Chem II: Final Exam Review (Part 2/4)

(quick links to all 4 parts)


Ch. 12 and 13 Practice Test (Solutions and Rates of Reaction) 

Chapter 14-16 Practice Test (Part 2: Chemical Equilibrium, Acids and Bases, and Acid-Base Equilibria)

Chapter 17 and 18 Practice Test (Part 3: Solubility and Complex-Ion Equilibria, and Thermodynamics and Equilibrium)

Chapters 19, 20, and 23 Practice Test (Part 4: Electrochemistry, Nuclear Chemistry, and Organic Chemistry)



Chapter 14-16 Practice Test

(Part 2: Chemical Equilibrium, Acids and Bases, 

and Acid-Base Equilibria)

1.       Define Equilibrium

a.        The rate of the forward rxn is equal to the rate of the reverse rxn.

b.       The concentration of reactants and products remains constant.

c.        Opposing processes occur at equal rates.

d.       Dynamic.

e.       Has nothing to do with collision.

f.         Examples

                                                               i.      R(evaporation) = R(condensation)

                                                              ii.      R(dissolution) = R(crystallization)

2.       The Equilibrium constant…

a.        Will not tell you the speed of the reaction.

b.       Will tell you the extent of the reaction.

c.        Predicts the direction of the reaction.

d.       Allows to calculate the equilibrium concentrations.

3.       Relate Qc to Kc to Direction of the reaction.

a.        Qc <Kc = reaction goes left to right.

b.       Qc>Kc = reaction goes right to left.

4.       If you reverse and cut in half a reaction…

a.        Kc = [1 / Kc^(1/2) ]

5.       If you increase pressure, volume…

a.        Expands towards the direction of fewer molecules of gas.

b.       Unless there are equal moles on each side.

c.        Adding H2O to a solution…

                                                               i.      …will not affect equilibrium.

                                                              ii.      …will dilute the Molarity.

6.       To find Equilibrium Pressure Kp…

a.        Kp = Kc(RT)^(∆n)

b.       ∆n = (c+d) – (a+b)

c.        R = 0.0821

d.       T = temp in K.

7.       Lewis Acids include:

a.        Something that will accept an electron pair.

b.       Group III (Boron, Aluminum, Galium)

8.       The _____ acid will have the strongest conjugate base.

a.        weakest

9.       Given Hydronium concentration, find pH:

a.        pH = –log (H3O)

10.    Given pH, find Hydronium concentration:

a.        [H3O] = 10^(-pH)

11.    Find Ka given initial concentration and pH:

a.        10^(-ph) = [H+]

b.       ([H+]^2) / Molarity = Ka

12.    What is Hydrolysis?

a.        When an ion reacts with water.

13. 

1.       Strong Acids and Bases:

a.        Are not Hydrolyzed

b.       Do not achieve equilibrium

c.        Strong Bases:

                                                               i.      Include Group IA (except H)

                                                              ii.      Include Group 2B (except Beryllium)

d.       Strong Acids:

                                                               i.      Include

e.       When added together, make a neutral solution.

2.       Ka will favor products if it is _greater_than _10^3_ and reactants if it is _lower_than_10^-3_.

3.       How do you determine if x in ([M]-x) can be neglected?

a.        If [A]/Ka > 100, then x may be neglected.

4.       How do you calculate % Dissociation?

a.        100 * (x dissociated Acid or Base / [M] initial concentration ) 

5.       About Buffers:

a.        Mix a weak acid or weak base and its conjugate.

b.       Water is not a buffer.

c.        Resist a change in pH when adding an acid or a base.

d.       Buffer Capacity is the limited amount the buffer can absorb before a significant change.

e.       Buffer Capacity depends on the number of moles.

6.       What is the Henderson Hasselbach Equation?

a.        pH = pka log ([Base]/[Acid])

b.       pOH = pkB log ([Acid]/[Base])

c.        pka = -log(ka)

d.       Ka = Kw/Kb

e.       Used to calculate the pH of a Buffer Solution.

7.       The second ionization constant:

a.        Is Ka₂ = A^2-

 

8.       What is the common ion effect? What’s it do? Give examples.

a.        The presence of an ion already present in the solution, based on the LeChatlier Principle.

b.       It causes:

                                                               i.      Suppression/Decrease of Ionization

                                                              ii.      Makes the degree of ionization smaller.

                                                            iii.      Shifts the reaction to the left.

                                                            iv.      Acid/Base is weakened.

c.        Eg. The solubility of silver chloride AgCl is reduced if a solution of NaCl is added.

9.       Amphoteric…

a.        Can act as an acid or a base, but doesn’t need a hydrogen.

b.       Eg. Al2O3

10.    Amphiprotic…

a.        Is Amphoteric but must have a Hydrogen in the formula.

b.       Eg. HSO4

11.    Arrhenius Acid:

a.        Releases H+ in water.

b.       Releases OH- in water.

12.    Br0wnstead/Lowry Acids:

a.        B/L A = proton donor

b.       B/L B = proton acceptor

13.    Lewis Acids and Bases…

a.        LA needs a lone pair of e-.

b.       LB has a lone pair of e- available.

14.    Acid Strength Depends on:

a.        Polarity – the more electronegative the NON-H part of a molecule, the stronger the acid, because H can break away easier.

b.       Bond Strength – Electron radius increases to the right and down. The larger the radius, the weaker the bond, so the stronger the acid.

c.        Oxoacids – The more O in an Oxoacid, the stronger it is.