Environmental Chemistry Lab
pH of Environmental Waters and Buffer
Capacity
Copyright © 2009 by DBS
Title
Objectives:
(i)
To learn proper use of the pH electrode by making pH
measurements
(i)
To determine the buffer capacity of natural waters
Introduction
• Natural waters contain a wide variety of both inorganic and
organic solutes
• Some of the more important solutes are inorganic and organic
acids and bases
• Organic: e.g. humic and tannic acids from OM
Role of pH in Water Quality
Brønsted-Lowry definition
• Acid is a proton donor
HCl + H2O → H3O+ + Cl• Base is a proton acceptor
NH3 + H2O → NH4+ + OHAcidic: H+ > OH-
Basic: OH- > H+
Where H3O+ = H+
Role of pH in Water Quality
Lewis definition
• Acid is an e- pair acceptor
H+ + :OH- → H2O
• Base is an e- pair donor
:NH3 + H2O → NH4+ + :OH-
pH Scale
pH = -log10 [H+]
[H+] = 10-pH
Or
pH = -log10 [H3O+]
Typically 0 – 14 (can go beyond this)
[H+] = [OH-] = 1.0 x 10-7 moles L-1
(pH = 7, neutral)
For each change of one pH unit [H+] changes x10
pH of Common Substances
Substance
pH
Battery acid
0.3
Lemon juice
2.4
Urine
4.8 - 7.5
Rainwater
5.5 - 6.0
Neutral water
7.00
Blood
7.35 - 7.45
Bleach
10.5
Ammonia
11.5
Typical pH Values
Reeve, 2002
Theory: Origin of Natural Acidity
• Some salts, such as ferric chloride (iron (III) chloride) affect
acidity:
Fe3+ (aq) + H2O(l) ⇌ [Fe(OH)]2+ (aq) + H+ (aq)
Fe3+ acts as Lewis acid
• Small highly charged metal ions produce acid solutions, nonmetal ions produce basic solutions
F- + H2O ⇌ HF + OH-
What is the pH of natural rain water?
Most acidity in natural waters is due to CO2
CO2: 370 ppm = 370 x 10-6 atm
CO2(g) + H2O ⇌ H2CO3 (aq)
Ka1 = 10-1.5
H2CO3(aq) ⇌ H+(aq) + HCO3-(aq)
Ka2 = 10-6.4
[H+] ~ [HCO3-]
Ka1 = [H2CO3]
[CO2]
Ka2 = [H+][HCO3-]
[H2CO3]
Ka1 = [H2CO3]
PCO2
Ka2 = [H+]2
[H2CO3]
10-1.5 = [H2CO3]
370 x 10-6
[H+]2 = Ka2[H2CO3] = 10-6.4 x 1 x 10-5 = 4 x 10-12
[H+] = 2 x 10-6
[H2CO3] = 1 x 10-5 M
pH = -log10[2 x 10-6] = 5.7
Theory: Buffers and Buffer Capacity
• Natural waters can neutralize both strong acids and bases:
HCO3-(aq) + HCl(aq) → H2O(l) + CO2 + Cl-(aq)
H2CO3(aq) + NaOH(aq) → H2O(l) + Na+(aq) + HCO3-(aq)
• Acts as a buffer – resists changes in pH
Theory: Buffers and Buffer Capacity
• Buffers contain a weak acid or base and a salt of the weak acid
or base (e.g. H2CO3:NaHCO3)
• Typical buffers change less than ± 0.1 unit on addition of acid or
base
Experimental
•
measure pH before and after the following:
1. Add 1 drop (0.050 mL) of concentrated HCl (12M) to one liter of DI
water (pH = 7)
[H3O+] = 0.050 mL x 1L / 1000 mL x 12 mol / L = 6.0 x 10-4 mol/L
pH = 3.22
2. Add 1 pellet of strong base ~ 0.10 g NaOH
[OH-] = (0.10 g x 1 mol / 40 g) / 1L = 0.0025 mol/L
pOH = 2.6, pH = 14 – pOH = 14 - 2.6 = 11.4
Henderson-Hasselbalch Equation
• calculating the pH of a buffer solution can be simplified by using an
equation derived from the Ka expression called the HendersonHasselbalch Equation
• the equation calculates the pH of a buffer from the Ka and initial
concentrations of the weak acid (HA) and salt of the conjugate base
(A-)
pH  pK a  log
[A - ]
[HA]
Deriving the Henderson-Hasselbalch Equation
[A - ][H3O ]
Ka 
HA 
 [HA] 

[H3O ]  K a  - 
 [A ] 
 [A [HA]
]  [HA]
[HA]
 




p]Ka
log
KK
a
 log    
log
 log[pH
H3OpH
pK
log

a 
a
- ][A
 - ]  
 [HA]
[A- ][A
       



pH  - log[H 3O ] pK a  -[HA]
log K a [A ]
-
 log

[A ]
 log
[HA]
Theory: Buffers and Buffer Capacity
• A buffer of specified pH can be made by adjusting the ratio
base: conjugate acid, [A-]/[HA]
• When[A-] = [HA], pH = pKa
• Buffer is said to be ‘centered’ at pKa
e.g. Acetic acid buffer Ka = 1.78 x 10-5 has center at pH = 4.74
Theory: Buffers and Buffer Capacity
• e.g. Acetic acid buffer Ka = 1.78 x 10-5 has center at pH = 4.74
• If 1 drop of 12 M HCl is added to 1 L of acetic acid buffer
containing 0.100 M acetic acid (HA) and 0.100 M sodium
acetate (A-)
What is the pH of a buffer that has 0.100 mol CH3COOH and 0.100 mol
NaCH3COO in 1.00 L that has 1 drop 12 M HCl (6.0 x 10-4 mol/L) added to it?
CH3COO- + H+ ⇌ CH3COOH
If the added chemical is a
base, write a reaction for
OH− with HA. If the added
chemical is an acid, write
a reaction for it with A−.
Construct a stoichiometry
table for the reaction
HA
A-
H+
mols Before
0.100
0.100
0
mols added
-
-
6.0 x 10-4
mols After
Part 1: stoichiometry
Original pH = 4.74
What is the pH of a buffer that has 0.100 mol CH3COOH and 0.100 mol
NaCH3COO in 1.00 L that has 1 drop 12 M HCl (6.0 x 10-4 mol/L) added to it?
CH3COO- + H+ ⇌ CH3COOH
HA
A-
H+
mols Before
0.100
0.100
0
mols added
-
-
6.0 x 10-4
0.100 + 6.0 x 10-4
0.100 - 6.0 x 10-4
0
mols After
Fill in the table, tracking
the number of moles for
each component
Part 1: stoichiometry
Original pH = 4.74
What is the pH of a buffer that has 0.100 mol CH3COOH and 0.100 mol
NaCH3COO in 1.00 L that has 1 drop 12 M HCl (6.0 x 10-4 mol/L) added to it?
CH3COO- + H+ ⇌ CH3COOH
HA
A-
H+
mols Before
0.100
0.100
0
mols added
-
-
6.0 x 10-4
0.101
0.099
0
mols After
Part 2: HH eqn.
Fill in the table, tracking
the number of moles for
each component
pH = pK + log [A-]/[HA]
= 4.74 + log (0.099/0.101) = 4.74
Theory: Buffers and Buffer Capacity
•
Buffer capacity is defined as the moles of strong acid (or base) needed to
change the pH of 1.00 L of a buffer by 1 pH unit
•
e.g. Acetic acid buffer system:
Center at 4.74, change by 1.0 pH units = 3.74:
pH = pKa + log [A-]/[HA]
3.74 = 4.74 + log [A-]/[HA]
log [A-]/[HA] = -1 or log [HA]/[A-] = 1, [HA]/[A-] = 10
[HA]/[A-] = (0.10 + x) / (0.10 – x) = 10
x = 0.082 mols
1 drop of 12 M HCl ~ 6.0 x 10-4 mols, therefore 137 drops = 0.082 mols
ISEs
•
•
•
Electrochemical potential - known
pH liquid inside the glass H+
sensitive membrane electrode vs.
unknown outside
Circuit is closed through the
solutions - internal and external and the pH meter
Electrodes generate a potential
(voltage) directly proportional to the
pH of the solution
– pH 7 potential is 0 V
– < 7 +ve V, > 7 –ve V
Analogy:
Battery where +ve is measuring
electrode, -ve is reference electrode
Flowing
• Internal KCl slowly flows to
the outside through the
junction (salt bridge)
• Must be refilled!
Gelled
• Slows leak but gets
contaminated
(shorter life-span)
Source: http://www.ph-meter.info
ISEs and Electrochemical Potential
Nernst equation
• Ecell = constant – 0.059 pH
(at 25 °C)
• Calibrated with buffer solutions of known pH
• Straight line plot of Ecell vs. pH
Activity Coefficient
•
Measured concentration (mol/L) of substance X (ax) is usually less than
the actual value
ax = γxCx
•
Where Cx = the actual concentration and γx = activity coefficient
(Note at low concentrations γx ~ 1, and ax ~ Cx)
•
Measure the pH of a 0.100 M solution of HCl (theoretical pH = 1…you
should obtain value > 1) why?
At 25 °C γH3O+ = 0.82
aHCl = (0.83)(0.10) = 0.083
pH = -log(0.083) = 1.08
Experimental
• Follow procedure p 55 Boehnke
• Use ultraclean glassware
• Must dilute 1 M HCl to make up 0.0100 M
• Take 10 mL using a 5 or 10 mL Eppendorf pipet dilute to 1 L to
make 0.01 M solution…stir well
• Graph and data analysis p 57
Questions and Further Thoughts
2.
Can you develop a general equation, based on concentrations of acid and conjugate base,
and pKa for calculating BC?
BC is the amount of strong acid or strong base required to change the pH of a buffer by 1.00
units
pH = log[A-]/[HA]) + pKa
(Henderson-Hasselbalch eqn.)
Max buffer capacity when [HA] = [A-]
Initial pH, pH0 = log(1) + pKa = pKa
Add NaOH to increase pH by 1.00
pH1 = log([A-]/[HA])1 + pKa
ΔpH = pH1 - pH0 = 1.00 = log([A-]/[HA])1
([A-]/[HA])1 = 10
(buffer center)
Questions and Further Thoughts
2.
Can you develop a general equation, based on concentrations of acid and conjugate base,
and pKa for calculating BC?
The greater the concentration of acid and base in the buffer the greater is its BC
Let C be the initial concentrations of acid (HA) and base (A-).
Addition of base: OH- + HA → H2O + Amol OH- added = BC
New conc.’s:
[HA]1 = C – BC
[A-] = C + BC
From previous slide ([A-]/[HA])1 = 10
Substituting in for [HA] and [A-] we have:
(C + BC) / (C – BC) = 10
BC = 9C/11
(for typical buffer C = 0.1 M, BC = 0.082 mol NaOH per liter of buffer)
BC is independent of pKa and proportional to concentration of acid or base
Text Books
•
•
•
•
•
•
•
•
Rump, H.H. (2000) Laboratory Manual for the Examination of Water, Waste Water and Soil.
Wiley-VCH.
Nollet, L.M. and Nollet, M.L. (2000) Handbook of Water Analysis. Marcel Dekker.
Keith, L.H. and Keith, K.H. (1996) Compilation of Epa's Sampling and Analysis Methods.
CRC Press.
Van der Leeden, F., Troise, F.L., and Todd, D.K. (1991) The Water Encyclopedia. Lewis
Publishers.
Kegley, S.E. and Andrews, J. (1998) The Chemistry of Water. University Science Books.
Narayanan, P. (2003) Analysis of environmental pollutants : principles and quantitative
methods. Taylor & Francis.
Reeve, R.N. (2002) Introduction to environmental analysis. Wiley.
Clesceri, L.S., Greenberg, A.E., and Eaton, A.D., eds. (1998) Standard Methods for the
Examination of Water and Wastewater, 20th Edition. Published by American Public Health
Association, American Water Works Association and Water Environment Federation.