Chemistry Study Cards

Here is a collection of study cards for my AP and General Chemistry classes. There are four cards per page. Each set of cards is saved as an Adobe Acrobat® file. Print each sheet on both the front and back of heavy paper to get two sets of study cards. Print one side of normal weight paper and fold paper lengthwise for single study cards.

Go to the Adobe Acrobat® website for a free software download if you do not have the plug-in.

A P C H E M I S T R Y Jump to General Chemistry

AP1-Introduction (20 cards on 5 pages) The Scientific Method, Observations and Measurements (Qualitative, Quantitative, Inferences), Significant Digits, Scientific Notation, Accuracy vs. Precision, Metric System, % and ppm, Unit Analysis, Temperature Scales, Mass vs. Weight, Potential Energy (PE) and Kinetic Energy (KE), Intensive vs. Extensive Properties, Calorimetry, Physical and Chemical Properties, Physical and Chemical Changes, Pure Substances, Elements, & Compounds; Homogeneous & Heterogeneous Mixtures; Separating Mixtures by Filtration, Distillation, and Chromatography; Early Laws: the Law of Definite Composition & the Law of Simple Multiple Proportions

AP2-Stoichiometry (24 cards on 6 pages) Formula Conventions, Stoichiometry Terms, Calculating Formula Mass, Mole Facts, Line Equations, Mole Relationships, Percentage Composition, Formula from % Composition, Equation Terms, Other Mole Problems and Conversions, Writing Formula Equations Things To Remember, Coefficients and Relative Volumes of Gases, Heart of the Problem, Mass-Mass Problems, Mass-Volume Problems, Mass-Volume-Particle Problems, Limiting Reactant Problems, How Much Excess Reactant is Left Over, Limiting Reactants, Theoretical Yield and Percentage Yield, Balancing Chemical Equations, Combustion Equations, Solutions -- Molar Concentration, Dilution Problems, Acid-Base Titrations

AP3-The Periodic Table (12 cards on 3 pages) The Subatomic Particles, Terms--Atomic Structure, Calculating Atomic Mass; Determining Numbers of Protons, Neutrons, and Electrons from the Isotopic Notation; Important People in the Development of the Atomic Theory; Metals, Nonmetals, and Metalloids; Rutherford's Gold Foil Experiment, The First Periodic Table, Families of the Periodic Table, Radioactivity Basics, J.J. Thomson & Cathode Ray Tubes

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Chemistry Study Cards

AP4-Electronic Structure (16 cards on 4 pages)

Wave Ideas You Should Know, Wave Calculations, The Balmer Series, The De Broglie Wavelength of Electrons, Standing Waves, Quantum Numbers (n, l, m, s), Three Rules for Filling Orbitals, Electron Configurations; The s-block, p-block... of the Periodic Table; Exceptions to the Filling Rules, Shapes of the Orbitals, Predicting the Atomic Size (radius), Trends in the Periodic Table, Explaining Sizes of Ions, The Lanthanide Contraction, Ionization Energy & Reactivity: Trends Across a Period & Down a Family, Ionization Energy: Trends in Successive Ionizations, Electron Affinity

AP5-Chemical Bonding--General (12 cards on 3 pages) Some Properties of Ionic and Molecular Compounds, Lewis Symbols of Atoms and Ions, The Ionic Bond, Noble and Pseudonoble Gas Configurations, Factors that Influence the Formation of Ionic Bonds, The Covalent Bond: Attractions and Repulsions, Groves' Electron Dot System: Multiple & Extended Valence Bonds; Bond Order: Bond Length, Strength, & Vibrational Frequency; Resonance, Coordinate Covalent Bonds, Electronegativity and Polar Bonds, Naming Ionic Compounds, Traditional and Stock Names, Naming Acids

AP6-Covalent Bonding & Molecular Structure Still gotta' do this one!

AP7-Chemical Reactions & the Periodic Table (16 cards on 4 pages) Solutions & Solubility, Weak and Strong Electrolytes, Ionic Reactions, Arrhenius Acids & Bases, Bronsted-Lowry Acids & Bases, Ions in Water, Some Metal Ions Make Water Acidic, Trends in Acid Strength, Lewis Acids & Bases, Oxidation Numbers, Balancing Redox Equations, Reactions in Basic Solutions, Metals as Reducing Agents, Activity Series of Metals, Non-Metals as Oxidizing Agents, Oxygen as an Oxidizing Agent, Combustion, Amphiprotic/Amphoteric & Leveling Effect

AP8-Ionic Reactions in Solution (12 cards on 3 pages) Driving Forces for Metathesis Reactions, Solubility Rules, Weak Electrolytes & Neutralization Reactions, Gas Formation During Metathesis, Preparation of Salts, Comparing Driving Forces; Weight Percent, ppm & ppb; Chemical Analyses: Precipitations, Combustions & Titrations; Titration Terminology, Acid-Base & Redox Titrations; Three Common Oxidizing Agents, Common Reducing Agents; equivalents, Equivalent Weights & Normality

AP9-The Gas Laws (12 cards on 3 pages) Boyle's Law, Boyle's Law Lab, Kelvin Temperature Scale, Charles' Law, Gay-Lussac's Law, The Combined Gas Law, The Ideal Gas Law, Dalton's Law of Partial Pressure, Why Do All Gases Cause the Same Pressure, Graham's Law, The Real Gas Law, Kinetic Molecular Theory, Pressure = Force/Area

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Chemistry Study Cards

AP10-States of Matter & IMF's (16 cards on 4 pages)

Comparing Gases, Liquid, and Solids, Surface Tension, KE Distributions & Evaporation, Molecular Crystals & IMF's, Hvap & IMF's, Vapor Pressures of Liquids, Boiling Point &

IMF's, Freezing Point, Melting Point, Hfusion, Common Crystal Structures and Units Cells,

Four Types of Crystals--A Summary, Crystal Types--Further Notes, Heating and Cooling Curves, Phase Diagrams--I, Phase Diagrams--II, Names of the Phase Changes, More Internet Resources

AP12-Chemical Thermodynamics (12 cards on 3 pages) Commonly Used Terms, Heat Capacity & Specific Heat, First Law of Thermodyanmics, Work & PV Work, Work is NOT a State Function, Reversible Processes, Change in Enthalpy, Hess's Law of Heat Summation, Bond Energy, Driving Forces: Entropy & Enthalpy, Second Law of Thermodynamics, Thermodynamics and Equilibrium

AP22-Nuclear Chemistry (16 cards on 4 pages) The People, Terms, Types of Radiation, Half-Life, Nuyclear Equations, Stabilizing Unstable Nuclei, Uses of Radioactivity, Fission and Fusion, Energy-Mass Conversion, What Happens During Beta and Positron Decay, Calculating Half-Lives, Radioactive Decay Series, Geiger-Muller Tubes, Smoke Detectors, and Brushes for Cleaning Negatives, Extending the Periodic Table

AP23-Organic Chemistry (16 cards on 4 pages) Historical Ideas, Alkane Series -- Saturated Hydrocarbons, Structural Formulas Can Be Misleading, Alkenes & Cis/Trans Isomerism; Alkynes, Alkadienes, and Cyclic Compounds; Naming Organic Compounds (IUPAC Rules), Common Errors in Drawing/Naming Structures, Optical Isomers / Chiral Compounds, Common Names You Should Know, Aromatic Compounds -- Benzene and Its Derivatives, Functional Groups I -- Alcohols & Ethers, Functional Groups II -- Aldehydes & Ketones, Functional Groups III -- Carboxylic Acids & Esters, Functional Groups IV -- Amines & Amides, Addition Polymerization, Condensation Polymerization

More To Come As soon as I get some more free time... stay tuned.

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Chemistry Study Cards

G E N E R A L C H E M I S T R Y Jump to AP Chemistry

GC1,2,3-Introduction, Measurement, & Prob. Solving (12 cards on 3 pages) The Scientific Method, Observations and Measurements (Qualitative, Quantitative, Inferences), Graphing--Great Graphs, Recongizing and Using Significant Digits, Scientific Notation, Accuracy vs. Precision, Metric System, % and ppm, Unit Analysis

GC4-Matter (8 cards on 2 pages) Mass & Weight; Pure Substances, Elements, Compounds & Mixtures; Homogeneous & Heterogeneous Mixtures, Separating Mixtures by Filtration, Distillation & Chromatography; Mass, Volume & Density; Intensive & Extensive Properties, Physical & Chemical Properties, Physical & Chemical Changes, Conservation of Mass, Symbols of the Elements, Relative Abundance of the Elements, Natural History of Airs Lab

GC5-Energy Study cards were never completed for this chapter... they are incorporated into other chapters.

GC6-Structure of the Atom (8 cards on 2 pages) Subatomic Particles, Terms, Calculating Atomic Mass; Determining Numbers of Protons, Neutrons & Electrons from Isotopic Notation; Important People, Early Experimental Observations, Rutherford's Gold Foil Experiment

GC7-Chemical Formulas (12 cards on 3 pages) Terms, Memorization Tips for Positive and Negative Ions, Formula Conventions, How Ions Form, Writing Ionic Formulas, More Ions & Oxidation Numbers, Determining Oxidation Numbers, Writing & Naming Acids, Stock Names vs. Traditional Names, Naming Molecular Compounds, Determining Ion Charges from Formulas

GC8-The Mathematics of Chemical Formulas (8 cards on 2 pages) Terms, Calculating Formula Mass, Mole Facts, Line Equations, Mole Relationships, Percent Composition, Formula from % Composition, Other Mole Problems

GC9-Chemical Equations (8 cards on 2 pages) Terms, Types of Reactions, Energy Changes, Showing Phases in Equations; Molecular, Ionic & Net Ionic Equations; Word Equations, Law of Conservation of Mass & Complete Combustion

GC10-The Mathematics of Chemical Equations (8 cards on 2 pages) Coefficients and Relative Volumes of Gases, Hearts of the Problem, Mass-Mass & Mass-Volume Problems, Limiting Reactant Problems, How Much Excess is Left, Baking Soda Lab Ideas

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Chemistry Study Cards

GC11-Phases of Matter (8 cards on 2 pages)

Pressure Definition & Units, Manometers & Measuring Pressure, heating Curve and Energy Changes, Names of Phase Changes & Energy, Kinetic Molecular Theory; Vapor Pressure, IMF's & Kinetic Energy; Boiling Point; Solids, Liquids & Gases

GC12-The Gas Laws (8 cards on 2 pages) Boyle's Law, Boyle's Law Lab, Kelvin Temperature Scale, Charles' Law, Gay-Lussac's Law, The Combined Gas Law, The Ideal Gas Law, Dalton's Law of Partial Pressure, Why Do All Gases Cause the Same Pressure, Graham's Law

GC13-Electron Configurations (4 cards on 1 page) Electron Energy Levels, Orbital Diagrams & Electron Configurations, Filling Orbitals, Valence Electrons

GC14-The Periodic Table (8 cards on 2 pages) Trends about the Periodic Table, People of the Periodic Table, Families (Groups) of the Table, Terms, Trends of the Periodic Table, Explaining the Size of Atoms and Ions, Explaining Ionization Energy, Isoelectronic Species, Clues for the Periodic Table

GC15-Chemical Bonding (8 cards on 2 pages) Three Types of Bonding, The Ionic Bond, The Covalent Bond, Lewis Dot Structures, The Bednarski Method of Drawing Dot Structures, Comparing Ionic & Molecular Substances, Electronegativity & Polar Bonds, Shapes & Polar Molecules

GC24-Organic Chemistry (12 cards on 3 pages) Historical Ideas, Alkane Series -- Saturated Hydrocarbons, Structural Formulas Can Be Misleading, Alkenes & Cis/Trans Isomerism; Alkynes, Alkadienes, and Cyclic Compounds; Naming Organic Compounds (IUPAC Rules), Common Errors in Drawing/Naming Structures, Aromatic Compounds -- Benzene and Its Derivatives, Functional Groups I -- Alcohols & Ethers, Functional Groups II -- Aldehydes & Ketones, Functional Groups III -- Carboxylic Acids & Esters, Functional Groups IV -- Amines & Amides --Four Extra Cards-- (4 cards on 1 page) Optical Isomers / Chiral Compounds, Common Names You Should Know, Addition Polymerization, Condensation Polymerization

GC26-Nuclear Chemistry (12 cards on 3 pages) The People, Terms, Types of Radiation, Half-Life, Nuyclear Equations, Stabilizing Unstable Nuclei, Uses of Radioactivity, Fission and Fusion, Energy -Mass Conversion, What Happens During Beta and Positron Decay

More To Come As soon as I get some more free time... stay tuned.

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Chemistry Study Cards

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1 • IntroductionThe Scientific Method

(1 of 20)

This is an attempt to state how scientists do science. It isnecessarily artificial. Here are MY five steps:

• Make observationsthe leaves on my plant are turning yellow

• State a Problem to be solvedhow can I get my plants healthy (non-yellow)

• Form a hypothesismaybe they need more water

• Conduct a controlled experimentwater plants TWICE a week instead of once a week

• Evaluate resultsif it works, good... if not, new hypothesis (sunlight?)

1 • IntroductionObservations and Measurements

Qualitative, Quantitative, Inferences(2 of 20)

Step 1 of the Scientific Method is Make Observations.These can be of general physical properties (color, smell,hardness, etc.) which are called qualitative observations.

These can be measurements which are called quantitativeobservations.

There are also statements that we commonly make based onobservations. “This beaker contains water” is an example.You infer (probably correctly) it is water because it is aclear, colorless liquid that came from the tap. Theobservations are that it is clear, it is colorless, it is a liquid,and it came from the tap.

1 • IntroductionSignificant Digits I

What do they mean?(3 of 20)

Consider: 16.82394 cmIn a measurement or a calculation, it is important to knowwhich digits of the reported number are significant.

That means… if the same measurement were repeated againand again, some of the numbers would be consistent andsome would simply be artifacts.

All of the digits that you are absolutely certain of plus onemore that is a judgment are significant.

If all the digits are significant above, everyone whomeasures the object will determine that it is 16.8239 cm, butsome will say …94 cm while others might say …95 cm.

1 • IntroductionSignificant Digits II

Some examples with rulers.(4 of 20)

1 2

a b c

(A composite ruler)a- No one should argue that the measurement is between 0.3and 0.4. Is it exactly halfway between (.35 cm)… or a littleto the left (.34 cm)? The last digit is the judgment of theperson making the measurement. The measurement has 2significant digits.b- The same ruler… so the measurement still goes to thehundredths place… 1.00 cm (3 significant digits).c- A ruler with fewer marks reads 1.6 cm (2 sig digits).

1 • IntroductionSignificant Digits III

Rules for Recognizing Sig. Digits(5 of 20)

In a number written with the correct number of sig. digits...• All non-zero digits are significant. 523 grams (3)• 0’s in the MIDDLE of a number are ALWAYS significant.

5082 meters (4) 0.002008 L (4)• 0’s in the FRONT of a number are NEVER significant.

0.0032 kg (2) 0.00000751 m (3)• 0’s at the END of a number are SOMETIMES significant.

• Decimal point is PRESENT, 0’s ARE significant2.000 Liters (4) 0.000500 grams (3)

• Decimal point is ABSENT, 0’s are NOT significant2000 Liters (1) 550 m (2)

NOTE: textbook values are assumed to have all sig. digits

1 • IntroductionScientific Notation

Useful for showing Significant Digits(6 of 20)

Scientific notation uses a number between 1 and 9.99 x 10 tosome power. It’s use stems from the use of slide rules.

Know how to put numbers into scientific notation:5392 = 5.392 x 103 0.000328 = 3.28 x 10–4

1.03 = 1.03 550 = 5.5 x 102

Some 0’s in numbers are placeholders and are not asignificant part of the measurement so they disappear whenwritten in sci. notation. Ex: 0.000328 above. In scientificnotation, only the three sig. digits (3.28) are written.

Scientific Notation can be used to show more sig. digits.Values like 550 ( 2 sig. digits) can be written 5.50 x 102 (3)

1 • IntroductionSignificant Digits IV

Significant Digits in Calculations(7 of 20)

When you perform a calculation using measurements, oftenthe calculator gives you an incorrect number of significantdigits. Here are the rules to follow to report your answers:

x and ÷: The answer has the same # of sig. digits as thenumber in the problem with the least number of sig. digits.example: 3.7 cm x 8.1 cm = 29.97 ≈ 30. cm2 (2 sig. digits)

+ and –: The last sig. digit in the answer is the largestuncertain digit in the values used in the problem.example: 3.7 cm + 8.1 cm = 11.8 cm (3 sig. digits)

Know how to ilustrate why these rules work.

1 • IntroductionAccuracy vs. Precision

(8 of 20)

Accuracy refers to how close a measurement is to someaccepted or true value (a standard).

Ex: an experimental value of the density of Al° is 2.69g/mL. The accepted value is 2.70 g/mL. Your value isaccurate to within 0.37%% error is used to express accuracy.

Precision refers to the reliability, repeatability, orconsistency of a measurement.

Ex: A value of 2.69 g/mL means that if you repeat themeasurement, you will get values that agree to thetenths place (2.68, 2.70, 2.71, etc.)± and sig. digits are used to express precision.

1 • IntroductionMetric System

(9 of 20)

We generally use three types of measurements:volume Liters (mL)length meters (km, cm and mm)mass grams (kg and mg)

We commonly use the prefixes:centi- 1/100thmilli- 1/1000thkilo- 1000

Occasionally you will encounter micro(µ), nano, pico, mega,and giga. You should know where to find these in chapter 1.Know that 2.54 cm = 1 inch and 2.20 lb = 1 kg

1 • Introduction% and ppm

(10 of 20)

Percentage is a mathematical tool to help compare values.Two fractions, 3/17 and 5/31 are difficult to compare:If we set up ratios so we can have a common denominator:

317

= x

100 =

17.65100

531

= x

100 =

16.13100

so… we can see that 317

> 531

.

There are 17.65 parts per 100 (Latin: parts per centum) or17.65 percent (17.65 %)… the % is a “1 0 0”

ppm (parts per million) is the same idea, (use 1,000,000

instead of 100) 317

= x

1 000 000 = 176,470 ppm

1 • IntroductionUnit Analysis

Converting between English and Metric Units(11 of 20)

Consider the metric/English math fact: 2.54 cm = 1 inchThis can be used as the “conversion factor”:

2.54 cm1 inch

or 1 inch

2.54 cm

You can convert 25.5 inches to cm in the following way:Given: 25.5 in

Desired: ? cm 25.5 in x 2.54 cm

1 in = 64.77 cm ≈ 64.8 cm

This is the required way to show your work. You have twojobs in this class, to be able to perform the conversions andto be able to prove that you know why the answer is correct.

1 • IntroductionTemperature Scales

(12 of 20)

The important idea is that temperature is really a measureof something, the average motion (kinetic energy,KE) of the molecules.

Does 0°C really mean 0 KE? nope... it simply means thefreezing point of water, a convenient standard.

We have to cool things down to –273.15°C before we reach0 KE. This is called 0 Kelvin (0 K, note: NO ° symbol.)

For phenomena that are proportional to the KE of theparticles (pressure of a gas, etc.) you must usetemperatures in K. K = °C + 273 °C = K – 273

1 • IntroductionMass vs. Weight

Theory, Measuring, Conversions(13 of 20)

mass is the amount of something...weight is how much gravity is pulling on the mass.(Weight will be proportional to the mass at a given spot.)

Mass is what we REALLY want to use... measured in grams.You use a balance to measure mass... you compare yourobject with objects of known mass.

Weight is measured with a scale (like your bathroom scaleor the scale at the grocery store). If there is no gravity, itdoesn’t work. Note: electronic balances are really scales!

You convert mass / weight using: 1 kg

2.205 lbs or

2.205 lbs1 kg

1 • IntroductionPotential Energy (PE) and

Kinetic Energy (KE)(14 of 20)

You can calculate the KE of an object: KE = 12mv2

m = mass, v = velocity [Note units: 1 J = 1 kg·m2·s–2]Temperature is a measure of the average kinetic energy.

PE = the potential to do work which is due to an object’sposition in a field. For example, if I hold a book 0.5 mabove a student’s head it can do some damage... 1.0 m aboveher/his head, more work can be done.

Important ideas:Objects tend to change from high PE to low PE (downhill).High PE is less stable than low PE.

1 • IntroductionMass, Volume, and Density

Intensive vs. Extensive Properties(15 of 20)

Extensive properties depend on the amount of substance.We measure these properties frequently... (mass &volume... mostly).

Intensive properties are independent of the size of thesample. These are useful for identifying substances...(melting point, boiling point, density, etc.)

It is interesting that an intensive property, density = mass

volumeis the ratio of two extensive properties... the size of thesample sort of “cancels out.” Be able to do density problems(3 variables) and know the usefulness of specific gravity.

1 • IntroductionCalorimetry

(16 of 20)

Heat is the total KE while temperature is the average KE.

A way to measure heat is to measure the temperature changeof a substance... often water. It takes 1 calorie of heatenergy (or 4.184 J) to heat 1 gram of H2O by 1 °C.

The specific heat of water = 1 cal

g·°C = 4.184

Jg·°C

heat = specific heat x mass H2O x ∆T

You can heat other substances as well, you just need toknow their specific heats. Notice that this is simplyheating or cooling a substance, not changing its phase.

1 • IntroductionPhysical and Chemical PropertiesPhysical and Chemical Changes

(17 of 20)

Equations to symbolize changes: reactants → products

Physical Properties can be measured from a sample of thesubstance alone... (density, MP, BP, color, etc.)

Chemical Properties are measured when a sample is mixedwith another chemical (reaction with acid, how does itburn in O2)

Physical Changes imply that no new substances are beingformed (melting, boiling, dissolving, etc.)

Chemical Changes imply the substance is forming newsubstances. This change is accompanied by heat, light,gas formation, color changes, etc.

1 • IntroductionPure Substances, Elements, & CompoundsHomogeneous & Heterogeneous Mixtures

(18 of 20)

Pure Substances

Matter Energy

CompoundsElements

Mixtures

HeterogeneousHomogeneous

This chart should help you sort out these similar terms.Be able to use chemical symbols to represent elements andcompounds. For example...CuSO4•5H2O, a hydrate, contains 21 atoms & 4 elements.

Memorize the 7 elements that exist in diatomic molecules:HONClBrIF or BrINClHOF or “H and the 6 that make a 7starting with element #7”

1 • IntroductionSeparating Mixtures by Filtration,Distillation, and Chromatography

(19 of 20)

Mixtures are substances the are NOT chemically combined...so if you want to separate them, you need to exploitdifferences in their PHYSICAL properties.

Filtration:some components of the mixture dissolve and some donot. The filtrate is what passes through the filter.

Distillation:some components vaporize at different temperatures orone component may not vaporize at all (e.g.: salt+water)complete separation may not be possible.

Chromatography:differences in solubility vs. adhesion to the substrate.Substaratemay be filter paper (paper chromatography),or other substances, GLC, TLC, HPLC, column, etc.

1 • IntroductionEarly Laws: the Law of Definite Composition

& the Law of Simple Multiple Proportions(20 of 20)

Definite Composition:samples of the same substance from various sources(e.g. water) can be broken down to give the same %’s ofelements. Calculation: percent composition

Multiple Proportions:samples of 2 substances made of the same 2 elements...(e.g. CO2 & CO or H2O and H2O2 or CH4 and C3H8)if you break down each to give equal masses of oneelement, the masses of the other element will be in asimple, whole-number ratio.Calculation: proportions to get equal amounts of oneelement and then simple ratios.

2•Stoichiometry: Chemical ArithmeticFormula Conventions

(1 of 24)

Superscriptsused to show the charges on ionsMg2+ the 2 means a 2+ charge (lost 2 electrons)

Subscriptsused to show numbers of atoms in a formula unitH2SO4 two H’s, one S, and 4 O’s

Coefficientsused to show the number of formula units2Br– the 2 means two individual bromide ions

Hydrates CuSO4 • 5 H2Osome compounds have water molecules included

2•Stoichiometry: Chemical ArithmeticStoichiometry Terms

(2 of 24)

stoichiometry study of the quantitative relationshipsin chemical formulas and equations.

atomic mass weighted average mass of an atom,found on the periodic table

formula mass sum of the atomic masses of theatoms in a formula

molecular mass sum of the atomic masses of theatoms in a molecular formula

gram molecular mass molecular mass written in grams

molar mass same as gram molecular mass

empirical formula formula reduced to lowest terms

2•Stoichiometry: Chemical ArithmeticCalculating Formula Mass

(3 of 24)

Formula or molecular mass is found by simply summingthe atomic masses (on the periodic table) of each atom in aformula.

H2SO41.01 + 1.01 + 32.06 + 16.0 + 16.0 + 16.0 + 16.0 = 98.08 u2(1.01) + 32.06 + 4(16.0) = 98.06 u or 98.06 g/mole

Generally, round off your answers to the hundredths ortenths place. Don’t round off too much (98.06 g/mol or98.1 g/mol is OK, but don’t round off to 98 g/mol)

UnitsUse u or amu if you are referring to one atom or molecule

2•Stoichiometry: Chemical ArithmeticMole Facts

(4 of 24)

A mole (abbreviated mol) is a certain number of things. Itis sometimes called the chemist’s dozen.A dozen is 12 things, a mole is 6.02 x 1023 things.

Avogadro’s Number1 mole of any substance contains 6.02 x 1023 molecules

Molar Volume (measured at P = 760 mmHg and T = 0 °C)1 mole of any gas has a volume of 22.4 Liters

Molar Mass (see gram formula mass)

1 mole

6.02 x 1023 molecules

1 mole22.4 L

1 mole

molar mass

2•Stoichiometry: Chemical ArithmeticLine Equations

(5 of 24)

A Line Equation is the preferred way to show conversionsbetween quantities (amount, mass, volume, and number) bycanceling units (moles, grams, liters, and molecules)

The line equation consists of the Given Value, the DesiredUnit, and the line equation itself.

Example: What is the mass of 135 Liters of CH4 (at STP)? Given: 135 L CH4 Desired: ? g CH4

135 L CH4 x 1 mol CH422.4 L CH4

x 16.0 g CH4 1 mol CH4

= 96.43 g CH4

2•Stoichiometry: Chemical ArithmeticMole Relationships

(6 of 24)

The “Mole Map” shows the structure of mole problems

Mass

Volume at STP

number ofatoms ormolecules

Mass

moles Volume at STP

number of atoms or molecules

➂ ➂➁ ➁

➀ ➀

1) 1 mol

molar mass 2)

1 mol22.4 L

3) 1 mol

6.02 x 1023 molecules

2•Stoichiometry: Chemical ArithmeticPercentage Composition (by mass)

(7 of 24)

Percentage Composition quantifies what portion (by mass)of a substance is made up of each element.

Set up a fraction: mass of element

mass of molecule

Change to percentage: 100 x mass of element

mass of molecule

Generally, round off your answers to the tenth’s place.

The percentage compositions of each element should add upto 100% (or very close, like 99.9% or 100.1%)

2•Stoichiometry: Chemical ArithmeticFormula from % Composition

(8 of 24)

Given the Percentage Composition of a formula, you cancalculate the empirical formula of the substance.Step 1 assume you have 100 g of substance so

the percentages become gramsStep 2 change grams of each element to moles

of atoms of that elementStep 3 set up a formula with the moles

example: C2.4 H4.8Step 4 simplify the formula by dividing moles

by the smallest value C2.42.4 H

4.82.4 = CH2

Step 5 If ratio becomes… 1:1.5 multiply by 21:1.33 or 1:1.66 multiply by 3

2•Stoichiometry: Chemical ArithmeticEquation Terms

(9 of 24)

equation condensed statement of facts about achemical reaction.

reactants → substances that exist before achemical rxn. Written left of arrow.

→ products substances that come into existence asa result of the reaction. Written to theright of the arrow.

word equation an equation describing a chemicalchange using the names of thereactants and products.

coefficients a number preceding atoms, ions, ormolecules in balanced chemicalequationns that showing relative #’s.

2•Stoichiometry: Chemical ArithmeticOther Mole Problems and Conversions

(10 of 24)

The gas density is often converted to molar mass:

Example :The gas density of a gas is 3.165 g/Liter (at STP). What isthe molar mass of the gas?

Knowing that 22.4 L is 1 mole, you can set up the ratio:3.165 g1 Liter

= molar mass

22.4 L

Other metric conversions you should know:1000 mL1 Liter

1 kg

1000 grams

2•Stoichiometry: Chemical ArithmeticWriting Formula Equations

Things To Remember(11 of 24)

Example: Write the formula equation of...sodium metal + water → sodium hydroxide + hydrogen gas

Na° + H2O → NaOH + H2

• metals often are written with the ° symbol to emphasizethat the metal is in the neutral elemental state, not an ion.

• some compounds have common names that you shouldjust know... water, H2O; ammonia, NH3; methane, CH4

• remember the seven diatomic elements so they can bewritten as diatomic molecules when they appear in theirelemental form. Other elemental substances are writtenas single atoms (e.g. sodium metal or helium gas, He)

2•Stoichiometry: Chemical ArithmeticCoefficients and Relative Volumes of Gases

(12 of 24)

Since every gas takes up the same amount of room (22.4 Lfor a mole of a gas at STP), the coefficients in an equationtell you about the volumes of gas involved.

Example: N2(g) + 3 H2(g) → 2 NH3(g)

+

2•Stoichiometry: Chemical ArithmeticHeart of the Problem

(13 of 24)

The “heart of the problem” conversion factor relates theGiven and the Desired compounds using the coefficientsfrom the balanced equation.

Example: N2 + 3 H2 → 2 NH3

♥ could be 3 moles H2

2 moles NH3

…which means that every time 2 moles of NH3 is formed, 3moles of H2 must react.

The format is always, moles of Desired moles of Given

2•Stoichiometry: Chemical ArithmeticMass-Mass Problems

Mass-Volume Problems(14 of 24)

Mass-Mass problems are probably the most common typeof problem. The Given and Desired are both masses (gramsor kg).

The pattern is:

Given x molar massof Given x ♥ x

molar massof Desired

In Mass-Volume problems, one of the molar masses is

replaced with 22.4 L1 mole

depending on whether the Given or

the Desired is Liters.

2•Stoichiometry: Chemical ArithmeticMass-Volume-Particle Problems

(15 of 24)

If the Given or Desired is molecules, then the Avogadro’s

Number conversion factor, 6.02 x 1023 molecules

1 mole is used

and the problem is a Mass-Particle or Volume-Particleproblem.

The units of the Given and Desired will guide you as towhich conversion factor to use:

Mass grams or kgVolume Liters or mLParticles molecules or atoms

2•Stoichiometry: Chemical ArithmeticLimiting Reactant Problems

(16 of 24)

In a problem with two Given values, one of the Given’s willlimit how much product you can make. This is called thelimiting reactant. The other reactant is said to be in excess.

Solve the problem twice using each Given… the reactantthat results in the smaller amount of product is the limitingreactant and the smaller answer is the true answer.

Example: N2 + 3 H2 → 2 NH3When 28.0 grams of N2 reacts with 8.00 grams of H2, whatmass of NH3 is produced?

(in this case, the N2 is the limiting reactant)

2•Stoichiometry: Chemical ArithmeticHow Much Excess Reactant is Left Over

(17 of 24)

Example: N2 + 3 H2 → 2 NH3When 28.0 grams of N2 reacts with 8.00 grams of H2, whatmass of NH3 is produced?

(in this case, the N2 is the limiting reactant)

To find out how much H2 is left over, do another lineequation:

Given: 28.0 g N2Desired: ? g H2

subtract the answer of this problem from 8.00 g H2

2•Stoichiometry: Chemical ArithmeticLimiting Reactants

(18 of 24)

It is difficult to simply guess which reactant is the limitingreactant because it depends on two things:

(1) the molar mass of the reactant and(2) the coefficients in the balanced equation

The smaller mass is not always the limiting reactant.

Example: N2 + 3 H2 → 2 NH31 mole (28 g N2) will just react with 3 moles (6.06 g H2)

so, if we react 28.0 g N2 with 8.0 g H2, only 6.06 g H2will be used up and 1.94 g of H2 will be left over.

In this case, N2 is the L.R. and H2 is in X.S.

2•Stoichiometry: Chemical ArithmeticTheoretical Yield and Percentage Yield

(19 of 24)

The answer you calculate from a stoichiometry problem canbe called the Theoretical Yield. Theoretically , youshould get this amount of product.

In reality , you often get less than the theoretical amount dueto products turning back to reactants or side reactions.The amount you actually get is called the Actual Yield.

Percentage Yield = Actual Yield

Theoretical Yield x 100

2•Stoichiometry: Chemical ArithmeticBalancing Chemical Equations

(20 of 24)

The balanced equation represents what actually occursduring a chemical reaction. Since atoms are not createdor destroyed during a normal chemical reaction, thenumber and kinds of atoms must agree on the left andright sides of the arrows.

__Na2CO3 + __HCl → __ NaCl + __H2O + __CO2

To balance the equation, you are only allowed to change thecoefficients in front of the substances... not change theformulas of the substances themselves.

Reduce the coefficients to the lowest terms.

Fractions may be used in front of diatomic elements.

2•Stoichiometry: Chemical ArithmeticCombustion Equations

(21 of 24)

The burning of fuels made of C, H, and O is calledcombustion. You need to memorize O2, CO2 and H2O

Example: The combustion of propane, C3H8, is written:

C3H8 + 5O2 → 3CO2 + 4H2O

Be careful when writing equations for alcohols, such asbutanol, C4H9OH• don’t forget to add the H’s (a total of 10 of them)• don’t forget to take account of the O atom in the alcohol

C4H9OH + 6O2 → 4CO2 + 5H2O

2•Stoichiometry: Chemical ArithmeticSolutions -- Molar Concentration

(22 of 24)

Many reactions are carried out in solution. Solutions areconvenient and speed up many reactions.

Concentration is often expressed as

Molarity ( M ) = moles of solute

Liters of solution

You can calculate the molarity of a solution when givenmoles (or grams) of a substance and its volume.

You can use the molarity of a solution as a conversion factor

0.150 M HCl ≈ 0.150 moles HCl

1 Liter HCl or

1 Liter HCl0.150 moles HCl

to convert moles to Liters and vice versa.Volumetric flasks are used to make solutions.

2•Stoichiometry: Chemical ArithmeticDilution Problems

(23 of 24)

You can calculate the moles of a solute using the volumeand molarity of the substance. Since diluting a solutionadds water and no solute, the moles of solute before andafter the dilution remains constant. So...

Vi · M i = Vf · M f

where “I” means “initial” and “f” means “final”

The units of volume or concentration do not really matter aslong as they match on the two sides of the equation.

2•Stoichiometry: Chemical ArithmeticAcid-Base Titrations

(24 of 24)

Acids form the H+ ion. Bases form the OH– ion.Acids + bases mix to form H2O (HOH) and a salt.

The moles of H+ = the moles of OH– in a neutralization.

An acid-base titration is the technique of carefullyneutralizing an acid with a base and measuring thevolumes used. An indicator (we used phenolphthalein)allows us to observe when the endpoint is reached.

If a monoprotic acid is neutralized with a base that only hasone OH– ion per formula unit, the simple formula:

Va · M a = Vb · M ballows you to determine the molarity of the unknown.

3 • The Periodic Table & Makeup of AtomsThe Subatomic Particles

(1 of 12)

Name Symbol Mass Charge Locationprotons p+ 1 u 1+ part of the nucleusneutron n° 1 u 0 part of the nucleus

electron e– 11837

u 1– normally at large

distances from thenucleus

J.J. Thompson is given credit for discovering electronsusing a Crookes tube and testing many different gases.Cathode rays were found to be beams of electrons.

Chadwick is given credit for the discovery of the neutron.

3 • The Periodic Table & Makeup of AtomsTerms I-- Atomic Structure

(2 of 12)

atoms the smallest particle of an element . Itconsists of a central nucleus and electronclouds outside the nucleus.

nucleus the dense central portion of an atom.

subatomic smaller than an atom. The proton,neutron, and electron are subatomic.

net charge the difference in the positive charge due toprotons and the negative charge due toelectrons in an atom.

nucleons the particles that make up the nucleus.

3 • The Periodic Table & Makeup of AtomsTerms II-- Atomic Structure

(3 of 12)

atomic number the number of protons in an atom. This #determines the identity of an element.

mass number the number of protons + neutrons

isotopes atoms with the same number of protons,but different numbers of neutrons. Atomswith the same atomic number, butdifferent mass numbers.

isotopic notation shorthand notation for a nucleus thatshows the mass #, atomic # and the

symbol. U-238 would be 238 92U

3 • The Periodic Table & Makeup of AtomsCalculating Atomic Mass

(4 of 12)

Any real sample of an element contains more than onenaturally occuring isotope. For instance, boron

isotope abundance mass # isotopic mass

boron-1010 5B 19.78% 10 mass = 10.013 u

boron-1111 5B 80.22% 11 mass = 11.009 u

The atomic mass is the weighted average of the isotopes.

at. mass = (19.78%)(10.013u) + (80.22%)(11.0009u)

100 or

at. mass = (0.1978)(10.013u) + (0.8022)(11.0009u)= 10.81 u

3 • The Periodic Table & Makeup of AtomsDetermining Numbers of Protons, Neutrons,

and Electrons from the Isotopic Notation(5 of 12)

Consider the following symbol: 3316S

2–

The 16 is the atomic number which is the number ofprotons.

The 33 is the mass number which is the mass of one of theisotopes. This mass is due to the protons and neutrons.

The number of neutrons is the mass number - the atomicnumber. 33 - 16 = 17 neutrons.

Since the charge is 2-, there are 2 more electrons thanprotons. In this case, there are 18 electrons.

3 • The Periodic Table & Makeup of AtomsImportant People in the Development of the

Atomic Theory(6 of 12)

Democritus [atomos]philosopher who argued that matter was discontinuous

John Dalton [billiard-ball model]experimented with gases… different substances aredifferent combinations of atoms

J.J. Thomson [plum-pudding model]experimented with gas-discharge tubes… atoms have + and– parts… the negative e–’s are the same for any atom

Ernest Rutherford [nuclear model/solar system model]most of the mass of the atom is concentrated in a tiny,positively-charged nucleus

Niels Bohr [quantized electron energy levels]the electrons have only certain allowed energy levels

3 • The Periodic Table & Makeup of AtomsMetals, Nonmetals, and Metalloids

(7 of 12)

The elements can be classified as metals, nonmetals, andmetalloids. Memorize the elements classified as metalloids(also called semi-metals or semiconductors).

Properties of metals include:ductility - the ability to pull a substance into a wiresectility - the ability to cut with a knifemalleability - the ability to pound substance into a sheetconductivity - the ability to carry an electrical current

Gold is the most malleable of all the metals.

3 • The Periodic Table & Makeup of AtomsRutherford’s Gold Foil Experiment

(8 of 12)

Ernest Rutherford’s classic gold foil experiment led to thenuclear model of the atom.

αα

most alpha'scame straightthrough here

a few bounced backat a large angle

gold foil

• the nucleus is tiny - because most of the alpha’s missedthe nucleus and went straight through the foil• the nucleus is positively charged - because the (+)charged alpha was repelled by the (+) charged nucleus• the nucleus is incredibly dense - because the nucleus wasable to bounce back at a very large angle

3 • The Periodic Table & Makeup of AtomsThe First Periodic Table

(9 of 12)

Meyer and Mendeleev are given credit for developing thefirst version of the periodic table. Mendelleev’s true claimto fame was that he actually predicted the existence ofseveral element that had not been discovered. He foundgaps in the table when he tried to organize the atoms and leftspaces for those elements (ekasilicon = "like silicon", etc.)

He predicted Ga, Ge, and Sc.

He also arranged elements in order of atomic number ratherthan the previous idea of atomic mass. Several of theelements change order... (like Te and I).

3 • The Periodic Table & Makeup of AtomsFamilies of the Periodic Table

(10 of 12)

Horizontal rows of the table are called periods.Vertical columns are called groups or families.

Memorize the names of some groups:IA - the alkali metalsIIA - the alkaline earth metalsVIIA - the halogens0 - the noble gases

Also know the transition metals, the inner transition metals(composed of the lanthanide series and the actinide series...the lanthanides are also called the rare earth metals)

Families IA - VIIIA are called the representative elements.

3 • The Periodic Table & Makeup of AtomsRadioactivity Basics

(11 of 12)

Radioactivity was discovered by Henri Becquerel (butnamed by Marie Curie).

“Becquerel rays” were found to consist of 3 types orradiation:alpha particles (α) a helium nucleus - 2 protons + 2 neutrons…easily stopped by paper or skin

beta particles (β) a high energy electron…stopped by Al foil (several thicknesses of foil)

gamma radiation (γ) a very high energy form of light (EMR)…the most penetrating and dangerous of the rays.

3 • The Periodic Table & Makeup of AtomsJJ Thomson & cathode ray tubes

(12 of 12)

Know the design of a cathode rays tubes. Realize thatcathode rays are really beams of electrons. The cathode raysare the same for any substance, but the canal rays (thepositive ions left after ionizing the gases) are different foreach gas.

Know how the bending of cathode rays can tell you thecharge-to-mass ratio (e/m) (but not the mass or the chargeof the electron).

Millikan’s oil drop experiment gave evidence for thecharge of the electron. Knowing this and the e/m ratio, youcan calculate the mass of the electron.

4 • Electronic Structure & the Per. TableWave Ideas You Should Know

(1 of 16)

EMR electromagnetic radiation … oscillating electric & magnetic fields at right angleswavelength (λ) the distance from crest to crest or

trough to trough.amplitude the distance from the equilibrium

point to the crest or trough.frequency (ν) the number of waves that pass a

point per second. (Hz, s–1, 1/s)continuous spectrum a “normal” rainbow that contains all

of the colors (ROYGBV).line spectrum a spectrum that only contains certain

bright lines that result from electrontransitions within an atom.

4 • Electronic Structure & the Per. TableWave Calculations

(2 of 16)

Exchanging wavelength, frequency, & energy of light:λ · ν = c c = 3.00 x 108 m/s [speed of light]

E = hν h = 6.626 x 10–34 J·s [Planck’s constant]

E = h cλ

Energy of Level “n” in the Hydrogen atom:

En = – An2 A = 2.18 x 10–18 J [Arrhenius constant?]

NOTE: You can learn the Balmer equation or the Rydbergequation listed on pp. 108 & 109. I suggest you use theabove equations to calculate the wavelengths or frequenciesof light emitted when electrons change energy levels.

4 • Electronic Structure & the Per. TableThe Balmer Series

(3 of 16)

When the electron “drops” to energy level n=2, visible lightis emitted. The bight line spectrum observed is called theBalmer Series of lines. [Memorize this info...]3 → 2 red4 → 2 blue-green5 → 2 blue-violet6 → 2 violet

All the lines result from an electron transition to level n=1are too high energy to be visible… UV [Lyman Series]All the lines that result from an electron transition to levelsn=3 [Paschen], n=4 [Brackett] and n=5 [Pfund] are tooLOW energy to be visible.

4 • Electronic Structure & the Per. TableThe De Broglie Wavelength of Electrons

(4 of 16)

Two equations can be combined into one to allow you tocalculate the wavelength of a particle:

E = mc2 E = h cλ

mc2 = h cλ

mc = hλ or more generally, mv =

hλ

λ = h

mv

h = Planck’s constant, m = mass of particle (in kg),v = velocity of particle (in m/s).

4 • Electronic Structure & the Per. TableStanding Waves

(5 of 16)

Standing waves are something that waves do…resulting from having repeating waves that always cancel atthe nodes and always add up at the antinodes.

We have seen standing wavesof strings: of sound:

of drumheads: + + +– –

We hypothesized the standing wavesof electrons: (4-dimensional vibrating wavicles)

the orbitals (s, p, d, f… probability waves)

4 • Electronic Structure & the Per. TableQuantum Numbers (n, l, m, s)

(6 of 16)

These are variables in some unseen equation:The Rules:

n = 1, 2, 3, ... determines the energy of the e–

l = 0 → (n–1) the type of orbital (subshell)0 ≈ s, 1 ≈ p, 2 ≈ d, 3 ≈ f

m = –l → +l which orientation of the orbital(x, y, z… for p orbitals)

s = +12

or –12

the “spin” of the electron

NOTE: m is also called ml and s is ms

4 • Electronic Structure & the Per. TableThree Rules for Filling Orbitals

(7 of 16)

The Pauli Exclusion Principle states that no two electronsin an atom may have the same four quantum numbers. Thistranslates to the idea that an orbital may contain no morethan two electrons.

The Aufbau Principle states that electrons occupy thelowest energy available orbital. You must memorize theorbital chart on page 120. [Aufbau = “building up”]

Hund’s Rule states that when you have several orbitals ofthe same energy (e.g. three p orbitals or five d orbitals) placeone electron in each orbital before doubling them up.

NOTE: Remember the Aufbau hotel analogy...

4 • Electronic Structure & the Per. TableElectron Configurations

(8 of 16)

This is a shorthand notation for the arrangement of electronin the orbitals. 1s2 means 2 electrons in the 1s orbital

Consider Arsenic, AsThis is the order in which the electrons fill…

As 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p3

however, we write the orbitals according to the distancefrom the nucleus (n=3’s then n=4’s, etc.)

As 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p3

long form: (shown above) show each subshellshort form: (show the last filled NRG level...noble gas core)

As [Ar] 3d10 4s2 4p3

4 • Electronic Structure & the Per. TableThe s-block, p-block... of the Periodic Table

Exceptions to the Filling Rules

(9 of 16)

The shape of the Periodic Table comes from the way theelectrons fill the orbitals.

s-block Families I and IIp-block Families III, IV, V, VI, VII, and VIIId-block The Transition Elementsf-block The Inner Transition Elements

Six elements do not follow the rules. They are in the d-block of the periodic table… the transition elements.

Cu, Ag, and Au... instead of a full s-orbital and almost filledd-orbital, they have a filled d and half-filled s-orbitalCr, Mo, W... instead of a full s-orbital and an almost half-filled d-orbital, they have d5 and s1

4 • Electronic Structure & the Per. TableShapes of the orbitals

(10 of 16)

Know the general shapes of thes (spherical) orbitals,

p (perpendicular) orbitals, and

d (diagonal) orbitals.

4 • Electronic Structure & the Per. TablePredicting the Atomic Size (radius)

Trends in the Periodic Table(11 of 16)

What: As you move ACROSS a period, the size of the atomDECREASES

Why: As you move ACROSS a period, the number ofprotons in the nucleus increases... so the proton-electron attraction increases... the size decreases.

NOTE: The increased e–-e– repulsion that one might expectis not important, because the outer electrons do notSHIELD the electrons from the nucleus..

What: As you move DOWN a family, the size of the atomINCREASES

Why: As the value of “n” increases, the average distance ofthe electron from the nucleus increases. The size ofthe atom IS the electron cloud... so the size increases

4 • Electronic Structure & the Per. TableExplaining The Sizes of IonsThe Lanthanide Contraction

(12 of 16)

The change in size of ions (compared to the neutral element)depends on the electron-electron repulsion.

If an atom gains electrons, the increased repulsion increasesthe size of the electron cloud. So... negative ions are largerthan the neutral atom.

Positive ions form by losing electrons (less repulsion) andget smaller.

The Lanthanide contraction makes the third row of thetransition elements about the same size as the second row...This causes these elements to be much more dense.

4 • Electronic Structure & the Per. TableIonization Energy

Trends Across a Period(13 of 16)

Following the decrease in size, the ionization energyGENERALLY increases as you move across a period. If theatom is smaller, the electron being removed is closer to thenucleus and therefore feels a stronger attraction.

The Jags: (why are some electrons EASIER to remove)In Family III, the electron being removed comes from the p-orbital rather than the s-orbital. The p-orbital electron is at ahigher energy and requires LESS energy to ionize.

In Family VI, the electron being removed comes from anorbital with TWO electrons. The e–- e– repulsion felt by theelectron allows it to be ionized with less added energy.

4 • Electronic Structure & the Per. TableIonization Energy & Reactivity

Trends Down a Family(14 of 16)

What: Down a family, Ionization energy DECREASESWhy: The trend here goes right with the size of the atom.

Since “n” increases, the distance of an electronfrom the nucleus increases and is easier to remove.

What: Down family I and II, the reactivity INCREASES.Why: These families lose electrons to become + ions.

The easier it is (lower ionization energy) the morereactive they are.

What: Down family VII, the reactivity DECREASES.Why: These elements GAIN electrons to become – ions.

Smaller atoms mean the attracted electrons will becloser to the nucleus... more effective attractions.

4 • Electronic Structure & the Per. TableIonization Energy

Trends in Successive Ionizations(15 of 16)

What: You can tell the family of an element by observingwhen the ionization energies get very large.For example: family III elementsAl + energy = Al+ + e– (3p electron)Al+ + energy! = Al2+ + e– (3s electron)Al2+ + energy!! = Al3+ + e– (3s electron)Al3+ + ENERGY = Al4+ + e– (2p electron!)

Why: The “easy” electrons to remove are in an orbitalwith a higher value of n. When n decreases, theaverage distance of the electron from the nucleusdecreases and the attractions between the protonand electrons increase.

4 • Electronic Structure & the Per. TableElectron Affinity

(16 of 16)

Electron affinity is the energy involved when an atom gainsan electron to become a negative ion.

F + e– → F– + energy

Elements in the upper right corner of the periodic table havethe greatest electron affinity (greatest attraction forelectrons).

The electron affinity may be + or –. Negative values meanenergy is released and also counts as a greater electronaffinity.

Electron affinity data is not complete, but it gives SOMEevidence for trends in the periodic table.

5 • Chemical Bonding: Gen ConceptsSome Properties of

Ionic and Molecular Compounds(1 of 12)

Compound Molecular IonicConducts as Solid NO NOConducts as Liquid NO YESConducts in Solution NO YESConducts as Gas NO YES

Hardness soft hard

MP / BP low high

Bonding covalent ionic

Examples He, CH4, CO2, NaCl, KI,C6H12O6 AgNO3

5 • Chemical Bonding: Gen ConceptsLewis Symbols of Atoms and Ions

(2 of 12)

Lewis symbols conisist of the atomic symbol surrounded byvalnece electrons. The four sides represent the four valenceorbitals. Atoms are usually shown in their excited states (II,III, IV)

Li Be B C N O F Ne

Ions include brackets. Positive ions show no valenceelectrons while negative ions usually have an octet.

[Li] [Mg] [ O ]2+ 2–+

5 • Chemical Bonding: Gen ConceptsThe Ionic Bond

Noble and Pseudonoble Gas Configurations(3 of 12)

Many ions can be explained because they have gained or lostelectrons and attain a noble gas configuration. For example:

P3– S2– Cl– Ar K+ Ca2+

all have the same electron arrangement: 1s2 2s2 2p6 3s2 3p6

A pseudonoble gas configuration is:1s2 2s2 2p6 3s2 3p6 3d10

This is found in Cu+ Zn2+ Ga3+ and Ge4+

Similar configurations are found in the next two periods.

The importance of this configuration is that there is morethan one reason why ions form what they do. Many ions arenot explained.

5 • Chemical Bonding: Gen ConceptsFactors that Influence

the Formation of Ionic Bonds(4 of 12)

Know the 5 steps that can be thought to occur when an ionicbond forms. Note whether each is exo- or endothermic...whether a larger energy helps or hinders the bond formation.

Overall: Li(s) + 1/2F2(g) → LiF(s)

1. heat of vaporization Li(s) + NRG → Li(g)2. heat of decomposition 1/2F2(g) + NRG → F(g)3. ionization energy Li(g) + NRG → Li+(g) + e–

4. electron affinity F(g) + e– → F–(g)5. lattice energy Li+(g) + F–(g) → LiF(s)

Large energy values for 1,2,3 hinder ionic bond formation.

5 • Chemical Bonding: Gen ConceptsThe Covalent Bond

Attractions and Repulsions(5 of 12)

The covalent bond between two atoms depends on thebalance of attractions between one atom’s + nucleus andthe other atom’s – electrons and the proton-protonrepulsions as well as electron-electron repulstions.

Distance between nuclei

PE

If two atoms have half-filled orbitals , the interactionsbalance at a small enough distanc e so the e–’s can be closeto both nucle i at the same time... this is a covalen t bond.

5 • Chemical Bonding: Gen ConceptsGroves’ Electron Dot System

Multiple & Extended Valence Bonds(6 of 12)

Count up your valence electrons.

Give every atom who “wants” and octet an octet.[the first 5 elements do not need octets... too small][Family I, II, and III do not form octets]

If you have drawn too man y electrons...“Take away a lone pair... take away a lone pair...

make these two atoms share”

If you have drawn too fe w electrons... place the extraelectrons on the central atom (extended valence shell)

5 • Chemical Bonding: Gen ConceptsBond Order: Bond Length, Strength, &

Vibrational Frequency(7 of 12)

Bond orde r is the number of pairs of electrons bonding twoatoms together.

single bond bond order = 1double bond bond order = 2triple bond bond order = 3

single bonds have the longest bond lengthsingle bonds have the weakest bond strengthsingle bonds have the lowest vibrational frequency

(think of single bonds as soft, springy springs...triple bonds are tight springs...sproinnnnng)

Bonds in resonance structure must be averaged... the S-Obond in SO2 has a bond order of 1.5. C-O in CO32– is 1.33

5 • Chemical Bonding: Gen ConceptsResonance

(8 of 12)

When you draw a Lewis structure (SO2, O3, CO32–, etc.) inwhich you must make a choice as to who gets a doublebond, the structure is actually a blend of two or threestructures.

We “say” that the structure “resonate s” or we say that thestructure contains contributions from each of the resonancestructures.

Resonance occurs simply because the electron-dot model(while very useful) is too limited to show how the electronsare being shared between the atoms... wait for π bonding.

5 • Chemical Bonding: Gen ConceptsCoordinate Covalent Bonds

(Preview: Lewis Acids)(9 of 12)

Coordinate covalent bond: When a covalent bond isformed by sharing a pair of electrons BUT the electron pairbelonged to only one of the atoms.

Classic Example: NH3 + BF3 → NH3BF3The bond between the N and the B is coordinate covalent.

The lone pair donor is called a Lewis Base.(this atom has a lone pair of electrons)

The lone pair acceptor is called a Lewis Acid.(this atom has an empty orbital)

“Have Pair Will Share” --Lewis Base

5 • Chemical Bonding: Gen ConceptsElectronegativity and Polar Bonds

(10 of 12)

You will be given a chart of electronegativity values.Memorize the most electronegative element s (F = 4.0)then oxygen (O = 3.5) and chlorine (Cl = 3.0). The noblegases have no electronegativity values… no bonds.Trend is large electronegativity in the upper right of theper. table and small in the lower left portion of the table.Classify the bond between any two atoms by subtractingtheir electronegativity values (∆e)

Non-polar covalent 0 < ∆e < 0.5Polar covalent 0.5 ≤ ∆e ≤ 1.7Ionic ∆e > 1.7

The more electronegative atom is more negative.Polar covalent bonds have partial charges δ+ and δ–

5 • Chemical Bonding: Gen ConceptsNaming Ionic Compounds

Traditional and Stock Names(11 of 12)

The Stock System of naming compounds is used…• when a positive ion has more than one possible chargeTraditional: mercurous, Hg22+ mercuric, Hg2+

Stock: mercury(I) mercury(II)Traditional: cuprous, Cu+ cupric, Cu2+

Stock: copper(I) copper(II)

• for molecular compounds where the elements have manydifferent oxidation states (i.e. N in NO2, NO, N2O, etc.)

Stock Name: Traditional Name:NO2 nitrogen(IV) oxide nitrogen dioxideNO nitrogen (II) oxide nitrogen monoxideN2O nitrogen(I) oxide dinitrogen monoxide

5 • Chemical Bonding: Gen ConceptsNaming Acids

(12 of 12)

Acids are ionic formulas in which the positive ion is H+.Use as many H+ ions as the charge on the negative ion.

Three rules for naming: if the anion ends with: the acid is named:

–ite ********ous acid–ate ********ic acid–ide hydro********ic acid

• Acids from sulfide, sulfite, and sulfate include a “ur” H2S is hydrosulfuric acid, not hydrosulfic acid• Acids from phosphate and phosphite include a “or” H3PO4 is phophoric acid, not phosphic acid

7 • Chemical Reactions & the Per. TableSolutions and Solubility

(1 of 16)

We learned the terms solute, solvent, and solution.Solubility (how MUCH solute will dissolve) is measured ing/100 mL of H2O or sometimes in M.This information may be given numerically or graphically.A saturated solution is one in which any additional soluteadded will simply settle on the bottom of the container.An unsaturated solution is any amount less than saturated.Supersaturated implies that the solution was saturated atsome higher temperature and then carefully cooled. Thisunstable situation can be changed with a “seed” crystal.Recall the supersaturated solution of NaC2H3O2 demo.The terms concentrated and dilute refer to the amount ofsolute and do not necessarily coincide with saturation.

7 • Chemical Reactions & the Per. TableElectrolytes: Weak and Strong

(2 of 16)

Solutions of acids, bases, and salts contain mobile ions andconduct electricity. These solutions are called electrolytes.Salts are ionic compounds that dissociate in water.Acids are actually molecular compounds (covalentlybonded) the become ions when dissolved in water.Only 8 acids are strong electrolytes and completelydissociate when dissolved. All others dissolve completely,but only partially dissociate into ions.Only 8 bases are strong electrolytes because they dissolvecompletely. All others have low solubility and remainsolids rather than dissolve. One common exception is theweak base NH4OH . It dissolves, but partially dissociates.

7 • Chemical Reactions & the Per. TableIonic Reactions

(3 of 16)

A common class of chemical reactions occurs when twoionic solutions are mixed. The double replacement ormetathesis reaction involves the formation of two newcombinations of ions.AgNO3 + NaCl → AgCl + NaNO3 (molecular equation)The new combinations may be more stable than the originaldue to low solubility (precipitate forms), weak electrolyte,gas formation, or complex ion formation.The reaction is written above as though the substances existas molecules. This is the easiest time to balance.The ionic equation shows strong electrolytes as separateions. The net equation eliminates “spectator ions”.

7 • Chemical Reactions & the Per. TableArrhenius Acids and Bases

(4 of 16)

Acid: a substance that increases the [H+] in solution.Base: a substance that increases the [OH–] in solution.Diprotic acids have more than one removable H. (H2SO4)However, only the first H is ever easily dissociated.Oxides of nonmetals (SO2) are acid anhydrides.Oxides of metals (Na2O)are basic anhydrides.

Just add water... to get the acid or base.SO2 + H2O → H2SO3 Na2O + H2O → 2 NaOHAcids and bases neutralize each other because the H+ andOH– ions form the very weak electrolyte... H2O (and a salt).Acid salts are the partially neutralized polyprotic acids.NaH2PO4 or Na2HPO4 or NaHSO4, etc... solid acids.

7 • Chemical Reactions & the Per. TableBrønsted-Lowry Acids and Bases

(5 of 16)

Acid: a proton (H+) donor. Base: a proton acceptor.This is a more general definition of acids and bases becauseit does not require the substance to be dissolved in water.Consider the following equations:The species that accepted the proton (base) can beconsidered a donor (conjugate acid) in the reverse reaction.

HF + H2O ⇔ H3O+ + F–

(acid) (base) (acid) (base)NH3 + H2O ⇔ NH4+ + OH–

(base) (acid) (acid) (base)Strong base≈weak conjugate acid, (good acceptor≈lousydonor). Conjugates differ by only a H+ (e.g. HF and F–)

7 • Chemical Reactions & the Per. TableIons in Water

Some Metal Ions Make Acidic Solutions(6 of 16)

Since water molecules are polar, they surround ions insolution (called hydration). When we write Na+(aq) andCl–(aq) we are implying this more complex situation.Some highly charged ions (Al3+, Cr3+, Fe3+) tend to tightlybind the water molecules. We can write them as complexions: Al(H2O)63+

The electron clouds are drawn toward the central ion andaway from the oxygen and therefore the O-H bond. Thisextra-polar O-H bond results in the H atom more readilyjoining with a passing water molecule... making the solutionacidic. [Note: this is a conjugate acid-base situation.]

Al(H2O)63+ + H2O ⇔ H3O+ + Al(H2O)5OH2+

7 • Chemical Reactions & the Per. TableTrend in Strengths of Acids and Bases

Three cases to explain(7 of 16)

Case 1: The more oxygens on an oxoacid, the stronger theacid. H2SO4 > H2SO3.. HClO4 > HClO3 > HClO2 > HClOWhy?: The electronegative oxygens draw electron densityaway from the central atom and therefore from the H-Obond... making it more polar. H leaves more easily.Case 2: The more electronegative the central atom, thestronger the oxoacid. H3PO4 < H2SO4 < HClO4

Why? Same as above... the more electronegative atom inthe center makes the O-H bond more polar.Case 3: Binary acid strength depends on the SIZE of theatom... HF < HCl < HBr < HI... not the electronegativity.Why?: The greater distance means a weaker attraction.

7 • Chemical Reactions & the Per. TableLewis Acids and Bases

(8 of 16)

Consider the formation of a coordinate covalent bond:Acid: electron pair acceptor [Note: proton donor]Base: electron pair donor [Note: proton acceptor]This definition is more general than the other two becausethis covers cases that don’t even involve hydrogen (protons).Classic case: HN3 + BF3 → NH3BF3 Other important cases:

oxide of metal + oxide of nonmetal → salt(base anhydride + acid anhydride)

CaO + SO2 → CaSO3(s)and the oft-confusing reactivity of CO2

OH– + CO2 → HCO3–

7 • Chemical Reactions & the Per. TableOxidation Numbers

(9 of 16)

Definition:Oxidation numbers are the apparent charges atoms have ifshared e–’s are assigned to the more electronegative atom.You can assign ox. #’s by studying a Lewis diagram or...The Rules:…ox. # of neutral atoms is 0In compounds……simple ions have ox. #’s equal to their charge.…F (-1), Family I (+1), II (+2), Al (+3)…O is usually (-2) except peroxide and OF2…H is usually (+1) except hydrides of Fam. I, II, Al…the sum of the ox. #’s of individual atoms equals thecharge on the entire species.

7 • Chemical Reactions & the Per. TableBalancing Redox Equations

Oxidation Number Change MethodStep-By-Step

(10 of 16)

1. Identify the ox. #’s of elements that change PER ATOM.One element will change UP (oxidation / lose e–’s) oneelement changes DOWN (reduction / gain e–’s).

2. Adjust coefficients and e– changes for situations wheremore than one atom MUST change:ex: 2HCl + K2Cr2O7 → KCl + 2CrCl3 + Cl2 + H2OCr: 2 x (3 e– per Cr) Cl: 2 x (1 e– per Cl)

3. Balance these changes in e–’s (e– gained = e– lost).4. Balance all elements except H and O.5. Balance O’s (add H2O’s as needed).6. Balance H’s (add H+’s as needed).7. If solution is basic, see card #12.

7 • Chemical Reactions & the Per. TableBalancing Redox Equations

Ion-Electron Method(Half-reaction Method)

Step-By-Step(11 of 16)

1. Identify the substances involved in the oxidation andreduction changes. Include substances so that allelements are represented (except H and O).

For each half-reaction…2. Balance all elements except H and O.3. Balance O’s (add H2O’s as needed).4. Balance H’s (add H+’s as needed).5. Balance charges (add e–’s to the more positive side).6. Balance e–’s in the two half-reactions.7. Combine the two half-reactions. Cancel substances that

show up on both sides of the equation (e–’s must cancel).8. If solution is basic, see card #12.

7 • Chemical Reactions & the Per. TableBalancing Redox EquationsReactions in Basic Solution

(12 of 16)

If the reaction occurs in a basic solution (usually statedclearly in the problem) then instead of H+’s and H2O’s, youutilize OH–’s and H2O’s.

An easy method is:- balance as though the reaction were in an acid solution.- add OH–’s to each side of the equation until all H+’s areturned into H2O’s.- cancel H2O’s that show on both sides of the equation.

7 • Chemical Reactions & the Per. TableMetals as Reducing Agents

(13 of 16)

In each case, the metal changes to the + ion: M → M+ + e–

Since the metal is oxidized, it is a reducing agent.Metals that most easily lose e–’s (those with low ionizationenergy and low electronegativity) make the best reducers.

M E M O R I Z E T H I S:

Some metals react with H2O [to H2(g) and OH–]Some react with non-oxidizing acids such as HCl, and cold,dilute H2SO4 (the H+ is the reacting species) [to H2(g)]Some react only with oxidizing acids:• dilute HNO3 [to colorless NO(g) + H2O],• conc. HNO3 [to red-brown NO2(g) and H2O], and• hot, conc. H2SO4 [to SO2(g) and H2O]

7 • Chemical Reactions & the Per. TableThe Activity Series of Metals

(14 of 16)

F O U R G R O U P S O F M E T A L S

1 - Most active - Families I and II - great reducing agentsreduction half-reaction: 2H2O + 2e– → H2 + 2OH–

2 - Most metals... Zn, Fe, Al, etc.reduction half-reaction: 2H+ + 2e– → H2

3 - Ag, Cu, Hgex. half-reaction: 2H+ + NO3– + e–→ NO2 + H2O

4 - Noble Metals - Au, Pt, Ironly changed by “aqua regia” [HNO3 + HCl forms Cl2]

A more active metal can reduce or displace the ion of a lessactive metal. Ex. Zn° is more active than Cu°, so…

Zn° + Cu2+ → Zn2+ + Cu° Cu° + Zn2+ → no reaction

7 • Chemical Reactions & the Per. TableNonmetals as Oxidizing AgentsOxygen as an Oxidizing Agent

Combustion as a Redox Reaction(15 of 16)

Elemental nonmetals (O, Cl, F, S, etc.) form negative ionsby gaining electrons (reduction) and are oxidizing agents.

The strongest oxidizing agents are those to the right of eachperiod (excluding noble gases) and those at the top of eachfamily. So we can predict that F>Cl>Br>I and O > S.

Example: Cl will displace Br–, but not F–

Cl2(g) + 2NaBr → Br2(l) + 2NaCl • Cl2 + NaF → N.R.

Oxygen, O2, is a common and powerful oxidizing agent.Corrosion of metals, formation of oxides, and combustionare all examples of redox with O2 as the oxidizer.

7 • Chemical Reactions & the Per. TableAmphiprotic/Amphoteric & Leveling Effect

(16 of 16)

Anything with a lone pair can act as a proton acceptor(base) Anything with a H atom can act as a proton donor(acid). See water on card 5 act as either an acid or a base.See acetic acid on page 241 act as both acid and base. Wesay these are amphiprotic or amphoteric substances.Leveling effect...You cannot tell which strong acid is strongest in water,because donating a proton to water is not a good enoughchallenge. Strong acids completely dissociate in waterbecause water is a pretty good acceptor of protons. We saythat water has a leveling effect. Acetic acid is amphiprotic,but a poor proton acceptor and therefore is a great test forwhich strong acid is the strongest...

8 • Ionic Reactions in SolutionDriving Forces for Metathesis Reactions

(1 of 12)

During a double replacement or metathesis reaction, twonew combinations of ions are produced. We identify fourreasons why these NEW combinations are more stablethan the original combos.

• a precipitate formsmemorize your solubility rules

• a gas forms which leaves the systemmemorize the list of gases that form

• a weak electrolyte formsmemorize the strong acid list so you will recognize weakacids, also H2O and NH4OH

• a complex ion formslearn the structure of complex ions and common ligands

8 • Ionic Reactions in SolutionPrecipitates as a Driving Force

The Solubility Rules(2 of 12)

Always Always Soluble compounds with alkali metal ions(Li+, Na+, K+, Cs+, Rb+), NH4+, NO3–, C2H3O2–, ClO3–

& ClO4–

Usually SolubleCl–, Br–, I– [except “AP/H”... Ag+, Pb2+, Hg22+]SO42– [except “CBS”: Ca2+, Ba2+, Sr2+ & “PBS”: Pb2+]

Usually NOT SolubleO2–, OH– [except alkali and “CBS” Ca2+, Ba2+, Sr2+ ]

Never SolubleCO32–, SO32–, S2–, PO43– [except NH4+ & alkali]

NOTE: some of these insoluble compounds WILL dissolvein acid solutions because of gas formation... useful idea!

8 • Ionic Reactions in SolutionWeak Electrolyte Formation as a Driving Force

Weak Acids and other Weak ElectrolytesNeutralization Reactions

(3 of 12)

Weak electrolytes are substances that break up into ions onlya LITTLE in solution... therefore, the two ions are MOSTLYin a combined state... not likely to re-form the reactants.

H2O, weak acids, NH4OHMemorize the 8 strong acids so you can recognize a weakacid when you see one...

HCl, HBr, HI, HNO3, H2SO4, HClO3, HClO4, HIO4

Acids (forming H+ ions) and bases (forming OH– ions)combine to form a salt (an ionic compound) and H2O... thevery weak electrolyte. Neutralization of the acid and baseoccur because the H+ & OH– ions are “tied up” as H2O.

8 • Ionic Reactions in SolutionGas Formation as a Driving Force

Gases That Commonly Form(4 of 12)

If you see the following substances formed duringmetathesis, realize that they will breakup into gases andleave the system (preventing re-formation of the reactants).

Watch For… It Turns Into…H2CO3 → CO2(g) + H2OH2SO3 → SO2(g) + H2OH2S → H2S(g) [rotten egg smell]NH4OH → NH3(g) + H2O2 HNO2 → NO(g) + NO2(g) + H2O

NOTE: these compounds are formed from acids withcarbonates, sulfites, sulfides, nitrites, and bases withammonium compounds.

8 • Ionic Reactions in SolutionPreparation of Salts

(5 of 12)

If you want to make the ionic compound, XY, you mixAY + XB to make XY + AB

Either XY or AB need to drive the reaction (ppt, gas, etc.)

You may need to do this in two steps... make a carbonate (orsulfite, sulfide, hydroxide or oxide) of the cation (+ ion) youneed and then react it with an acid that has the proper anion.

The Practical Side:Keep in mind how you could recover the product you want...could you filter the product mixture? Do you want what isin the filter paper or what is in the filtrate? If you need thefiltrate, you need to be careful not to have excess ions in it.

8 • Ionic Reactions in SolutionComparing Driving Forces

(6 of 12)

Gas formation is a very strong driving force... evencompounds that exist as insoluble solids will react (slowly)to form gases because gases leave the system and CANNOTre-form reactants.

CaCO3(s) + 2HCl → CO2(g) + H2O + Ca2+ + 2 Cl–

The tendency to form H2O is very strong. Insoluble oxideswill react with acids.

ZnO(s) + 2HCl → Zn2+ + 2Cl– + H2OSometimes, one insoluble solid can change into another evenMORE insoluble solid... but you need more than thesolubility rules to predict this (you need Ksp’s).

AgCl(s) + Br– → AgBr(s) + Cl–

8 • Ionic Reactions in SolutionMore Concentration Units

Weight Percent, ppm, and ppb(7 of 12)

Percent means “parts per 100”

96% means 96100

, that is, 96 out of every 100.

Weight Percent (w/w) means 96g100g

… (w/v) means 96g

100mL

ppm means “parts per million”

96 ppm means 96

1 000 000 , 96 out of every 1 million.

ppb means “parts per billion”

96 ppb means 96

1 000 000 000, 96 out of every BILLION!

8 • Ionic Reactions in SolutionChemical Analyses

Precipitations, Combustions,and Titrations

(8 of 12)

Real chemistry often deals with testing what is in aparticular reaction mixture, environmental sample, etc.Stoichiometry is used to analyze the compositions.

• You can precipitate an ion you are interested in, filter theprecipitate and then determine from its mass the amountof compound in the original sample.

• You can burn a sample and collect the combustionproducts (CO2 & H2O) to determine the amount of C andH in the original sample.

• You can carefully measure the volumes of solutions usedduring a titration. The endpoint must have some sort ofindicator to allow you to recognize when the correctamounts of reactants have been added.

8 • Ionic Reactions in SolutionTitration Terminology

Acid-Base and Redox Titrations(9 of 12)

A titration is a volumetric analysis because you carefullymeasure the volume of titrant, dispensing it from a buret.When you have added just enough titrant to completely reactwith the sample, you have reached the endpoint. This isusually apparent because of the color change of someindicator molecule (such as phenolphthalein). The endpointcan also be tracked because of changes in pH or changes involtage due to the amount of some ion.

Acid-Base & Redox titrations follow the formula: V·N=V·Nwhere the N indicates [H+] or [OH–] in acid-base titrationsand [Oxidizer]· e– gained or [Reducer] · e– lost in redoxtitrations.

8 • Ionic Reactions in SolutionThree Most Common Oxidizing Agents

and what they turn into(10 of 12)

purple permanganate ionacid solution MnO4– + 8H+ 5e– → Mn2+ + 4H2O

Mn2+ ion is colorlessneutral/basic MnO4– + 2H2O + 3e– → MnO2(s) + 4OH–

MnO2(s) is a black solidyellow chromate / orange dichromate ion depends on [H+]

2CrO42– + 2H+ →← Cr2O72– + H2Oacid solution Cr2O72– + 14H+ + 6e– → 2Cr3+ + 7H2Oslightly basic CrO42– + 4 H2O + 3e– → Cr(OH)3 + 5OH–

Cr(OH)3 is a solidvery basic CrO42– + 2H2O + 3e– → CrO2– + 4OH–

8 • Ionic Reactions in SolutionCommon Reducing Agents

and what they turn into(11 of 12)

Tin(II) (a gentle reducing agent)Sn2+ → Sn4+ + 2 e–

Sulfites and Bisulfitesacidic solution: HSO3– + H2O → SO42– + 3H+ + 2e–

basic solution: SO32– + 2OH– → SO42– + H2O + 2e–

Thiosulfate ion (also called “hypo” in photography)strong oxidizer: S2O32– + 5H2O → 2SO42– + 10H+ + 8e–

half-reaction w/I2: 2S2O32– → S4O62–+ 2e–

complete: I2 + 2S2O32– → 2I– + S4O62–

excess I–: I2 + I– → I3–

I2 + starch →← starch•I2 complex (blue-black)

8 • Ionic Reactions in SolutionEquivalents, Equivalent Weights,

and Normality(an old-fashioned, but useful idea)

(12 of 12)

equivalents = H+, OH–, or electrons gained or lostn = # of equivalents in the balanced chemical equationexample: I2 + 2S2O32– → 2I– + S4O62– n = 2equivalent weight = molar mass ÷ n…mass of a chemical that provides 1 mole of equivalents.

Normality, N, = n · M, = moles equivalents

Liter solution

This idea is useful in acid-base and redox titrations becauseit takes into account the differences of acids and bases oroxidizers and reducers. This concept allows the use of thesimple formula: V·N = V·N

9 • Properties of GasesBoyle’s Law (P and V)

(1 of 12)

General: When P↑, V↓ (inversely proportional)Formula: P·V = constant or P1V1 = P2V2

Restrictions: P1 and P2 must be in the same unitsV1 and V2 must be in the same units

Convert pressures using conversion factors using the factthat 1 atm = 760 mmHg = 760 torr = 101.3 kPa = 14.7 psi

psi = lbin2

Example: 730 mmHg x 101.3 kPa760 mmHg

= 97.3 kPa

9 • Properties of GasesBoyle’s Law Lab

(2 of 12)

Graphically:

P

V

P

1/V

In our lab, we had to add the atmospheric pressure to ourmeasurements because tire gauges only measure thepressure ABOVE atmospheric pressure.

Consistent (“ good”) data form a straight line (P vs. 1V

).

9 • Properties of GasesKelvin Temperature Scale

(3 of 12)

K = °C + 273 °C = K – 273Examples: 0 °C + 273 = 237 K

25 °C + 273 = 298 K100 °C + 273 = 373 K

300 K – 273 = –27 °C

The Kelvin scale is used in gas law problems because thepressure and volume of a gas depend on the kinetic energyor motion of the particles.

The Kelvin scale is proportional to the KE of theparticles… that is, 0 K (absolute zero) means 0 kineticenergy. 0 °C is simply the freezing point of water.

9 • Properties of GasesCharles’ Law (V and T)

Gay-Lussac’s Law (P and T)(4 of 12)

Charles’ LawGeneral: When T↑, V↑ (directly proportional)

Formula:VT

= constant or V1T1

= V2T2

Restrictions: T must be in KelvinsV1 and V2 must be in the same units

Gay-Lussac’s LawGeneral: When T↑, P↑ (directly proportional)

Formula:PT

= constant or P1T1

= P2T2

Restrictions: T must be in KelvinsP1 and P2 must be in the same units

9 • Properties of GasesThe Combined Gas Law

(5 of 12)

Formula:P·VT

= constant or P1·V1

T1 =

P2·V2T2

Restrictions: T must be in KelvinsV1 and V2 must be in the same unitsP1 and P2 must be in the same units

STP (“standard temperature and pressure”) is often used asone of the two conditions

T = 0 °C = 273 K P = 1 atm = 760 mmHg = 101.3 kPa

Each of the three gas laws is really a special case of thislaw.

Example: If T1 = T2, the law becomes P1V1 = P2V2

9 • Properties of GasesThe Ideal Gas Law

(6 of 12)

Formula: P·V = n·R·T or PV = nRTwhere P = pressure

V = volumen = number of molesR = the ideal gas constantT = temperature (in Kelvins)

The value of R depends on the P and V units used.

R = PVnT

so you can use the molar volume info to calculate R

R = (101.3 kPa)(22.4 L)

(1 mole)(273 K) = 8.31

L·kPamol·K

R = 62.4 L·mmHgmol·K

= 0.0821 L·atmmol·K

9 • Properties of GasesDalton’s Law of Partial Pressure

(7 of 12)

When you have a mixture of gases, you can determine thepressure exerted by each gas separately. This is calledthe partial pressure of each gas.

Since each gas has the same power to cause pressure (seecard #8) the partial pressure of a gas depends on howmuch of the mixture is composed of each gas (in moles)

Example: Consider air, a mixture of mostly O2 and N2moles O2

moles total =

PO2Ptotal

moles N2

moles total =

PN2Ptotal

Also: Ptotal = PO2 + PN2This idea is used when a gas is collected over water

Patm = Pgas + PH2O PH2O is found on a chart

9 • Properties of GasesWhy Do All Gases Cause the Same Pressure?

(8 of 12)

The gas laws work (to 3 significant digits) for all gases…that is, all gases have the same power to cause pressure.

At the same temperature, the KE of each gas is the same.KE = 1/2 mass·velocity2… if two particles have different

masses, their velocities are also different. So…

SMALL particles move FAST mv2

LARGE particles move SLOWLY mv2

We can use this idea with numbers as well: (Graham’s Law)KEA = KEB mAvA2 = mBvB2

[another version of this formula is on the next card]

9 • Properties of GasesGraham’s Law of Effusion

(9 of 12)

mAvA2 = mBvB2 can also be used as the equation…

rate of effusion of A rate of effusion of B =

MBMA

Notice that the A is in the numerator in the ratio of therates and in the denominator in the radical.

“Effusion” is similar to diffusion. It means to escapethrough a small opening.

The ratio of the rates (or velocities) of CH4 (mass=16 u) to

SO2 (mass=64 u) is 6416

= 4 = 2

9 • Properties of GasesThe Real Gas Law

(10 of 12)

Ideal gases have no volume & no attractions for eachother. Luckily, real gases act pretty much like ideal gases atroom temperature and pressure. The most ideal of real gasesis He.

The REAL GAS Law is:

(P + n2aV2 )(V - nb) = nRT

where:a corresponds to the attractions between real gas particlesb corresponds to the size of the real gas particle

9 • Properties of GasesKinetic Molecular Theory

(11 of 12)

Explaining the behavior of gases involves the kineticmolecular theory. Here are the main ideas:

• all particles are in constant, random motion• temperature is a measure of the average kinetic energy• pressure is due to collisions of gas particles with the walls

of the container• increased temperature causes more collisons as well as

harder collisions• some particles are moving fast, some are moving slowly

9 • Properties of GasesPressure = Force ÷ Area

(12 of 12)

P = FA

Pressure is proportional to the force pushing and inverselyproportional to the area over which that force pushes.

P = FA

P = FA

10 • States of Matter & IMFs Comparing Gases, Liquids, and Solids

(1 of 16)

Properties of matter depend on the model that gas particles are spread out; liquids are close together, but random; and solids are close together and arranged in a crystal lattice.

gas liquid solid

Volume and shape, compressibility, and the ability of substances to diffuse depend on these models. Gases have no set shape or volume. Liquids have a constant volume, but no set shape. Solids have constant volume and shape.

10 • States of Matter & IMFs Surface Tension

(2 of 16)

Surface tension- a measure of the amount of energy needed to expand the surface area of a liquid.

An interior molecule is surrounded by molecules to which it is attracted… no net attraction. A surface molecule feels a net attraction toward the interior. To move a molecule to the surface (i.e. increase the surface area), energy must be used, work must be done. The potential energy of the liquid is increased.

Substances tend toward the lowest potential energy so liquids tend toward the minimum surface area. A sphere is the smallest surface area for given volume.

10 • States of Matter & IMFs KE Distributions and Evaporation

(3 of 16)

In any sample of liquid, the distribution of KE varies. Particles to the right of the line (the “threshold energy” have enough KE to escape the IMF’s holding them in the liquid.

Increasing the temperature (average KE) of the liquid moves the curve to the right. The line depends on the IMF of the liquid.

Only particles at the surface of the liquid may escape (evaporate.)

10 • States of Matter & IMFs Molecular Crystals and IMFs

(4 of 16)

Substances that exist as molecules (as opposed to ionic, metallic, or covalent network crystals) are in three groups:

nonpolar molecules London Forces

polar molecules Dipole-dipole attractions

polar with H-O-, H-N-, or H-F H-bonding

Weak IMF. Due to polarizable e- clouds & temp. attraction

+ end of one molecule attracting - end of other molecule

Strong dipole because of high electroneg. /small size of O, N, & F

In London Forces--larger atoms and larger molecules have stronger London forces due to more sites or more polarizable electron clouds .

10 • States of Matter & IMFs Hvap and IMFs

(5 of 16)

Heat of Vaporization, ∆∆ Hvap, can be thought of as the energy needed to vaporize a mole of a liquid. It can be used as a conversion factor in a calculation of heat during a phase change. [Calorimetry is used for temperature changes between the phase changes.]

Ex: kJ 250.OH mol 1

kJ 40.6OH g 18.0OH mol 1

OH 111g22

22 =××

It can also serve as a indicator of the strength of the IMF (intermolecular forces of attraction) in the liquid.

Ex: CH4 (9.20 kJ/mol) vs. C3H8 (18.1 kJ/mol) Larger molecule… greater IMF… greater Hvap

10 • States of Matter & IMFs Vapor Pressures of Liquids

(6 of 16)

In a closed container, the number of particles changing from liquid →→ vapor will eventually equal the number of particles changing from vapor →→ liquid.

The amount of vapor when this balance is reached depends on the IMF and the KE of the liquid (& not on the volume of the container). The pressure exerted by this vapor is called the equilibrium vapor pressure (VP) of the liquid.

As temperature increases, the VP increases . (This is important for why/when a liquid boils.)

VP is another indicator of the strength of the IMF. The stronger the IMF, the smaller the VP.

10 • States of Matter & IMFs Boiling Point and IMFs

(7 of 16)

Boiling occurs when the vapor pressure (VP) of a substance = the air pressure above the liquid.

You can boil a liquid by increasing the VP of the liquid (heating) or by lowering the pressure above the liquid.

The temperature at which a liquid reaches 760 mmHg is called the “normal boiling point” of the liquid.

---Again, BP is an indicator of IMF.--- ↑↑ boiling points (BP) ~ ↑↑ IMF’s ~ ↓↓ vapor pressures.

Altitude (low air pressure) lowers the boiling temperature of water in an open container (increases cooking time). Pressure cookers ↑ BP by ↑ the pressure above the liquid.

10 • States of Matter & IMFs Freezing Point, Melting Point, ∆∆ Hfusion

(8 of 16)

Freezing Point and Melting Point are the same temperature, just opposite directions. That is, a substance will freeze and melt at the same temperature.

When you fuse two metals, you MELT them… thus the term fusion means melting. ∆∆ Hfusion is the energy needed to melt a mole of solid into a mole of liquid.

As with ∆Hvap, ∆Hfus can be used as a conversion factor as well as an indicator of IMF strength.

Note: Freezing is more complicated than vaporization because the process of forming a crystal causes some subtle considerations which we will not deal with in this course.

10 • States of Matter & IMFs Common Crystal Structures and Unit Cells

(9 of 16)

Three common crystal structures are Simple Cubic; Body-Centered Cubic (which means there is an atom, “a body,” in the center of the cubic structure; and Face-Centered Cubic (with an atom in the center of each side or “face”).

From Dr. John Gelder’s

Solid State Chemistry Page

A unit cell is the theoretical arrangement of atoms that, if repeated, will recreate the crystal.

This topic, although interesting, is no longer on the AP curriculum and will not be dealt with here. One resource is a page by OSU chemist, Dr. John Gelder. (See Card 16)

10 • States of Matter & IMFs Four Types of Crystals -- Summary

(10 of 16)

Crystal Molecular Ionic Covalent Metallic lattice points:

molecules or atoms

+ and - ions atoms positive ions

IMF’s London, dipole,

H-bonding

attraction between +

& - ions

covalent bonds

attr. between + ions & “sea

of e-‘s” Props. soft, low

MP, non-conduct

hard, brittle, high MP, (l) (aq) conduct

v. hard, high MP,

nonconduct (graphite)

high luster, conductor,

variable MP, soft/hard

Ex I2 H2O HI NaBr C SiC WC Na° Fe° Cu°

10 • States of Matter & IMFs Crystal Types -- Further Notes

(11 of 16)

Students are often confused between molecular crystals in which covalent bonds hold the molecules together (but the IMF = London forces, dipole-dipole attractions or hydrogen bonding) and covalent crystals in which covalent bonds hold the crystal together (the IMF = covalent bonds).

Substances can conduct electricity for two reasons: freely moving ions or delocalized electrons.

Ionic compounds have freely ions in the liquid state and when dissolved in water. Metals have delocalized electrons -- the “sea of electrons.” Graphite has a chicken-wire shaped π-bond above & below each sheet of sp2-hybridized C atoms, allowing it to conduct.

10 • States of Matter & IMFs Heating and Cooling Curves

(12 of 16)

Time (min)KE

PE(s)

(l)

(g)

(s) & (l)

(l) & (g)

KE

KE

PEKE = kinetic energy changes which are times when the heat energy speeds up the molecules.

PE = potential energy changes which are times when the heat energy separates the molecules from solid to liquid or liquid to gas.

10 • States of Matter & IMFs Phase Diagrams-I

(13 of 16)

The phase diagram shows the phases of a substance at all temperatures and pressures.

Moving across the diagram gives you the MP and then BP of a substance.

There is a point above which it is no longer possible to liquefy a substance (the critical point, C). Moving vertically you can see the effect of pressure on the phase of the substance. This is a diagram for a substance like CO2, in which the liquid can be compressed into the solid. (Unlike H2O.)

10 • States of Matter & IMFs Phase Diagrams-II

(14 of 16)

Water’s phase diagram is unique because the liquid phase is less dense than the solid phase. To maximize hydrogen bonding, the solid must expand.

The B-D boundary of the phase diagram has a negative slope.

The “triple point” is the temperature and pressure in which the solid, liquid, and vapor phases of a substance can co-exist. I visualize this as a boiling glass of ice water.

By increasing the pressure, dry ice can melt. By decreasing the pressure, solid water can sublime.

10 • States of Matter & IMFs Name of the Phase Changes

(15 of 16)

NRG is REQUIRED NRG is RELEASED solid →→ liquid liquid →→ solid melting or fusion freezing

liquid →→ gas gas →→ liquid vaporization condensation evaporation or boiling

solid →→ gas gas →→ solid sublimation solidification

The energy involved in the phase change is calculated using heat of fusion (solid →→ liquid or liquid →→ solid) heat of vaporization (liquid →→ gas or gas →→ liquid)

10 • States of Matter & IMFs More Internet Resources

(16 of 16)

Searching the Internet, I found an interesting set of topic reviews. These are from Purdue University (Indiana)

I was looking at the topic, LIQUIDS, but there are many topics to choose from. chemed.chem.purdue.edu/genchem/topicreview/ The unit cell is a frame from an online movie file on Dr. John Gelder’s Solid State Chemistry page. (OK State Univ.) www.okstate.edu/jgelder/solstate.html

12 • Chemical Thermodynamics Commonly Used Terms

(1 of 12)

thermodynamics Study of NRG changes & flow of NRG. Tells whether a reaction is possible.

system That portion of the universe on which we focus.

surroundings Everything outside the system. adiabatic A change without heat transfer between

the system and its surroundings. isothermal A change that occurs at constant

temperature. state functions Properties that depend on the initial and

final state, not on how the change was made. ex: ∆H, temp, but not work

12 • Chemical Thermodynamics Heat Capacity and Specific Heat

(2 of 12)

Energy changes can be measured using calorimetry. Often this involves heating water under controlled conditions (a bomb calorimeter).

Three closely related terms are: heat capacity is the amount of heat needed to change a system by 1°C. molar heat capacity is the amount of heat needed to change a mole of a substance by 1°C. specific heat is the amount of heat needed to change 1 gram of a substance by 1°C. (Water: 1 cal/g°C = 4.184 J/g°C)

Note that heat capacity is an extensive property whereas the other two are intensive.

12 • Chemical Thermodynamics First Law of Thermodynamics

(3 of 12)

The first law: “if a system undergoes some series of changes that ultimately brings it back to its original state, the net energy change is zero.”

∆E = Efinal - Einitial ∆E = 0 when Efinal = Einitial

The usefulness of this idea is that the internal energy, ∆E, depends only on the init ial and final state, not on how you get there. Every path you take from Einitial to Efinal takes the same amount of energy. (no perpetual motion)

There are two ways for a system to exchange energy with the surroundings, heat and work. ∆E = q - w [q = heat absorbed by system, w = work done by system]

12 • Chemical Thermodynamics Work and PV Work

(4 of 12)

From physics, work = force × distance = F × d. In chemistry there is electrical work (e-‘s through a wire) and P∆∆ V work as a gas expands.

Use pressure = force/area and area × distance = volume to derive work = P∆∆ V from work = F ×× d.

Note that units of work are units of energy. Energy and work are two forms of energy. Doing work on a system increases the potential energy of the system.

P∆V work is in L⋅⋅ atm 1 L⋅atm = 24.2 cal = 101.3 J

PV = nRT… work can be calculated as work = ∆∆ nRT for chemical reactions where the # of moles of gas change, ∆n.

12 • Chemical Thermodynamics Work is NOT a State Function

(It depends on HOW you do the work.) (5 of 12)

Q: Gas in a piston at a pressure of 10 atm is allowed to expand from 1 L to 10 L. How much work done?

A: It depends on the resisting force. If it expands against 0 atm pressure, P∆V = 0 atm × 9 L = 0 L⋅⋅ atm Expanding against 1 atm pressure, Vfinal is 10 L P∆V = 1 atm × 9 L = 9 L⋅⋅ atm Expanding in two steps, first, against 2 atm, Vfinal = 5 L P∆V = 2 × 4 L = 8 L⋅atm then in a second step against 1 atm, Vfinal = 10 L P∆V = 1 × 5 L = 5 L⋅atm… Total = 13 L ⋅⋅ atm

12 • Chemical Thermodynamics Reversible Processes

(6 of 12)

From the P∆V work example, you can see that the more steps (and smaller increments) you use to get the work from the system, the more work you can get. There is an upper limit to how much work can be derived from any system. That maximum work is called the Gibb’s Free Energy, ∆G. The theoretical maximum work can be achieved if the steps are so small that they can go either way, if they are reversible.

12 • Chemical Thermodynamics ∆∆ H = ∆∆ E + P∆∆ V

(7 of 12)

If a chemical reaction occurs in a closed container ∆∆ V = 0 and so P∆∆ V work = 0. ∆∆ E = q - w becomes ∆∆ E = q.

In a system at constant pressure (a common situation) the energy change involves both heat (q) & work (P∆V). ∆∆ E = q - P∆∆ V [P∆V = work done BY the system] or q = ∆∆ E + P∆∆ V q at constant pressure is called ∆∆ H.

∆∆ H = ∆∆ E + P∆∆ V or ∆∆ E = ∆∆ H - P∆∆ V

The work (P∆V) is generally insignificant unless the # of moles of gas (∆n = nfinal - ninitial) is changing. P∆∆ V = ∆∆ nRT

Recall that if ∆n is negative (# moles decreasing) work is negative (work is being done ON the system).

12 • Chemical Thermodynamics Hess’s Law of Heat Summation

(8 of 12)

If several reactions add up to give an overall reaction, the ∆∆ H’s of the reactions will add up to the overall ∆∆ H.

Standard Heats of Formation, ∆∆ Hf°° , are useful for this purpose. This is the energy involved in making a mole of a substance from its elements at 25°C and 1 atm pressure.

This law is often written as: ∆∆ H°° reaction = ΣΣ ∆∆ Hf°° products - ΣΣ ∆∆ Hf°° reactants

Note: ∆Hf° for elements is 0. If you are NOT using heats of formation, you need to write out the equations to see how they combine.

12 • Chemical Thermodynamics Bond Energies

(9 of 12)

Bond energy is the amount of energy needed to BREAK a certain bond. You can determine the approximate energy change in a chemical reaction by summing the bond energies of the reactants and subtracting the bond energies of the products. Note: this is opposite to the Hess’s Law (products - reactants) because bond energy involves breaking bonds whereas ∆H°f involve forming bonds.

12 • Chemical Thermodynamics Two Big Driving Forces of the Universe

Enthalpy (∆∆ H) and Entropy (∆∆ S) (10 of 12)

Things in the world tend toward lowest energy (-∆H) and also tend toward greatest disorder (+∆S).

More disorder can be recognized as: • greater # of moles of gas formed • gas > liquid > solid and (aq) > (s) • greater volume formed • mixed molecules (HI) formed from diatomic molecules

example: H2(g) + I2(g) → 2HI(g) +∆S Note: ∆S° can be calculated using Hess’s Law, but S° of

elements is NOT 0. Also, S° is often reported in J/mol⋅K and cal/mol⋅K, not kJ and kcal… watch your units!

12 • Chemical Thermodynamics The Second Law of Thermodynamics

∆∆ G = ∆∆ H - T∆∆ S (11 of 12)

The second law defines a value called Gibb’s Free Energy symbolized as ∆G. This represents the theoretical maximum work that can be done by a system.

A reaction is spontaneous when ∆G is negative (∆G<0). The reverse reaction is spontaneous when ∆G is positive. When ∆G = 0, the reaction is at equilibrium.

∆∆ H ∆∆ S Spontaneous… + … at all temperatures … at LOW temperatures + + … at HIGH temperatures + the reverse reaction is spontaneous

12 • Chemical Thermodynamics ∆∆ G (thermodynamics) and Keq (equilibrium)

(12 of 12)

We have been calculating ∆G° values, that is, ∆G under standard conditions. Once a reaction begins, the concentrations change and the value of ∆G is different.

If ∆∆ G°° < 0, then the reaction will proceed in a forward direction. As it does, however, the value of ∆G will increase until the forward and reverse reactions are balanced and ∆∆ G = 0. This is called equilibrium.

If ∆∆ G°° > 0, then the reverse reaction will proceed and the value of ∆G will decrease until equilibrium is reached.

∆G° gives information about where we are in relation to equilibrium. There is a formula that combines ∆∆ G and Keq.

22 • Nuclear Chemistry The People

(1 of 16)

• Wilhelm Roentgen (1845-1923) discovered X-rays, a high energy form of light. (1895)

• Henri Becquerel (1852-1909) found that uranium ores emit radiation that can pass through objects (like x-rays) and affect photographic plates. (1896)

• Marie Sklodowska Curie (1867-1934) Marie and Pierre worked with Becquerel to understand radioactivity. The three shared a Nobel Prize in Physics in 1903. Marie won a second Nobel Prize in Chemistry in 1911 for her work with radium and its properties.

• E. O. Lawrence invented the cyclotron which was used at UC Berkeley to make many of the transuranium elements.

22 • Nuclear Chemistry Terms I-- Radioactivity

(2 of 16)

radioactivity the spontaneous breakdown of atomic nuclei, accompanied by the release of some form of radiation (also called radioactive decay)

half-life time required for half of a radioactive

sample to decay transmutation one element being converted into another

by a nuclear change nuclides isotopes of elements that are identified by

the number of their protons and neutrons

22 • Nuclear Chemistry Terms II--Radioactivity

(3 of 16)

decay series the sequence of nuclides that an element changes into until it forms a stable nucleus

radioactive using half-life information to determine dating the age of objects. C-14/C-12 is common

for organic artifacts. Uranium is common for rocks.

nuclear fission large nucleus breaking down into pieces of

about the same mass nuclear fusion two or more light nuclei blend to form one

or more larger nuclei

22 • Nuclear Chemistry Types of Radiation

(4 of 16)

Alpha particles are the same as a helium nucleus, 42 He, with

a mass of 4 amu. It travels about 1/10th the speed of light and is the most easily stopped of the three particles (a sheet of paper will stop them). It is the least dangerous. Beta particles are high speed electrons,

01− e, with a mass of

0.00055 amu and travel at nearly the speed of light. They can be stopped by a sheet of aluminum. It is more penetrating and therefore more dangerous than alpha. Gamma rays are extremely high energy light, γ, with no mass, and are the most penetrating (several cm’s of lead are needed to stop them). They can cause severe damage.

22 • Nuclear Chemistry Half-Life Problems

(5 of 16)

In each half-life problem there are basically four variables: • total time • half-life • starting amount • ending amount

64g 32g 16g 8g 4g 2g 1g 0.5g 0.25g

Question: If you have 0.25 g of a radioactive substance with a half life of 3 days, how long ago did you have 64 grams?

Answer: Draw the chart to determine the number of half-lives to get from the ending amount to the starting amount… each half-life is worth 3 days…24 days.

22 • Nuclear Chemistry Half-Life (6 of 16)

Half-Life The time it takes for half of a radioactive substance to decay.

The decay graph has a characteristic shape:

time

#

The time it takes for the amount of substance or the activity of the substance to drop to half is the same WHEREVER you start on the graph. This is a first-order reaction. Half-lives can range from microseconds to thousands of years and is characteristic of each substance.

22 • Nuclear Chemistry Nuclear Equations

(7 of 16)

Memorize the symbols for the important particles alpha beta positron neutron

42 He 0

1− e 0 1+ e 1

0 n Decay means the particle is on the right side of the equation: example: alpha decay of U-238

23892 U → 4

2 He + 23490 Th

The 234 and 90 are calculated… the Th is found on the periodic table (find the element with atomic # = 90). Several neutrons can be shown together and written as…

3(10 n) and would be counted as 3

0 n in the equation.

22 • Nuclear Chemistry How Each Type of Decay Can Stabilize an

Unstable Nucleus (8 of 16)

Certain values of p+’s and n°’s in the nucleus are stable. A nucleus can be unstable (radioactive) for 3 reasons: • the nucleus has too many protons compared to neutrons solution: positron decay (change a proton into a neutron and a positive electron… …a positron)

• the nucleus has too many neutrons compared to protons solution: beta decay (change a neutron into a proton and a negative beta particle)

• the nucleus is too big (too many protons and neutrons) solution: alpha decay (lose 2 p+ and 2 n°)

22 • Nuclear Chemistry Uses of Radioactivity

(9 of 16)

Radioactive Dating: In every living thing there is a constant ratio of normal C-12 and radioactive C-14. You can calculate the time needed to change from what is expected to what is actually found.

Radioisotopes: Many substances can be radioactive and then followed as they move through the body.

Fission Reactors: Current nuclear reactors use fission reactions to produce heat which is used to turn water into steam and drive turbine engines that produce electricity.

The Sun and Stars are powered by nuclear fusion… this is related to the fact that the most abundant element in the universe is hydrogen… followed by helium.

22 • Nuclear Chemistry Fission and Fusion Reactions

(10 of 16)

U-235 is “fissionable” which means it can be split when bombarded by neutrons.

23592 U + 1

0 n → 14156 Ba + 92

36 Kr + 310 n + energy

The fact that each splitting nucleus can emit neutrons that can split other nuclei is the basis for the “chain reaction.”

“Breeder reactors” use different isotopes.

Fusion in the Sun involves several steps that can be summed up as: 4( 1

1 H) → 42 He + 2 0

1 e + energy Thermonuclear devices use isotopes of hydrogen (deuterium and tritium): 2

1 H + 31 H → 4

2 He + 10 n + energy

22 • Nuclear Chemistry Energy–Mass Conversions

(11 of 16)

Einstein’s famous equation, E = mc2 , is the basis for explaining where the energy associated with nuclear changes comes from.

When a nuclear change occurs, the mass of the products is slightly less than the mass of the reactants. This loss in mass is called the mass defect.

E = the energy m = the mass defect c = the speed of light, 3.00 x 108 m/s

1 kg of mass converted into energy would be equivalent to burning 3 billion kg of coal!

22 • Nuclear Chemistry What Happens During

Beta and Positron Decay (12 of 16)

During beta decay, 1 neutron changes into 1 proton + 1 negative beta particle (The atomic # increases by one due to the new proton. The mass # is unchanged… a neutron is gone. To maintain electrical neutrality, a negative beta particle is also formed.)

Example: 23592 U → 0

1− e + 23593 Np

During positron decay, 1 proton changes into 1 neutron + 1 positron particle (The atomic # decreases by one due to the loss of a proton. Since it changed into a neutron, the mass # is unchanged.) Example: 235

92 U → 0 1+ e + 235

91 Pa

22 • Nuclear Chemistry Calculating Half-Lives

(13 of 16)

When a problem involves whole numbers of half-lives, divide by 2 to determine the amounts involved. For other situations, the following equations are useful:

ln t

0

[A][A]

= kt and the special case for half-life, t½, where by

definition, [A]t = ½ [A]0 ln 2 = 0.693= kt½ [A] is the concentration (or activity) of the radioactive substance, t = time, k = the rate constant (the same that is in Rate Laws). Note: if you know the half-life, you can calculate the rate constant and vice-versa.

22 • Nuclear Chemistry Radioactive Decay Series

(14 of 16)

Once a nucleus decays, the daughter isotope is often unstable as well. Many decays may occur before a stable nucleus is formed.

A classic example is U-238 that decays through 14 steps into stable Pb-206. Each step has a characteristic decay particle and half-life.

This characteristic decay series is the method used to verify the identity of newly formed atoms.

The fact that daughter products can be even more radioactive than the parent isotope adds to the problem of nuclear waste and its storage/disposal.

22 • Nuclear Chemistry Geiger-Muller Tubes, Smoke Detectors, and

Brushes for Cleaning Negatives (15 of 16)

An useful characteristic of decay particles are that they ionize the air they pass through by striking atoms and knocking off electrons.

Geiger counters use this idea. As radioactive particles pass through a chamber with two electrodes, ionized particles migrate to the + and - electrodes and complete the circuit.

Smoke Detectors use a tiny piece of radioactive Am to keep a circuit flowing due to ionized particles. Smoke particles attract ionized particles, break the circuit, & set off the alarm.

Brushes are kept ionized by tiny bits of radioactive material to more easily attract tiny bits of dust.

22 • Nuclear Chemistry Extending the Periodic Table

(16 of 16)

Uranium, Z=92, is the largest naturally-occurring element. Larger atoms were manufactured. Elements 93 and 94 were formed in atomic bomb tests and identified by Seaborg. Glenn Seaborg and Al Ghiorso at UC Berkeley were able to use E. O. Lawrence’s cyclotron to make larger atoms (95-103).

Some of these new elements have uses in the medical field as well as helping to further the understanding of the nucleus. For many of the larger elements, however, only a few atoms or even one atom formed. They were identified by their characteristic decay series.

As of July 2000, 118 is the largest element.

23 • Organic Chemistry Historical Ideas

(1 of 16)

Chemicals from living things were thought to contain a “vital force” that could not be duplicated in the lab. This changed with Friedrich Wöhler who mixed cyanic acid (HCNO) with ammonium hydroxide making ammonium cyanate (NH4CNO).

C

O

NH2 NH2

urea

He usually allowed the salt solution to evaporate overnight, but tried heating it to hurry the process. The result was a crystal that he recognized as urea (H2NCONH2).

The modern view of organic chemistry is the chemistry of carbon compounds. C is the key element. It can form four bonds and that are very strong bonds due to its small size.

23 • Organic Chemistry Alkane Series -- Saturated Hydrocarbons

(2 of 16)

The alkanes (paraffins) follow the formula: CnH2n+2: These molecules contain ONLY single bonds. They are said to be “saturated” with hydrogens.

Memorize these prefixes also used with alkenes & alkynes. CH4 methane C6H14 hexane C2H6 ethane C7H16 heptane C3H8 propane C8H18 octane C4H10 butane C9H20 nonane C5H12 pentane C10H22 decane

Given a formula, you can tell that it contains only single bonds because it fits the alkane formula.

As the molecules increase in size, they tend to be liquids and

23 • Organic Chemistry Structural Formulas Can Be Misleading

(3 of 16)

CH4, can be drawn using a structural formula. This can be misleading. The molecule is not flat with bond angles of 90°. You must be aware of the 3-D structure and the 109.5 °° bond angles.

C

H

H

H

H

C

Cl

H H

Cl

C

H

H Cl

Cl

For example, there is only one isomer of dichloromethane, but you can draw it at least two ways.

Building models of the molecules is an important part of strengthening this skill.

23 • Organic Chemistry Alkenes and cis- /trans- Isomerism

(4 of 16)

Alkenes contain 1 double bond. The formula is CnH2n. They are said to be “unsaturated” (like unsaturated fats). The double bond can be broken and more hydrogens added.

C

H

C

H H

H

ethene

Since double bonds cannot easily rotate (due to the pi bonding) cis- and trans- isomers can be formed.

Example: 1,2-dichloroethene can be built two ways.

C

H

C

Cl Cl

H cis -1,2-dichloroethene

(a polar molecule)

CH

C

Cl H

Cl trans-1,2-dichloroethene

(a nonpolar molecule)

23 • Organic Chemistry Alkynes, Alkadienes, and Cyclic Hydrocarbons

(5 of 16)

Alkynes contain 1 triple bond (unsaturated). Formula: CnH2n-2.

CH C H ethyne (acetylene)

The triple bond is linear, so no cis/trans isomerism occurs.

Alkadienes are molecules with two double bonds. They have the same formula as the alkynes, CnH2n-2.

Example: C4H6 is named 1,3-butadiene because the double bonds start on carbons #1 and #3.

CH2C

CCH2

H

H

CH2

CH2

CH2

Cyclic compounds contain rings having the same formula as the alkenes, CnH2n. Example: cyclopropane, C3H6.

23 • Organic Chemistry Naming Organic Compounds

(Organic Nomenclature Using IUPAC Rules) (6 of 16)

The basic idea is to name the molecule after the longest continuous chain of carbon atoms. Side groups are listed with #’s to indicate the C atom to which they are attached.

Side Groups : -Cl chloro -Br bromo -I iodo -CH3 methyl -C2H5 ethyl -C3H7 propyl, etc.

di- = 2 groups tri- = 3 groups tetra- = 4 groups

2,2,3-tribromobutane (not 2,3,3-) Note that we # the carbons from whichever end results in the smallest numbers.

C C C C

Br

H

Br

Br

H

H

H

H

H

H

23 • Organic Chemistry Common Errors in Drawing/Naming Structures

1-methylsomething & 2-ethylsomething (7 of 16)

While drawing the isomers of pentane, C5H12, students draw this structure, naming it 2-ethylpropane. (a chain of 3 C’s with an ethyl group) The longest chain is four C’s, and should be named 2-methylbutane.

C C C H

H

CH2

H

H

H

H

H

CH3

A similar error is to draw and name “1-methylsomething”.

C C O

H

H

H

H

H 5 bonds to C

Two more tips… double check that each C has four and only four bonds. Also, remember that N and O atoms have lone pairs of e-‘s although they are seldom drawn. (Impt. for steric #!)

23 • Organic Chemistry Optical Isomers

Chiral Compounds (8 of 16)

CH CH2

CH2

CH3

CH3

CH

2

CH3

3-methylhexane

Some molecules have the ability to rotate polarized light.

These molecules can be recognized by a C atom (the chiral carbon) bonded to four different groups .

This carbon is bonded to H, methyl, ethyl, & propyl groups.

You can build two versions of this molecule that are “nonsuperimposable mirror images of each other.” One will rotate light clockwise, one counterclockwise.

In biology, these are called dextro- and levo- (D and L) forms.

23 • Organic Chemistry Common Names You Should Know About

(9 of 16)

ethene is also called ethylene propene is also called propylene

CH2

CH3

CHCH3

CH3

2-methylbutane is also called isopentane. “Iso-“ means the same… the same two methyl groups come branch from C #2. 2-methylpentane is isohexane, etc.

CH3C

CH3

CH3

CH3

2,2-dimethylpropane is called neopentane. These common names show up occasionally in names… such as in isopropyl alcohol.

23 • Organic Chemistry Aromatic Compounds

Benzene and its Derivatives (10 of 16)

CH

CHCH

CH

CHCH

CH

CHCH

CH

CHCH

two resonance structures

Benzene, C6H6, is unique. It can be drawn as shown, but the actual structure involves a circular pi bond (sp2 orbitals & delocalized e-‘s).

Benzene is also shown with a circle as the pi bond.

The carbon #’s can be used for substituted benzene. Example: dichlorobenzene 1,2- is known as the ortho- position 1,3- is known as the meta- position 1,4- is known as the para- position

5

4

3

2

1

6

Paradichlorobenzene: the main ingredient in some moth balls.

23 • Organic Chemistry Functional Groups I Alcohols and Ethers

(11 of 16)

Alcohols General formula: R-O-H [R ≈ Rest of molecule]

C1C3 C

2

C 1

C1

C atoms are classified as primary (1), secondary (2), or tertiary (3) by the number of C atoms it is bonded to. A primary alcohol has the -OH group bonded to a primary carbon, etc.

This is not a base because the -OH is covalent, not ionic. Naming: group + “alcohol” (e.g. ethyl alcohol or ethanol)

Ethers General formula: R-O-R’ [R’ can = R, but not H] Naming: two groups + “ether” diethyl ether was the 1st effective surgical and dental anesthetic.

CH3

CH2

OCH2

CH3

23 • Organic Chemistry Functional Groups II

Aldehydes and Ketones (12 of 16)

Aldehydes Ketones General formula:

RC

H

O

RC

R'

O

Naming: names end in “al”

or “aldehyde” methanaldehyde (formaldehyde)

names end in “one”

propanone (acetone)

Aldehydes and ketones both have a C=O group (carbonyl group). Aldehydes have it on an end carbon. Ketones have it on a middle carbon. Reactions: Primary alcohols can be oxidized into aldehydes. Secondary alcohols into ketones .

23 • Organic Chemistry Functional Groups III

Carboxylic Acids and Esters (13 of 16)

Carboxylic Acids Esters General formula:

RC

OH

O

RC

OR

O

Naming: names end - “oic

acid” ethanoic acid (acetic acid)

names end - “ate” ethyl acetate (acetic acid + ethyl alcohol)

Reactions: Acids can be made by oxidizing aldehydes. Esters are formed (“esterification”) from a carboxylic acid & an alcohol. Water is removed (a “condensation” reaction).

Esters often have pleasant, agreeable odors (e.g. banana.)

23 • Organic Chemistry Functional Groups IV

Amines & Amides (14 of 16)

Amines Amides General formula:

R

N H H

RC

NH2

O

Naming: names contain

“amino” or end in “amine”

aminomethane (methylamine)

names end in “amide”

acetamide

The N may have 1 or 2 or all 3 H atoms replaced with groups. The lone pair on the N atom makes these molecules basic. Your body needs certain amines “vital amines” ≈ “vitamins.”

23 • Organic Chemistry Polymers I

Monomers & Addition Polymerization (15 of 16)

Monomer = one part Polymer = many parts

C C

H

H H

H

“ethylene”

One kind of polymer is made up of monomers that contain a double bond. The double bond can break and we can ADD to it… “Addition polymerization.”

C C

H

H*

H*

H

+ C C

H

H H

H

→ C C

H

H

C

H

C*

H

H

H H

*H

Different monomers form different polymers. This polymer would be called polyethylene. Replace on H on the monomer with Cl and you can make polyvinyl chloride, “PVC.”

23 • Organic Chemistry Polymers II

Copolymers & Condensation Polymerization (16 of 16)

Another polymerization involves condensation reactions.

C

OH

OC

OH

OR

a di-acid

COH

COH

R

H

H

H

H a di-alcohol (a glycol)

Esters form from an acid and an alcohol. Using a di-acid and a di-alcohol, you can make a continuous chain by removing water molecules. The resulting polymer is called a polyester.

Soda bottles are made from a polyester, polyethylene terephthalate ester (PETE).

Nylon (a polyamide) can be made from a di-amine & a di-acid.

1 • IntroductionThe Scientific Method

(1 of 12)

This is an attempt to state how scientists do science. It isnecessarily artificial. Here are MY five steps:

• Make observationsthe leaves on my plant are turning yellow

• State a Problem to be solvedhow can I get my plants healthy (non-yellow)

• Form a hypothesismaybe they need more water

• Conduct a controlled experimentwater plants TWICE a week instead of once a week

• Evaluate resultsif it works, good... if not, new hypothesis (sunlight?)

1 • IntroductionObservations and Measurements

Qualitative, Quantitative, Inferences(2 of 12)

Step 1 of the Scientific Method is Make Observations.These can be of general physical properties (color, smell,hardness, etc.) which are called qualitative observations.These can be measurements which are called quantitativeobservations.

There are also statements that we commonly make based onobservations. “This beaker contains water” is an example.

You infer (probably correctly) it is water because it is aclear, colorless liquid that came from the tap. Theobservations are that it is clear, it is colorless, it is a liquid,and it came from the tap. Recognize the difference.

1 • IntroductionGraphing -- Great Graphs

(3 of 12)

5 Steps to a Great Graph• Descriptive Title• Subtitle (dependent variable vs. independent variable)• Label axes with variable name and units used• Number axes so most of graph paper is used (should the point 0,0 be included in graph)• Draw an appropriate line

straight line (if data looks like it is directly proprotional)smooth curve (if data looks like it follows a trend)dot-to-dot (if data looks like the two variables areunrelated to each other... like Dow Jones averages)

Also: if more than one line is on a graph, provide a Legend

2 • MeasurementSignificant Digits I

What do they mean?(4 of 12)

Consider: 16.82394 cmIn a measurement or a calculation, it is important to knowwhich digits of the reported number are significant.

That means… if the same measurement were repeated againand again, some of the numbers would be consistent andsome would simply be artifacts.

All of the digits that you are absolutely certain of plus onemore that is a judgment are significant.

If all the digits are significant above, everyone whomeasures the object will determine that it is 16.8239 cm, butsome will say …94 cm while others might say …95 cm.

2 • MeasurementSignificant Digits II

Some examples with rulers.(5 of 12)

1 2

a b c

(A composite ruler)a- No one should argue that the measurement is between 0.3and 0.4. Is it exactly halfway between (.35 cm)… or a littleto the left (.34 cm)? The last digit is the judgment of theperson making the measurement. The measurement has 2significant digits.b- The same ruler… so the measurement still goes to thehundredths place… 1.00 cm (3 significant digits).c- A ruler with fewer marks reads 1.6 cm (2 sig digits).

2 • MeasurementSignificant Digits III

Rules for Recognizing Sig. Digits(6 of 12)

• All non-zero digits are significant.523 grams (3) 972,366 seconds (6)

• 0’s in the MIDDLE of a number are ALWAYS significant.5082 meters (4) 0.002008 L (4)

• 0’s in the FRONT of a number are NEVER significant.0.0032 kg (2) 0.00000751 m (3)

• 0’s at the END of a number are SOMETIMES significant.• Decimal point is PRESENT, 0’s ARE significant

2.000 Liters (4) 0.000500 grams (3)• Decimal point is ABSENT, 0’s are NOT significant

2000 Liters (1) 550 m (2)

NOTE: textbook values are assumed to have all sig. digits

2 • MeasurementSignificant Digits IV

Significant Digits in Calculations(7 of 12)

When you perform a calculation using measurements, oftenthe calculator gives you an incorrect number of significantdigits. Here are the rules to follow to report your answers:

x and ÷: The answer has the same # of sig. digits as thenumber in the problem with the least number of sig. digits.example: 3.7 cm x 8.1 cm = 29.97 ≈ 30. cm2 (2 sig. digits)

+ and –: The last sig. digit in the answer is the largestuncertain digit in the values used in the problem.example: 3.7 cm + 8.1 cm = 11.8 cm (3 sig. digits)

Know how to ilustrate why these rules work.

2 • MeasurementAccuracy vs. Precision

(8 of 12)

Accuracy refers to how close a measurement is to someaccepted or true value (a standard).

Ex: an experimental value of the density of Al° is 2.69g/mL. The accepted value is 2.70 g/mL. Your value isaccurate to within 0.37%% error is used to express accuracy.

Precision refers to the reliability, repeatability, orconsistency of a measurement.

Ex: A value of 2.69 g/mL means that if you repeat themeasurement over and over, you will get values thatagree to the tenths place (2.68, 2.70, 2.71, etc.)± uncertainty and sig. digits are used to expressprecision

2 • MeasurementMetric System

(9 of 12)

We generally use three types of measurements:volume Liters (mL)length meters (km, cm and mm)mass grams (kg and mg)

We commonly use the prefixes:centi- 1/100thmilli- 1/1000thkilo- 1000

Occasionally you will encounter micro(µ), nano, pico, mega,and giga. You should know where to find these in Ch. 1.Know that 2.54 cm = 1 inch and 2.20 lb = 1 kg

2 • Measurement% and ppm

(10 of 12)

Percentage is a mathematical tool to help compare values.Two fractions, 3/17 and 5/31 are difficult to compare:If we set up ratios so we can have a common denominator:

317

= x

100 =

17.65100

531

= x

100 =

16.13100

so… we can see that 317

> 531

.

There are 17.65 parts per 100 (Latin: parts per centum) or17.65 percent (17.65 %)… the % is a “1 0 0”ppm (parts per million) is the same idea, (use 1,000,000instead of 100)

317

= x

1 000 000 = 176,470 ppm

3 • Problem SolvingScientific Notation

Useful for showing Significant Digits(11 of 12)

Scientific notation uses a number between 1 and 9.99 x 10 tosome power. It’s use stems from the use of slide rules.

Know how to put numbers into scientific notation:5392 = 5.392 x 103 0.000328 = 3.28 x 10–4

1.03 = 1.03 550 = 5.5 x 102

Some 0’s in numbers are placeholders and are not asignificant part of the measurement so they disappear whenwritten in sci. notation. Ex: 0.000328 above. In scientificnotation, only the three sig. digits (3.28) are written.

Scientific Notation can be used to show more sig. digits.Values like 550 ( 2 sig. digits) can be written 5.50 x 102 (3)

3 • Problem SolvingUnit Analysis

Converting between English and Metric Units(12 of 12)

Consider the metric/English math fact: 2.54 cm = 1 inchThis can be used as the “conversion factor”:

2.54 cm1 inch

or 1 inch

2.54 cm

You can convert 25.5 inches to cm in the following way:Given: 25.5 in

Desired: ? cm 25.5 in x 2.54 cm

1 in = 64.77 cm ≈ 64.8 cm

This is the required way to show your work. You have twojobs in this class, to be able to perform the conversions andto be able to prove that you know why the answer is correct.

4 • MatterMass & Weight -- Two Properties of Matter

(1 of 8)

mass is the amount of something...weight is how much gravity is pulling on the mass.(Weight will be proportional to the mass at a given spot.)

Mass is what we REALLY want to use... measured in grams.You use a balance to measure mass... you compare yourobject with objects of known mass.

Weight is measured with a scale (like your bathroom scaleor the scale at the grocery store). If there is no gravity, itdoesn’t work. Note: electronic balances are really scales!

You convert mass / weight using: 1 kg

2.205 lbs or

2.205 lbs1 kg

4 • MatterPure Substances, Elements, & CompoundsHomogeneous & Heterogeneous Mixtures

(2 of 8)

Pure Substances

Matter Energy

CompoundsElements

Mixtures

HeterogeneousHomogeneous

This chart should help you sort out these similar terms.Be able to use chemical symbols to represent elements andcompounds. For example...CuSO4•5H2O, a hydrate, contains 21 atoms & 4 elements.

Memorize the 7 elements that exist in diatomic molecules:HONClBrIF or BrINClHOF or “H and the 6 that make a 7starting with element #7”

4 • MatterSeparating Mixtures by Filtration,Distillation, and Chromatography

(3 of 8)

Mixtures are substances the are NOT chemically combined...so if you want to separate them, you need to exploitdifferences in their PHYSICAL properties.

Filtration: some components of the mixture dissolve &some do not. Filtrate is what passes through the filter.

Distillation: some components vaporize at differenttemperatures or one component may not vaporize at all(e.g.: salt+water) complete separation may not bepossible.

Chromatography: differences in solubility vs. adhesion tothe substrate. Substratemay be filter paper (paperchromatography), or other substances, GLC, TLC,HPLC, column, etc.

4 • MatterMass, Volume, and Density

Intensive vs. Extensive Properties(4 of 8)

Extensive properties depend on the amount of substance.We measure these properties frequently... (mass &volume... mostly).

Intensive properties are independent of the size of thesample. These are useful for identifying substances...(melting point, boiling point, density, etc.)

It is interesting that an intensive property, density = mass

volumeis the ratio of two extensive properties... the size of thesample sort of “cancels out.” Be able to do density problems(3 variables). See Sample Problems on pages 72 & 73

4 • MatterPhysical and Chemical PropertiesPhysical and Chemical Changes

(5 of 8)

Equations to symbolize changes: reactants → products

Physical Properties can be measured from a sample of thesubstance alone... (density, MP, BP, color, etc.)

Chemical Properties are measured when a sample is mixedwith another chemical (reaction with acid, how does itburn in O2)

Physical Changes imply that no new substances are beingformed (melting, boiling, dissolving, etc.)

Chemical Changes imply the substance is decomposinginto new substances or mixing with another chemical toform new substances. This change is accompanied by heat,light, gas formation, color changes, etc.

4 • MatterConservation of Mass

Symbols of the Elements(6 of 8)

Conservation means the quantity does not change during areaction. If you carefully measure the reactants before areaction and the products after the reaction, no mass isgained or lost. This is called Conservation of Mass.

Know your symbols of the elements (make Flash Cards).Be careful with the spelling of:

Cl, chlorine F, fluorine Ni, nickel

Recall that many of the symbols come from the Latin name.

4 • MatterRelative Abundance of Elements

(7 of 8)

In the atmosphere:78% nitrogen gas, N221% oxygen gas, O2<1% argon gas, Ar

In the earth’s crust :Most of the crust is made up of SiO2 (quartz, sand, glass)

46.7% oxygen (mostly combined with silicon)27.7% silicon (mostly combined with oxygen)8.1% aluminum (in combined form)5.0% iron (in combined form)

In the universe:almost all hydrogen gas, H2... then He (fusion product)

4 • MatterNatural History of Airs Lab

The Chemistry of the Airs, the Recipes, and theTests for the Three Gases

(8 of 8)

Name: Dephlogisticated Air - Oxygen gas, O2Recipe: Catalytic Decomposition of hydrogen peroxide

(H2O2) by yeastTest: Glowing Splint Test (oxygen supports combustion)

Name: Inflammable Air - Hydrogen gas, H2Recipe: Drano™ (NaOH) + Al° → Na+ + AlO3– + H2Test: Burning Splint Test (H2 + O2 → H2O + energy)

Name: Fixed Air - Carbon Dioxide gas, CO2Recipe: chalk (CaCO3) + vinegar (acetic acid) (HC2H3O2)

→ H2O + CO2 + Ca(C2H3O2)2Test: Limewater (Ca(OH)2 + CO2 → CaCO3(s) + H2O

Mn, manganese vs. Mg, magnesium

Refer to page 80 for a nice listing with the Latin names.

6 • Structure of the AtomThe Subatomic Particles

(1 of 8)

Name Symbol Mass Charge Locationprotons p+ 1 u 1+ part of the nucleusneutron n° 1 u 0 part of the nucleus

electron e– 11837

u 1– normally at large

distances from thenucleus

J.J. Thompson is given credit for discovering electronsusing a Crookes tube and testing many different gases.Cathode rays were found to be beams of electrons.

Cavendish is given credit for the discovery of the neutron.

6 • Structure of the AtomTerms I-- Atomic Structure

(2 of 8)

atoms the smallest particle of an element . Itconsists of a central nucleus and electronclouds outside the nucleus.

nucleus the dense central portion of an atom.subatomic smaller than an atom. The proton,

neutron, and electron are subatomicparticles.

net charge the difference in the positive charge due toprotons and the negative charge due toelectrons in an atom.

nucleons the particles that make up the nucleus.

6 • Structure of the AtomTerms II-- Atomic Structure

(3 of 8)

atomic number the number of protons in an atom. Thisnumber determines the identity of anelement.

mass number the number of protons + neutrons

isotopes atoms with the same number of protons,but different numbers of neutrons. Atomswith the same atomic number, butdifferent mass numbers.

isotopic notation shorthand notation for a nucleus thatshows the mass #, atomic # and the

symbol. U-238 would be 238 92U

6 • Structure of the AtomCalculating Atomic Mass

(4 of 8)

Any real sample of an element contains more than onenaturally occuring isotope. For instance, boron

isotope abundance mass # isotopic mass

boron-1010 5B 19.78% 10 mass = 10.013 u

boron-1111 5B 80.22% 11 mass = 11.009 u

The atomic mass is the weighted average of the isotopes.

at. mass = (19.78%)(10.013u) + (80.22%)(11.0009u)

100 or

at. mass = (0.1978)(10.013u) + (0.8022)(11.0009u) =10.81 u

6 • Structure of the AtomDetermining Numbers of Protons, Neutrons,

and Electrons from the Isotopic Notation(5 of 8)

Consider the following symbol: 3316S

2–

The 16 is the atomic number which is the number ofprotons.

The 33 is the mass number which is the mass of one of theisotopes. This mass is due to the protons and neutrons.

The number of neutrons is the mass number - the atomicnumber. 33 - 16 = 17 neutrons.

Since the charge is 2-, there are 2 more electrons thanprotons. In this case, there are 18 electrons.

6 • Structure of the AtomImportant People in the Development of the

Atomic Theory(6 of 8)

Democritus [atomos]philosopher who decided that matter was discontinuous

John Dalton [billiard-ball model]experiments with gases… different substances aredifferent combinations of atoms

J.J. Thomson [plum-pudding model]experiments with gas-discharge tubes… atoms havepositive and negative parts… the negative electrons arethe same from atom to atom

Ernest Rutherford [nuclear model/solar system model]most of the mass of the atom is concentrated in a tiny,positively-charged nucleus

Niels Bohr [Bohr Model (quantized e- energy levels/orbits)]the electrons had only certain allowed energy levels

6 • Structure of the AtomEarly Experimental Observations That Would

Later Be Explained By The Atomic Theory(7 of 8)

The Law of Conservation of Massthe mass of all the reactant molecules = the mass of all theproduct molecules

The Law of Definite Compositionthe percentage composition of any sample of a substanceis the same

The Law of Multiple Proportionswhen two compounds made of the same two elements(such as CO and CO2) are broken down to give the samemass of one element… the masses of the other elementwill be in simple whole-number ratio.

6 • Structure of the AtomRutherford’s Gold Foil Experiment

(8 of 8)

Ernest Rutherford’s classic gold foil experiment led to thenuclear model of the atom.

αα

most alpha'scame straightthrough here

a few bounced backat a large angle

gold foil

• the nucleus is tiny - because most of the alpha’s missedthe nucleus and went straight through the foil• the nucleus is positively charged - because the (+)charged alpha was repelled by the (+) charged nucleus• the nucleus is incredibly dense - because the nucleus wasable to bounce back at a very large angle

7 • Chemical FormulasFormula and Compound Terms

(1 of 12)

anion another name for a negative ioncation another name for a positive ionbinary compound contains two elementsternary compound contains three or more elementsionic compound made of a positive & a negative ionmolecular compound atoms share electrons…not ioniccovalent compound same as molecular compoundchemical formula shows # & kind of atoms (includes

molecular, empirical and structural)molecular formula shows # & kind of atoms in moleculeempirical formula simplest whole # ratio of atomsstructural formula show how atoms are connectedmonatomic: one atom diatomic: two atomspolyatomic: many atoms

7 • Chemical FormulasMemorization Tips - Negative Ions

(2 of 12)

all negative ions (anions) end in “-ide”, “-ate”, or “-ite”

”–ides”• these are single atoms• exceptions: hydroxide (OH–) and cyanide (CN–)• you can tell charge from position on the periodic table.• Family VII (F, Cl, Br, I ) all form 1– ions.• Family VI (O, S) all form 2– ions.

“–ates”• these contain several oxygen atoms. You just have tomemorize them… there is no rule about how many oxygens.

“–ites” contain one less O than the –ates… same charge.

7 • Chemical FormulasMemorization Tips - Positive Ions

(3 of 12)

metals form + ions “–ous” ions < “–ic” ions • Family I (Li, Na, K, etc.) all form 1+ ions • Family II (Be, Mg, Ca, Sr, Ba, Ra) all form 2+ ions • Family III (Al) forms a 3+ ion

mercury mercurous, Hg22+ mercuric, Hg2+

mercury(I) mercury(II)copper cuprous, Cu+ cupric, Cu2+

copper(I) copper(II)tin stannous, Sn2+ stannic, Sn4+

tin(II) tin(IV)iron ferrous, Fe2+ ferric, Fe3+

iron(II) iron(III)

7 • Chemical FormulasFormula Conventions

(4 of 12)

Superscriptsused to show the charges on ionsMg2+ the 2 means a 2+ charge (lost 2 electrons)

Subscriptsused to show numbers of atoms in a formula unitH2SO4 two H’s, one S, and 4 O’s

Coefficientsused to show the number of formula units2Br– the 2 means two individual bromide ions

Hydrates CuSO4 • 5 H2Osome compounds have water molecules included

7 • Chemical FormulasHow Ions Form

(5 of 12)

Positive ions form by LOSING one or more electrons.Negative ions form by GAINING one or more electrons

PROTONS are not gained or lost from the nucleus except innuclear reactions that require MUCH more energy than isusually available.

Metals become POSITIVE ions.Non-Metals become NEGATIVE ions.Semi-Metals sometimes become ions and sometimes shareelectrons as molecular compounds.

7 • Chemical FormulasWriting Ionic Formulas

(6 of 12)

The positve ion is written first.

The total positive charge must match the total negativecharge in the compound.

Use parentheses when you need several polyatomic ions…Al2(SO4)3 is correct Al2(Cl)3 is incorrect

Be careful of OH– ions…Ba(OH)2 is correct BaOH2 is incorrect

Reduce subscripts in final formula except with Hg22+

SnS2 is correct Sn2S4 is incorrectmercurous chloride, Hg2Cl2 is correct

7 • Chemical FormulasOther Ions (beyond the 40)

and how they relate to oxidation numbers(7 of 12)

bicarbonate, HCO3– bisulfate, HSO4–

bisulfide, HS– biphosphate, HPO42–

perchlorate ClO4– oxidation number of Cl = +7chlorate ClO3– oxidation number of Cl = +5chlorite ClO2– oxidation number of Cl = +3hypochlorite ClO– oxidation number of Cl = +1chlorine Cl2 oxidation number of Cl = 0chloride Cl– oxidation number of Cl = –1

chlorate ClO3–

bromate BrO3– similar to chlorateiodate IO – similar to chlorate

7 • Chemical FormulasOxidation Numbers

What they are and how you find them(8 of 12)

The oxidation number is the “apparent charge” on an atom.

The oxidation number of any substance in its elemental formis defined as 0.

Example: the oxidation number of H in H2 is 0 the oxidation number of H in H2O is +1

The oxidation numbers of each of the atoms in a substanceadd up to the charge on the substance…

CO32– C + O + O + O = –2x + –2 + –2 + –2 = –2 ∴ x = +4

CH4 C + H + H + H + H = 0 x + 4(+1) = 0 ∴ x = –4

7 • Chemical FormulasWriting and Naming Acids

(9 of 12)

Acids are ionic formulas in which the positive ion is H+.

Use as many H+ ions as the charge on the negative ion.

Three rules for naming: if the anion ends with: the acid is named:

–ite ********ous acid–ate ********ic acid–ide hydro********ic acid

• Acids from sulfide, sulfite, and sulfate include a “ur” H2S is hydrosulfuric acid, not hydrosulfic acid• Acids from phosphate and phosphite include a “or” H3PO4 is phophoric acid, not phosphic acid

7 • Chemical FormulasStock Names vs. Traditional Names

(10 of 12)

The Stock System of naming compounds is used…• when a positive ion has more than one possible charge

(i.e. cuprous, cupric, etc.)Traditional: mercurous, Hg22+ mercuric, Hg2+

Stock: mercury(I) mercury(II)Traditional: cuprous, Cu+ cupric, Cu2+

Stock: copper(I) copper(II)

• for molecular compounds where the elements have manydifferent oxidation numbers (i.e. N in NO2, NO, N2O, etc.)

Stock Name: Traditional Name:NO2 nitrogen(IV) oxide nitrogen dioxideNO nitrogen (II) oxide nitrogen monoxide

7 • Chemical FormulasNaming Molecular Compounds

(Traditional Method)(11 of 12)

The first element is named using the name of the element.The second element always end in “–ide.”

Indicate the number of atoms using the prefix…1 mono- 6 hexa-2 di- 7 hepta-3 tri- 8 octa-4 tetra 9 nona-5 penta 10 deca-

If the first element has only one atom, don’t use the mono-If the second element is oxygen, drop the vowel…

monoxide, not monooxidetetroxide, not tetraoxide

7 • Chemical FormulasDetermining Ions from Formulas

(12 of 12)

Given the formula of an ionic compound, you can determinethe original ions.

Most of these problems are obvious if you have the ionsmemorized…

NaCl Na+ Cl–

K2SO4 K+ SO42–

Some need a little “detective work”…CuS Cu+ or Cu2+?

since S is 2– (memorized) Cu must be 2+.

You can also figure out ions you’ve never seen before…Ga(NO3)3 must be Ga3+ since NO3– (memorized)

8 • Mathematics of Chemical FormulasStoichiometry Terms

(1 of 8)

stoichiometry study of the quantitative relationshipsin chemical formulas and equations.

atomic mass weighted average mass of an atom,found on the periodic table

formula mass sum of the atomic masses of theatoms in a formula

molecular mass sum of the atomic masses of theatoms in a molecular formula

gram atomic mass atomic mass written in gramsgram formula mass formula mass written in gramsgram molecular mass molecular mass written in grams

molar mass same as gram formula mass

8 • Mathematics of Chemical FormulasCalculating Formula Mass

(2 of 8)

Formula or molecular mass is found by simply summingthe atomic masses (on the periodic table) of each atom in aformula.

H2SO41.01 + 1.01 + 32.06 + 16.0 + 16.0 + 16.0 + 16.0 = 98.08 u2(1.01) + 32.06 + 4(16.0) = 98.06 u or 98.06 g/mole

Generally, round off your answers to the hundredths ortenths place. Don’t round off too much (98.06 g/mol or98.1 g/mol is OK, but don’t round off to 98 g/mol)

UnitsUse u or amu if you are referring to one atom or molecule

8 • Mathematics of Chemical FormulasMole Facts

(3 of 8)

A mole (abbreviated mol) is a certain number of things. It issometimes called the chemist’s dozen.A dozen is 12 things, a mole is 6.02 x 1023 things.

Avogadro’s Number1 mole of any substance contains 6.02 x 1023 molecules

Molar Volume (measured at P = 760 mmHg and T = 0 °C)1 mole of any gas has a volume of 22.4 Liters

Molar Mass (see gram formula mass)1 mole of any substance is its gram formula mass

1 mole6.02 x 1023 molecules

1 mole22.4 L

1 mole

molar mass

8 • Mathematics of Chemical FormulasLine Equations

(4 of 8)

A Line Equation is the preferred way to show conversionsbetween quantities (amount, mass, volume, and number) bycanceling units (moles, grams, liters, and molecules)

The line equation consists of the Given Value, the DesiredUnit, and the line equation itself.

Example: What is the mass of 135 Liters of CH4 (at STP)? Given: 135 L CH4 Desired: ? g CH4

135 L CH4 x 1 mol CH422.4 L CH4

x 16.0 g CH4 1 mol CH4

= 96.43 g CH4

8 • Mathematics of Chemical FormulasMole Relationships

(5 of 8)

The “Mole Map” shows the structure of mole problems

Mass

Volumeat STP

number ofatoms ormolecules

Mass

moles Volumeat STP

number ofatoms ormolecules

➂ ➂➁ ➁

➀ ➀

1) 1 mol

molar mass 2)

1 mol22.4 L

3) 1 mol

6.02 x 1023 molecules

8 • Mathematics of Chemical FormulasPercentage Composition (by mass)

(6 of 8)

Percentage Composition quantifies what portion (by mass)of a substance is made up of each element.

Set up a fraction: mass of element

mass of molecule

Change to percentage: 100 x mass of element

mass of molecule

Generally, round off your answers to the tenth’s place.

The percentage compositions of each element should add upto 100% (or very close, like 99.9% or 100.1%)

8 • Mathematics of Chemical FormulasFormula from % Composition

(7 of 8)

Given the Percentage Composition of a formula, you cancalculate the empirical formula of the substance.

Step 1 assume you have 100 g of substance sothe percentages become grams

Step 2 change grams of each element to molesof atoms of that element

Step 3 set up a formula with the molesexample: C2.4 H4.8

Step 4 simplify the formula by dividing molesby the smallest value C

2.42.4 H

4.82.4 = CH2

Step 5 If ratio becomes…1:1.33 or 1:1.66 multiply by 3

8 • Mathematics of Chemical FormulasOther Mole Problems and Conversions

(8 of 8)

The gas density is often converted to molar mass:

Example:The gas density of a gas is 3.165 g/Liter (at STP). What isthe molar mass of the gas?

Knowing that 22.4 L is 1 mole, you can set up the ratio:3.165 g1 Liter

= molar mass

22.4 L

Other metric conversions you should know:1000 mL1 Liter

1 kg

1000 grams

9 • Chemical EquationsEquation Terms

(1 of 8)

equation condensed statement of facts about achemical reaction.

reactants → substances that exist before a chemicalreaction. Written on the left of the arrow.

→ products substances that come into existence as aresult of the reaction. Written to the rightof the arrow.

word equation an equation describing a chemical changeusing the names of the reactants andproducts.

coefficients a number preceding atoms, ions, ormolecules in balanced chemical equationsthat showing relative #’s.

9 • Chemical EquationsTypes of Reactions and other Terms

(2 of 8)

synthesis A + B → Cdecomposition AB → A + Bsingle replacement AB + C → A + BCdouble replacement AB + CD → AD + CB

precipitate solid that is formed during a reactionspectator ions ion that undergoes no chemical

change during a reactionmolecular equation equation with reactants and products

written as whole moleculesionic equation equation with soluble salts written as

individual ionsnet ionic equation equation with spectator ions removed

9 • Chemical EquationsEnergy Changes

(3 of 8)

EXOTHERMIC• reaction gives off energy• energy is written as a product on the right side of arrow• reaction mixture generally gets warmer or must be cooled

[combustion] CH4 + 2 O2 → CO2 + 2 H2O + energy[freezing] H2O(l) → H2O(s) + energy

ENDOTHERMIC• reaction requires or takes in energy• energy is written as a reactant on the left side of the arrow• reaction mixture takes warmth from surroundings or must be warmed... for example

[electrolysis of water] 2 H2O + energy → 2 H2 + O2

9 • Chemical EquationsShowing Phases in Equations

(4 of 8)

We have seen the phases of matter in earlier chapters. Seepage 212 for atomic pictures.

(s) solid phase may be used to show a ppt.(l) liquid phase(g) gaseous phase

(aq) aqueous phase -- solid or gas dissolved in water(ppt) precipitate -- solid (s) formed during a reaction

use Appendix D or solubility rules to predict whena product is a precipitate.

(l) vs (aq) sugar(l) would be melted sugarsugar(aq) would be sugar water (dissolved)

9 • Chemical EquationsMolecular, Ionic, Net Ionic Equations

(5 of 8)

Consider the compounds: silver nitrate + sodium chromateAg+ NO3– Na+ CrO42–

molecular equation [balance at this stage](use double replacement pattern to predict the products)

2 AgNO3 + Na2CrO4 → Ag2CrO4(s) + 2 NaNO3

ionic equation [use sol. rules to determine (aq) or (s)]2 Ag+ + 2 NO3– + 2 Na+ + CrO42– →

Ag2CrO4(s) + 2 Na+ + 2 NO3–

net ionic equation [remove spectator ions]2 Ag+ + CrO42– → Ag2CrO4(s)

9 • Chemical EquationsWriting Word Equations

Things To Remember(6 of 8)

Example: Write the word equation of...SiO2 + 4 HF → SiF4 + 2 H2O

silicon dioxide + hydrofluoric acid→ silicon tetrafluoride + water

• molecular compounds must be named using mono-, di-,tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca-.

• watch for acids (ionic compounds ... positive ion is H+)acid naming rules apply (-ide = hydro---ic acid, etc.)

• ionic compounds do NOT use di-, tri-, etc. unless theyare part of the ion name (e.g. dichromate, Cr2O72–)ionic cmpds are named as the positive and negative ion.Stock names may be used.

9 • Chemical EquationsWriting Formula Equations

Things To Remember(7 of 8)

Example: Write the formula equation of...sodium metal + water → sodium hydroxide + hydrogen gas

Na° + H2O → NaOH + H2

• metals often are written with the ° symbol to emphasizethat the metal is in the neutral elemental state, not an ion.

• some compounds have common names that you shouldjust know... water, H2O; ammonia, NH3; methane, CH4

• remember the seven diatomic elements so they can bewritten as diatomic molecules when they appear in theirelemental form. Other elemental substances are writtenas single atoms (e.g. sodium metal or helium gas, He)

9 • Chemical EquationsMiscellaneous:

Law of Conservation of Massand Complete Combustion Reactions

(8 of 8)

The law of conservation of mass can be shown bycomparing the total masses of reactants and products.:

Example: Show that the law of conservation of massapplies to the balanced equation...

C3H8 + 5 O2 → 3 CO2 + 4 H2O 44 g + 5(32 g) = 3(44 g) + 4(18 g)

Combustion (burning) implies a fuel and three chemicals:O2, CO2, and H2O. Example: combustion of C3H8 above.

Careful when balancing: C2H5OH... Notice: 6H’s and an O.

Use fractions to show odd #’s of O atoms, 32 O2 = 3 atoms

10 • Mathematics of Chemical EquationsCoefficients and Relative Volumes of Gases

(1 of 8)

Since every gas takes up the same amount of room (22.4 Lfor a mole of a gas at STP), the coefficients in an equationtell you about the volumes of gas involved.

Example: N2(g) + 3 H2(g) → 2 NH3(g)

+

10 • Mathematics of Chemical EquationsHeart of the Problem

(2 of 8)

The “heart of the problem” conversion factor relates theGiven and the Desired compounds using the coefficientsfrom the balanced equation.

Example: N2 + 3 H2 → 2 NH3

♥ could be 3 moles H2

2 moles NH3…which means that every time 2 moles of NH3 is formed, 3moles of H2 must react.

The format is always, moles of Desired moles of Given

10 • Mathematics of Chemical EquationsMass-Mass Problems

Mass-Volume Problems(3 of 8)

Mass-Mass problems are probably the most common typeof problem. The Given and Desired are both masses (gramsor kg).

The pattern is:

Given x molar massof Given x ♥ x

molar massof Desired

In Mass-Volume problems, one of the molar masses is

replaced with 22.4 L1 mole

depending on whether the Given or

the Desired is Liters.

10 • Mathematics of Chemical EquationsMass-Volume-Particle Problems

(4 of 8)

If the Given or Desired is molecules, then the Avogadro’s

Number conversion factor, 6.02 x 1023 molecules

1 mole is used

and the problem is a Mass-Particle or Volume-Particleproblem.

The units of the Given and Desired will guide you as towhich conversion factor to use:

Mass grams or kgVolume Liters or mLParticles molecules or atoms

10 • Mathematics of Chemical EquationsLimiting Reactant Problems

(5 of 8)

In a problem with two Given values, one of the Given’s willlimit how much product you can make. This is called thelimiting reactant. The other reactant is said to be in excess.

Solve the problem twice using each Given… the reactantthat results in the smaller amount of product is the limitingreactant and the smaller answer is the true answer.

Example: N2 + 3 H2 → 2 NH3When 28.0 grams of N2 reacts with 8.00 grams of H2, whatmass of NH3 is produced?

(in this case, the N2 is the limiting reactant)

10 • Mathematics of Chemical EquationsHow Much Excess Reactant is Left Over

(6 of 8)

Example: N2 + 3 H2 → 2 NH3When 28.0 grams of N2 reacts with 8.00 grams of H2, whatmass of NH3 is produced?

(in this case, the N2 is the limiting reactant)

To find out how much H2 is left over, do another lineequation:

Given: 28.0 g N2Desired: ? g H2

subtract the answer of this problem from 8.00 g H2

10 • Mathematics of Chemical EquationsLimiting Reactants

(7 of 8)

It is difficult to simply guess which reactant is the limitingreactant because it depends on two things:

(1) the molar mass of the reactant and(2) the coefficients in the balanced equation

The smaller mass is not always the limiting reactant.

Example: N2 + 3 H2 → 2 NH31 mole (28 g N2) will just react with 3 moles (6.06 g H2)

so, if we react 28.0 g N2 with 8.0 g H2, only 6.06 g H2will be used up and 1.94 g of H2 will be left over.

In this case, N2 is the L.R. and H2 is in XS.

10 • Mathematics of Chemical EquationsLab 16 (Baking Soda Lab) Ideas

(8 of 8)

The reaction: NaHCO3 + HCl → H2O + CO2(g) + NaCl

The Point:We can calculate the mass of NaCl expected from any

starting mass of NaHCO3We can also experimentally measure the mass of NaCl left

over after reacting NaHCO3 and HClThese two values should match. Percent error is a good

way to report how good your experimental value is.

Techniques:We used excess HCl. The CO2 escapes as a gas. The H2O

is driven off by evaporation. Drying thoroughly is impt.

11 • Phases of MatterPressure Definition and Units

(1 of 8)

Pressure is defined as force/area which is measured in

units of lbin2 (psi) or

Nm2 (Pa) , atm , mmHg (torr) or kPa

At sea level, the air pushes down with a pressure of 1 atm = 14.7 psi = 760 mmHg = 760 torr = 101.3 kPa

You can convert from one unit to the other with conversion

factors such as 760 mmHg101.3 kPa

or 101.3 kPa760 mmHg

Example: 500 mmHg x 101.3 kPa760 mmHg

= 66.6 kPa

11 • Phases of MatterManometers and Measuring Pressure

(2 of 8)

A manometer is used to measure pressure.

Pgas

If difference in height is 155 mmand atmospheric pressure = 100 kPa,

convert the difference in the mercury levels to kPa and subtract from the air pressure because you can SEE that the

air pressure is greater than the gas pressure.

A barometer is a special case of a manometer.

11 • Phases of MatterHeating Curve and Energy Changes

(3 of 8)Time (min)

KEPE(s)

(l)

(g)

(s) & (l)

(l) & (g)

KE

KE

PEKE = kinetic energy changes which are times when the heat energy speeds up the molecules.

PE = potential energy changes which are times when the heat energy separates the molecules from solid to liquid or liquid to gas.

11 • Phases of MatterNames of Phase Changes and Energy

(4 of 8)

NRG is REQUIRED NRG is RELEASEDsolid → liquid liquid → solidmelting or fusion freezing

liquid → gas gas → liquidvaporization condensationevaporation or boiling

solid → gas gas → solidsublimation solidification

The energy involved in the phase change is calculated usingheat of fusion (solid → liquid or liquid → solid)heat of vaporization (liquid → gas or gas → liquid)

11 • Phases of MatterAssumptions of the Kinetic Molecular Theory

(5 of 8)

• Gases are tiny particles separated by large areas of emptyspace.

• Molecules are in constant, random motion.• Pressure results from collisions of the gas molecules with

the walls of the container.• Molecules of an “ideal” gas show no attraction or repulsion

for each other. “Real” gases have intermolecular forces ofattraction (IMF’s) that are strongest in solids and weakestin gases.

• Some molecules are moving fast, some are moving slowly,temperature is a measure of the average KE…corresponds to the average speed.

(Relate this to the KE distribution curve Fig 11-13.)Phase changes are the balance between KE and IMF’s.

11 • Phases of MatterVapor Pressure, IMF’s, and KE

(6 of 8)

Vapor pressure is the push exerted by the particles of vaporthat escape from a liquid (or solid). It is a good measureof the IMF’s of a substance.

Large vapor pressure = small IMF’s(particles can easily escape the liquid to become vapor)

Small vapor pressure = large IMF’s(particles are tightly held as a liquid)

Greater temperature (KE) allows more particles to escapeagainst the IMF’s and so as temp ↑, vapor pressure ↑.

Vapor pressure = power to form bubbles.

11 • Phases of MatterBoiling Point

(7 of 8)

Boiling occurs when the vapor pressure of a substance =the air pressure above the liquid.

You can make a liquid boil by increasing the vaporpressure of the liquid (heating it up) or by lowering theair pressure above the liquid. Both were demonstrated.

Use a chart of vapor pressure values or a graph todetermine the vapor pressure needed to boil at varioustemperature and pressure conditions..

↑ boiling points (BP) ~ ↑IMF’s ~ ↓vapor pressures.

Altitude (low air pressure) lowers the boiling temperature ofwater in an open container (increases cooking time).

11 • Phases of MatterSolids, Liquids, and Gases

(8 of 8)

For most substances, the solid is the most dense form ofmatter because the atoms are packed together tightly.

SOLID LIQUID GASWater is unique: when the solid forms, the molecules

spread out and form a stable, but less dense pattern, ice.

SOLID WATER

12 • The Gas LawsBoyle’s Law (P and V)

(1 of 8)

General: When P↑, V↓ (inversely proportional)Formula: P·V = constant or P1V1 = P2V2

Restrictions: P1 and P2 must be in the same unitsV1 and V2 must be in the same units

Convert pressures using conversion factors using the factthat 1 atm = 760 mmHg = 760 torr = 101.3 kPa = 14.7 psi

psi = lbin2

Example: 730 mmHg x 101.3 kPa760 mmHg

= 97.3 kPa

12 • The Gas LawsBoyle’s Law Lab

(2 of 8)

Graphically:

P

V

P

1/V

In our lab, we had to add the atmospheric pressure to ourmeasurements because tire gauges only measure thepressure ABOVE atmospheric pressure.

Consistent (“ good”) data form a straight line (P vs. 1V

).

12 • The Gas LawsKelvin Temperature Scale

(3 of 8)

K = °C + 273 °C = K – 273Examples: 0 °C + 273 = 237 K

25 °C + 273 = 298 K100 °C + 273 = 373 K

300 K – 273 = –27 °C

The Kelvin scale is used in gas law problems because thepressure and volume of a gas depend on the kinetic energyor motion of the particles.

The Kelvin scale is proportional to the KE of theparticles… that is, 0 K (absolute zero) means 0 kineticenergy. 0 °C is simply the freezing point of water.

12 • The Gas LawsCharles’ Law (V and T)

Gay-Lussac’s Law (P and T)(4 of 8)

Charles’ LawGeneral: When T↑, V↑ (directly proportional)

Formula:VT

= constant or V1T1

= V2T2

Restrictions: T must be in KelvinsV1 and V2 must be in the same units

Gay-Lussac’s LawGeneral: When T↑, P↑ (directly proportional)

Formula:PT

= constant or P1T1

= P2T2

Restrictions: T must be in KelvinsP1 and P2 must be in the same units

12 • The Gas LawsThe Combined Gas Law

(5 of 8)

Formula:P·VT

= constant or P1·V1

T1 =

P2·V2T2

Restrictions: T must be in KelvinsV1 and V2 must be in the same unitsP1 and P2 must be in the same units

STP (“standard temperature and pressure”) is often used asone of the two conditions

T = 0 °C = 273 K P = 1 atm = 760 mmHg = 101.3 kPa

Each of the three gas laws is really a special case of thislaw.

Example: If T1 = T2, the law becomes P1V1 = P2V2

12 • The Gas LawsThe Ideal Gas Law

(6 of 8)

Formula: P·V = n·R·T or PV = nRTwhere P = pressure

V = volumen = number of molesR = the ideal gas constantT = temperature (in Kelvins)

The value of R depends on the P and V units used.

R = PVnT

so you can use the molar volume info to calculate R

R = (101.3 kPa)(22.4 L)

(1 mole)(273 K) = 8.31

L·kPamol·K

R = 62.4 L·mmHgmol·K

= 0.0821 L·atmmol·K

12 • The Gas LawsDalton’s Law of Partial Pressure

(7 of 8)

When you have a mixture of gases, you can determine thepressure exerted by each gas separately. This is calledthe partial pressure of each gas.

Since each gas has the same power to cause pressure (seecard #8) the partial pressure of a gas depends on howmuch of the mixture is composed of each gas (in moles)

Example: Consider air, a mixture of mostly O2 and N2moles O2

moles total =

PO2Ptotal

moles N2

moles total =

PN2Ptotal

Also: Ptotal = PO2 + PN2This idea is used when a gas is collected over water

Patm = Pgas + PH2O PH2O is found on a chart

12 • The Gas LawsWhy Do All Gases Cause the Same Pressure?

Graham’s Law(8 of 8)

The gas laws work (to 3 significant digits) for all gases…that is, all gases have the same power to cause pressure.

At the same temperature, the KE of each gas is the same.KE = 1/2 mass·velocity2… if two particles have different

masses, their velocities are also different. So…

SMALL particles move FAST mv2

LARGE particles move SLOWLY mv2

We can use this idea with numbers as well: (Graham’s Law)KEA = KEB mAvA2 = mBvB2

[another version of this formula is on pg 323 of the text]

13 • Electron ConfigurationsElectron Energy Levels

(1 of 4)

The electrons in an atom exist in various energy levels.

When an electron moves from a lower energy level to ahigher energy level, energy is absorbed by the atom.When an electron moves from a higher to a lower energylevel, energy is released (often as light).

Neils Bohr was able to determine the energy levels ofhydrogen by the visible light energy that is released whenthe electron drops from 3 → 2 (red light), 4 → 2 (blue-green), 5 → 2 (blue-violet) and 6 → 2 (violet).

Transitions to level n = 1 are too high energy to see (UV).Transitions to level n = 3 are too low energy to see (IR).

13 • Electron ConfigurationsShowing Electron Arrangments:

Orbital Diagrams and Electron Configurations(2 of 4)

1s

2s2p

3s3p

4s4p 3d

A shorthand notation is the electron configuration:1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6, etc.

13 • Electron ConfigurationsFilling Orbitals

(3 of 4)

Three rules define how the orbitals fill:

The Pauli Exclusion PrincipleEach orbital can be occupied by no more than twoelectrons.

The Aufbau PrincipleThe electrons occupy the lowest energy orbitalsavailable. The “Ground State” for an atom is whenevery electron is in its lowest energy orbital.

Hund’s RuleWhen more than one orbital exists of the same energy(p, d, and f orbitals), place one electron in each orbitalbefore doubling them up.

13 • Electron ConfigurationsValence Electrons

(4 of 4)

The valence electrons are the outermost electrons… thosefarthest from the nucleus. They have the largest principalquantum number, n.

These electrons occupy the s and p orbitals in the highestenergy level. These orbitals are called the valence orbitals.

The columns of the periodic table are labeled, I, II, III, IV,V, VI, VII and VIII (ignoring the transition and rare earthelements). This label tell you the number of valenceelectrons of every element in that column (except He.)

The valence electrons are important in how atoms bond.Note that filled energy levels match up with the noble gases.

14 • The Periodic TableTerms about the Periodic Table

People of the Periodic Table(1 of 8)

period a horizontal row of the tablegroup / family a vertical column of the table

periodic law properties repeat if you arrange elementsby atomic mass. The modern periodiclaw arranges elements by atomic number

Döbereiner triads of similar elementsNewlands Law of Octaves for similar elementsMeyer close to a modern periodic table

Dimitri first to publish, predicted missingMendeleev elements, his table was very detailed

14 • The Periodic TableFamilies (Groups) of the Periodic Table

(2 of 8)

hydrogen a family by itself because it acts like bothgroup I and group VII

alkali metals Family I… forms 1+ ionsalkaline earth Family II… forms 2+ ions metalshalogens Family VII (salt formers)… forms 1– ionsnoble gases Family VIII… He, Ne, Ar are inert

representative Families I – VIII elements s-block (I & II) and p-block (III – VIII)main transition elements with unfilled inner shells elements d-block (often form colored compounds)inner transition f-block (lanthanoids and actinoids) elements

14 • The Periodic TableTerms used in the Trends

(3 of 8)

ionization energyenergy needed to remove an electron from an atomex: Li + energy → Li+ + e–

atomic radius size of an atomionic radius size of an ion

negative ions get larger; positive ions get smaller

metallic character (compared to nonmetals)low ionization energies… form positive ionslow electronegativities high lustereasily deformed (malleable and ductile)good conductors of heat and electricity

electronegativity… the “pull” or attraction for electrons

14 • The Periodic TableTrends of the Periodic Table

(4 of 8)

Trends you should know and be able to explain:

ionization energy increaseselectronegativity increases

atomic radius increasesmetallic character increases

14 • The Periodic TableExplaining The Size of Atoms and Ions

(5 of 8)

The size of an atom depends on the electron cloud… theaverage distance of the valence e–’s from the nucleus.

Three important factors:the e– - e– repulsion … making the atom largerthe p+ - e– attractions… making the atom smallerthe principal quantum number, n…as n increases, the average distance of the valence e–’sfrom the nucleus increases… making the atom larger

• across a period… more p+’s… more attraction… smaller• down a family… n increases… e–’s farther… larger• + ion… lose e–’s… less repulsion… smaller

–

14 • The Periodic TableExplaining Ionization Energy

(6 of 8)

Ionization energy trends follow the size trends:As atoms or ions get larger, the electron being removed is

farther from the nucleus… the attractions are less… theenergy needed is less… the ionization energy is less.

Ionization energy greatly increases when you startremoving electrons from an inner shell (n decreases).

Moving across a period, two other factors come into play:• “p” e–’s are higher energy… require less energy to

remove than “s” e–’s with the same quantum number, n• e–’s in filled orbitals are easier to remove than e–’s in

singly occupied orbitals because of e– - e– repulsions.

14 • The Periodic TableIsoelectronic Species

(7 of 8)

The elements before and after a noble gas form ions bygaining or losing electrons until they have the same

electron configuration as the noble gas.

N3– O2– F– Ne Na+ Mg2+ Al3+

all have the electron configuration: 1s2 2s2 2p6

We say these ions and atoms are isoelectroniciso means “same” and electronic means “electrons”

Using ideas from Study Card #5, we see that if the electronsare the same but there are increasing numbers ofprotons, the increased attractions cause the sizes todecrease.

14 • The Periodic TableClues from the Periodic Table

(8 of 8)

The group or family number (I through VIII) tells you…• the number of valence electrons of the elements in each

family (except He which only has 2 valence electrons)• the ion that commonly forms• the bonding capacity (the number of bonds it will form)

I II III IV V VI VII VIII# of valence e– 1 2 3 4 5 6 7 8ion formed 1+ 2+ 3+ * 3– 2– 1– none

bonding capacity 1 2 3 4 3 2 1 0

IUPAC family # 1 2 13 14 15 16 17 18

* Family IV usually shares e–’s rather than forming ions

15 • Chemical Bonding Three Types of Bonding

(1 of 8)

There are three general classes of bonds that form between atoms. You can predict which will form by classifying the atoms as metals or nonmetals:

metal + metal metallic bond Au-Ag alloy metal + nonmetal ionic bond MgCl2 nonmetal + nonmetal covalent bond SO2 or CH4

Some compounds can contain both ionic and covalent bonds such as K2SO4... the sulfate ion is held together with covalent bonds... the potassium ions are ionically bonded to the sulfate ions.

Acids are exceptions... they are ionic only when dissolved.

15 • Chemical Bonding The Ionic Bond

(2 of 8)

Many ions can be explained because they have gained or lost electrons and attain a noble gas configuration.

For example: P3– S2– Cl– Ar K+ Ca2+ all have the same electron arrangement: 1s2 2s2 2p6 3s2 3p6

The importance of this configuration is that this is one reason why ions form. After these ions form, they stick together in a crystal lattice because opposites attract: + - + - + - + - - + - + - + - + There are other reasons why some + - + - + - + - ions (ex: Cu+ or Zn2+) form. - + - + - + - +

15 • Chemical Bonding The Covalent Bond

(3 of 8)

The covalent bond between two atoms depends on the balance of attractions between one atom’s + nucleus and the other atom’s – electrons and the proton-proton repulsions as well as electron-electron repulstions.

Distance between nuclei

PE

If two atoms have half-filled orbitals, the interactions balance at a small enough distance so the e–’s can be close to both nuclei at the same time... this is a covalent bond.

15 • Chemical Bonding Lewis Electron Dot Structures

(4 of 8)

Lewis symbols consist of the atomic symbol surrounded by valence electrons. The four sides represent the four valence orbitals. Atoms are usually shown in their excited states. (Families II, III, & IV can also be in their “ground state.”)

Ions include brackets and charges. Positive ions show no valence electrons while negative ions show an octet.

[Li] [Mg] [ O ]2+ 2–+

15 • Chemical Bonding Drawing Electron Dot Structures The (Chris) Bednarski Method

(5 of 8)

Example: CO2 Draw the Lewis symbols for each atom. Connect the unpaired electrons. Clean up your drawing.

O C O

O C O

O C O

15 • Chemical Bonding Comparing Ionic & Molecular Substances

(6 of 8)

Compound Molecular Ionic Conducts as Solid NO NO Conducts as Liquid NO YES Conducts in Solution NO YES Conducts as Gas NO YES

Hardness soft hard

MP / BP low high

Bonding covalent ionic

Examples He, CH4, CO2, NaCl, KI, C6H12O6 AgNO3

15 • Chemical Bonding Electronegativity and Polar Bonds

(7 of 8)

You will be given a chart of electronegativity values. Memorize (F = 4.0) (O = 3.5) and (Cl = 3.0). The noble gases have no values… no bonds. Large electronegativity in the upper right of the per. table and small in the lower left portion of the table. Classify the bond between any two atoms by subtracting their electronegativity values (∆e) Non-polar covalent 0 < ∆e < 0.5 Polar covalent 0.5 � ∆e � 1.7 Ionic ∆e > 1.7 The more electronegative atom is more negative . Polar covalent bonds have partial charges δ+ and δ–

15 • Chemical Bonding Shapes and Polar Molecules

(8 of 8)

Use VSEPR theory to predict the shape of molecules. The Steric Number (the # of lone pairs + bonded atoms) relates the shape of the electron pairs around a central atom.

[1=linear, 2=linear, 3=trigonal planar, 4=tetrahedral]

If each shape is symmetrical, the bond dipoles will cancel resulting in a nonpolar molecule.

If a shape has lone pairs of electrons on the central atom, the shape is often unsymmetrical, the molecule is polar.

Polar molecules and ions dissolve well in polar solvents while nonpolar molecules dissolve in nonpolar solvents.

“Like Dissolves Like”

24 • Organic Chemistry Historical Ideas

(1 of 12)

Chemicals from living things were thought to contain a “vital force” that could not be duplicated in the lab. This changed with Friedrich Wöhler who mixed cyanic acid (HCNO) with ammonium hydroxide making ammonium cyanate (NH4CNO).

C

O

NH2 NH2

urea

He usually allowed the salt solution to evaporate overnight, but tried heating it to hurry the process. The result was a crystal that he recognized as urea (H2NCONH2).

The modern view of organic chemistry is the chemistry of carbon compounds. C is the key element. It can form four bonds and that are very strong bonds due to its small size.

24 • Organic Chemistry Alkane Series -- Saturated Hydrocarbons

(2 of 12)

The alkanes (paraffins) follow the formula: CnH2n+2: These molecules contain ONLY single bonds. They are said to be “saturated” with hydrogens.

Memorize these prefixes also used with alkenes & alkynes. CH4 methane C6H14 hexane C2H6 ethane C7H16 heptane C3H8 propane C8H18 octane C4H10 butane C9H20 nonane C5H12 pentane C10H22 decane

Given a formula, you can tell that it contains only single bonds because it fits the alkane formula.

As the molecules increase in size, they tend to be liquids and

24 • Organic Chemistry Structural Formulas Can Be Misleading

(3 of 12)

CH4, can be drawn using a structural formula. This can be misleading. The molecule is not flat with bond angles of 90°. You must be aware of the 3-D structure and the 109.5 °° bond angles.

C

H

H

H

H

C

Cl

H H

Cl

C

H

H Cl

Cl

For example, there is only one isomer of dichloromethane, but you can draw it at least two ways.

Building models of the molecules is an important part of strengthening this skill.

24 • Organic Chemistry Alkenes and cis- /trans- Isomerism

(4 of 12)

Alkenes contain 1 double bond. The formula is CnH2n. They are said to be “unsaturated” (like unsaturated fats). The double bond can be broken and more hydrogens added.

C

H

C

H H

H

ethene

Since double bonds cannot easily rotate (due to the double bond) cis- and trans- isomers can be formed.

Example: 1,2-dichloroethene can be built two ways.

C

H

C

Cl Cl

H cis -1,2-dichloroethene

(a polar molecule)

CH

C

Cl H

Cl trans-1,2-dichloroethene

(a nonpolar molecule)

24 • Organic Chemistry Alkynes, Alkadienes, and Cyclic Hydrocarbons

(5 of 12)

Alkynes contain 1 triple bond (unsaturated). Formula: CnH2n-2.

CH C H ethyne (acetylene)

The triple bond is linear, so no cis/trans isomerism occurs.

Alkadienes are molecules with two double bonds. They have the same formula as the alkynes, CnH2n-2.

Example: C4H6 is named 1,3-butadiene because the double bonds start on carbons #1 and #3.

CH2C

CCH2

H

H

CH2

CH2

CH2

Cyclic compounds contain rings having the same formula as the alkenes, CnH2n. Example: cyclopropane, C3H6.

24 • Organic Chemistry Naming Organic Compounds

(Organic Nomenclature Using IUPAC Rules) (6 of 12)

The basic idea is to name the molecule after the longest continuous chain of carbon atoms. Side groups are listed with #’s to indicate the C atom to which they are attached.

Side Groups : -Cl chloro -Br bromo -I iodo -CH3 methyl -C2H5 ethyl -C3H7 propyl, etc.

di- = 2 groups tri- = 3 groups tetra- = 4 groups

2,2,3-tribromobutane (not 2,3,3-) Note that we # the carbons from whichever end results in the smallest numbers.

C C C C

Br

H

Br

Br

H

H

H

H

H

H

24 • Organic Chemistry Common Errors in Drawing/Naming Structures

1-methylsomething & 2-ethylsomething (7 of 12)

A common error when drawing isomers of pentane, C5H12, is naming this structure as 2-ethylpropane. (a chain of 3 C’s with an ethyl group) The longest chain is really four C’s, & should be named 2-methylbutane.

C C C H

H

CH2

H

H

H

H

H

CH3

A similar error is to draw and name “1-methylsomething”.

C C O

H

H

H

H

H 5 bonds to C

Two more tips… double check that each C has four and only four bonds. Also, remember that N and O atoms have lone pairs of e-‘s although they are seldom drawn. (Impt. for steric #!)

24 • Organic Chemistry Aromatic Compounds

Benzene and its Derivatives (8 of 12)

CH

CHCH

CH

CHCH

CH

CHCH

CH

CHCH

two resonance structures

Benzene, C6H6, is unique. It can be drawn as shown, but the actual structure involves a circular pi bond (sp2 orbitals & delocalized e-‘s).

Benzene is also shown with a circle as the pi bond.

The carbon #’s can be used for substituted benzene. Example: dichlorobenzene 1,2- is known as the ortho- position 1,3- is known as the meta- position 1,4- is known as the para- position

5

4

3

2

1

6

Paradichlorobenzene: the main ingredient in some moth balls.

24 • Organic Chemistry Functional Groups I Alcohols and Ethers

(9 of 12)

Alcohols General formula: R-O-H [R ≈ Rest of molecule]

C1C3 C

2

C 1

C1

C atoms are classified as primary (1), secondary (2), or tertiary (3) by the number of C atoms it is bonded to. A primary alcohol has the -OH group bonded to a primary carbon, etc.

This is not a base because the -OH is covalent, not ionic. Naming: group + “alcohol” (e.g. ethyl alcohol or ethanol)

Ethers General formula: R-O-R’ [R’ can = R, but not H] Naming: two groups + “ether” diethyl ether was the 1st effective surgical and dental anesthetic.

CH3

CH2

OCH2

CH3

24 • Organic Chemistry Functional Groups II

Aldehydes and Ketones (10 of 12)

Aldehydes Ketones General formula:

RC

H

O

RC

R'

O

Naming: names end in “al”

or “aldehyde” methanaldehyde (formaldehyde)

names end in “one”

propanone (acetone)

Aldehydes and ketones both have a C=O group (carbonyl group). Aldehydes have it on an end carbon. Ketones have it on a middle carbon. Reactions: Primary alcohols can be oxidized into aldehydes. Secondary alcohols into ketones .

24 • Organic Chemistry Functional Groups III

Carboxylic Acids and Esters (11 of 12)

Carboxylic Acids Esters General formula:

RC

OH

O

RC

OR

O

Naming: names end - “oic

acid” ethanoic acid (acetic acid)

names end - “ate” ethyl acetate (acetic acid + ethyl alcohol)

Reactions: Acids can be made by oxidizing aldehydes. Esters are formed (“esterification”) from a carboxylic acid & an alcohol. Water is removed (a “condensation” reaction).

Esters often have pleasant, agreeable odors (e.g. banana.)

24 • Organic Chemistry Functional Groups IV

Amines & Amides (12 of 12)

Amines Amides General formula:

R

N H H

RC

NH2

O

Naming: names contain

“amino” or end in “amine”

aminomethane (methylamine)

names end in “amide”

acetamide

The N may have 1 or 2 or all 3 H atoms replaced with groups. The lone pair on the N atom makes these molecules basic. Your body needs certain amines “vital amines” ≈ “vitamins.”

24 • Organic Chemistry -- Extra Optical Isomers

Chiral Compounds (1 of 4)

CH CH2

CH2

CH3

CH3

CH

2

CH3

3-methylhexane

Some molecules have the ability to rotate polarized light.

These molecules can be recognized by a C atom (the chiral carbon) bonded to four different groups.

This carbon is bonded to H, methyl, ethyl, & propyl groups.

You can build two versions of this molecule that are “nonsuperimposable mirror images of each other.” One will rotate light clockwise, one counterclockwise.

In biology, these are called dextro- and levo- (D and L) forms.

23 • Organic Chemistry -- Extra Common Names You Should Know About

(2 of 4)

ethene is also called ethylene propene is also called propylene

CH2

CH3

CHCH3

CH3

2-methylbutane is also called isopentane. “Iso-“ means the same… the same two methyl groups come branch from C #2. 2-methylpentane is isohexane, etc.

CH3C

CH3

CH3

CH3

2,2-dimethylpropane is called neopentane. These common names show up occasionally in names… such as in isopropyl alcohol.

24 • Organic Chemistry -- Extra Polymers I

Monomers & Addition Polymerization (3 of 4)

Monomer = one part Polymer = many parts

C C

H

H H

H

“ethylene”

One kind of polymer is made up of monomers that contain a double bond. The double bond can break and we can ADD to it… “Addition polymerization.”

C C

H

H*

H*

H

+ C C

H

H H

H

→ C C

H

H

C

H

C*

H

H

H H

*H

Different monomers form different polymers. This polymer would be called polyethylene. Replace on H on the monomer with Cl and you can make polyvinyl chloride, “PVC.”

24 • Organic Chemistry -- Extra Polymers II

Copolymers & Condensation Polymerization (4 of 4)

Another polymerization involves condensation reactions.

C

OH

OC

OH

OR

a di-acid

COH

COH

R

H

H

H

H a di-alcohol (a glycol)

Esters form from an acid and an alcohol. Using a di-acid and a di-alcohol, you can make a continuous chain by removing water molecules. The resulting polymer is called a polyester.

Soda bottles are made from a polyester, polyethylene terephthalate ester (PETE).

Nylon (a polyamide) can be made from a di-amine & a di-acid.

26 • Nuclear ChemistryThe People

(1 of 12)

• Wilhelm Roentgen (1845-1923) discovered X-rays, ahigh energy form of light. (1895)

• Henri Becquerel (1852-1909) found that uranium oresemit radiation that can pass through objects (like x-rays)and affect photographic plates. (1896)

• Marie Sklodowska Curie (1867-1934) Marie and Pierreworked with Becquerel to understand radioactivity. Thethree shared a Nobel Prize in Physics in 1903. Marie wona second Nobel Prize in Chemistry in 1911 for her workwith radium and its properties.

• E. O. Lawrence invented the cyclotron which was used atUC Berkeley to make many of the transuranium elements.

26 • Nuclear ChemistryTerms I-- Radioactivity

(2 of 12)

radioactivity the spontaneous breakdown of atomicnuclei, accompanied by the release ofsome form of radiation (also calledradioactive decay)

half-life time required for half of a radioactivesample to decay

transmutation one element being converted into anotherby a nuclear change

nuclides isotopes of elements that are identified bythe number of their protons and neutrons

emission the particle ejected from the nucleus.

26 • Nuclear ChemistryTerms II--Radioactivity

(3 of 12)

decay series the sequence of nuclides that an elementchanges into until it forms a stable nucleus

radioactive using half-life information to determinedating the age of objects. C-14/C-12 is common.

nuclear fission large nucleus breaking down into pieces ofabout the same mass

nuclear fusion two or more light nuclei blend to form oneor more larger nuclei

26 • Nuclear ChemistryTypes of Radiation

(4 of 12)

Alpha particles are the same as a helium nucleus, 42He, with

a mass of 4 amu. It travels about 1/10th the speed of lightand is the most easily stopped of the three particles (a sheetof paper will stop them). It is the least dangerous.

Beta particles are high speed electrons, 0-1e, with a mass of

0.00055 amu and travels at nearly the speed of light. It canbe stopped by a sheet of aluminum. It is more penetratingand therefore more dangerous than alpha.Gamma rays are extremely high energy light, γ, with nomass, and are the most penetrating (several cm’s of lead areneeded to stop them). They can cause severe damage.

26 • Nuclear ChemistryHalf-Life Problems

(5 of 12)

In each half-life problem there are basically four variables:• total time • half-life• starting amount • ending amount

64g 32g 16g 8g 4g 2g 1g 0.5g 0.25g

Question:If you have 0.25 g of a radioactive substance with a halflife of 3 days, how long ago did you have 64 grams?

Answer: Draw the chart to determine the number of half-lives to get from the ending amount to the starting amount…each half-life is worth 3 days…24 days.

26 • Nuclear ChemistryHalf-Life(6 of 12)

The time it takes for half of a radioactive substance to decay.The graph has a characteristic shape:

time

#

The time it takes for the amount or the activity of thesubstance to drop to half is the same WHEREVER you starton the graph.

Half-lives can range from microseconds to thousands ofyears and is characteristic of each substance.

26 • Nuclear ChemistryNuclear Equations

(7 of 12)

Memorize the symbols for the important particlesalpha beta positron neutron42He

0-1e

0+1e

10n

Decay means the particle is on the right side of the equation:example: alpha decay of U-238

238 92U →

42He +

234 90Th

The 234 and 90 are calculated… the Th is found on theperiodic table (find the element with atomic # = 90).Several neutrons can be shown together and written as…

3(10n) and would be counted as

30n in the equation.

26 • Nuclear ChemistryHow Each Type of Decay Can Stabilize an

Unstable Nucleus(8 of 12)

Certain values of p+’s and no’s in the nucleus are stable. Anucleus can be unstable (radioactive) for 3 reasons:• the nucleus has too many protons compared to neutrons

solution: positron decay(change a proton into a neutron and a positive electron…

…a positron)

• the nucleus has too many neutrons compared to protonssolution: beta decay

(change a neutron into a proton and a negative beta particle)

• the nucleus is too big (too many protons and neutrons)solution: alpha decay (lose 2 p+ and 2 n°)

26 • Nuclear ChemistryUses of Radioactivity

(9 of 12)

Radioactive Dating: In every living thing there is aconstant ratio of normal C-12 and radioactive C-14. Youcan calculate the time needed to change from what isexpected to what is actually found.Radioisotopes: Many substances can be radioactive andthen followed as they move through the body.

Fission Reactors: Current nuclear reactors use fissionreactions to produce heat which is used to turn water intosteam and drive turbine engines that produce electricity.The Sun and Stars are powered by nuclear fusion… this isrelated to the fact that the most abundant element in theuniverse is hydrogen… followed by helium.

26 • Nuclear ChemistryFission and Fusion Reactions

(10 of 12)

U-235 is “fissionable” which means it can be split whenbombarded by neutrons.

235 92U +

10n →

141 56Ba +

9236Kr + 3

10n + energy

The fact that each splitting nucleus can emit neutrons thatcan split other nuclei is the basis for the “chain reaction.”“Breeder reactors” use different isotopes. See page 774.

Fusion in the Sun involves several steps that can be

summed up as: 4(11H) →

42He + 2

01e + energy

Thermonuclear devices use isotopes of hydrogen

(deuterium and tritium): 21H + 31 H →

42He +

10n + energy

26 • Nuclear ChemistryEnergy–Mass Conversions

(11 of 12)

Einstein’s famous equation, E = mc2, is the basis forexplaining where the energy associated with nuclear changescomes from.

When a nuclear change occurs, the mass of the products isslightly less than the mass of the reactants. This loss in massis called the mass defect.

E = the energym = the mass defectc = the speed of light, 3.00 x 108 m/s

As stated in your text, 1 kg of mass converted into energywould be equivalent to burning 3 billion kg of coal!

26 • Nuclear ChemistryWhat Happens During

Beta and Positron Decay(12 of 12)

During beta decay,1 neutron changes into 1 proton + 1 negative beta particle(The atomic # increases by one due to the new proton. Themass # is unchanged… a neutron is gone. To maintainelectrical neurtality, a negative beta particle is also formed.)

Example: 235 92U →

0–1e +

235 93Np

During positron decay,1 proton changes into 1 neutron + 1 positron particle(The atomic # decreases by one due to the loss of a proton.Since it changed into a neutron, the mass # is unchanged.)

Example: 235 92U →

0+1e +

235 91Pa