Transcript No Slide Title

pH and Buffering
 Aim
 to know the logarithmic scale of pH
 to understand how weakly dissociating acids can buffer the pH
of an aqueous environment
 to know the importance of the carbonate - bicarbonate
buffering system
pH, The master variable
–
Consumed and produced
Enzyme/biological optima
Biological activity (enzyme activity)
–
4
5
6
7
pH
8
9
10
Dissociation of Water
OH H   K


H 2O
By Convention
therefore
w
 1014
[H2O] = 1
[OH-] [H+] = 10-14
So, if [H+] is known, [OH-] is also known
if [H+] = 10-5,
then [OH-] =10-9
Dealing in [H+] is cumbersome
Deal in pH (minus the log of the hydrogen ion concentration)
pH = - log[H+]
if [H+] = 0.1 M or 10-1 M, then pH = 1
pH is a log scale
[H+]
pH
10-7
7
10-7
10-6
6
10-8
10-5
5
10-9
10-3
3
10-11
10-11
11
10-3
[OH-]
Measurement of pH


pH meter and glass electrode
– quick
– easy
– accurate
– portable
Indicators
– titrations
phenolphthalein: pink  colourless below pH 8.3
methyl orange: red  yellow above pH 4.3
Weak acids and strong acids

An acid is substance produces H+ in water
H2SO4  2H+ + SO42-

A base produces OH- and/or accepts H+
NaOH  Na+ + OH-



A strong acid dissociates completely
1 mole HCl  1 mole H+ + 1 mole Cl1 mole H2SO4  2 mole H+ + 1 mole SO42A weak acid dissociates only partially
1 mole CH3COOH  0.0042 mole H+ + 0.0042 mole CH3COOThe concentration of hydrogen ions [H+] is therefore not always the same
as the concentration of the acid
Buffers


Chemicals which resist pH change
–
Acetic acid
Acetate
CH3COOH  CH3COO- + H+
–
Carbonate
Bicarbonate
CO32- + H+ 
HCO3-
Amphoteric chemicals
e.g. Proteins and amino acids
(have both +ve and -ve charged groups on the same molecule)
–

Buffering range of a buffering chemical is indicated
by its pKa
pKa is the pH at which the buffering chemical
is half dissociated:
for
HA  H+ + Awhen [HA] = [H+] = [A-], then pH = pKa
therefore buffering greatest when pH = pKa

Buffering capacity is given by the amount of
buffering chemical present
Carbonate-Bicarbonate Buffering
Major buffering in aquatic systems
CO2 (g)  CO2 (aq)
CO2 (aq) + H2O  H2CO3
(carbonic acid)
Difficult to distinguish between the two forms in water.
[H2CO3*] = [CO2] + [H2CO3]
H2CO3* is a proxy for “dissolved CO2 plus carbonic acid”
"Carbonic acid" dissociates to form bicarbonate
H2CO3*  HCO3- + H+
pKa = 6.3
Bicarbonate dissociates to form carbonate
HCO3-  CO32- + H+
pKa = 10.3
Carbonate can also come from the dissolution of carbonate
containing minerals:
MgCO3, Ca CO3
MgCO3  Mg2+ + CO32CaCO3 + CO2(aq) + H2O  Ca2+ + 2 HCO3-
Carbonate / bicarbonate system in a particular water depends on its
contact with air (CO2) and carbonate minerals.
For a closed system with no minerals or CO2 input, the species are:
1.0
HCO3-
H2CO3
CO32-
0.8
0.6
0.4
0.2
0
4
5
6
7
pKa
6.3
8
pH
9
10
pKa
10.3
11
12
References
Sawyer, McCarty, Parkin(1994)
Chemistry for Environmental Engineering
 Snoeyink, V.L. and Jenkins, D. (1980) Water
chemistry, Wiley.
 Stum, J and Morgan, J.J. (1981) Aquatic
Chemistry, Wiley Interscience.
 Loewenthal, R.E. and Marais, G.V.R (1976)
Carbonate Chemistry of Aquatic Systems,
Butterworths.
