AP Exam Review

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AP Exam Review

1. Thermochemistry

2. Acids-Bases Review Session 3. Equilibrium

4. Kinetics

5. Electrochemistry 6. Colligative Properties

I. Thermo A. Energy

1. Definition and forms 2. Units

B. Difference between temp and heat C. System vs. Surroundings

D. What are state functions? E. What is the standard state? F. Laws of Thermodynamics

1. 1st Law

2. 2nd Law

3. 3rd Law

G. Enthalpy

1. Definition 2. How measured

3. Endothermic vs. Exothermic 4. Solve problems

5. Hess’s Law

6. Heats of formation 7. Specific Heat 8. Heat Capacity

9. Calorimetry Lab with nested coffee cups

a. Set-Up – procedure; weigh before water added, thermometer b. Equations

H. Entropy

1. Amount of disorder; more positive, more disorder 2. ΔSuniv = ΔSsurr + ΔSsystem

I. Gibbs Free Energy 1. Spontanaeity

2. How solve ΔG = ΔH - TΔS J. Catalysts

II. Acids Bases

A. Arrhenius Theory – examples

B. Brønsted-Lowry Theory – examples (and one not Arrhenius)

C. Lewis Theory – examples (and coordination complexes, how not Arrhenius or Brønsted) D. Amphoteric

E. Strong acids and bases and their conjugates F. Calculate pH and pOH

1. How to do it without a calculator 2. Strong Acids and Bases

3. Weak Acids and Bases using Ka and Kb


1. From weak acids and bases 2. Calculate pH

3. Calculate pH changes when strong acids or bases are added H. Salts

1. Prediction of pH range 2. Calculate its pH I. Titrations – Labs

1. Proper set-up of lab – procedure (rinse buret, void tip, etc.)

a. How to get correct molarity of solutions (M = mol/L; convert grams to moles,


2. Endpoint-Equivalence Point 3. Strong-Strong titrations 4. Weak-Strong titrations 5. Use of indicators J. Anhydrides

K. Oxyacids

III. Equilibrium –

A. Definition of Equilibrium

1. Forward and reverse reactions continue at the same rate B. Solve problems using LeChatelier’s

1. 3 ways to effect equilibrium

a. Temperature (which is the only one that affects K) b. Pressure (inverse of volume)

c. Concentration 2. Practice Problems

C. Set-up of an equilibrium expression

1. Pure Solids and Pure liquids (including water) do NOT appear in the equilibrium


2. aA + bB  cC + dD 3. Law of Mass Action


4. Manipulation of the constant

a. A + ½ B2  AB ; K1 = 2.3 x 103 b. 2A + B2  2AB ; K2 = K12

c. AB  A + ½ B2 ; K2 = 1/K1 d. Ktotal = K1 • K2 • K3

5. Value of K (Kc; ‘c’ for concentration)

a. K > 1; product formation is favored ([react] < [prod]) b. K < 1; reactant formation is favored ([react] > [prod]) c. K = 1; ([react] = [prod])

6. ICE Tables

a. ICE = Initial, Change, Equilibrium

b. Used to identify the CONCENTRATIONS of each species at equilibrium. 7. Practice Problems

D. Reaction Quotient

1. Q – measure of proportions of reactants and products in equation. (Probably not at



b. K > Q; not equilibrium, product formation favored ([react] < [prod]) c. K < Q; not equilibrium, reactant formation favored ([react] > [prod]) d. K = Q; at equilibrium ([react] = [prod])

2. Practice Problems

E. Calculate and interpret the value for Kp 1. Kp = Kc(RT)n – for gases only

a. R = 0.0821 L•atm/mol•K b. T = temperature in Kelvin c. n = moles of gas

IV. Kinetics – What you need to Know A. Reaction Rate

1. Change in concentration (or pressure) over time

a. Graph showing reactant conc decr. over time and product conc. incr. over time 2. Rate affected by: (its all about the collisions – Collision Theory)

a. Concentration (pressure for gaseous species) b. Surface Area

c. Temperature d. Catalyst

B. Rate Constant and the Rate Law Expression 1. Rate Law Expression

a. A + B  C + D b. Rate = k[A]m[B]n

i. Rate = rate of reaction

ii. k = rate constant (specific to an experiment at a given temp) iii. m and n = order of reaction (0,1, or 2)

2. Zero Order

a. No change in rate with change in concentration 3. 1st Order

a. Proportional change in rate with change in concentration 4. 2nd Order

a. Change in rate doubles the concentration change (quadruples?) 5. Integrated Rate Law

a. Finding concentration at time (t)

Reaction Order IntegratedRate Law Linear Plot Slope

0 [A] = [A0] – kt [A] vs t -k

1 ln[A] = ln[A0] – kt ln[A] vs t -k

2 1/[A] = 1/[A0] + kt 1/[A] vs t k

C. Activation Energy

1. Solve using Arrhenius equation a. k = Ae-Ea/RT

b. lnk = -Ea/RT + lnA 2. Interpret Graphs


1. Catalysts used in a reaction but not used up, intermediates not there before and not there


2. Heterogeneous vs. homogenous

a. Homogeneous is a catalyst in the same phase as the reactants b. Heterogeneous is a catalyst in a different phase than the reactants F. Mechanisms

1. Slow step is the rate determining 2. 2A + 2B  2C + D2

a. A + B  E (Fast with equal rates) b. E + A  E2 + C (Slow)

c. E2 + B  D2 + C (Fast)

d. 2A + 2B  2C + D2 Correct Mechanism 3. Use of a catalyst

a. Cl(g) + O3(g)  ClO(g) + O2(g) b. O(g) + ClO(g)  Cl(g) + O2(g)

c. O(g) + O3(g)  2 O2(g)

V. Electrochemistry

A. Oxidation State of elements and compounds

1. Rules

B. Difference between oxidation and reduction

1. Oxidation – loss of electrons at anode (both vowels) 2. Reduction – gain in electrons at cathode (both consonants) C. What is being oxidized and reduced

1. Oxidizing Agent – is being reduced 2. Reducing Agent – is being oxidized D. Balance Redox reactions

1. In acidified solutions

a. Add water to balance oxygens b. Add H+ to balance hydrogens 2. In basic solutions

a. Add OH- to both sides, enough to cancel H+ b. Then cancel waters

3. Complex Redox Rxns a. Common Oxidizers

i. High Charged Metal Cations (Sn4+  Sn2+) ii. MnO4¯ & MnO2 Mn2+

iii. Cr2O72-  Cr3+

iv. Halogens (Diatomic)  Ion (Cl2  2Cl¯) b. Common Reducers

i. Free Metals (Zn  Zn2+) ii. H2  2H+

iii. C2O42¯  CO2

c. Both

i. H2O2 – Acidic to H2O, Basic to O2 + H2O E. Problems with Galvanic Cells


i. An ions attracted to anode

b. Reduction – gain in electrons at cathode (both consonants) i. Cati ons attracted to cathode

c. Salt Bridge – connect the two in order to complete the circuit; U-tube with porous

plugs on each end

i. KNO3 as an example 2. Cell diagram

ANODE | ionic species in solution || ionic species in solution | CATHODE in anodic half-cell in cathodic half-cell

3. Cell Potential

a. Eºcell = Ecathode - Eanode

Example: Find Ecell for the galvanic cell whose overall redox reaction under standard conditions is:

3 Zn(s) + 2 Cr3+(aq)  3 Zn2+(aq) + 2 Cr(s)

First, split the reactions up into the oxidation and reduction halves. Oxidation at anode: Zn(s)  Zn2+(aq) + 2e- E = +0.76 V

Reduction at cathde: Cr3+(aq) + 3e-  Cr(s) E = -0.74 V

In order to combine these half-reactions to give the overall redox reaction, note that we must first multiply the zinc oxidation reaction by 3 and the Cr3+ reduction equation by 2. Recall from balancing

redox reactions that we want the electrons to cancel out.

NOTE THAT YOU DO NOT MULTIPLY E BY ANYTHING!!! 3Zn(s)  3Zn2+(aq) + 6e- E = +0.76 V

2Cr3+(aq) + 6e-  3Cr(s) E = -0.74 V

---3 Zn(s) + 2 Cr3+(aq)  3 Zn2+(aq) + 2 Cr(s) E

cell = +0.02 V

F. Draw and label a simple voltaic cell G. Draw and label an electrolytic cell H. Solve problems with the Nernst equation

1. Gº = -nFEºcell

a. n = number of moles of electrons transferred

b. F = Faradays constant = 96,500 C/mol e- = 96,500 J/V

2. at 25 ºC

VI. Colligative Properties A. Properties

1. Vapor Pressure (Raolt’s Law for Ideal Solutions – totally miscible) a. Vapor Pressure = XP

b. X = mole fraction of solvent (mol solvent/total moles) c. P = Vapor Pressure of PURE solvent

d. Total Pressure found by just adding the vapor pressures 2. Boiling Point Elevation


b. ΔTb = change in boiling point, increase c. i = van’t Hoff factor

i. i = (# of part. after dissoc.)/(# of formula units) ii. CaCl2 Ca+ + 2 Cl-; I = 3/1 = 3

d. m = molality (moles of solute/kg of solvent) 3. Freezing Point Depression

a. ΔTf = i*Kf*m

b. ΔTf = change in freezing point, decrease c. i = van’t Hoff factor

4. Osmotic Pressure

a. π = iMRT

b. π = osmotic pressure c. M = Molarity

d. R = 0.0821 – gas constant e. T = Temp in Kelvin f. i = van’t Hoff factor

VII. Misc.

A. Gases/Solutions 1. Phase Diagram

a. Which part of the phase is a solid, liquid, or gas?

b. Which part of the diagram shows the phase changes, melting/freezing,


2. Solubility of gases into solution B. Ligands/Complex Ions

1. [Al(H2O)6]3+, [Fe(H2O)6]3+

2. Use of ‘outermost’ level empty bonding orbital 3. Lewis Base most of the time, lone pair donor

a. Find the charge of the ligand, double the charge for the cation and add the charges

altogether to find the charge of the polyatomic ion.

C. Calculations