1. CH 104: ACID-BASE PROPERTIES OF AQUEOUS SOLUTIONS
In the Arrhenius theory an acid produces H+ in aqueous solution and a base produces OH– in aqueous solution.
The more general Brønsted-Lowry theory defines an acid as a H+ donor and a base as a H+ accepter.
Svante Arrhenius Johannes Brønsted Thomas Lowry

2. HYDROGEN, HYDRONIUM, AND HYDROXIDE IONS
The symbol H+(aq) is convenient to use; however, it is not accurate. Hydrogen ion (H+) is a proton without an electron. It is hydrated in water and exists as hydronium ion (H3O+(aq)).
The self-ionization of water:
Preferred: 2H2O(l) = H3O+(aq) + OH–(aq)
Accepted: H2O(l) = H+(aq) + OH–(aq)
H3O+ (or H+) is acid.
Hydroxide ion (OH–) is base.

3. pH AND HYDRONIUM ION CONCENTRATION
The pH scale measures acidity. It typically ranges from 0 to 14. The acidity is neutral at pH 7. Values less than pH 7 are increasingly acidic. Values greater than pH 7 are increasingly basic.

4. pH AND HYDRONIUM ION CONCENTRATION
If pH = 0.0, then [H3O+] = 1 M = 1x10–0 M
If pH = 1.0, then [H3O+] = 0.1 M = 1x10–1 M
If pH = 2.0, then [H3O+] = 0.01 M = 1x10–2 M
If pH = 3.0, then [H3O+] = M = 1x10–3 M
If pH = 4.0, then [H3O+] = M = 1x10–4 M
If pH = 5.0, then [H3O+] = M = 1x10–5 M
If pH = 6.0, then [H3O+] = M = 1x10–6 M
If pH = 7.0, then [H3O+] = M = 1x10–7 M
If pH = 8.0, then [H3O+] = M = 1x10–8 M
If pH = 9.0, then [H3O+] = M = 1x10–9 M
If pH = 10.0, then [H3O+] = M = 1x10–10 M
If pH = 11.0, then [H3O+] = M = 1x10–11 M
If pH = 12.0, then [H3O+] = M = 1x10–12 M
If pH = 13.0, then [H3O+] = M = 1x10–13 M
If pH = 14.0, then [H3O+] = M = 1x10–14 M
pH 2.0, 0.01 M, and 1x10–2 M each have 1 significant figure. For pH and other logarithms, the numbers to the right of the decimal are significant. The numbers to the left of the decimal are NOT significant. The 0 in pH 2.0 is significant. The 2 in pH 2.0 is NOT significant, it defines the 2 in 1x10–2 M.

5. THE pH OF COMMON SUBSTANCES

6. THE pH OF ACID PRECIPITATION
This is the distribution of precipitation pH for North America.
The combustion of sulfur-containing coal from Midwestern power plants is a major cause of acid precipitation.
2S + 3O2 + 2H2O → 2H2SO4
The prevailing winds carry this acid from the Midwest to the East.

7. MATHEMATICS, ACIDS, AND BASES
All concentrations are in moles per liter.
(1) The Ion Product of Water = Kw = 1.0x10–14 = [H3O+][OH–]
Rearranging Equation 1.
(2) [H3O+] = (1.0x10–14) / [OH–]
(3) [OH–] = (1.0x10–14) / [H3O+]
“p” is the negative base 10 logarithm.
(4) pH = –log10[H3O+] = log10(1 / [H3O+])
(5) pOH = –log10[OH–] = log10(1 / [OH–])

8. MATHEMATICS, ACIDS, AND BASES
Taking the “p” of Equation 1 and rearranging.
Kw = 1.0x10–14 = [H3O+][OH–]
pKw = = pH + pOH
(6) pH = – pOH
(7) pOH = – pH
Taking the antilogarithm of Equation 4.
pH = –log10[H3O+]
(8) [H3O+] = 10(–pH)
Taking the antilogarithm of Equation 5 and inserting Equation 7.
pOH = –log10[OH–]
[OH–] = 10(–pOH)
(9) [OH–] = 10(pH – 14.00)

9. MATHEMATICS, ACIDS, AND BASES
In summary,
(1) The Ion Product of Water = Kw = 1.0x10–14 = [H3O+][OH–]
(2 and 8) [H3O+] = (1.0x10–14) / [OH–] = 10(–pH) = 10(pOH – 14.00)
(3 and 9) [OH–] = (1.0x10–14) / [H3O+] = 10(–pOH) = 10(pH – 14.00)
(4 and 6) pH = –log10[H3O+] = log10(1 / [H3O+]) = – pOH
(5 and 7) pOH = –log10[OH–] = log10(1 / [OH–]) = – pH
Complete this table.
[H3O+]
[OH–] pH
pOH
3.5x10–4 M
5.0x10–10 M
5.38
7.58
2.9x10–11 M
3.46
10.54
2.0x10–5 M
4.70
9.30
4.2x10–6 M
2.4x10–9 M
8.62
3.8x10–7 M
2.6x10–8 M
6.42

10. STRONG ACIDS AND STRONG BASES
A strong acid or a strong base in distilled water will almost completely ionize.
Strong acid: HCl(g) + H2O(l) → H3O+(aq) + Cl–(aq)
Strong base: NaOH(s) + H2O(l) → Na+(aq) + OH–(aq)
Common strong acids and strong bases.
a H2SO4 ionizes in 2 steps. The first ionization goes to completion. The second ionization does not go to completion.
Acids
Bases
HCl
HBr HI
HClO4
HNO3
H2SO4a
NaOH
KOH
RbOH
CsOH
Ca(OH)2
Sr(OH)2
Ba(OH)2

11. WEAK ACIDS AND WEAK BASES
Most acids and most bases are weak. That is, most acids and most bases in distilled water do not completely ionize.
Weak acid: HC2H3O2(l) + H2O(l) = H3O+(aq) + C2H3O2–(aq)
A weak acid (HC2H3O2) is in equilibrium with its conjugate base (C2H3O2–).
Ionization Constant = Ka = [H3O+][C2H3O2–] / [HC2H3O2] = 1.74x10–5 at 25° C.
Weak base: NH3(aq) + H2O(l) = NH4+(aq) + OH–(aq)
A weak base (NH3) is in equilibrium with its conjugate acid (NH4+).
Ionization Constant = Kb = [NH4+][OH–] / [NH3] = 1.74x10–5 at 25° C.

12. BUFFERS AND THE HENDERSON-HASSELBALCH EQUATION
A buffer is a solution that resists drastic changes in pH when an acid or base is added.
Furthermore, a buffer resists drastic changes in pH when it is diluted.
Buffers are used to control pH. For example, human blood is buffered at pH 7.4±0.1. The ability of blood to carry oxygen depends on the pH being within this range.
A buffer is a mixture of a weak acid and a salt of its conjugate base, or a weak base and a salt of its conjugate acid.
For example, a mixture of acetic acid (HC2H3O2) and sodium acetate (NaC2H3O2) is a common buffer.
What is the conjugate base of acetic acid?
Acetate (C2H3O2–).

13. BUFFERS AND THE HENDERSON-HASSELBALCH EQUATION
This Henderson-Hasselbalch equation shows how a buffer of a weak acid and its salt resists drastic changes in pH.
(A similar Henderson-Hasselbalch equation would show how a buffer of a weak base and its salt resists drastic changes in pH.)

14. BUFFERS AND THE HENDERSON-HASSELBALCH EQUATION
Therefore, the buffering of a weak acid and its salt depends on the relative concentrations of its conjugate base (A–) and its unionized acid (HA).
If a small amount of strong acid is added, it will combine with A– to make HA. If the change in [A–]/[HA] is small, the change in pH will be small.
Conversely, if a small amount of strong base is added, it will react with HA to make A–. If the change in [A–]/[HA] is small, the change in pH will be small.
What has a larger buffering capacity (a larger resistance to changes in pH)? A solution with [A–] = M and [HA] = M. Or a solution with [A–] = 1 M and [HA] = 1 M.
The solution with [A–] = 1 M and [HA] = 1 M has a larger buffering capacity.

15. MEASURING pH BY GLASS ELECTRODE
The glass membrane of a pH electrode is made out of silicate groups with exchangeable hydrogen ions, Si-O-H+. These Si-O– groups are attached to the glass membrane. These H+ ions are in equilibrium with the surface of the glass membrane and the sample.
Glass membrane-Si-O-H+(s) =
Glass membrane-Si-O–(s) + H+(aq)
If the number of H+ ions in the sample is large, then the number of H+ ions on the glass membrane is large and the electrode voltage is small. Conversely, if the number of H+ ions in the sample is small, then the number of H+ ions on the glass membrane is small and the electrode voltage is large. This voltage is converted to a pH value.

16. MEASURING pH BY GLASS ELECTRODE

17. MEASURING pH BY PAPER
A wide variety of dyes are used to make pH paper. These dyes change color with pH.

18. SAFETY
Give at least 1 safety concern for the following procedure.
Using acids and bases.
These are irritants. Wear your goggles at all times. Immediately clean all spills. If you do get either of these in your eye, immediately flush with water.
Your laboratory manual has an extensive list of safety procedures. Read and understand this section.
Ask your instructor if you ever have any questions about safety.

19. SOURCES
Christian, G.D Analytical Chemistry, 3rd ed. New York, NY: John Wiley & Sons, Inc.
Harris, D.C Quantitative Chemical Analysis, 5th ed. New York, NY: W.H. Freeman Company.
Hill, J.W., D.K. Kolb Chemistry for Changing Times, 11th ed. Upper Saddle River, NJ: Pearson Prentice Hall.
McMurry, J., R.C. Fay Chemistry, 4th ed. Upper Saddle River, NJ: Prentice Hall.
Park, J.L ChemTeam: Photo Gallery Menu. Available: [accessed 9 October 2006].
Petrucci, R.H General Chemistry Principles and Modern Applications, 4th ed. New York, NY: Macmillan Publishing Company.