Download Week #6: Acid - Base Properties of Aqueous Solutions

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
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.
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.
pH AND HYDRONIUM ION CONCENTRATION
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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+] = 0.001 M = 1x10–3 M
If pH = 4.0, then [H3O+] = 0.0001 M = 1x10–4 M
If pH = 5.0, then [H3O+] = 0.00001 M = 1x10–5 M
If pH = 6.0, then [H3O+] = 0.000001 M = 1x10–6 M
If pH = 7.0, then [H3O+] = 0.0000001 M = 1x10–7 M
If pH = 8.0, then [H3O+] = 0.00000001 M = 1x10–8 M
If pH = 9.0, then [H3O+] = 0.000000001 M = 1x10–9 M
If pH = 10.0, then [H3O+] = 0.0000000001 M = 1x10–10 M
If pH = 11.0, then [H3O+] = 0.00000000001 M = 1x10–11 M
If pH = 12.0, then [H3O+] = 0.000000000001 M = 1x10–12 M
If pH = 13.0, then [H3O+] = 0.0000000000001 M = 1x10–13 M
If pH = 14.0, then [H3O+] = 0.00000000000001 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.
THE pH OF COMMON SUBSTANCES
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.
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–])
MATHEMATICS, ACIDS, AND BASES
Taking the “p” of Equation 1 and rearranging.
Kw = 1.0x10–14 = [H3O+][OH–]
pKw = 14.00 = pH + pOH
(6) pH = 14.00 – pOH
(7) pOH = 14.00 – 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)
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+]) = 14.00 – pOH
(5 and 7) pOH = –log10[OH–] = log10(1 / [OH–]) = 14.00 – pH
• Complete this table.
[H3O+]
[OH–]
2.9x10–11 M
3.5x10–4 M
2.0x10–5 M
5.0x10–10 M
pH
3.46
4.70
4.2x10–6 M
2.4x10–9 M
5.38
3.8x10–7 M
2.6x10–8 M
6.42
pOH
10.54
9.30
8.62
7.58
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.
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Acids
Bases
HCl
HBr
HI
HClO4
HNO3
H2SO4a
NaOH
KOH
RbOH
CsOH
Ca(OH)2
Sr(OH)2
Ba(OH)2
H2SO4 ionizes in 2 steps. The first ionization goes to completion. The
second ionization does not go to completion.
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.
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–).
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.)
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–] = 0.001 M and [HA] = 0.001 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.
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.
MEASURING pH BY GLASS ELECTRODE
MEASURING pH BY PAPER
• A wide variety of dyes are used to make pH paper. These
dyes change color with pH.
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.
SOURCES
• Christian, G.D. 1986. Analytical Chemistry, 3rd ed. New York, NY: John
Wiley & Sons, Inc.
• Harris, D.C. 1999. Quantitative Chemical Analysis, 5th ed. New York, NY:
W.H. Freeman Company.
• Hill, J.W., D.K. Kolb. 2007. Chemistry for Changing Times, 11th ed.
Upper Saddle River, NJ: Pearson Prentice Hall.
• McMurry, J., R.C. Fay. 2004. Chemistry, 4th ed. Upper Saddle River, NJ:
Prentice Hall.
• Park, J.L. 2004. ChemTeam: Photo Gallery Menu. Available:
http://dbhs.wvusd.k12.ca.us/webdocs/Gallery/GalleryMenu.html
[accessed 9 October 2006].
• Petrucci, R.H. 1985. General Chemistry Principles and Modern
Applications, 4th ed. New York, NY: Macmillan Publishing Company.