Acids, Bases and the Common Ion Effect HF + H 2 O H 3 O + + F - H + + Cl - Consider the following acid equilibrium of a weak acid: K a = [H 3 O + ][F - ] [HF] HCl (aq) What happens when we add some strong acid to the mixture? HCl completely dissociates, adding free H 3 O + to solution. This is a common ion to the weak acid equilibrium. By LeChatelier’s principle, we predict the HF dissociation should be driven left, suppressing the dissociation. More quantitative a) Determine [F - ] in a solution of 0.500M HF. b) Determine [F - ] in a solution of 0.500M HF and 0.10M HCl. K HF = 6.6x10 -4 . [ ] [ ] H F K K K C HF HF HF HF + − = = − + + 2 4 2 a) A solution of a weak acid. Let’s use the quadratic. [F - ] = 0.0178M
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Acids, Bases and the Common Ion Effect - … Acids, Bases and the Common Ion Effect HF + H 2O H 3O+ + F-H+ + Cl-Consider the following acid equilibrium of a weak acid: K a = [H 3O+][F-]
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Acids, Bases and the Common Ion Effect
HF + H2O H3O+ + F-
H+
+Cl-
Consider the following acid equilibrium of a weak acid:Ka = [H3O+][F-]
[HF]
HCl(aq)
What happens when we add some strong acid to the mixture?
HCl completely dissociates, adding free H3O+ to solution. This is a common ion to the weak acid equilibrium.
By LeChatelier’sprinciple, we predict the HF dissociation should be driven left, suppressing the dissociation.
More quantitativea) Determine [F-] in a solution of 0.500M HF. b) Determine [F-] in a solution of 0.500M HF and 0.10M HCl.KHF = 6.6x10-4.
[ ] [ ]H F K K K CHF HF HF HF+ −= =− + +2 4
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a) A solution of a weak acid. Let’s use the quadratic.
[F-] = 0.0178M
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Continued
Initialweak acid 0.50 0 0strong acid 0 0.10 0
Change -x + x + xEquilib 0.50-x 0.10 +x x
HF + H2O H3O+ + F-
b) Use ICE table. Let HF and HCl dissociate separately.
Ka = [H3O+][F-] [HF]
Ka = (0.10 +x)(x) / (0.50-x)
Assumption: ionization of HF is suppressed. Assume x<<0.1; x<<0.5
Ka ~ (0.10)(x) / (0.50)
ContinuedRearrange: x = [F-] = (0.50M)Ka/ 0.10M = 0.0033 M
Comparison a) no common ion, [F-] = 0.0178 Mb) common ion, [F-] = 0.0033 M
5 X less!
This is one example of the common ion effect.
When a strong acid supplies the common ion, H3O+, the ionization of a weak acid is suppressed.
When a strong base supplies the common ion, OH-, the ionization of a weak base is suppressed.
Acid/Base buffer – a system that resists changes to pH caused by addition of excess base or acid.
Chemical buffer – a chemical system that resists change
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Acid Base Buffer SystemsA mixture of a weak acid and it’s conjugate base
O
O
H O
O
Benzoic acid benzoate
CH3
OHO
acetic acid acetate
CH3
O O
A mixture of a weak base and it’s conjugate acidNCH3
CH3 CH3 NH
+
CH3
CH3 CH3
trimethylamine trimethylammonium
Add as sodium salts for example
Add as chloride salt for example
An acid/base buffer system
K [H ][A ][HA]a =+ −
What happens if we add acid and conjugate base in equal concentrations? i.e. equal moles of HA and NaA
HA + H2O H3O+ + A-
Initialweak acid 0.10 0 0conj base 0 0 0.10
Change -x +x + xEquilib 0.10-x +x 0.10+x
K [H ][0.10 + x][0.10- x]a =+
But due to presence of A-, ionization of HA is suppressed more than usual. x<<0.10
[H+] = Ka
pH = pKa!
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Henderson-Hasselbalch Equation
K [H ][A ][HA]a =+ −
Solve for [H+]
-log [H+] = -log Ka –log [HA] + log [A-]
[H ] K [HA][A ]a+
−=
Take –log of both sides
-log [H+] = -log Ka + log [A-]/[HA]
pH pK log [A ][HA]a
-= + pH pK log
CCa
AHA
-= +
From previous ICE table, we see: [HA] = CHA – x[A-] = CA
- + x But in general, CHA = CA
-. If x<<CHA
Buffer ratiosFor correct buffering, there should be significant amounts of HA and A-.Just as importantly, the ratio should be such that:
0.10 CC
10.0HA
A≤ ≤
−
pH pK 1buffer a= ±
Substituting this constraint into the H-H equation
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How to prepare a buffer1. What pHbuffer do you require?
2. Find an acid/conj base system from tables that has pKa within 1 unit of the required pH.
3. Use the H-H equation to determine the exact ratio of the acid and the base you will need.
4. What buffer capacity do you need…0.1M, 0.01M, 0.001M? Answer depends on how much acid or base could potentially impact the system that you are trying to buffer.
5. Prepare solutions taking into account 3) and 4).
ExamplePrepare 1L of buffer with pH = 3.5, with a total concentration, [A-] + [HA] = 0.01M.
Since we are preparing 1L, we need to add 0.003 moles of citric acid and 0.007 moles of monosodium citrate.
massHA = 0.003 mol x 192.1 g/mol = 0.576g
massNaA = 0.007 mol x 214.1 g/mol = 1.50g
Dissolve to make 1L of solution.OR
Dissolve 0.01mol x 192.1g/mol = 1.92g of citric acid in about 1L of water, and neutralize with NaOH until pH = 3.5! (Total volume must be 1L.)
Examples of Buffers in NaturepH of blood is 7.40+0.05. pH is maintained by a series of buffer systems, H2CO3/HCO3
-, phosphate and proteins. pH regulation in the body is critical since functioning of enzymes is highly pH dependent.
Alkalosis – raising of pH results from hyperventilation, or exposure to high elevations (altitude sickness).
Acidosis – lowering of pH in blood by organ failure, diabetes or long term protein diet.
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What causes alkalosis at high elevation?…a
multiple equilibrium hypothesis.
The bicarbonatebuffer is essential
for controllingblood pH
CO2(aq) + H2O H2CO3(aq)
H2CO3(aq) + H2O H3O+ + HCO3-
CO2(g)
[HCO ][H CO ]
~ 1032 3
−
pKa = 6.4+
OH-
2 H2O
pH increases !
in the lungs
in the blood
At high elevation,P drops,
hyperventilation!
Acid Base TitrationsAcid base titrations are examples of volumetric techniques used to analyze the quantity of acid or base in an unknown sample.
Acid + Base H2O + salt
This is done by detecting the point at which we have added an equal number of equivalents of base to the acid. This is the equivalence point. For neutralization of an unknown monoprotic acid (A) with a base (B), we have at equivalence:
moles of acid = moles of base CAVA = CBVB
We detect the equivalence point with a pH meter or by identifying the end-point with an acid base indicator.
A
equivBA V
VCC =
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Acid Base Titrations - some terminologyequivalence point: the point at which
moles of acid = moles of baseend-point: the experimental approximation of
the equivalence pointtitration error: the difference between the
equivalence point and the end-pointtitrant: the solution that is added in a
measured quantitystandard solution: a solution of known
concentration
IndicatorsAcid-base indicators are highly coloured weak acids or
bases.HIndic Indic- + H+
colour 1 colour 2Colour transition occurs for 0.1< [Indic-]/[HIndic]<10
They may have more than one colour transition. Example. Thymol blue
Red Yellow BlueOne of the forms may be colourless - phenolphthalein
(colourless to pink)
1pK[HIn]
][InlogpKpH inda
indatransition ±=+=
−
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IndicatorsSelection of an indicator.
The colour transition range for an indicator is pKa + 1. Choose an indicator that has a pKa close to the pH at the equivalence point.
phenolphthalein
Indicator Examples
bromthymol blue
methyl red
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Acid Base Titrations16_11
2
4
6
8
10
14
12pH
0 5 10 15 20 25 30 35 40 45 50
Volume of NaOH added (mL)
Volume NaOHadded (mL)
05
1015202122232425262728293035404550
1.001.181.371.601.952.062.202.382.697.00
11.2911.5911.7511.8711.9612.2212.3612.4612.52
pH
Pink
ColorlessPhenolphthalein
Equivalence point
Bromcresol green
Blue
Yellow
Titration of strong acid with strong base
Low initial pH, large pH change at equivalence point
ExampleWhat is the pH at 0mL, 10mL, 25.0mL and 35mL in titration of 25mL of 0.100 M strong acid (HCl) with 0.100M NaOH?