CHEMISTRY 162
FINAL EXAM CONCEPTUAL REVIEW
1. Spontaneous processes
a. Processes that do not require a constant input of energy
i. Aging
ii. Water flowing down hill
iii. Air coming out of a tite with a hole in it.
b. Entropy and Enthalpy two driving forces promoting spontaneous changes
i. Free energy is the difference between enthalpy and entropy called Gibbs
free energy
ii. When this difference is negative the chemical process is spontaneous.
iii. Positive difference nonspontaneous
iv. And a zero difference is equilibrium
v. Some spontaneous changes are temperature dependent; you should be
able to predict and calculate this knowing values of enthalpy and entropy
2. Entropy units j/K-mole
a. Entropy is a measure of disorder
i. Since there are more possible states of disorder, then the probability of
processes having disorder are favored and this is why spontaneous
changes having disorder are favored.
1. solid → gas, favored
2. Solid → dissolving favored, only if interparticle forces are weak.
Strong interparticle forces, such as hydrogen bonding can cause
more order in solution formation.
3. Microstates are used to predict order. A microstate is an area
where a particle may be found. The larger the volume the more
possible locations.
ii. Second law of thermodynamics, entropy tends toward a maximum, or
the entropy of the universe tends toward a maximum
iii. Third law of thermodynamics the entropy of a perfect crystal is zero.
iv. Calculating entropy
3. Kinetics (Rates of chemical reactions)
a. Units concentration/time
b. Average rate and how to calculate
c. Instantaneous rate and how to calculate
d. Rate Laws
i. Differential rate law
ii. Integral rate law
iii. Rate constant, only temperature change will cause a change in the rate
constant.
iv. Reaction order
v. Reaction half-life
vi. Pseudo order reactions
e. Factors affecting rate
i. Temperature
ii. Concentration
iii. Surface area
iv. Catalyst
v. Nature of reactant
f. Arrhenius equation
i. Energy of activation
ii. Frequency factor
iii. Two point Arrhenius equation
1. Eliminates the y-intercept
2. Derived by subtracting two equations at different temperatures.
g. Reaction mechanisms
i. Elementary steps
ii. Slowest step determines the rate law
iii. Catalysts
iv. Reactive intermediates
h. Reaction Quotient Q
i. Q>K, means shift left
ii. Q<K, means shift right
iii. Q=K only at equilibrium
4. Equilibrium
a. Rate forward reaction equal to the rate of the reverse reaction
b. Equilibrium constants Kc, Kp
c. Equilibrium expression: products/reactants
i. Solids and liquids do not play the equilibrium game
ii. Relation between kp and kc
iii. How does the equilibrium constant change by
1. Doubling the coefficients of the expression
2. Reversing the expression
d. LeChatelier’s Principle
i. Changing concentrations
ii. Pressure changes
iii. Temperature changes
1. Constant changes if temp goes up
2. Does not change if heat is added isothermally
iv. Effect of catalyst
e. Thermodynamics and equilibrium
i. ΔG= 0 for equilibrium
ii. ΔG = ΔG° +RT lnQ at equilibrium ΔG =0, so ΔG° = -RT lnK
iii. Different temperatures different K, ΔG° same or different
5. Acids Bases and Salts
a. Weak acids/bases are in equilibrium
b. Salts are not in equilibrium, unless made from weak acid or base
i. Weak acid + strong base creates a basic salt
ii. Weak base + strong acid creates and acidic salt
iii. Weak acid + weak base creates neutral, acidic, or basic salt.
c.
d.
e.
f.
g.
h.
i.
j.
k.
l.
1. If Ka of the weak acid is larger than the Kb, then the salt is acidic
2. If Ka of the weak acid is smaller than the Kb, then the salt is basic
3. If Ka = Kb then a neutral salt is formed
Conjugate acids of weak bases are relatively strong (Compared to the acid)
Conjugate bases of weak acids are relatively strong (Compared to the base)
Cannot make strong acids or bases
Auto ionization of water and kw
Water ionization does not affect pH/pOH unless the dilute solutions have
concentrations on the order of 10-14
Acid strength and structure
i. Ion size in a group
ii. Electronegativity in a period
iii. Relative to the ratio of H to O
pH of salt solutions, KaKb = 10-14 be sure to use the correct one. Very common
mistake on midterm #2
Buffers and the Henderson Hesselbach
i. Molarity or moles since same volume
ii. Only works for both HA and A- present
iii. Be sure to determine excess acid or base
iv. pH and pOH for strong acids and bases depend only their concentration
Solubility equilibria
Titration
i. End point
ii. Equivalence point
iii. Analate
iv. Titrant
v. Indicator