unit 0 basics of chemistry - ms. diorio's ap chemistry...

Unit 0Basics of Chemistry

Ms. DiOrio

Rm 109

Ch. 1: Chemical Foundations

Lab Safety and Techniques

Lab Safety

• Waft

• Always wear:• Safety Goggles

• Hair Tied Back

• Closed toed shoes

• Apron or gloves when necessary

• Never point heated substance towards another student

• Rinse acid or base spill on skin immediately with water!

• Notify instructor of spills, breakage, fire, etc.

• Check glassware to see if cool• Always wash hands when finished with

lab• NO HORSEPLAY!

Lab Techniques

• Massing on Balance

• Lighting Bunsen Burner

• Reading Glassware

• Titrations

• Preparation of solution from solid

• Preparation of solution from concentrated solutions

Lab Glassware

• Buchner funnel

• Filtration flask

• Volumetric flask

• Volumetric pipet

• Mohr pipet

Lab Equipment

• Drying oven

• Spectrophotometer

• Hot/stir plate

• Desiccator

Measurement

SI Units

• Systéme International d’Unités

• A different base unit is used for each quantity

Metric System

• Prefixes convert the base units into units that are appropriate for the item being measured

Mass vs. Weight

Mass• Measures the amount of material• Balance UsedWeight• The force of the object due to gravity• Scale Used

• **You can be weightless but you cannot be massless

Temperature

• A measure of the average kinetic energy of the particles in a sample

• Measured in Kelvin, Celsius, and Fahrenheit

Temperature

Celsius• Based on the properties of water (0℃ melting point and 100℃ boiling point)Kelvin• Based on the properties of gases• There are no negative Kelvin temperaturesFahrenheit• No used in scientific measurements

Temperature Conversions

MEMORIZE!!!!

℉ =95 ℃ + 32

℃ =59 (℉ − 32)

𝐾 = ℃ + 273.15

Volume

• The most commonly used metric units for volume are the liter (L) and the milliliter (mL)• A liter is a cube 1 dm long on each side

• A milliliter is a cube 1 cm long on each side

Derived Units

• Made by combining the seven SI units to make new units

• Ex: Volume, density, speed, force, pressure, and energy

Measurement Derived Unit Symbol

Density Grams/milliliter g/mL

Concentration Moles/liter mol/L

Area Meters cubed m2

Volume Decimeter squared dm3 ( = Liter )

Speed (velocity) Meters/second m/s

Acceleration Meters/second sq. m/s2

Density

• A physical property used to help characterize a substance

• Units of g/cm3 or g/mL

Uncertainty in Measurements

• Different measuring devices have different uses and different degrees of accuracy and precision

Significant Figures

• Digits that are measured

• Determines how precise and instrument is

• When measuring à go one place beyond markers (except with electronic equipment)

• When calculating à round to sig figs at the end of calculation• You do not want to overstate the accuracy of answers

• Avoiding rounding error throughout calculations

Significant Figures

1. All nonzero digits are significant

2. Zeroes between two significant figures are themselves significant

3. Zeroes at the beginning of a number are never significant

4. Zeroes at the end of a number are significant if a decimal point is written in the number

Atlantic-Pacific Rule

1410

25085932

2 1

1 -

2 1 -

5 4 3 2 1

2 sig figs

1 sig fig

2 sig figs

5 sig figs

1.4.080.002508.593005400.

- 12 sig figs

1 sig fig

3 sig figs

6 sig figs

4 sig figs

1 2 3 4 5 6

- - - 1 2 3

1 2 3 4

1 2

Significant Figures

Multiplication / Division

• Answers are rounded to the least number of sig figs in any of the numbers used in the calculation

Addition / Subtraction

• Answers are rounded to the least significant decimal place

**Remember, carry all digits through until the very end of a multistep calculation

Sig Fig Review

v 0.0234v 0.0230v 0.12v 1.000

v 1.2v 100.5v 5050v 5000

3 Sig Figs

3 Sig Figs

2 Sig Figs

4 Sig Figs

2 Sig Figs

4 Sig Figs

3 Sig Figs

1 Sig Fig

Prefix Conversion Steps

1. Write down what is given in scientific notation

2. Put the units given in the denominator3. Put the units wanted in the numerator4. Put a 1 with the larger unit and the positive 10x

difference with the smaller unit5. Solve

1. Exponents in numerator are added, exponents in denominator are subtracted

Prefix Conversion Practice

v Convert 467g into Mg

vConvert 0.0056 pm into cm

4.67 x 102 g 1 Mg106 g

5.6 x 10-3 pm 1 cm109 pm

= 4.67 x 10-4 Mg

= 5.6 x 10-12 cm

Dimensional Analysis

• Converts one quantity to another

• Most commonly utilized conversion factors• Ex: 1 in = 2.54 cm

Dimensional Analysis Steps

1. Start with the number and units that are given

2. Set up a conversion with the units given in the denominator to cancel out

3. Put units you want in the numerator

4. Solve

Dimensional Analysis with Cubes

• Don’t forget, if a unit is cubed your conversion must also be cubed

Convert 4.5 mm3 to nm3

4.5 mm3 (106 nm)3

(1 mm)3 = 4.5 x 1018 nm3

Challenge Question

• What is the length in micrometers of a 2.4 g cube of gold whose density is 19.3 g/mL? 2.4 g 1 mL 1 cm3

19.3 g 1 mL= 0.124 cm3

Length = 0.124 cm31

= 0.499 cm 104 nm1 cm

= 0.499 x 104 nm = 5.0 x 103 nm

Error Analysis

Error

Accuracy

• Refers to the proximity of a measurement to the true value of a quantity

Precision

• Refers to the proximity of several measurements to each other

Error

• The difference between a measured value and a true value

Error & Accuracy vs. Precision

• If the significant digits of a measurement (expressed in however many sig figs) agree with the true value, then the measurement is accurate

• A poorly used instrument which delivers 6 significant digits may be precisebut will not be helpful in obtaining an accurate value

• In general, the greater the number of measurements, the closer the estimate and the smaller the error

Types of Error

Random• Non-reproducible and not the same magnitude for any two measurements (except

by chance)• Often tend to cancel out if enough samples are taken• Repetition of work is the best way to minimize the effects on accuracyDeterminate• Affects each measurement in the same way and to the same extent• Includes systematic, proportional, and constant error

Types of Determinate ErrorTypes of Error Explanation Can be reduced by…

SystematicOften occurs in one

directionDifferent magnitude

Calibration of equipment

Proportional Constant relative error such as incorrect concentration

Using controls with known values

ConstantReproducible error

Seen with fixed amount of reagents used

Using a “blank”

Error

All measurements have some amount of error.• This is due to both human error (estimation)

and error built into an instrument• Most instruments are labeled with the

uncertainty

**If not stated, assume uncertainty is ±1 in the last significant digit

Instrument Uncertainty (±)

Balance, top loading 0.01 g

Balance, analytical 0.001 g

Thermometer 0.2

10 mL graduated cylinder 0.1 mL




100 mL gas measuring tube 0.2 mL

50 mL buret 0.05 mL

1 mL pipet 0.006 mL

10 mL pipet 0.02 mL

25 mL pipet 0.03 mL

100 mL volumetric flask 0.08 mL

250 mL volumetric flask 0.12 mL

Logic of Uncertainty in Calculations

Addition and Subtraction:

23.2 ± 0.3 mL + 5.6 ± 0.1 mL =

Smallest Possible Value: Largest Possible Value:

Range:

Report as:

28.4 mL 29.2 mL0.8 mL28.8 ± 0.4 mL

Uncertainty with Addition and Subtraction

• The previous example only works if the direction of both errors is the same. To correct for the sign of the uncertainty, it is more correct to use the following equation:

Example:

𝑈𝑛𝑐𝑒𝑟𝑡𝑎𝑖𝑛𝑡𝑦𝑖𝑛𝑟𝑒𝑠𝑢𝑙𝑡 = (𝑢𝑛𝑐𝑒𝑟𝑡𝑎𝑖𝑛𝑡𝑦> )?+(𝑢𝑛𝑐𝑒𝑟𝑡𝑎𝑖𝑛𝑡𝑦? )? + ⋯+ (𝑢𝑛𝑐𝑒𝑟𝑡𝑎𝑖𝑛𝑡𝑦A )?�

23.2 ± 0.3 mL + 5.6 ± 0.1 mL = 𝑈𝑛𝑐𝑒𝑟𝑡𝑎𝑖𝑛𝑡𝑦 = (0.3)?+(0.1)?�

= ± 0.3 mL

28.8 ± 0.3 mL

Uncertainty with Multiplication and Division

• The error propagation with these operations is not as straight forward

• Use the following equation:

𝑈𝑛𝑐𝑒𝑟𝑡𝑎𝑖𝑛𝑡𝑦𝑖𝑛𝑟𝑒𝑠𝑢𝑙𝑡 = 𝐶𝑎𝑙𝑐𝑢𝑙𝑎𝑡𝑒𝑑𝑟𝑒𝑠𝑢𝑙𝑡𝑥(𝑢𝑛𝑐𝑒𝑟𝑡𝑎𝑖𝑛𝑡𝑦> )?

(𝑚𝑒𝑎𝑠𝑢𝑟𝑒𝑚𝑒𝑛𝑡> )?+

(𝑢𝑛𝑐𝑒𝑟𝑡𝑎𝑖𝑛𝑡𝑦? )?

(𝑚𝑒𝑎𝑠𝑢𝑟𝑒𝑚𝑒𝑛𝑡? )?+ ⋯+

(𝑢𝑛𝑐𝑒𝑟𝑡𝑎𝑖𝑛𝑡𝑦A )?

(𝑚𝑒𝑎𝑠𝑢𝑟𝑒𝑚𝑒𝑛𝑡A )?�

Q-Test: Rejection of Data

• Sometime one value in a set of data appears different in a non-random way The q-test is an objective statistical test for deciding if a result should be rejected as an “outlier”

• Only the survivors of the q-test are used in the mean, standard deviation, and t-test

Q-Test

1. Arrange the data in order2. Calculate the difference (d) between the suspect value

and the nearest neighbor3. Calculate the range (r) between the suspect value and

the farthest neighbor4. Calculate the experimental value for Qexp

QIJK = dr

5. Compare Qexp to Qtable

6. Reject the data point if Qexp > Qtable

Confidence Level90% 95% 99%

n Q Q Q3 0.89 0.94 0.994 0.68 0.76 0.895 0.56 0.64 0.786 0.48 0.56 0.78

7 0.43 0.43 0.64For n>7, use value for 7

Q-Test Practice

• Should any of the data be thrown out at a 95% confidence level?

Trial Volume (mL)1 14.82 12.63 13.94 13.25 14.1

Highest Value: 14.8

Lowest Value: 12.6

QIJK = dr

Confidence Level90% 95% 99%

n Q Q Q3 0.89 0.94 0.994 0.68 0.76 0.89

5 0.56 0.64 0.786 0.48 0.56 0.78

7 0.43 0.43 0.64For n>7, use value for 7

QIJK = 14.8−14.114.8 − 12.6

= 0.32

0.32 < 0.64

QIJK = 13.2−12.614.8 − 12.6

= 0.27

0.27 < 0.64

Standard Deviation (s or 𝜎)

• A method for calculating the precision of a series of measurements

• Where x is the measurement and x is the mean (measurement-mean) and n is the number of samples

**The smaller the standard deviation, the more precise the data

Calculate the Standard Deviation

Mean:

Standard Deviation:

Trial Volume (mL)1 45.62 46.23 45.74 45.2

45.675

=(45.6 − 45.675)?+(46.2 − 45.675)?+(45.7 − 45.675)?+(45.2 − 45.675)?

4 − 1�

s = 0.411

One Sample T-Test

• Used to estimate the reliability of the derived result from a small set of data• Sets up a “confidence interval” for a result (the probability that the true value is

close to the mean (𝜇))

• x = mean• t = value from table• s = standard deviation• n = number of trials

𝑥 ± 𝑡(𝑠𝑛�)

**Report results using this formula

T-Table

• Confidence level or % probability that the actual value lies within the interval from the mean

T-Test Practice

• Use a T-Test at a 95% confidence level to find the mean and range of uncertainty.

Trial Volume (mL)1 45.62 46.23 45.74 45.2

𝑥 ± 𝑡(𝑠𝑛�)

45.7 ± 3.18(0.4114�)

𝟒𝟓. 𝟕 ± 𝟎. 𝟕𝒎𝑳

The Beauty of Modern Technology

• Standard Deviation and T-Tests can also be calculated in excel or with a TI graphing calculator

Excel

=STDEV.S(range)

TI- Press [2nd], [LIST], - Scroll to MATH and select 7:stdDev(- Input in brackets within parenthesesseparated by commas

The Beauty of Modern Technology

• Standard Deviation and T-Tests can also be calculated in excel or with a TI graphing calculator

Excel

=T.TEST(range)

TI

- Input Data to L1- Press [STAT] - Scroll to TESTS and select 2:T-Test- Scroll down to Calculate

Summary

• % Error – simple calculation to determine accuracy (known actual value)

• Uncertainty in Measurements – table of values for instruments, equations for add/sub and mult/div

• Q-Test – determines if data can be thrown out of statistical analysis

• Standard Deviation – complex determination of error and analysis of mean

• T-Test – used in small data sets to determine value of mean in relation to true value

Ch. 2: Atoms, Molecules and Ions

Classification of Matter

Matter

• We define matter as anything that has mass and takes up space

Atoms of an element Molecules of an element Molecules of a compound Mixture of elements and a compound

Matter

• Atoms are the building blocks of matter

• Each element is made of the same kind of atom

• A compound is made of two or more different kinds of elements

States of Matter

Properties and Changes of Matter

Types of Properties

• Physical Properties• Can be obtained without changing a substance into another substance

• Boiling point, density, mass, volume, etc.

• Chemical Properties• Can only be observed when a substance is changed into another substance

• Flammability, corrosiveness, reactivity with acid, etc.

Types of Properties

• Intensive Properties• Are independent of the amount of the substance that is present

• Density, boiling point, color, etc.

• Extensive Properties• Depend upon the amount of the substance present

• Mas Volume, energy, etc.

Types of Changes

• Physical Changes• These are changes in matter that do not change the composition of a substance

• Changes of state, temperature, volume, etc.

• Chemical Changes• Chemical changes result in new substances

• Combustion, oxidation, decomposition, etc.

Separation of Mixtures

Evaporation

• Evaporation uses difference in the boiling points of substances to separate a mixture into its components• The substance with the

lower boiling point is lost to the atmosphere

Distillation

• Distillation uses difference in the boiling points of substances to separate a homogeneous mixture into its components

• All components are retained

Filtration

• In filtration, solid substances are separated from liquids and solutions

Types of Filtration

• Decanting• Involves pouring off the top layer of liquid

• The simplest way to separate a liquid and solid and should only be used for qualitative work

• Gravity Filtration• Uses the natural downward flow of the liquid through a porous medium (usually paper) to

separate components

• Vacuum Filtration• Creates a partial vacuum in the collection flask to speed up separation

• Uses Buchner funnel and filtration flask

Chromatography

• Separates substances on the basis of difference in polarity

• Types of chromatograph differ by their mobile and stationary phases

Absorbent vs. Adsorbent

• Absorbent• The ability to carry a liquid in small chambers

• Ex: sponge

• Related to volume and diffusion

• Adsorbent• When particles get caught in small pores by weak Van der Waals forces (i.e. London

Dispersion) to form a layer• Ex: Chromatography

• Related to surface area and adhesion

Chromatography

Substances with similar structures (polar or nonpolar) will be attracted to each other

Types of Chromatography

• Thin Layer Chromatography (TLC)• Simple method using a gel or silica as the stationary phase and solvent as mobile phase

• Paper Chromatography• Similar to TLC but using filter paper as the stationary phase

Uses for Chromatography

• Identify unknowns based on comparison

• Determine the purity of a substance

• Monitor the progress of a reaction

Calculating Rf Values

Retention Factor (Rf)

• Value to help determine the relative distance of compounds in TLC

𝑅Y = 𝑑𝑖𝑠𝑡𝑎𝑛𝑐𝑒𝑡𝑟𝑎𝑣𝑒𝑙𝑒𝑑𝑏𝑦𝑐𝑜𝑚𝑝𝑜𝑢𝑛𝑑𝑑𝑖𝑠𝑡𝑎𝑛𝑐𝑒𝑡𝑟𝑎𝑣𝑒𝑙𝑒𝑑𝑏𝑦𝑠𝑜𝑙𝑣𝑒𝑛𝑡

• Best for compare compounds on the same stationary phase• Can compare if TLC has constant solvent system, stationary

phase, amount of material spotted, and temperature.

Calculating Rf Values

**Always mark in pencil the origin line (before starting) and the solvent front (after sample is run before solvent dries)

History of Atomic Theory

Democritus (~400 BC)

• First suggested the existence of atoms

• All matter is made up of small indestructible particles

John Dalton (1803-1807)

• First Atomic Theory based on proof.

1. All matter is made up of atoms

2. Atoms of the same element are the same, atoms of different elements are different

3. Atoms combine in small whole number ratios to form compounds

4. Atoms are rearranged in a chemical reaction

Thomson’s Experiment


• Passing an electric current makes a beam appear to move from the negative to the positive end


• By adding an electric field, he found that the moving pieces were negative

J.J. Thomson (1897)

• Used cathode ray tubes to discover the electron

• Measured the charge to mass ratio of the electron

• Since atoms are neutral species, he suggested that negative electrons are floating in a positively charged medium

Millikan’s Oil Drop Experiment

• From the mass of the drop and the charge of the plates, he calculated the mass of an electron

9.1 x 10-31 kg

Radioactivity

• Bequerel (~1890s)• Discovered radioactivity by accident

• Curie (~1900s)• Isolated radioactivity

• Rutherford (~1900)• Discovered three types:• Alpha – Helium nucleus (+2 charge, large mass)• Beta – High speed electron• Gamma – High energy light

Rutherford’s Gold Foil Experiment


What he expected


Because he thought the mass was evenly distributed in the atom


What he got

Earnest Rutherford (1908-1913)

1. Atom contains a dense nucleus

2. The atom is mostly empty space

Niels Bohr (~1910s)

• Electrons have certain energies which allow them to stay in certain orbits around the nucleus

• Still used to explain energy levels in atoms

Modern View of the Atom

• The atom is mostly empty space

• Two regions• Nucleus – protons and neutrons

• Electron Cloud – region where you might find an electron

Quantum Model (Electron Cloud Model)

The Periodic Table

Elements

Atomic Number

Atomic Mass

Atomic Symbol

Periodic Table

Metals

Nonmetals

Metalloids

Alkali Metals

Alkali Earth Metals

Transition Metals

Halogens

Noble Gases

Inner Transition Metals

Common Charges

Chemical Bondingand Nomenclature

Chemical Bonds

Covalent Bonding• Sharing electrons• 2 non-metals• Makes molecules

Ionic Bonding• Gaining/losing electrons• Metal and non-metal (or polyatomic)• Held together by opposite charges• Makes compounds

The forces that hold atoms together

Types of Formulas

• Chemical Formula• The number and type of atoms in a

molecule

• Structural Formula• Shows the connects but not

necessarily the shape

• Skeletal Structure• Abbreviated from structural

• C and H are not written (C at bend)

• Ball and Stick Model• Shows the three dimensional shape

with accurate angles

Butane

Empirical vs. Molecular Formula

Empirical Formula

• The reduced formula of a compound

Molecular Formula

• The unreduced or actual formula of a compound

**The empirical and molecular formulas can be the same!**Empirical formula can be found from the molecular formula, but data is

needed to go from empirical to molecular

Molecular àEmpirical

• Divide all subscripts in the MF by the greatest common factor

MF Si4O8 C25H45 CO2 C8H10N4O2 Ba3(PO4)2

EF SiO2 C5H9 CO2 C4H5N2O Ba3(PO4)2

Empirical à Molecular

Formula Molar Mass GCF

MF

EF

Determine the molecular formula of a compound with an empirical formula of NH2 and a formula mass of 32.06 amu.

**EF formula and MF molar mass given1. Calculate the molar mass of the EF2. Divide MF molar mass by EF molar

mass to find the greatest common factor (GCF)

3. Multiply all subscripts in the EF by the GCF

NH2

N2H4 32.06

16.032

1

23

Ions

Cations

• Positive ions

• Lose electrons

Anions

• Negative ions

• Gain electrons

Atoms of groups of atoms with a chargeIonic solids are called salts

Polyatomic Ions

• Groups of atoms covalently bonded that have an overall charge

• Bromine and iodine are just like chlorine• Bromate (BrO3

-1), perbromate (BrO4-1)

• Adding hydrogen to polyatomics• Sulfate (SO4

-2), hydrogen sulfate or bisulfate (HSO4-1)

Hypochlorite ClO-1 -2Chlorite ClO2

-1 -1Chlorate ClO3

-1 basePerchlorate ClO4

-1 +1

Polyatomic Ions to Memorize

Naming Ionic Compounds

From formula to name:

1. Name cationa. Transition metals must have roman numeral or -ic/-ous

2. Name aniona. Monatomic ion à end in –ide

b. Polyatomic ion à leave name as is

Transition Metals

• Have multiple possible ion charges• Exceptions: Ag+1 Zn+2 Cd+2

Element (ous) (ic) NameIron (Fe) +2 +3 Ferr-Lead (Pb) +2 +4 Plumb-Tin (Sn) +2 +4 Stann-Mercury +1 +2 Mercur-Copper +1 +2 Cupr-

Examples

Name the following:

• CaS

• AlPO4

• FeS

• CuI3

• KClO4

• YBrO2

• Cr(ClO)6

Calcium sulfide

Aluminum phosphate

Iron (II) sulfide OR Ferrous sulfideCopper (III) iodide

Potassium perchlorate

Yttrium (II) bromite

Chromium (VI) hypochlorite

Naming Ionic Compounds

From name to formula:

1. Determine charge of cation and anion

2. Criss-cross charges as positive subscripts1. Put () around polyatomics with charge on the outside

3. Reduce subscripts if needed

Examples

Write the formula for the following:• Tungsten (II) nitrite• Stannous chromate• Potassium permanganate• Zinc phosphate• Sodium oxide• Aluminum hydrogen sulfate

W(NO2)2

SnCrO4

KMnO4

Na2OZn3(PO4)2

Al(HSO4)3

Naming Covalent Compounds


1. Write first element with prefixa. Exception à never write mono with

the first element

2. Write second element with prefix and end in -ide

Covalent Prefixesmono- 1 hexa- 6

di- 2 hepta- 7tri- 3 octa- 8

tetra- 4 nona- 9penta- 5 deca- 10

Examples

Name the following:

• CO2

• CO

• CCl4

• N2O4

• XeF6

• N4O4

• P2O10

Carbon dioxide

Carbon monoxide

Carbon tetrachloride

Dinitrogen tetraoxide

Xenon hexafluoride

Tetranitrogen tetraoxide

Diphosphorus decaoxide

Naming Covalent Compounds


1. Prefixes indicate subscript on each element

2. If no prefix on first element, assume 1

3. DO NOT REDUCE!

Covalent Prefixesmono- 1 hexa- 6

di- 2 hepta- 7tri- 3 octa- 8

tetra- 4 nona- 9penta- 5 deca- 10

Examples

Write the formula for the following:

• Dinitrogen pentachloride

• Carbon disulfide

• Triboron heptafluoride

• Phosphorus trioxide

N2Cl5CS2

B3Fl5PO3

Acids

• Substances that produce H+ when dissolved in water

• All acids begin with H

• Two types of acids:• Oxyacids (contains oxygen, usually a polyatomic)

• Non-oxyacids (can be binary acids à H with a monatomic ion, or non-oxygen polyatomics)

Naming Oxyacids


1. Name polyatomic / anion1. Change -ate to –ic acid

2. Change –ite to –ous acid

2. Do NOT include hydrogen in the name


1. Look for acid at end of name

2. Determine polyatomic1. If –ic acid, then use –ate polyatomis

2. If –ous acid, then use –ite polyatomis

3. Use # of H to balance out charge of polyatomic

Naming Non-oxyacids


1. Write hydro-

2. Name anion1. Change –ide to –ic acid


1. Name contains hydro-

2. Determine –ide anion formula

3. Use # of H to balance out charge

Examples

Name the following:

• H2CrO4

• HCl

• H2S

• HMnO4

• HCN*

• HNO2

Chromic acid

Hydrochloric acid

Hydrosulfuric acid

Permanganic acid

Hydrocyanic acid ORHydrogen cyanide

Nitrous acid

Examples

Write the formula for the following:• Hydrofluoric acid• Dichromic acid• Carbonic acid• Hydrophosphoric acid• Perchloric acid• Phosphorous acid

HFH2Cr2O7

H2CO3

HClO4

H3P

H3PO3

Hydrates

• Salts containing coordinated water

• Name must include water

Naming Hydrates


1. Write name of salt

2. Write hydrate with covalent prefix to indicate # of water molecules


1. Determine formula for salt

2. Put a dot, then xH2O1. X is prefix from name for # of

water molecules

Examples

Write the formula for the following:

• Calcium chloride dihydrate

• Chromium (III) nitrate hexahydrate

Write the name for the following:

• CuSO4 • 5H2O

• Na2O • 3H2O

CaCl2 • 2H2O

Cr(NO3)3 • 6H2O

Copper (II) sulfate pentahydrate

Sodium oxide trihydrate

Organic Nomenclature

Organic Molecules

• Hydrocarbon• Organic molecules containing just

hydrogen and carbon

• Substituent• General term for a functional group or

R group

• Functional Group• Group containing oxygen, nitrogen,

and/or sulfur which gives the molecule different properties/reactivities

• R group• A hydrocarbon chain of undetermined

length attached to another hydrocarbon or functional group

• Aromatics• Molecules which contain alternating

double bonds, usually cyclical

• Ex: Benzene

Covalently bonded molecules containing hydrogen, carbon, oxygen, nitrogen, and/or sulfur

Simple Organic Naming


1. Use organic prefixes for # of carbons

2. End in suffix based on # hydrogens / types of bonds


1. Use prefix to determine # of carbons

2. Based on suffix, use formula to determine # of hydrogen

Organic Prefixes

Meth- 1 Hex- 6Eth- 2 Hept- 7Prop- 3 Oct- 8But- 4 Non- 9Pent- 5 Dec- 10 **Use cyclo before prefix if

hydrocarbon is in a ring

Organic Suffix

Ending C-C Bond Formula-ane Single bond CnH2n+2

-ene Double bond CnH2n

-yne Triple bond CnH2n-2

Functional GroupsFunctional Group Name Example

Aldehyde ______al Ethanal

Alcohol _____ol Ethanol

Carboxylic Acid ____oic acid Ethanoic

acid

Ketone _____one Propanone

Functional GroupsFunctional Group Name Example

Ether R-yl R’-yl ether Dimethyl ether

Amine _____ylamine Methylamine

Ester R’-yl _____oate Methyl propanoate

Amide _____amide Ethanamide

Examples

Write the names for the following:

• CH3CH2NH2

• -

• CH3CH2COOCH2CH2CH2CH3

• -

ethylamine

pentanone

butyl propanoate

ethyl methyl ether

Examples

Draw the structures for the following:

• Pentanol

• Butanoic acid-

• Ethyl propyl ether

• Cyclopentanamide

Challenge Problem

IUPAC Nomenclature

• Systematic method to indicate where something is occurring in an organic structure

• Using a numbering system of the “parent” chain

IUPAC Nomenclature


1. Find the longest hydrocarbon chain (This is the parent chain)

2. Number the chain so that substituents coming off the parent chain get the lowest numbered carbons

3. Write the #, substituent, -, then parent chain1. If substituent is a functional group, use functional group naming with the #

Examples

1-chlorobutane

2-methylhexane

2-butene

3-hexanol

1-bromocyclopentane

4-ethyloctane

1

1

1

1

1

1

2

2

2

2

2

2

3

3

3

3

3

3

4

4

4

4

4

4

5

5

5

5

6

6

678

Parent chain

Substituent

• Instead of putting the number for functional groups in front of the name, they can be placed immediately before the prefix ending

Alternate Naming

2-butene but-2-ene

• Instead of putting the number for functional groups in front of the name, they can be placed immediately before the prefix ending

Alternate Naming

3-hexanol hexan-3-ol

Multiple Substituents

• When the parent chain has multiple substituents, they are named ALPHABETICALLY rather than in number order

4-ethyl-2-methylhexane

1 2 3 4 5 6

ethylmethyl

Multiple Substituents

• This is where the alternate numbering comes in handy

4-ethyl-2-methylhex-3-ene

1 2 3 4 5 6

ethylmethyl

unit 0 basics of chemistry - ms. diorio's ap chemistry...

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