This document provides information about buffer solutions including definitions, examples, calculations, and the Henderson-Hasselbalch equation. It discusses how buffers resist changes in pH when small amounts of strong acid or base are added. Example calculations are provided to determine the pH of buffer solutions using given concentrations of weak acids/bases and their conjugate species.

1. IB Chemistry Power Points
Topic 18
Acids and Bases
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Buffer
Solutions

2. Buffer Solutions
 DEFINITION: A buffer solution contains a weak acid mixed
with its conjugate base (or weak base and conjugate acid)
 Buffers resist changes in pH when a small amount of a
strong acid or base is added to it.
HA ∏ H+ + A-

3.  If a small amount of a strong acid (H+) is added eqm
shifts to the left as [H+] increases so system adjusts to
increase [HA] and reduce [H+] again.
HA ∏ H+ + A-

4.  A small amount of a strong base will react with H+ to
form H2O and eqm will shift to the right to increase
[H+] again.
HA ∏ H+ + A-

5. Making Buffer Solutions
 An example of a weak acid is ethanoic acid. This could be
mixed with sodium ethanoate which will provide ethanoate
ions (conjugate base).
CH3COOH(aq) ∏ CH3COO-(aq)+H+(aq)
 An example of a weak base is ammonia. This could be mixed
with ammonium chloride to provide ammonium ions
(conjugate acid).
NH3(aq) + H2O(aq) ∏ NH4+(aq) +OH-(aq)
In order for a buffer to work well the concentration of the acid/base and
its salt must be much higher than the strong acid/base added.

6. Optimum Buffer
 A buffer is most effective when the concentration
of weak acid and its salt (the conjugate base) are
equal and the pH is equal to pKa.
 In practice it will work reasonably well with similar
concentrations and the effective buffer range of
any weak acid/ base is pKa 1.

7. Blood has buffering capacity
 Blood must maintain a pH of 7.4 so its enzymes can
work. If 0.01 mol of H+ or OH- is added to blood it
only changes pH by 0.1 unit.
 The eqm is:
CO2(aq) + H2O(l) ) ∏ H+(aq) + HCO3-(aq)

8. Buffer Calculation #1
Calculate the pH of a 1.00 dm3 buffer solution made by
dissolving 0.50 mol of sodium ethanoate into a 0.075 mol dm-3
ethanoic acid solution.
1. Identify the weak acid / conjugate base OR weak base /
conjugate acid pair. Determine their concentrations.

9. Buffer Calculation #1
Calculate the pH of a 1.00 dm3 buffer solution made by
dissolving 0.50 mol of sodium ethanoate into a 0.075 mol dm-3
ethanoic acid solution.
1. Identify the weak acid / conjugate base OR weak base /
conjugate acid pair. Determine their concentrations.
weak acid [CH3COOH] = 0.075 mol dm-3
conjugate base [CH3COO-] = 0.50 mol dm-3

10. Buffer Calculation #1
Calculate the pH of a 1.00 dm3 buffer solution made by
dissolving 0.50 mol of sodium ethanoate into a 0.075 mol dm-3
ethanoic acid solution.
2. Write the equilibrium equation and expression

11. Buffer Calculation #1
Calculate the pH of a 1.00 dm3 buffer solution made by
dissolving 0.50 mol of sodium ethanoate into a 0.075 mol dm-3
ethanoic acid solution.
2. Write the equilibrium equation and expression
CH3COOH ∏ H+ + CH3COO-
[H + ][CH 3COO - ] - pK
Ka = = 10 a
[CH 3COOH ]

12. Buffer Calculation #1
Calculate the pH of a 1.00 dm3 buffer solution made by
dissolving 0.50 mol of sodium ethanoate into a 0.075 mol dm-3
ethanoic acid solution.
ICE table for clarity
CH3COOH ∏ H+ + CH3COO-
I C E
HA 0.075 -x 0.075 - x
H+ 0 +x x
A- 0.50 +x 0.50 + x

13. Buffer Calculation #1
Calculate the pH of a 1.00 dm3 buffer solution made by
dissolving 0.50 mol of sodium ethanoate into a 0.075 mol dm-3
ethanoic acid solution.
3. In this case, assume the equilibrium concentrations of the
weak acid and the salt anion are assumed to be the same as
the given information (very little change when equilibrium is
established.

14. Buffer Calculation #1
Calculate the pH of a 1.00 dm3 buffer solution made by
dissolving 0.50 mol of sodium ethanoate into a 0.075 mol dm-3
ethanoic acid solution.
3. In this case, assume the equilibrium concentrations of the
weak acid and the salt anion are assumed to be the same as
the given information (very little change when equilibrium is
established.
[H + ][0.50]
Ka = = 10-4.76
[0.075]
æ 0.075´10-4.76 ö
pH = -log[H + ] = -log ç ÷
è 0.50 ø

15. Buffer Calculation #1
Calculate the pH of a 1.00 dm3 buffer solution made by
dissolving 0.50 mol of sodium ethanoate into a 0.075 mol dm-3
ethanoic acid solution.
pH = 5.6

16. Henderson-Hasselbalch Equation

17. Henderson-Hasselbalch Equation
æ [A- ] ö
pH = pK a + log ç ÷
è [HA] ø
æ 0.50 ö
pH = 4.76 + log ç ÷
è 0.075 ø
pH = 5.58

18. Buffer Calculation #2
Calculate the mass of ammonium chloride that would need to be
dissolved into 1.00 dm3 of 0.100 mol dm-3 NH3 solution to create a buffer
with a pH of 9.00. Assume no change in overall volume.
1. Identify the weak acid / conjugate base OR weak base /
conjugate acid pair. Determine their concentrations.

19. Buffer Calculation #2
Calculate the mass of ammonium chloride that would need to be
dissolved into 1.00 dm3 of 0.100 mol dm-3 NH3 solution to create a buffer
with a pH of 9.00. Assume no change in overall volume.
1. Identify the weak acid / conjugate base OR weak base /
conjugate acid pair. Determine their concentrations.
weak base [NH3] = 0.100 mol dm-3
conjugate acid [NH4+] = ?

20. Buffer Calculation #2
Calculate the mass of ammonium chloride that would need to be
dissolved into 1.00 dm3 of 0.100 mol dm-3 NH3 solution to create a buffer
with a pH of 9.00. Assume no change in overall volume.
2. Write the equilibrium equation and expression

21. Buffer Calculation #2
Calculate the mass of ammonium chloride that would need to be
dissolved into 1.00 dm3 of 0.100 mol dm-3 NH3 solution to create a buffer
with a pH of 9.00. Assume no change in overall volume.
2. Write the equilibrium equation and expression
NH3+ H2O ∏ NH4+ + OH-
[NH 4+ ][OH - ] - pKb
Kb = = 10
[NH 3 ]

22. Buffer Calculation #2
Calculate the mass of ammonium chloride that would need to be
dissolved into 1.00 dm3 of 0.100 mol dm-3 NH3 solution to create a buffer
with a pH of 9.00. Assume no change in overall volume.
3. In this case, assume the equilibrium concentrations of the
weak base is the same as the given information (very little
change when equilibrium is established. Calculate [OH-]
from desired pH
(14 pH )
[ NH 4 ][10 ] 4.75
kb 10
[0.100]
0.100 10 4.75
[ NH 4 ] 0.1778 mol dm-3
10 (14 9)

23. Buffer Calculation #2
Calculate the mass of ammonium chloride that would need to be
dissolved into 1.00 dm3 of 0.100 mol dm-3 NH3 solution to create a buffer
with a pH of 9.00. Assume no change in overall volume.
3. In this case, assume the equilibrium concentrations of the
weak base is the same as the given information (very little
change when equilibrium is established. Calculate [OH-]
from desired pH

24. Buffer Calculation #2
Calculate the mass of ammonium chloride that would need to be
dissolved into 1.00 dm3 of 0.100 mol dm-3 NH3 solution to create a buffer
with a pH of 9.00. Assume no change in overall volume.
mass = 9.50 g

25. Buffer Calculation #3
A buffer can also be made by mixing excess weak acid/base with a lesser
amount of strong base/acid. For example, calculate the pH of a buffer
formed when 25 cm3 of 0.075 mol dm-3 HCl is added to 40 cm3 of a 0.150
mol dm-3 ammonia solution.
1. Do the stoichiometry to determine what the concentrations
are AFTER neutralization.
-

26. Buffer Calculation #3
A buffer can also be made by mixing excess weak acid/base with a lesser
amount of strong base/acid. For example, calculate the pH of a buffer
formed when 25 cm3 of 0.075 mol dm-3 HCl is added to 40 cm3 of a 0.150
mol dm-3 ammonia solution.
1. Do the stoichiometry to determine what the concentrations
are AFTER neutralization.
HCl + NH3  NH4+ + Cl-

27. Buffer Calculation #3
A buffer can also be made by mixing excess weak acid/base with a lesser
amount of strong base/acid. For example, calculate the pH of a buffer
formed when 25 cm3 of 0.075 mol dm-3 HCl is added to 40 cm3 of a 0.150
mol dm-3 ammonia solution.
1. Do the stoichiometry to determine what the concentrations
are AFTER neutralization.
HCl + NH3  NH4+ + Cl-
(0.025 0.075)
NH 4 HCl limiting (why?)
0.065
NH 4 0.02885
(0.040 0.150) (0.025 0.075)
NH 3
0.065
NH 3 0.06346

28. Buffer Calculation #3
A buffer can also be made by mixing excess weak acid/base with a lesser
amount of strong base/acid. For example, calculate the pH of a buffer
formed when 25 cm3 of 0.075 mol dm-3 HCl is added to 40 cm3 of a 0.150
mol dm-3 ammonia solution.
2. Write the equilibrium equation and expression

29. Buffer Calculation #3
A buffer can also be made by mixing excess weak acid/base with a lesser
amount of strong base/acid. For example, calculate the pH of a buffer
formed when 25 cm3 of 0.075 mol dm-3 HCl is added to 40 cm3 of a 0.150
mol dm-3 ammonia solution.
2. Write the equilibrium equation and expression
NH3+ H2O ∏ NH4+ + OH-
[NH 4+ ][OH - ] - pKb
Kb = = 10
[NH 3 ]

30. Buffer Calculation #3
A buffer can also be made by mixing excess weak acid/base with a lesser
amount of strong base/acid. For example, calculate the pH of a buffer
formed when 25 cm3 of 0.075 mol dm-3 HCl is added to 40 cm3 of a 0.150
mol dm-3 ammonia solution.
3. In this case, assume the equilibrium concentrations of the
weak base and cation are the same as determined by
stoichiometry (very little change when equilibrium is
established.

31. Buffer Calculation #3
A buffer can also be made by mixing excess weak acid/base with a lesser
amount of strong base/acid. For example, calculate the pH of a buffer
formed when 25 cm3 of 0.075 mol dm-3 HCl is added to 40 cm3 of a 0.150
mol dm-3 ammonia solution.
3. In this case, assume the equilibrium concentrations of the
weak base and cation are the same as determined by
stoichiometry (very little change when equilibrium is
established.
[0.02885][OH - ]
Kb = = 10-4.75
[0.06346]
æ 0.06346 ´10-4.75 ö
pOH = -log ç ÷
è 0.02885 ø

32. Buffer Calculation #3
A buffer can also be made by mixing excess weak acid/base with a lesser
amount of strong base/acid. For example, calculate the pH of a buffer
formed when 25 cm3 of 0.075 mol dm-3 HCl is added to 40 cm3 of a 0.150
mol dm-3 ammonia solution.
pH= 9.6