AP Chemistry
Name:__________________
Chapter 14 and 15: Acids and Bases
Date:_____________
Chapter 14 and 15 Summary:
Arrhenius definition of acids and bases:
Acid: a chemical compound that increases the concentration of hydrogen ions (H +) in
aqueous solution
Base: a substance that increases the concentration of hydroxide ions (OH -) in
aqueous solution
Brønsted- Lowry definition of acids and bases:
Acid: a molecule or an ion that is a proton donor
Base: a molecule or an ion that is a proton acceptor
HNO3 + H2O → H3O+ + NO3acid
base
conjugate
acid
conjugate
base
Example 1: Give the formulas for the conjugate base of H2SO4 and the conjugate acid of CH3NH2.
Lewis definition of acids and bases:
Acid: atom, ion or molecule that accepts an electron pair to form a covalent bond
Base: atom, ion or molecule that donates an electron pair to form a covalent bond
Strong And Weak Acids:
type of acid, HA
reversibility of
reaction
Ka value
ions existing when
acid, HA, dissociates
in H2O
strong
not reversible
(→)
very large
H+ and A- only, no HA
weak
reversible
(⇌)
small
H+, A-, HA
1
YOU SHOULD KNOW:
Strong Acids: HCl, HBr, HI, HClO4, HNO3, H2SO4
Strong Bases: group I and II hydroxides (e.g. NaOH)
Most common base: NH3 (weak)
Acetic acid (diluted to make vinegar): HC2H3O2 or CH3COOH (weak)
Example 2: List the acids in order of increasing strength (use the K a values in the appendix of your
textbook):
HCN, HCl, HClO2, HNO2.
Example 3: Arrange the following species in order of increasing base strength: NO 2-, ClO2-, CN-, Cl-
Calculating pH of Strong Acids and Bases
pH = -log [H+]
pOH = -log [OH-]
pH + pOH = 14
Example 4: Calculate the pH of 0.010 M HCl
Example 5: The pH of a Sr(OH)2 solution is 13.50. Calculate the concentration of Sr(OH)2.
**number of decimal places in a pH measurement is equal to the number of decimal places in the pH
Calculating pH of Weak Acids
Weak: not all acid dissociates
1.
Write the reaction of the acid with water
2. Set up an ICE chart
3. Write the equilibrium expression for Ka
4. Plug in equilibrium values and solve for x
5. Calculate pH from [H3O+] = x
Example 6: Calculate the pH of 0.25 HCN
2
Calculating % dissociation
𝑥
[𝐻𝐴]0
x 100%
Example 7: The percent dissociation of an acid, HA, which is 1.00 M is 2.5%. Calculate the K a of the
acid.
Calculating pH of Weak Bases
Same steps as pH of weak acids except you find pOH and then in the last step get pH.
Example 8: The pH of a 0.20 M solution of H2NNH2 is 11.38. Calculate Kb for H2NNH2.
3
Polyprotic Acids

Donate more than one H+

Dissociate my losing one H+ at a time
Example 9: Calculate the [H+] of a 0.20 M solution of H3AsO4. Also calculate the concentrations of
H3AsO4, H2AsO4-, HAsO42-, AsO43- .
4
Acid –Base Properties of Salts
Use the following flowcharts
Example 10: Determine whether an aqueous solution of KC2H3O2 is acidic, basic or neutral.
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Calculating the pH of Salts
1.
Decide whether the salt is acidic, basic or neutral.
2. If basic, the hydrolysis reaction will produce an acid
3. Write equation for the anion with water to form the acid and OH- ions
4. Write equilibrium, expression of the above reaction
Example 11: Determine the pH of a 0.100 M aqueous solution of NaCN. The K a for HCN is 5.8 x 10-10.
KaKb = K w
Kw = 1.0 x 10-14
6
Buffers
Buffer: a solution that resists change in pH when acids or bases are added.
The components of a buffer are summarized in the next table.
Buffer
Components
weak acid + salt containing the
conjugate base
weak base + salt containing the
conjugate acid
excess weak acid + strong base
excess weak base + strong acid
Examples
HCN and NaCN
CH3NH2 and CH3NH3Cl
2 mol of HCN + 1 mol NaOH react to yield
1 mol HCN and 1 mol NaCN
HCN + NaOH → HCN NaCN
2 mol NH3 + 1 mol HCl react to yield 1 mol
NH3 and 1 mol NH4Cl
NH3 + HCl → NH3 + NH4Cl
Calculating the pH of a Buffer

similar to equilibrium calculations, but 2 initial conc., one for each part of the buffer pair

same steps as calculating pH of weak acids (in below example)
Example 12: Calculate the pH of a solution that is 0.60 M HF and 1.00 M KF. K a for HF is 7.2 x 10-4.
1.
Write the reaction of the acid with water
2. Set up ICE chart
3. Write equilibrium expression for Ka
4. Plug in equilibrium values and solve for x
5. Calculate pH from [H3O+] = x
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Preparation of a Buffer
2 ways to solve
1.
ICE chart
2. Henderson-Hasselbalch Equation: pH = pKa + log([conjugate base]/[acid])
pKa = -log Ka
Example 13: Calculate the mass of NaC2H3O2 required toprepare a buffer of pH 4.55 when added to
0.500 L of 0.67 M acetic acid. (Assume no change in volume.) K a = 1.8 x 10-5 for HC2H3O2.
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Addition of Strong Acid to a Buffer
Solve problem the same way as Example 9, with the new first steps
1.
Calculate initial moles of weak acid and conjugate acid present in the buffer
Write these amounts under the reaction above the ICE table
2. Calculate the moles of strong acid, H+, added
The added strong acid will react with the basic part of the buffer
3. Make a new line under the initial moles called “add mol H +”
- shows amount of conjugate base decreasing upon H + added
+ shows the amount of acid increasing
4. Calculate the new initial moles of weak acid and its conjugate base, after the conjugate base
has reacted with the strong acid, by adding or subtracting the number of moles H + added
5. Fill out ICE chart computing initial concentration of weak acid and conjugate base by
dividing moles of weak acid and conjugate base by the total volume (volume of a buffer +
volume of added strong base, in L)
6. Fill in ICE chart and plug values into the Ka expression (or use Henderson-Hasselbalch) as in
normal buffer problems
Example 14: Calculate the pH when 100.0 mL of 0.50 M HCl are added to the buffer consisting of
20.0 g HC2H3O2 and 18.0 g of NaC2H3O2 dissolved in 5.00 x 102 mL of water.
Equation:
Initial Moles:
Add mol H+:
New Initial Moles:
I
C
E
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Addition of Strong Base to a Buffer
Similar steps to example 14
Example 15: Calculate the pH when 100.0 mL 0.50 M NaOH are added to the buffer consisting of
20.0 g of HC2H3O2 and 18.0 g of NaC2H3O2 dissolved in 5.00 x 102 mL of water.
Equation:
Initial Moles:
Add mol H+:
New Initial Moles:
I
C
E
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Titrations
Titration: Used to determine the concentration of an unknown solution. The calculations involve
stoichiometric principles 
2 main types:
Titrations
1
2
Strong acid
titrated with
strong base
Species in solution
affect the pH
(in addition to H2O)
Weak base titrated
with a strong acid
Species in solution
affect the pH
(in addition to H2O)
Before
equivalence point
Use excess [H+]
ion to calculate pH
pH < 7
At equivalence
point
H+
H 2O
OH-
Salt hydrolysis
Excess [H+] ion
pH < 7
HB+
H+
pH = 7
Weak base +
conjugate acid
(buffer)
pH > 7
B, HB+
After equivalence
point
Use excess [OH-]
ion to calculate the
pH > 7
1. Titration of a strong acid with a strong base
moles H+ = moles OH-
At the equivalence point:
Example 16: Consider the titration of 40.0 mL of 0.200 M HBr by 0.100 m KOH. Calculate
the pH of the resulting solution when the following volumes of KOH have been added
a.
10.0 mL
b. 80.0 mL
c.
100.0 mL
Problem
mol H+
mol OH-
mol ion in XS
Total volume
(acid + base)
pH
a
b
c
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Calculation of the pH before the equivalence point:
Halfway to the equivalence point:
At the equivalence point:
After the equivalence point:
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2. Titration of a weak base with a strong acid
(see table)
Characteristics of Titration curves
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