unit 1 atoms, molecules and ions chemistry ii starter monday august 24th draw a picture of the atom....

Unit 1Atoms, Molecules and Ions

Chemistry II

Starter Monday August 24th• Draw a picture of the atom. Label the

following: nucleus, electron cloud, proton, neutron, and electron

• Under your picture, construct a table showing the relationship in mass, charge, and location for the three subatomic particles. For table, make your columns as follows: type of particle, mass, charge, location. Make your rows: protons, neutrons and electrons

Atoms/Bonding

Forces:

Atoms, Molecules, and Ions

4 force of nature

1)Strong- between protons and neutrons

2)Weak- hold neutrons together

3)Electromagnet- between protons and electrons

4)Gravity- between matter

Extension

• Give an example of each of these forces in action. You can either give a description in paragraph form for each force or draw a visual aid/picture/model that shows how nature uses these forces.

Starter Aug 27th • In same table from Monday, add the

following particles to your rows: beta particles, alpha particles and positrons, then based on any prior knowledge or research fill in the columns done on Monday.

• Research the terms leptons, hadrons, fermions, bosons, muons and gluons…What is their significance within the atom.

What determines these forces?• Proton/neutron ratio- the bigger the ratio the

more likely an atom can overcome its strong force- radioactive decay

• Electronegativity- the pull of the nucleus on its surroundings determines its behavior

• Size of atom- greater atomic size, greater shielding

• Size of matter- as matter behaves on the macroscale, the larger the size the more pull it has on other matter

Proton/Neutron Ratio

• The ratio between number of protons and neutrons inside of the atom.

• The greater the difference between the two the less stable the atom

Band of Stability• Unstable nuclei

naturally break down.

• As this break down occurs radioactivity is given off.

Radioactive Decay

• Alpha decay

• Beta decay

• Gamma decay

• Positron emission

• Electron capture

Alpha Particles ()

Radium

R226

88 protons138 neutrons

Radon

Rn222

Note: This is themass number, whichis the number ofprotons plus neutrons


+ nnp

p

He)

2 protons2 neutrons

The alpha-particle is a Helium nucleus.

It’s the same as the element Helium, with no electrons!

Beta Particles ()

CarbonC14

6 protons8 neutrons

NitrogenN14

7 protons7 neutrons

+ e-

electron(beta-particle)

A neutron “converted” into a proton, and an electron was ejected.

This occurs when the atom’s neutron/proton ratio falls above the band ofstability.

Gamma particles ()In much the same way that electrons in atoms can be in an excited state, so can a nucleus.

NeonNe20


(in excited state)


(lowest energy state)

+

gamma

NeonNe20

A gamma is a high energy light particle.

It is NOT visible by your naked eye because it is not in the visible part of the EM spectrum.

A gamma is a high energy light particle.

It is NOT visible by your naked eye because it is not in the visible part of the EM spectrum.

Positron Emission

+ e+

A proton emits a positron from it, causing it to be converted into a neutron and therefore decreasing the atomic number by 1, while the mass number stays the same.

Happens when the neutron/proton ratio falls below the band of stability.

38K 38Ar

Electron Capture

e-

An electron is is pulled into a proton becoming a neutron. There atomic number decreases by 1, but the mass number stays the same.

+37 Ar37Cl

Nuclear binding energy

• Quantitative measure of nuclear stability.

• The energy required to break up a nucleus into its component protons and neutrons.

• Mass defect- difference between the mass of an atom and the sum of the protons, neutrons and electrons.

Nuclear Decay Series

Uranium has an atomic number greater than83. Therefore it is naturally radioactive.

Most abundant isotope

Alpha Particle

Predicting the Product of a Nuclear Reaction- AlphaWhat product is formed when radium-226 undergoes alpha decay?

What product is formed when polonium-212 undergoes alpha decay?

Predicting products of nuclear decay- beta

Write a reaction for the beta decay of Potassium-40.

Homework:

• Draw the complete nuclear series that decays Uranium-238 into Lead-206.

The series goes as follows: α, β, β, α, α, α, α, α, β, α, β, β, α, and β

Rate of Decay• How frequently atoms emit radiaton or how long does it take for atoms to decay?• Cannot predict when a particular entity will decay.

•But can predict when a number in a large sample will decay. Called a radioactive half-life

• Decay rate- •some atoms break down extremely quickly- Uranium-231•Some atoms break down slowly- Carbon-14

Half-Life• “Half-life” (h) = time it takes for half the atoms of a radioactive substance to decay.•Example:

•we had 20,000 atoms of a radioactive substance. If the half-life is 1 hour, how many atoms of that substance would be left after:

10,000 50% or 1/2

5,000 25% or 1/4

2,500 12.5% or 1/8

1 hour

2 hours

3 hours

Time #atoms

remaining% of atomsremaining

Half-life equation• Nt = No X (.5)n

Nt = amount of sample remaining

No = initial amount of sample

n= # of half-lifes or total time divided by time in one half life (i.e. n= t/h)

Or

• Mt = Mo X (.5)t/h

Mt = mass of sample remaining Mo = mass of initial sample

Logs

• Sometimes in order to calculate half-lifes it is important you remember how to do log functions.

• Log xy = y Log x

• We will revisit radioactive decay later in the year after we discuss kinetics and look at natural logs.

Problem:A sample of Iodine-131 had an original

mass of 16g. How much will remain in 24

days if the half life is 8 days?

Problem:• What is the ½ life of a sample if after 40

years 25 grams of an original 400 gram sample is left ?

27

Problem:• You have 400 mg of a radioisotope with a

half-life of 5 minutes. How much will be left after 30 minutes?

28

problem:• Cobalt-60 is a radioactive isotope used

in cancer treatment. Co-60 has a half-life of 5 years. If a hospital starts with a 1000 mg supply, how many mg will need to be purchased after 10 years to replenish the original supply?

Homework:

• Find an article on the use of radioactive decay in medicine. Research this and determine the risks and benefits of this material and it’s use. Bring in your thoughts to present to the class.

Electronegativity

The difference in electronegativity controls atomic behavior

• If the difference is small between two elements, they can share their electrons with each other.

• If great enough, one atom will pull harder on the electrons than the other, creating a postively charged cation and a negatively charged anion.

Bond Energy

• Bond energy: amount of work needed to pull two atoms completely apart.

• Bond dissociation energy: energy required to break a chemical bond.

Ionic Bonding

Element Protons Electrons Valence Electrons

Oxidation number(charge)

Potassium

Fluorine

Oxygen

Aluminum

Using your periodic table fill in the chart with the appropriate information

Naming these ions

• Metals– Named as is

• Example: – Hydrogen atom becomes a Hydrogen ion(Shown on previous slide)

• Nonmetals– Write first syllable of element with the ending –ide

• Example: – Oxygen becomes oxide– Fluorine becomes fluoride(shown on previous slide)

Metal Ions (cations)

• 3 types– Metal with specific charge

• Written as is • Example Li+1 is a lithium ion

– Metals with varying charges• Written with roman numeral • Example Cu+2 is Copper II Cu+3 is Copper III

– Polyatomic ion • Written as is • NH4

+ or ammonium

Nonmetals (anions)

• Monoatomic ions– 1st syllable of name with –ide ending– Oxygen is Oxide

• Polyatomic ions– Written as is

– SO4-2 is sulfate

– SO3-2 is sulfite

Naming Binary Compounds

• NaCl

• Al2O3

• Fr2O

• Mg3P2

Polyatomic Ions

A group of atoms with an overall charge. Result of a covalent bond.NH4

+ _______ OH- __________

NO3- ________ NO2

- ________

CO32- _____________

HCO3- ____________________

More Polyatomic Ions

Sulfur

SO42- sulf______ SO3

2- sulf_____

HSO4- hydrogen sulf____

HSO3- hydrogen sulf____

Phosphate

PO43- phosphate PO3

3- ____________

HPO42- _______________________________

H2PO4- _______________________________

Naming Ternary CompoundsContain at least 3 elementsName the nonmetals as a polyatomic ionExamples:

NaNO3 Sodium nitrate

K2SO4 Potassium sulfate

Al(HCO3)3 Aluminum bicarbonate

or

Aluminum hydrogen carbonate

Learning Check

Name the following compounds:

1)Na2ClO3

2)CaSO3

3)Mg3(PO4)2

4)Be(HCO3)2

5)CaCO3

6)(NH4)3PO4

7)LiNO3

Naming Binary Covalent Compounds

Two nonmetalsName each element End the last element in -ideAdd prefixes to show more than 1 atomDo not put the prefix mono on the first element

Prefixesmono 1 penta 5di 2 hexa 6tri 3 hepta 7tetra 4 octa 8

Learning Check

A. aluminum nitrate

1) AlNO3 2) Al(NO)3 3) Al(NO3)3

B. copper(II) nitrate

1) CuNO3 2) Cu(NO3)2 3) Cu2(NO3)

C. Iron (III) hydroxide

1) FeOH 2) Fe3OH 3) Fe(OH)3

D. Tin(IV) hydroxide

1) Sn(OH)4 2) Sn(OH)2 3) Sn4(OH)

Covalent Bonding

Learning Check

Indicate whether a bond between the following would be 1) Ionic 2) covalent

____ A. sodium and oxygen

____ B. nitrogen and oxygen

____ C. phosphorus and chlorine

____ D. calcium and sulfur

____ E. chlorine and bromine

What happens if they both have a high electronegativity?

• If the electronegativity difference is 1.7 or less, then the elements will share electrons. Not necessarily equally but share nonetheless…

• Two ways of sharing:– Polar- sharing unequally (.3-1.7)

• 5% to 50% ionic character

– Non-polar- sharing equally (.3 or less)• Only about 5% or less ionic character

Covalent Bonds

Occur between 2 nonmetals

Electrons are shared

single bond shares one pair electrons (2)

double bond shares two pairs electrons (4)

triple bond shares three pairs electrons (6)

Other single bonds:

Two nonmetal atoms form a covalent bond because they have less energy after they bonded

H + H H : H = HH = H2

• Diatomic MoleculesGases that exist as diatomic moleculesare H2, N2, O2, F2, Cl2, Br2, I2

Learning CheckFill in the blanks to complete the following names of covalent compounds.

CO carbon ______oxide

CO2 carbon _______________

PCl3 phosphorus _______chloride

CCl4 carbon ________chloride

N2O _____nitrogen _____oxide

Learning Check

A.P2O5 1) phosphorus oxide2) phosphorus pentoxide3) diphosphorus pentoxide

B. Cl2O7 1) dichlorine heptoxide2) dichlorine oxide3) chlorine heptoxide

C. Cl2 1) chlorine2) dichlorine3) dichloride

Naming Acids

• Acids- substance that yields H+ in water.

• Oxoacids: acids containing hydrogen, oxygen and another central element like C, N, S, Cl or P

Naming Acids

1.Addition of an O atom to the “ic” acid, produces the prefix “per” and the suffix “ic”

2. Removal of an O atom to an “ic” acid, produces an “ous” acid.

3. Removal of two O atoms from an “ic” acid produces a prefix “hypo” and suffix “ous” acid

Example

anion oxoacidperchlorate= ClO4

-1 perchloric acid= HClO4 chlorate = ClO3

-1 chloric acid= HClO3

chlorite = ClO2-1 chlorous acid= HClO2

hypochlorite= ClO-1 hypochlorous acid= HClO

Homework: Determine the anion and oxoacid for the set of ions produced by each of the following central atoms.1) nitrogen2) phosphorus3) manganese

Polar Bonds

• Nonpolar bonds: – .3 difference or less: – When the atoms in a bond are the same, the

electrons are shared equally.

• Polar bonds:– Between .3 and 1.7 difference– Unequal sharing of electrons

How to show a bond is polar• Isn’t a whole charge just a partial chargemeans a partially positivemeans a partially negative

• The Cl pulls harder on the electrons• The electrons spend more time near the Cl

H Cl

Polarity

• For the below compounds, determine whether each of the bonds inside of the molecule or formula unit is ionic,polar or non-polar

• NaCl

• CH4

• CCl4• H2S

Polar Molecules

Molecules with ends

Polar Molecules

• Molecules with a positive and a negative end• Requires two things to be true The molecule must contain polar bonds

This can be determined from differences in electronegativity.

Symmetry can not cancel out the effects of the polar bonds.

Must determine geometry first.

HHO

-

+

Water is asymmetrical

+

ClH

Space filling model“Electron-Cloud” model

-+

Is it polar?

• HF

• H2O

• NH3

• CCl4

• CO2

Intermolecular Forces

What holds molecules to each other

• Attractions between molecules– van der Waals forces

– Hydrogen “bonding”

– Dipersion Forces

– Dipole-Dipole

Intermolecular attractions

van der Waals

• The force that weakly attracts one molecule to another

• Non-polar molecules can exist in liquid and solid phases because van der Waals forces keep the

molecules attracted to each other

• Exist between CO2, CH4, CCl4, CF4, diatomics and monoatomics

van der Waals periodicity

• increase with molecular mass. – Greater van der Waals force?

• F2 Cl2 Br2 I2

• increase with closer distance between molecules– Decreases when particles are farther away

Hydrogen Bonding

H is shared between

2 atoms of

OXYGEN NITROGEN FLUORINE

of 2 different molecules

Hydrogen bonding

• Are the attractive force caused by hydrogen bonded to F, O, or N.

• F, O, and N are very electronegative so it is a very strong dipole.

• The hydrogen partially share with the lone pair in the molecule next to it.

• The strongest of the intermolecular forces.

Hydrogen Bonding

HH

O+ -

+

H HO+-

+

Hydrogen bonding

HH

O H HO

HH

O

H

H

OH

HO

H

HO HH

O

Dispersion forces

• The pull of electrons on the surroundings

• Depend on the number of electrons

• More electrons stronger forces

• Bigger molecules more electrons

• Fluorine is a gas

• Bromine is a liquid

• Iodine is a solid

Dipole interactions

• Occur when polar molecules are attracted to each other.

• Slightly stronger than dispersion forces.

• Opposites attract but not completely hooked like in ionic solids.

Dipole interactions

• Occur when polar molecules are attracted to each other. (The charge of one end on another.)

• Slightly stronger than dispersion forces.

• Opposites attract but not completely hooked like in ionic solids.

H F

H F

Dipole Interactions

MOLECULAR SHAPES

OFCOVALENT

COMPOUNDS

VSepR tHEORY

What Vsepr means

Electrons orient themselves as far apart as possible, from each other.

This leads to molecules having specific shapes.

Things to remember

•Atoms bond to form an Octet

•Bonded electrons take up less space then un-bonded/unshared pairs of electrons.

Linear

•Number of Bonds = 1-2

•Shared Pairs of Electrons =1-2

•Bond Angle = 180°

EXAMPLE:

BeF2

Trigonal Planar

•Number of Bonds = 3

•Shared Pairs of Electrons = 3

•Unshared Pairs of Electrons = 0


EXAMPLE:

GaF3

Bent #1




•Bond Angle = < 120°

EXAMPLE:

H2O

Bent #2


•Number of Shared Pairs of Electrons = 2

•Number of Unshared Pairs of Electrons = 1

•Bond Angle = >120°

EXAMPLE:

O3

Tetrahedral




•Bond Angle = 109.5°

EXAMPLE:

CH4

Trigonal Pyramidal




•Bond Angle = <109.5°

EXAMPLE:

NH3

Trigonal bIPyramidal




•Bond Angle = <120°

EXAMPLE:

NbF5

OCTAHEDRAL





EXAMPLE:

SF6

Percent Composition

Molar Mass Practice• Calculate the molar mass of the following:

Hydrogen

Fe

Iron II sulfate

Percent Composition• Percent Composition – the percentage by mass of

each element in a compound

Percent = _______PartWhole

x 100%

Percent compositionof a compound or =molecule

Mass of element in 1 mol____________________Mass of 1 mol

x 100%

Percent CompositionDetermine the percentage composition of sodium carbonate (Na2CO3)?

Molar Mass Percent Composition

% Na =46.0 g106 g

x 100% =43.4 %

% C =12.0 g106 g

x 100% =11.3 %

% O =48.0 g106 g

x 100% =45.3 %

Na = 2(23.00) = 46.0C = 1(12.01) = 12.0O = 3(16.00) = 48.0 MM= 106 g

Percent CompositionDetermine the percentage composition of ethanol (C2H5OH)?

% C = 52.13%, % H = 13.15%, % O = 34.72%

_______________________________________________

Determine the percentage composition of sodium oxalate(Na2C2O4)?

% Na = 34.31%, % C = 17.93%, % O = 47.76%

Percent CompositionCalculate the mass of bromine in 50.0 g of Potassium bromide.

1. Molar Mass of KBr

K = 1(39.10) = 39.10 Br =1(79.90) =79.90

MM = 119.0

79.90 g ___________119.0 g

= 0.6714

3. 0.6714 x 50.0g = 33.6 g Br

2.

Empirical and Molecular Formulas

Formulas

Empirical Formula – formula of a compound that expresses lowest whole number ratio of atoms.

Molecular Formula – actual formula of a compound showing the number of atoms present

Percent composition allow you to calculate the simplest ratio among the atoms found in compound.

Examples:

C4H10 - molecular

C2H5 - empirical

C6H12O6 - molecular

CH2O - empirical

Calculating Empirical FormulaWhen a 2.000 g sample of iron metal is heated in air, it reacts with oxygen to achieve a final mass of 2.573 g. Determine the empirical formula.

2.000 g Fe 1 mol Fe

55.85 g Fe= 0.03581 mol Fe

0.573 g O 1 mol O

16.00 g= 0.03581 mol Fe

Fe = 2.000 g O = 2.573 g – 2.000 g = 0.5730 g

1 : 1

FeO

Calculating Empirical FormulaThe most common form of nylon (Nylon-6) is 63.38% carbon, 12.38% nitrogen, 9.80% hydrogen and 14.14% oxygen. Calculate the empirical formula for Nylon-6.

Step 1:

In 100.00g of Nylon-6 the masses of elements present are 63.38 g C, 12.38 g n, 9.80 g H, and 14.14 g O.

Step 2:

63.38 g C 1 mol C

12.01 g C= 5.302 mol C

12.38 g N 1 mol N

14.01 g N= 0.8837 mol N

9.80 g H 1 mol H

1.01 g H= 9.72 mol H

14.14 g O 1 mol O

16.00 g O= 0.8838 mol O

Calculating Empirical FormulaThe most common form of nylon (Nylon-6) is 63.38% carbon, 12.38% nitrogen, 9.80% hydrogen and 14.14% oxygen. Calculate the empirical formula for Nylon-6.

Step 3:

5.302 mol C

0.8837= 6.000 mol C

0.8837 mol N

0.8837= 1.000 mol N

9.72 mol H

0.8837= 11.0 mol H

0.8838 mol O

0.8837= 1.000 mol O

6:1:11:1

C6NH11O

Calculating Molecular FormulaA white powder is analyzed and found to have an empirical formula of P2O5. The compound has a molar mass of 283.88g. What is the compound’s molecular formula?

Step 1: Molar Mass of EF

P = 2 x 30.97 g = 61.94gO = 5 x 16.00g = 80.00 g

141.94 g

Step 2: Divide MF Molar Mass by EF Molar Mass

238.88 g141.94g

= 2

Step 3: Multiply

(P2O5)2 =

P4O10

Calculating Molecular FormulaA compound has an experimental molar mass of 78 g/mol. Its empirical formula is CH. What is its molecular formula?

C = 12.01 gH = 1.01 g 13.01 g

78 g/mol

13.01 g/mol= 6

(CH)6 =

C6H6

Reactions and Stoichiometry

What is a chemical reaction composed of?

1) Contains reactants and products

2) Formulas must be written correctly with symbols and subscripts

3) Law of conservation of matter requires that coefficients be used to ensure that atoms

Describing chemical reaction

• The way atoms are joined is changed• Atoms aren’t created or destroyed.• May involve a catalyst• Can be described several ways

In a sentence • Solid Copper reacts with chlorine gas to

form aqueous copper (II) chloride.In a word equation or formula equation

• Copper(s) + chlorine(g) copper(II) chloride(aq)

• Cu(s) + Cl2(g) CuCl2(aq)

All chemical reactions are accompanied by a change in energy.

Exothermic - reactions that release energy to their surroundings (usually in the form of heat)

ΔH (enthalpy) is negative – energy leaving system

Endothermic - reactions that need to absorb heat from their surroundings to proceed.

ΔH (enthalpy) is positive – energy coming into the system

Reaction Energy

•Spontaneous Reactions - Reactions that proceed immediately when two substances are mixed together. Not all reactions proceed spontaneously.

Some require…

Reaction Energy

•Activation Energy – the amount of energy that is required to start a chemical reaction.

•Once activation energy is reached the reaction continues until you run out of material to react.

What is a catalyst?• Does not cause a reaction to occur,

but speeds up the rate which a reaction occurs

• Can be in the form of the following: – Energy- light, heat– Chemicals– Enzymes are biological or protein

catalysts.

Summary of Symbols

Reactions cont.

Synthesis Reactions• Also called combination reactions• 2 elements, or compounds combine

to make one compound.• A + B AB• Na (s) + Cl2 (g) NaCl (s)

• Ca (s) +O2 (g) CaO (s)

• SO3 (s) + H2O (l) H2SO4 (s) • We can predict the products if they

are two elements.• Mg (s) + N2 (g)

Mg3N2 (s)

Decomposition Reactions• decompose = fall apart• one compound (reactant) falls apart into

two or more elements or compounds.• Usually requires energy

• AB A + B

• NaCl Na + Cl2

• CaCO3 CaO + CO2

electricity

Replacement Reactions • Require a solution for reaction, because a

dissolvable ion is produced that allows the reaction to occur.

Remember: solvent- substance that dissolves substances into component ions

solute- substance that is dissolved into its component ions

Single Replacement• single displacement• One element replaces another• Reactants= an element and a compound.• Products= a different element and a

different compound.

• A + BC AC + B

• Based on the activity series (in other words the element replacing must be more active.

Activity Series

• Series that determines the reactivity of various metals and nonmetals in a solution.

• Controls whether a reaction will occur.

Activity Series

Foiled again –Foiled again –Aluminum loses to CopperAluminum loses to Copper

Element Reactivity

LiRbKBaCaNaMgAlMnZnCrFeNiSnPbH2

CuHgAgPtAu

Halogen Reactivity

F2

Cl2Br2

I2

Potassium reacts with Water

P O W !

Double Replacement (metathesis reaction)

Two ionic compounds or acids react

Requires aqueous solution

AB + CD AD + CB

AgNO3 + NaCl AgCl + NaNO3

ZnS + 2HCl ZnCl + H2S

Formation of a solid AgCl

AgNO3(aq) + KCl(aq) KNO3 (aq) + AgCl(s)

Single and Double Replacement Reactions

Double-replacement reaction

CaCO3 + 2 HCl CaCl2 + H2CO3

General form: AB + CD AD + CB

Single-replacement reaction

Mg + CuSO4 MgSO4 + Cu

General form: A + BC AC + B

Net Ionic Equations

Ionic equations

• Ionic equation: a complete equation showing all dissolved ions in a solution and any insoluble substance

Net ionic equation

• Spectator ions- dispersed ions on both sides… They do not contribute to the overall reaction and are therefore not truly considered a part of the reaction.

• Net ionic equation- shows reaction without the spectator ions

A Net Ionic Equation

• 2KI + Pb(NO3)2 → 2KNO3 + PbI2

Write a net ionic equation for the below reaction:

• 2AgNO3 + CaCl2 → 2AgCl + Ca(NO3 )2

Combustion

• A reaction in which a compound (often carbon) reacts with oxygen

• CH4 + O2 CO2 + H2O

• C3H8 + O2 CO2 + H2O

• C6H12O6 + O2 CO2 + H2O

• The charcoal used in a grill is basically carbon. The carbon reacts with oxygen to yield carbon dioxide. The chemical equation for this reaction is C + O2 CO2

Combustion of a Hydrocarbon

GENERAL FORMULA: CH + O2 CO2 + H2O

Many homes get heat from propane (C3H8) heaters. Write a balanced chemical equation for the completecombustion of propane gas.

C3H8(g) + O2(g) CO2(g) + H2O(g)

C3H8(g) + O2(g) CO2(g) + H2O(g)5 3 4 + energy

Incomplete Combustion of a Hydrocarbon

C3H8(g) + O2(g) CO2(g) + H2O(g)5 3 4 + energy

C3H8(g) + O2(g) 3 CO (g) + H2O(g)3 + energy

C3H8(g) + O2(g) C (g) + H2O(g)2 3 4 + energy

Ideal Stoichiometry

Too ‘rich’ (not enough oxygen – too much fuel)

SOOT

4

Combustion of Iron

• Formation of Rust

• Thermite Reaction• underwater welding• Temp. = ~3500oC

Fe2O3 + 2 Al 2 Fe + Al2O3 + 199 kcal

4 Fe + O2 2 Fe2O3

Combustion of Copper

• Copper burns with a green color

• Copper forms a patina (oxide) – green in color

• CuO2

– black in color• CuOStatue of Liberty is covered with

copper that has oxidized to formcopper (II) oxide, CuO2.

Acid/Base Reaction

• An acid and a base react to form a salt and water.

• Always in aqueous solution

• Acid (H+) + Base (OH-) → Salt + H2O

NaOH + HCl → NaCl + H2O

NH4OH + H2SO4 → (NH4)2SO4 + H2O

Acids• from the Latin word acere “sour”• taste sour (but you wouldn’t taste an acid to see)• change litmus paper red• corrosive to some metals (reacts to create hydrogen

gas – H2)• Donates a hydrogen ion (H+) to another substance• Create a hydrogen ion (H+) or hydronium ion (H3O+)

when dissolved in water

HCl H+ + Cl-

HydrochloricAcid

Hydrogenion

Chlorideion

Examples: hydrochloric acid, vinegar, lemon juice, rainwater

H2O Notice howthe hydrogenion is releasedwhen the acid

is in water

Bases (Alkalis)• taste bitter • feel slippery or soapy• change litmus paper blue• react with oils and grease- soaps• Accept a hydrogen ion (H+) • create a hydroxide ion (OH-) when dissolved in

water

Examples: sodium hydroxide, Drano, Tums, baking soda

NaOH Na+ + OH-

SodiumHydroxide

Sodiumion

Hydroxideion

H2O

Notice howthe hydroxideion is releasedwhen the baseis in water; this

ion can accept a hydrogen ion

(H+)

Neutralization Reaction

• occurs when acids and bases react with each other to produce water and salt– acids release a hydrogen ion (H+) and bases

release a hydroxide ion (OH-) water (H2O)– the negative ion from the acid joins with the

positive ion of a base salt

HCl + NaOH H2O + NaClHydrochloric

Acid(acid)

Sodium Hydroxide

(base)

WaterSodium Chloride

(salt)

Both the salt and water are neutral substances; therefore, that is why this is referred to as a neutralization reaction.

Acid, Base, or Neutralization?

Zn + 2H+ Zn2+ + H2

NH3 + H2O NH4+ + OH-

HClO + LiOH LiClO + H2O

HCl + H2O H3O+ + Cl-

Acid – because H2 gas was given off

Acid – because H3O+ is present in the products

Base – because OH- is present in the products

Neutralization – because of the salt and water in the products

Some Definitions

Arrhenius acids and bases– Acid: Substance that, when dissolved

in water, increases the concentration of hydrogen ions (protons, H+).

– Base: Substance that, when dissolved in water, increases the concentration of hydroxide ions.

The Brønsted-Lowry acid donates a proton (H+ ion),

while the Brønsted-Lowry base (H+ ion) accepts it.

Brønsted-Lowry Acids and Bases:

Which is the acid and which is the base in each of these rxns?

A Brønsted–Lowry acid…

…must have a removable (acidic) proton.

HCl, H2O, H2SO4

A Brønsted–Lowry base…

…must have a pair of nonbonding electrons.

NH3, H2O

Types of Proton acceptors and donators:

• Monoprotic acid- Donates 1 proton (H+)– HCl, HF, HI, HClO3

• Diprotic acid- Donates 2 protons (2H+)– H2S, H2SO4

• Triprotic acid- Donates 3 protons (3 H+)– H3PO4

If it can be either…

...it is amphiprotic.

HCO3–

HSO4 –

H2O

How to recognize which type

• Look at the reactants• Element(E), Compound(C)

• E + E • C• E + C• C + C• Acid + Base

• Look at the Products• CO2 + H2O

RedoxSynthesisDecompositionSingle replacement

Double replacementAcid/Base reaction

Combustion

Redox Reactions

• A reaction that shows the flow (transfer) of electrons between atoms in the substances.

OILRIG • Oxidation- the loss of electrons from an

atom• Reduction- the gain of electrons from an

atom• Shown as a change in oxidation states.

Oxidation States

• The charge of an element with a chemical reaction.

• Determine based on what the element is attached to and the charge of other elements that do not change.

- See oxidation states review from earlier in this unit

Oxidation State rules1) Fluorine is always -1. 2) Free elements have an oxidation state of 0. 3) Ions oxidation state = formal charge4) Oxygen has an oxidation state of -2 except in peroxide (O2

-2 = state is -1). 5) Hydrogen’s oxidation state is +1 unless attached to something less electronegative (metals in binary compounds).6. neutral molecule= sum of oxidation state= 07. polyatomic ion= sum of oxidation state= charge

Homework

• Complete oxidation state homework, bring in ready to discuss tomorrow.

Examples

Synthesis

Decomposition

DecompositionSingle replacement

Single replacement

Double replacement

Double replacement

H2 + O2

H2O AgNO3 + NaCl

Zn + H2SO4 HgO

KBr +Cl2

Mg(OH)2 + H2SO3

Examples

Acid/Base

Decomposition

Single replacementSynthesis

Acid/Base

Single replacement

Double replacement

HNO3 + KOH

CaPO4 AgBr + Cl2

Zn + O2 HgO + Pb

HBr + NH4OH

Cu(OH)2 + KClO3

SummaryAn equation:• Describes a reaction• Must be balanced because to follow Law

of Conservation of Energy• Can only be balanced by changing the

coefficients.• Has special symbols to indicate state, and

if catalyst or energy is required. • Can describe 5 different types of

reactions.

Stoichiometry

Stoichiometry is…• Greek for “measuring elements”Pronounced “stoy kee ahm uh tree”

• Defined as: calculations of the quantities in chemical reactions, based on a balanced equation.

• There are 4 ways to interpret a balanced chemical equation

#1. In terms of Particles

• An Element is made of atoms• A Molecular compound (made

of only nonmetals) is made up of molecules (Don’t forget the diatomic

elements)

• Ionic Compounds (made of a metal and nonmetal parts) are made of formula units

Example: 2H2 + O2 → 2H2O

• Two molecules of hydrogen and one molecule of oxygen form two molecules of water.

• Another example: 2Al2O3 Al + 3O2

2 formula units Al2O3 form 4 atoms Al

and 3moleculesO2

Now read this: 2Na + 2H2O 2NaOH + H2

#2. In terms of Moles

• The coefficients tell us how many moles of each substance

2Al2O3 Al + 3O2

2Na + 2H2O 2NaOH + H2

• A balanced equation is a Molar Ratio- We will look at this next.

#3. In terms of Mass

• The Law of Conservation of Mass applies

• We can check mass by using moles.

2H2 + O2 2H2O

2 moles H2

2.02 g H2

1 mole H2

= 4.04 g H2

1 mole O2

32.00 g O2

1 mole O2

= 32.00 g O2

36.04 g H2 + O236.04 g H2 + O2

+

reactants

In terms of Mass (for products)

2H2 + O2 2H2O

2 moles H2O18.02 g H2O1 mole H2O

=36.04 g H2O

36.04 g H2 + O2= 36.04 g H2O

The mass of the reactants must equal the mass of the products.

36.04 grams reactant = 36.04 grams product

#4. In terms of Volume

• At STP, 1 mol of any gas = 22.4 L

2H2 + O2 2H2O (2 x 22.4 L H2) + (1 x 22.4 L O2) (2 x 22.4 L H2O)

NOTE: mass and atoms are ALWAYS conserved - however, molecules, formula units, moles, and volumes will not necessarily be conserved!

67.2 Liters of reactant ≠ 44.8 Liters of product!

159

Mole Ratios

Ratio between two of the substances in a

balanced equation

Derived from coefficients of any two

substances in an equation.

160

Writing Mole Factors

4 Fe + 3 O2 2 Fe2O3

Fe and O2

4 mol Fe and 3 mol O2

3 mol O2 4 mol Fe

Fe and Fe2O3

4 mol Fe and 2 mol Fe2O3

2 mol Fe2O3 4 mol Fe

161

O2 and Fe2O3

3 mol O2 and 2 mol Fe2O3

2 mol Fe2O3 3 mol O2

162

Learning Check

3 H2(g) + N2(g) 2 NH3(g)

A. A mol factor for H2 and N2 is

1) 3 mol N2 2) 1 mol N2 3) 1 mol N2

1 mol H2 3 mol H2 2 mol H2

B. A mol factor for NH3 and H2 is

1) 1 mol H2 2) 2 mol NH3 3) 3 mol N2

2 mol NH3 3 mol H2 2 mol NH3

163

3 H2(g) + N2(g) 2 NH3(g)

A. A mol factor for H2 and N2 is

2) 1 mol N2

3 mol H2

B. A mol factor for NH3 and H2 is

2) 2 mol NH3

3 mol H2

Answers:

Converting Moles

2Al2O3 Al + 3O2

– each time we use 2 moles of Al2O3 we will also make 3 moles of O2

2 moles Al2O3

3 mole O2

or2 moles Al2O3

3 mole O2

Molar ratios can also be known as conversion factors. We could use them to solve calculations.

Example:• How many moles of O2 are

produced when 3.34 moles of Al2O3 decompose?

2Al2O3 Al + 3O2

3.34 mol Al2O3

2 mol Al2O3

3 mol O2 = 5.01 mol O2

If you know the amount of ANY chemical in the reaction, you can find the amount of ALL the other chemicals!

Conversion factor from balanced equation

166

4 Fe + 3 O2 2 Fe2O3

How many moles of Fe2O3 are produced when

6.0 moles O2 react?

6.0 mol O2 x mol Fe2O3 = 4.0 mol Fe2O3

mol O2

Practice:

167

4 Fe + 3 O2 2 Fe2O3

How many moles of Fe are needed to react with 12.0 mol of O2?

1) 3.00 mol Fe

2) 9.00 mol Fe

3) 16.0 mol Fe

More Practice:

168

4 Fe + 3 O2 2 Fe2O3

12.0 mol O2 x mol Fe = 16.0 mol Fe

mol O2

4

3

Answer:

More Practice

4 Fe + 3 O2 2 Fe2O3

How many grams of O2 are needed to produce

0.400 mol of Fe2O3?

1) 38.4 g O2

2) 19.2 g O2

3) 1.90 g O2

169

170

0.400 mol Fe2O3 x 3 mol O2 x 32.0 g O2

2 mol Fe2O3 1 mol O2

= 19.2 g O2

Answer:

171

Balance equation

Convert starting amount to moles

Use coefficients to write a mol-mol ratio

Convert moles of desired to grams

Converting Mass

172

The reaction between H2 and O2 produces

13.1 g of water. How many grams of O2

reacted?

Write the equation

H2 (g) + O2 (g) H2O (g)

Balance the equation

2 H2 (g) + O2 (g) 2 H2O (g)

Example:

173

Organize data

2 H2 (g) + O2 (g) 2 H2O (g)

? g 13.1 g

Plan g H2O mol H2O mol O2 O2

Setup

13.1 g H2O x 1 mol H2O x 1 mol O2 x 32.0 g O2

8.0 g H2O 2 mol H2O 1 mol O2

= 11.6 g O2

Mass-Mass Problem:

6.50 grams of aluminum reacts with an excess of oxygen. How many grams of aluminum oxide are formed?

4Al + 3O2 2Al2O3

=6.50 g Al

? g Al2O3

1 mol Al

26.98 g Al 4 mol Al

2 mol Al2O3

1 mol Al2O3

101.96 g Al2O3

(6.50 x 1 x 2 x 101.96) ÷ (26.98 x 4 x 1) = 12.3 g Al2O3

are formed

Another example:

• If 10.1 g of Fe are added to a solution of Copper (II) Sulfate, how many grams of solid copper would form?

2Fe + 3CuSO4 Fe2(SO4)3 + 3Cu

Answer = 17.2 g Cu

176

More Practice:

Acetylene gas C2H2 burns in the oxyactylene

torch for welding. How many grams of C2H2 are

burned if the reaction produces 75.0 g of CO2?

2 C2H2 + 5 O2 4 CO2 + 2 H2O

75.0 g CO2 x _______ x _______ x _______

177

2 C2H2 + 5 O2 4 CO2 + 2 H2O

75.0 g CO2 x 1 mol CO2 x 2 mol C2H2 x 26.0 g C2H2

44.0 g CO2 4 mol CO2 1 mol C2H2

= 22.2 g C2H2

Answer:

Volume-Volume Calculations:

• How many liters of CH4 at STP are required

to completely react with 17.5 L of O2 ?

CH4 + 2O2 CO2 + 2H2O

17.5 L O2 22.4 L O2 1 mol O2

2 mol O2

1 mol CH4

1 mol CH4 22.4 L CH4

= 8.75 L CH4

22.4 L O2 1 mol O2

1 mol CH4 22.4 L CH4

Notice anything relating these two steps?

Avogadro told us:

• Equal volumes of gas, at the same temperature and pressure contain the same number of particles.

• Moles are numbers of particles• You can treat reactions as if they

happen liters at a time, as long as you keep the temperature and pressure the same. 1 mole = 22.4 L @ STP

Shortcut for Volume-Volume?

• How many liters of CH4 at STP are required to completely react with 17.5 L of O2?

CH4 + 2O2 CO2 + 2H2O

17.5 L O2 2 L O2

1 L CH4 = 8.75 L CH4

Note: This only works for Volume-Volume problems.

181

If the amounts of two reactants are given, the reactant used up first determines the amount of product formed.

Limiting reagents- the reactant that is used up in a chemical reaction.

Excess reagents- the reactant that is left over after chemical reaction.

Limiting Reagants

182

Hints for LR Problems

1. For each reactant amount given, calculate the

moles (or grams) of a product it could produce.

2.The reactant that produces the smaller amount of product is the limiting reactant.

3. The number of moles of product produced by the limiting reactant is ALL the product possible. There is no more limiting reactant left.

Limiting Reagents - Combustion

How do you find out which is limited?

• The chemical that makes the least amount of product is the “limiting reagent”.

• You can recognize limiting reagent problems because they will give you 2 amounts of chemical

• Do two stoichiometry problems, one for each reagent you are given.

• If 10.6 g of copper reacts with 3.83 g sulfur, how many grams of the product (copper (I) sulfide) will be formed?

2Cu + S Cu2S

10.6 g Cu 63.55g Cu 1 mol Cu

2 mol Cu 1 mol Cu2S

1 mol Cu2S

159.16 g Cu2S

= 13.3 g Cu2S

3.83 g S 32.06g S 1 mol S

1 mol S 1 mol Cu2S

1 mol Cu2S

159.16 g Cu2S

= 19.0 g Cu2S

= 13.3 g Cu2S

Cu is the Limiting

Reagent, since it

produced less product.

Another example:• If 10.3 g of aluminum are

reacted with 51.7 g of CuSO4 how much copper (grams) will be produced?

2Al + 3CuSO4 → 3Cu + Al2(SO4)3

the CuSO4 is limited, so Cu = 20.6 g

• How much excess reagent will remain? Excess = 4.47 grams

The Stoichiometric Concept of :

188

Actual yield is the amount of product actually recovered from an experiment

Theoretical (possible) yield is the maximum amount of product that could be produced from

the reactant.

Percent Yield is the actual yield compared to the maximum (theoretical yield) possible.

Types of Yield:

Details on Yield• Percent yield tells us how “efficient” a

reaction is.

• Percent yield can not be bigger than 100 %.

• Theoretical yield will always be larger than actual yield!– Why? Due to impure reactants; competing

side reactions; loss of product in filtering or transferring between containers; measuring

190

Percent Yield Calculation

What is the percent yield of cookies?

Percent Yield = Actual Yield (g) recovered X 100 Possible Yield (g)

% cookie yield = 48 cookies x 100 = 80% yield

60 cookies

Example:

• 6.78 g of copper is produced when 3.92 g of Al are reacted with excess copper (II) sulfate.

2Al + 3 CuSO4 Al2(SO4)3 + 3Cu

• What is the actual yield?

• What is the theoretical yield?

• What is the percent yield?

= 6.78 g Cu

= 13.8 g Cu

= 49.1 %

Solution Stoichiometry

• Concentration: amount of solute present in a given quantity of solvent or solution.

Molarity= quantitative description of concentration measured as moles of solute per liters of solution

mol of soluteL of solution

M =

Molarity (M)

• Concentration of solution; or its strength

• Can change with temperature.

Preparing a molar solution1) Accurately weigh the correct amount of solute and transfer to a volumetric flask through a funnel.

2) Add a small amount of water through the funnel and swirl the flask. Continue to do this until all of the solid is dissolved.

3) Then, add the remainder of water to the flask until it reaches exactly the volume mark used.

Dilution

• Preparing a less concentrated solution from a more concentrated one.

M1V1 = M2V2

Example problem

• Describe how you would prepare 5.00 X102 mL of a 1.75 M H2SO4 solution, starting with a 8.61 M stock solution of H2SO4.

Gravimetric Analysis

• Analysis technique based on the measurement of mass involving formation, isolation and mass determination of a precipitate.

Highly accurate technique that allows you produce a specific material using a known precipitation reaction and its predicted stoichiometry.

Example:

• A .5562 g sample of an ionic compound containing chloride ions and an unknown metal is dissolved in water and treated with excess AgNO3. If 1.0882 g of AgCl precipitate forms, what is the percent mass of the Cl in the original compound?

unit 1 atoms, molecules and ions chemistry ii starter monday august 24th draw a picture of the atom....

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number of protons

mass number

atomic number

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following particles

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