AP Chemistry

Atomic Structure:

Atomic Number = # of protons = # of electrons in a neutral atom

Mass Number = protons and neutrons (Isotopes)

Mass #-238

U Protons=92 Electrons 92

Atomic # 92

Neutrons=238-92=146

Electron Configurations:

1. Order of filling: 1s22s

22p

63s

23p

64s

23d

104p

65s

24d

105p

66s

24f

145d

106p

67s

2

proper way of writing: 1s22s

2p

63s

2p

6d

104s

2p

6d

10f14

5s2p

6d

10f14

…

2. Atoms gain or lose electrons to obtain a filled octet

a) when transition metals lose electrons, they lose from the “s” sublevel first

ex. Fe: 1s22s

22p

63s

23p

64s

23d

6 Fe

3+ : 1s

22s

22p

63s

23p

63d

5

b) metals get oxidized (lose electrons) to form cations (pos.) while nonmetals get

reduced (gain electrons) to form anions (neg.).

3. Quantum #‟s n = energy level = (1,2,3…)

L = sublevel = (0 → n-1)

mL = orbital = - L → + L

ms = spin # = ± ½ (clockwise or counterclockwise)

4. Oxidation states vs. group # (Group # = highest possible oxidation state)

Group: I = +1 IV = +4 (Except Carbide =C-4

) VII = -1

II = +2 V = -3 VIII = 0

III = +3 VI = -2

Transition metals have multiple oxidation states but (+2) is the most common b/c of “s”

sublevel being lost first.

*Hund‟s rule: diamagnetic: No unpaired electrons

Pauli Exclusion Principle paramagnetic: unpaired electrons

Shielding Effects (penetration) ferromagnetic: Fe, Co, Ni

Heisenberg Uncertainty Principle

Periodic Trends:

I. Atomic Size(radius) : On Periodic Table Increases ←

↓ a) decreases left to right b/c electrons are in same energy level therefore they do not

add size but nuclear charge increases pulling electron cloud in more tightly

b) increases going down in a group b/c adding more energy levels

Greater difference in size between energy levels 1,2, & 3 than between 4,5,6, & 7

b/c energy levels are not evenly spaced.

(Transition metals are all nearly the same size)

II. Ionization Energy = inversely related to size (the bigger the atom, the lower the

ionization energy)

a) exceptions to trend occur when electrons are in filled or ½ filled sublevel

Between groups IIA and IIIA - Increased shielding by "s" electrons

b) Between groups VA and VIA - Increased electron ↔ electron

repulsions because of electrons beginning to pair up in "p" orbitals

c) ions become very stable (high I.E.) once they obtain a noble gas config.

III. Electron Affinity: Opposite of Ionization Energy

Groups IIA, IIB, & VIII have least attraction (most negative value) for extra electrons b/c

they are already stable

Intermolecular Forces 4 types of substances

1. Ionic: metal w/ nonmetal

a) High melting & boiling points b/c of high lattice energy binding ions together

1. Coulumb‟s Law: ΔH = Kq1q2 (q1 & q2 = charges on ions)

(lattice energy) r

2. As the product of the charges increases, the electrostatic attraction

increases, therefore higher melting and boiling points

3. The bigger the ions, the lower the electrostatic attractions, therefore the

lower melting and boiling points.

b) Do not conduct electricity in the solid state b/c the ions are held rigidly in

place. They do conduct in the liquid (molten) and aqueous states b/c the ions

dissociate and are free to move.

c) Brittle and therefore neither malleable nor ductile

d) Form large crystals-no molecules

2. Metallic: Like metal atoms bonded together

a) conduct electricity in the solid and liquid states b/c of the “sea of free-floating

valence electrons.”

b) Malleable and ductile b/c (see part a)

c) Vary wide variety of melting points and boiling points.

From Hg = liquid to W = solid w/ highest melting point (Generally the melting point

is greater than 300o C and therefore solids at room temperature)

d) Insoluble in H2O

e) Mixture of metals is called an “Alloy”

3. Covalent Network (Large crystals of nonmetals covalently bonded to one another).

a) only Cdiamonds, Cgraphite, SiO2, SiC

\Allotropes/ (Quartz)

b) Do not conduct electricity (except Graphite)

c) Diamonds = sp3 3d crystal-very hard

Graphite = sp2 2d crystal-forms loose layers

d) insoluble in H2O

e) High melting and boiling points and therefore solids at room temperature

4. Covalent Molecular: nonmetal w/ nonmetal

a) nonmetals bond with each other to form molecules

b) molecules are held together by weak intermolecular forces called van der

Waals forces

1. London Forces: Increases with the number of electrons in the molecule. (size of

the molecule)

a) The weakest of intermolecular forces

b) The only force of attraction between nonpolar molecules

c) Also called: Dispersion, Induced Dipole, Instantaneous Dipole

2. Dipole-Dipole: The attraction between oppositely charged portions of

polar molecules.

a) The greater the dipole moment (a measurement of the molecule‟s

polarity) the stronger the D.P. –D.P. attraction.

b) Stronger than London forces

3. Hydrogen Bonds: H atom must be attached directly to a N, O, or F.

a) Strongest of the 3 intermolecular forces but covalent network and

ionic bonds are still much stronger

Bonding

*Lewis Dot Structure: #valence electrons = group number

*Atoms want to achieve a stable octet noble gas configuration

*exceptions to rule: Be, Al, B - form stable molecules w/ 6 e- around central atom

*Resonance- occurs when molecule can have more than 1 possible structure (always has

double bonds/triple bonds)

Ex. O=N-O ↔ O-N=O (bond order = 1½ )

*all the bonds are actually the same length (average of the 2 bonds)

*Double bonds and Triple Bonds

bond length: single(longest) → double → triple(shortest)

Diatomic elements: H2, O2, N2, Cl2, F2, I2, Br2

*all form single bonds except: O2 (double bond) O=O

N2 (triple bond) N≡N

*Sigma and Pi bonds-

first bond between any two atoms=sigma bond (strongest)

second and third bonds = pi bonds

bond order = strength of bond (single, double, triple)

*Polar vs. Nonpolar

polar: bond that acquires positive and negative ends(has lone pair, asymmetrical)

nonpolar: shared pair of electrons is equally shared, no net displacement

(symmetrical- no lone pair on central atom) *No dipole moment

*Likes Dissolve Likes

Polar dissolves only into polar (*Water is Polar)

Ionic dissolves into polar

Nonpolar dissolves only into nonpolar

*Electronegativity-measure of ability of an atom to attract electrons to itself

most electronegative element: F, least: Fr

*structural isomers- appear to have same structure but are different

ex. C2H5OH and CH3OCH3

^

^

Hydrogen Bond No H bond

*optical isomer- molecule where one of its atoms has 4 different atoms or groups attached to it:

*VSEPR theory (valence shell electron pair repulsion)

-bonding and nonbonding pairs repel to obtain the farthest distances between each other

molecular shapes:

O H

: N≡C-H S H-C-H

O O H

linear trigonal planar tetrahedral trigonal-bipyramidal octahedral

*Hybridization: s, p, and d orbitals can be mixed to form new sets of orbitals

Hybrid Orbitals

Required (Could Have)

2 sp linear AX2 triple, double, or single bonds

3 sp2 trigonal planar AX3 double bonds, or single bonds

4 sp3 tetrahedral AX4 single bonds

5 sp3d trigonal bipyramidal AX5 single bonds

6 sp3d

2 octahedral AX6 single bonds

*Add up exponents to see how many bonds

*Some atoms cannot have sp3d hybridization because they don‟t have access to a “d” orbital ex:

NF5 doesn‟t exist, but PF5 does

Phases of Matter: Phase Diagrams:

1. Triple point-vapor pressure of solid & liquid are equal; all phases exist at equilibrium

2. Critical Point (temp.)-temp. at which liquid will form gas at any pressure minimum

amt. of pressure needed = critical pressure)

3. Normal melting/freezing pt.(1 atm)

4. Normal boiling/condensing pt.(1 atm)

Water

*Slope of solid to liquid line is (-) b/c liquid is more dense than solid For everything else, all slopes are (+)

solidliquid = melting

liquidsolid = freezing

liquidgas = boiling

gasliquid = condensation

solidgas = sublimation

gas solid = crystallization

Differences in Boiling Points

-for covalent compounds, always look for Hydrogen bonding

-also look for London Forces (vary w/ # of electrons ex. bigger molecules have higher B.P.)

-For ionic compounds: Coulombs Law ΔH= KQ1Q2

r

lattice energy-energy holding oppositely charged ions together

-Heat of Vaporization-energy needed to vaporize something at its boiling point (latent heat)

-varies with strength of attractive force

- Heat of Fusion - energy needed to melt a substance at its melting point. (Latent Heat)

Gases Gas Laws: 1 atm = 760mm Hg (torr)

1. Boyle‟s Law: P1V1 = P2V2

Pressure & volume are inversely proportional

2. Charles‟ Law: V1/T1 = V2/T2

Volume is directly proportional to temperature

3. Gay Lussac‟s Law: P1/T1 = P2/T2

Pressure is directly proportional to temperature (T is in Kelvin)

4. Avagadro‟s Law: N1/V1 = N2/V2

Moles of gas is directly proportional to volume

5. Ideal Gas Law: PV = nRT

R = universal gas constant = 0.0821 (L*atm)/(k*mol)

Another form: MM = (dRT)/P

(molar mass = MM, density = d)

6. Dalton‟s Law of Partial Pressure: PT = P1 + P2 + P3 …

Partial Pressure of a gas = X(Ptotal)

(X = mole fraction of gas)(P = total pressure)

PT = XAP˚ A + XBP˚ B

(ideal solution with more than one volatile compound)

STP = 1atm, 273K

1 mol gas = 22.4L gas

R = 0.0821 (L*atm)/(mol*K) = 8.314 Joules/(mol*K)

Ideal Gases versus Real Gases

2 assumptions of ideal gas:

o no attractive forces between molecules

o molecules are infinitely small

Real Gases behave ideally at:

o High temperature

o Low pressure

Thermodynamics

- thermodynamics – science of heat & work

- ΔH = enthalpy change (heat transferred)

- ΔHrxn = (Σ ΔH°f products – Σ ΔH°f reactants)

- ΔS = entropy (measure of randomness)

increase in entropy =>

solid → liquid → gas

(least entropy) (most entropy)

- ΔG = ΔH – TΔS

When ΔH = TΔS

(and both ΔH and ΔS are positive/negative)

Then ΔG = O (meaning at equilibrium)

- ΔG = - RTlnK

R = 8.314

K = equilibrium constant

Negative value Positive value

ΔH Exothermic Endothermic

ΔS Entropy decrease Entropy increase

ΔG Spontaneous Non Spontaneous

Less than 1 Greater than 1

K Non Spontaneous

(reactant favored)

Spontaneous

(product favored)

Kinetics

- reaction mechanisms – detailed pathways taken by atoms and molecules as a reaction proceeds

- factors that affect reaction rates:

1.) concentration of reactants

2.) temperature

3.) catalysts

4.) orientation of molecules

- Rate Equation (Law): rate = K[A]x[B]

y

K = rate constant; „x‟ and „y‟ can only be determined by experimental data

Δrate = (Δ[conc]) x

◊ Rate Law dependent upon ◊ Slowest Step ◊ of reaction mechanisms

- Order of Reaction is sum of the exponents (x + y …)

- If concentration of reactant is increased, the rate is usually increased (unless the exponent is 0 – not

involved in rate law)

▪ catalyst lowers activation energy

▪ temp increase allows more molecules to overcome EA

▪ 2 criteria for a mechanism:

1.) steps must add up stoichiometricly.

2.) rate law of rate def‟n step = rate law.

Slow Rates Fast Rates

▪ strong bonds

in reactant

molecules

▪ catalyst

▪ high temp.

▪ high [reactant]

▪ low activation

energy

Equilibrium

▪ K does not change unless temp. changes

Equilibrium expression for:

aA + bB ↔ cCdD

K = [C]c[D]

d ([products]/[reactants])

[A]a[B]

b

▪ only gases and aqueous included in equilibrium expression

Kc = equilibrium constant of concentration

Kp = equilibrium constant of pressure

Kp = Kc(RT) Δn

Δn = total moles gaseous product – total moles gaseous reactant

R = 8.31

K = 1 at equilibrium when [prod.] = [reactants], RARE

K > 1 product favored (spontaneous)

K < 1 reactant favored (non spontaneous)

Q = Reaction Quotient

Q > K reactant favored (non spontaneous)

Q < K product favored (spontaneous)

Q = K at equilibrium

Le Chatelier‟s Principle – a change in any of the factors that determine equilibrium conditions of a system cause

the system to change in such a manner to counteract the effect of the change

▪ factors that effect equilibrium:

1. change in concentration

2. change in temp

3. change in volume

4. addition of catalyst

2A + B ↔ C + D Shifts

1.) Add A / Remove D →

2.) Remove B / Add C ←

3.) Increase Volume ← (Shifts to side w/ greater moles of gas)

4.) Decrease Volume → (Shifts to side w/ fewer moles of gas)

If endothermic reaction, X + H → Y

1.) Add heat, Shifts right

2.) Remove heat, Shifts left

If Exothermic X → Y + H

1.) Add heat, Shifts left

2.) Remove heat, Shifts right

Electrochemistry

Oxidation Reduction

- lose electrons - gain electrons

- metals get oxidized - non metals get reduced

- occurs at ANODE - occurs at CATHODE

- cations formed - anions formed

- a.k.a. reducing agent - a.k.a. oxidizing agent

- has a negative Ecell - has a positive Ecell

- E˚ cell = Ered - Eox

˚ indicates standard conditions (25˚C, 1.0 atm)

▪ if Ecell ( + ) spontaneous

( - ) non spontaneous

- Common Oxidizing Agents:

KMnO4 → Mn2+

K2Cr2O7 → Cr3+

Concentrated HNO3 → NO2

Dilute HNO3 → NO

H2O2 → H2O + O2

Getting reduced → ← Getting oxidized

Salt Bridge – device used for maintaining balance of ion charges in the cell

compartments

- 1 Coulomb = current (amps) * time (seconds)

- 1 Faraday = 96,500 coulombs = 1mol e-

- Nernst Equation: E = E˚ - [(0.0257V)/n]*lnQ at 25˚C

Cathode(-) Anode(+)