October 1, 2004

Review Notes, Exam 2

Exam covers Chapter 13 Equilibrium, Chapter 15 Aqueous Equilibrium and the portion of Chapter 16 concerned with Ksp and the solubility of ionic salts. A review of Ch 14 might be appropriate, since that's where the H+ / H₃O+ / OH- / H₂O relationship is introduced.

The exam will not include any calculations for diprotic or polyprotic acids, although you should understand the concepts involved.

The only equation provides is for quadratic equation

• -- if ax₂ + bx + c =0 then
• x = (1/2a) { -b + (b²-4ac)^1/2}
• In practice I won't ask you to finish the solution for any quadratic equation if it arises.

1. You will not be expected to know values for any specific equilibrium constant-- except that for Kw. You should know that Kw = 1.0 x 10^-14.
2. You should know that HCl and HNO₃ are strong acids and that NaOH, KOH and Ca(OH)₂ are strong bases. Strong means 100% dissociated in water, forming H₃O⁺ or OH^- ions.
3. First, you must be able to write the equilibrium expression associated with any stated chemical equation. (Products on top, reactants on the bottom. Gases as Pi if a Kp is used or [gas] if Kc. Solutions use [species], the molar concentration; Solids and pure liquids, including water as a solvent, are omitted in writing K.)
4. You must also be able to write the chemical equation for many common reaction types, illustrated below by specific examples.

• weak acid H-A + H₂O ---> H₃O⁺ + A- Ka
• weak base NH₃ + H₂O --> NH₄+ + OH^- Kb
• salt solution MX₂ (solid) ---> M²⁺ + 2 X^- Ksp
• water H₂O --> H+ + OH^- Kw (or 2 H2O --> H₃O⁺ +OH^- )
• diprotic acid H₂CO₃ + H₂O ---> H₃O⁺ + HCO₃ - Ka1
• HCO₃ - + H₂O ---> H₃O⁺ + CO₃ -2 Ka2
• evaporation H₂O (liq) ---> H₂O(g) Kp In the case of acid-base you must know the definition of pH (pOH, Pka) and be able to evaluate [H₃O+] if you are given the pH (or pH for a given [H₃O+].)
  • You must also know that pH + POH = pKw or [H₃O⁺][OH^-] = 1.0 x 10^-14
  1. If the conc/pressure of all species is given, you should be able to use Q to determine which direction the mixture will change to reach equilibrium. (Applied to gases, to solutions, to mixtures of ions that form a soal to low solubility.)
  2. If all conc/pressure (or pH) are given at equilibrium, you should be able to evaluate K.
  3. If given K and starting conditions you should be able to find final conditions. This typically involves a table of initial values, x= change of one variable, other changes expressed with x, final concentrations= intital + change. Substitute into Keq expression.
• You must understand the Bronstead Lowry concept of acids and complementary bases
  • e.g., NH₃ (base) + H₂O (acid)--> NH₄+ (acid) + OH^- (base)
  • NH₃ (proton acceptor) + H₂O (proton donor)--> NH₄+ (acid) + OH^- (base)
  • A complementary acid base pair (like NH₄+ and NH₃ or Acetic acid/ acetate ion ) have
    • Ka * Kb = Kw
  • There are a few ways of expressing a solution -- degree of dissociation of a weak acid
• You should know that Keq generally varies with temperature-- Keq increases as the temperature rises if the reaction is endothermic; K decreases as temperature rises for freactions that are exothermic. (You won't be asked to compute K at another temperature)
• LeChatlier's Law allows one to predict how an equilibrium mixture adjusts when conditions change.

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  Sample Problems

  Assume that 1.0 % of acetic acid vapor dimerizes at 2.5 atm pressures.

  2 CH₃CO₂H (g) ---> {CH₃CO₂H}₂ (g)

  Evaluate Keq

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  A 0.050 M solution of a weak acid has a pH of 4.55. What is Ka for the acid?

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  The salt QMF₃ --> Q³⁺ (aq) + 3F^- (ag) and has a solubility of 2.1 x 10-4 moles per liter.

  Evaluate Ksp

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  Ksp for Ba(OH)₂ is 3 x 10^-3. If we mix equal volumes of 0.05M BaCl2 and 0.020 M NaOH, will we get a precipitate to form?

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  NH₄+ (or NH₄Cl ) acts like a weak acid

  Kb for ammonia is 1.75 x 10-5

  what is Ka for ammonium ion?

  what is the pH of a 0.10 M solution of ammonium chloride

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  What is the pH of a solution that is 0.10 M Acetic acid; Ka = 1.8 x 10^-5

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  Kc = 138 for the reaction Fe³⁺ + SCN- ----> Fe(SCN)²⁺

  If we start with 2.0 x 10-3M Fe³⁺ and 0.050 M SCN-, how much Fe(SCN)²⁺ forms?

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  What is the pH of a solution that is 0.030M Acidic acid and 0.020 M potassium acetate?

  Ka for acetic acid is 1.8 x 10^-5.

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  Concepts--

  • Why is sodium oxalate more soluble in 0.1 M HCl than it is in water?
  • Why is PbCl₂ less soluble in 0.1M HCl than it is in water?
  • Will adding a drop of HCl (1M) to 5 ml of 1 M solution of acetic acid
    • a. raise the pH significantly
    • b. lower the pH significantly
    • c, have relatively little effect on the pH
  • Will adding a drop of 1M Acetic acid to 5 ml of 1 M HCl
    • (same three choices)
  • If you add 1 drop (1/20 ml) of 1 M NaOH to 5.0 ml of water, what will be the pH?
  • If you added that NaOH to a solution that is 0.1M Acetic acid and 0.10 Sodium acetate,
    • why would the pH change be very small
  • Adding sodium thiosulfate to water forms a solution that is several degrees cooler than the starting water.
    • Is sodium thiosulfate more soluble in hot water or in cold water?
  • If you slowly add 0.1M NaOH to a solution of HCl and acetic acid could you...
    • produce a solution that is mainly NaCl and Acetic acid?
    • produce a solution that is mainly HCl and sodium acetate?
    • a solution that has lost half of the HCl and half of the acetic acid?
    • a solution where OH- and acetate ion both exist at modest concentrations?
    • a solution that is pH 7.00 ?
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  • The freezing point of a 1.0 M solution of glucose (a sugar) is -1.86 ^oC
    • The freezing point of a 1.0 M solution of a weak acid is 2.01 ^oC
    • why the difference?
    • how could this information be used to find Ka for the acid
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  • CaCO ₃ (solid limestone) can be heated in air to completely decompose to CaO + CO₂ (g)
    • If CaO is left exposed to normal air, it picks up CO₂ and becomes CaCO₃ again.
    • can you explain and use this information to tell something about the reaction
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