The Daily Constant: answers

Saturday, May 7, 2011

Chem II: Final Exam Review (Part 3/4)

(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 17 and 18 Practice Test

(Part 3: Solubility and Complex-Ion Equilibria,

and Thermodynamics and Equilibrium)

1. Write the Ksp expression and x-formula for the following salts:

a. Ca(OH)₂

i. K_sp = [Ca²⁺][OH^-]²

ii. K_sp = (x) * [(2x)²]

b. Na₃PO₄

i. K_sp = [Na⁺]³[PO₄^3-]

ii. K_sp = [(3x)³][x]

2. What are Ksp and Qsp, and what is the difference between them?

a. Ksp = Product Solubility Constant, it tells you the solubility of a salt in water at equilibrium. It may be used to tell if a precipitate will form upon at equilibrium.

b. Qsp = Reaction Quotient, Ratio of Products over reactants at instant of mixing. Can also be expressed as Qip to determine if a precipitate will form upon the instant of mixing.

c. If Qip > Ksp, then the solution is super-saturated (precipitate forms)

d. If Qip = Ksp, then the solution is saturated and at equilibrium.

e. If Qip < Ksp, then it is unsaturated, and there is no precipitate.

3. 45mL of 0.015M CaCl₂ is added to 55mL of 0.010 M Na₂SO₄.

a. What is the Qip?

i. 3.71 x 10^-5

b. What is the Ksp?

i. 2.4 x 10^-5

c. Will a precipitate form, and why or why not?

i. Qip > Ksp, so a Precipitate will form.

4. The solubility of magnesium oxalate, MgC₂O₄ in water is 0.0093 mol/L. Calculate K_sp

a. 8.65 x 10^-5

5. The solubility of silver chromate, Ag₂CrO₄ in water is 0.022

10⁰ g/L. Calculate K_sp.

a. 1.16 x 10^-12

6. What is a complex ion, how are the formation and dissociation constants for it expressed? Give an acidic and basic example.

a. A complex ion is an anion that is composed of a metal and one or more ligands.

b. A ligand is a molecule or anion bound to a metal through coordination of lone pairs.

c. The Formation Constant is Kf and the Dissociation constant Kd = 1/Kf.

d. Typically the Kf for a complex ion is very large and the reaction goes to completion.

e. Co(H₂O)₆²⁺ - CoCl₄^2- are both examples of a complex ion (Co) bound to ligands.

f. The complex ion acts as a Lewis Acid, accepting elections.

g. The ligands act as a Lewis Base, donating elections.

h. Solubility is increased because the formation of it removes a product.

7. The solubility of magnesium fluoride, MgF₂, in water is 0.016 g/L. What is the solubility (in grams per liter) of magnesium fluoride in 0.014 M sodium fluoride, NaF?

a. 2.2 x 10^-5g/L

8. What is the solubility of MgF₂ in water? The K_sp for MgF₂ is 7.4 x 10^-11

a. 2.6 x 10^-4M

9. What are the Thermodynamic State Functions?

a. E, G, S, P, T, H, V, Qp “Every Good Student Picks The Hot Valentine Qp.”

10. Which Thermodynamic variables are not State Functions?

a. q (heat) and w (work).

11. Why study Thermodynamics?

a. It can tell us if a reaction is possible, the entropy, energy, and whether or not it is spontaneous or not. This saves us time, energy, money, and frustration.

12. Why study Kinetics?

a. It can tell us how fast a reaction happens, how we got there, the reaction mechanism, and collision theory.

13. What is ∆H, and what can it tell us, and how?

a. Enthalpy, also known as q_p,and is a state function.

b. If ∆H is negative, heat is a product and the reaction is exothermic.

c. If ∆H is positive, heat is a reactant, and the reaction is endothermic.

14. Are the following situations more likely to be Spontaneous or NonSpontaneous?

a. Flow of heat from a hot place to a cold place? (Spontaneous)

b. State change from liquid to solid? (Non-Spontaneous)

15. What is Gibb’s Free Energy Equation and what must be remembered?

a. ∆G = ∆H - T∆S

b. Remember ∆S is in Joules, not KiloJoules.

c. ∆G = 0 at equilibrium.

16. How can Gibb’s Equation be used to determine K , and what must be remembered?

a. ∆G = -RT ln(K)

b. Ln(K) = ∆G / -RT

c. Convert ∆G to Joules, as R is our constant.

d. ∆G = 0 at equilibrium.

17. What are the units of measurement in the following?

a. ∆G = kJ/mol

b. ∆S = J/K

c. ∆H = kJ/mol

d. R = J/K*mol

18. What is the 1^st Law of Thermodynamics and its associated formula?

a. Energy can neither be created nor destroyed.

b. Change in Internal Energy (∆u) is equal to the heat of the reaction plus the work done to the system.

c. ∆u = q+w (Internal Energy = Thermal + Kinetic Energy.)

19. What is the 2^nd Law of Thermodynamics and its associated formula?

a. Entropy will increase if a process is spontaneous.

b. ∆S > (q/T) is a spontaneous reaction

c. ∆S = (q/T) is an equilibrium reaction

d. ∆S < (q/T) is a non-spontaneous reaction

20. What is the 3^rd Law of Thermodynamics and its associated formula?

a. Entropy will increase with temperature.

b. The Entropy of a perfectly ordered crystalline substance at 0 K is 0.

21. Which anions are soluble in acidic solutions? Why? Give examples.

a. Very Basic Anions will be more soluble in acid than water.

b. F-, CO₃^2-, S^2-, OH^-, PO+4^3-

c. Weak Basic Anions will not be affected much by pH.

2. Under what conditions do the following indicate a spontaneous reaction?

a. K ? (If value is greater than 1.0 x 10³ it is spontaneous and favors products)

b. ∆G (if value is negative)

c. ∆S (if value is positive)

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

(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.