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EQUILIBRIUM REVIEW #2 (Ka/Kb)
Ka or Kb
Here are some things in common with most Ka/Kb problems.
1. Many of them will ask for you to determine pH or pOH. To do this, you obviously need to know [H+] or [OH-]. If
you have a strong acid or strong base, you can determine [H+] or [OH-] directly from the acid or base
concentration. If you have a weak acid or base, you must set up an ICE table and determine [H+] or [OH-], then
pH or pOH.
You may be asked to find a % ionization (dissociation – same thing).
To find the % dissociation you take [H+] / [HA] X 100
To find the percent dissociation for a weak base: [OH-] / [B] x 100
2. Buffers are very common to Ka and Kb. This means knowing the following.
(a) A buffer is a combination of a weak acid and its conjugate base or a weak base and its conjugate acid.
(b) To determine the pH of a buffer, use the Henderson-Hasselbalch equation: pH = pKa + log B/A. The
ratio B/A is important. It can either be a ratio of moles or molarities. This means that if you dilute a
buffer you do not change the pH of the buffer – but you change the buffers capacity for consuming H +
and OH-.
(c) The purpose of a buffer is to resist a change in pH. Thus a buffer must have a base component and an
acid component.
(d) Expect stoichiometry if a buffer is present. You may be asked to determine the final pH of a buffer when
a certain number of moles of H+ or OH- are added. The very first thing you do write a net ionic equation
and set up an I(initial)C(change)F(final) table. At completion – ask WHAT DO I HAVE?????????? That
way you can find pH.
(e) If you have a buffer –use H & H to find pH
(f) If you have excess strong acid, use pH = -log [H+]
(g) If you have excess strong base – just use pOH = -log[OH-], then pH = 14- pOH
2.
Titrations are also very common. Know the following:
(a) Parts of a pH (titration) curve
(b) How to calculate pH at various points in a strong acid-strong base titration: Before equiv. pt, excess
strong acid determines pH, at equiv. pt,- pH is 7, beyond equiv pt. – excess strong base determines pH
(find pOH first, then subtract from 14)
(c) How to calculate pH at various points in a weak acid./strong base titration (or weak base/ strong acid
titration): before equiv pt. – buffer is formed and determines pH, at equiv. pt – hydrolysis occurs (write
hydrolysis reaction to determine whether pH at equiv. pt. is acidic or basic and then do a pH calculation),
beyond equiv. pt – excess strong base (or strong acid) determines the pH.
(d) How to pick an appropriate indicator for a titration. (Need to know pH at equiv. pt to determine this.)
3.
Always remember that for a conjugate acid/base pair that Ka x Kb = Kw. They will use this very frequently – give
you Ka but what you need is Kb.
4.
Be able to identify Bronsted acids and bases (conjugate acid-base pairs as well), and Lewis acids and bases.
5.
Be able to determine strength of an acid based on its Ka value or structure.
PRACTICE PROBLEM
CH3NH2 + H2O <===> CH3NH3+ + OH¯
1)
Methylamine, CH3NH2, is a weak base that reacts according to the equation above. The value of the ionization constant, K b,
is 5.25 x 10¯4. Methylamine forms salts such as methylammonium nitrate, (CH 3NH3+) (NO3¯).
(a) Calculate the hydroxide ion concentration, [OH¯], of a 0.225-molar solution of methylamine.
(b) Calculate the pH of a solution made by adding 0.0100 mole of a solid methylammonium nitrate to
120.0 milliliters of a 0.225-molar solution of methylamine. Assume that no volume change occurs
.
(c) How many moles of either NaOH or HCl (state clearly which you choose) should be added to the
solution in(b) to produce a solution that has a pH of 11.00? Assume that no volume change occurs.
(d) A volume of 100. milliters of distilled water is added to the solution in (c). How is the pH of the
solution affected? Explain.
PRACTICE PROBLEM
1) The acid ionization constant, Ka, for propanoic acid, C2H5COOH, is 1.3 x 10-5.
(a) Calculate the hydrogen ion concentration, [H+], in a 0.20-molar solution of propanoic acid.
(b) Calculate the percentage of propanoic acid molecules that are ionized in the solution in (a).
(c) What is the ratio of the concentration of propanoate ion, C2H5COO-, to that of propanoic acid in a buffer
solution with a pH of 5.20 ?
(d) In a 100-milliliter sample of a different buffer solution, the propanoic acid concentration is 0.35-molar and the sodium
propanoate concentration is 0.50-molar. To this buffer solution, 0.0040 mole of solid NaOH is added. Calculate the pH
of the resulting solution.
HOMEWORK: EQUILIBRIUM
1. Answer the following questions that relate to the chemistry of halogen oxoacids.
(a) Use the information in the table below to answers part (a) (i).
Acid
HOCl
HOBr
Ka at 298K
2.9 x 10-8
2.4 x 10-9
(i) Which of the acids is stronger, HOCl or HOBr? Justify you answer in terms of K a.
(ii) Draw a complete Lewis electron-dot diagram for the acid that you identified in part (a)(i).
(iii) Hypoiodous acid has the formula HOI. Predict whether HOI is a stronger acid or a weaker acid
than the acid you identified in part (a)(i). Justify your prediction in terms of chemical bonding.
(b) Write the equation for the reaction that occurs between hypochlorous acid and water.
(c) A 1.2 M NaOCl solution is prepared by dissolving solid NaOCl in distilled water at 298 K. The hydrolysis
reaction OCl- (aq) + H2O (l)  HOCl (aq) + OH- (aq) occurs.
(i) Write the equilibrium-constant expression for the hydrolysis reaction that occurs between OCl- (aq) and
H2O (l).
(ii) Calculate the value of the equilibrium constant at 298 K for the hydrolysis reaction.
(iii) Calculate the value of [OH-]- in the 1.2 M NaOCl solution at 298K.
(d) A buffer solution is prepared by dissolving some solid NaOCl in a solution of HOCl at 298 K. The pH of
the buffer solution is determined to be 6.48.
(i) Calculate the value of [H3O+] in the buffer solution.
(ii) Indicate which of HOCl (aq) or OCl- (aq) is present at the higher concentration in the buffer solution.
Support your answer with a calculation.