Chemical equilibria
Buffer solutions
Definition
The buffer solutions have the ability to resist changes in pH when smaller amounts of acid or base is added.
Importance
They are applied in the chemical practice (reactions performed at an optimal pH), in biochemistry and in physiology.
Composition of the buffer solutions
They contain a weak acid and its salt (or its conjugate base) e.g. CH3COOH and CH3COONa.
They can be prepared from a weak base and its salt (or conjugate acid) e.g. NH3 solution and NH4Cl.
[CH3COOH] = cacid
Acidic buffers I
CH3COOH(aq) CH3COO- + H+
K = [ ]
] ][
[
3 3
COOH CH
H COO
CH = 1.8 · 10-5 (at 25 0C)
The buffer is composed of acetic acid (cacid), it ionizes only slightly.
In order to prepare a buffer, sodium acetate is added to the acetic acid solution. The sodium acetate (csalt), dissociates completely.
The acetate concentration was increased, the undissociated acid concentration was also increased.
[CH3COO-] =csalt
Acidic buffers III
Ka = csalt ·
cacid
H ]
[ [H+] =
salt acid a
c c K
[H+] = Ka ·
] [
] [
A
HA
–log [H+] = –log Ka + log
] [
] [
HA A
pH = pKa + log
] [
] [
HA A
According to the law of mass action:
A general expression for the acidic buffers
The equation above is the Henderson-Hasselbalch equation describing the pH value of an acidic buffer system, it depends only on the ratio of the components.
Basic buffers
OH
-+ NH
4+NH
3(aq) + H
2O(l)
K
b=
] [
] ][
[
3 4
NH OH NH
[OH-] = Kb
] [
] [
4 3
NH
NH
E.g. it consists of NH3 solution and NH4Cl
pOH = pK
b+ log
] [
] [
B HB
Henderson-Hasselbalch equation for basic buffers
The features of the buffer solutions I
The pH value of a certain buffer is determined by salt/weak acid ratio (acidic buffer)
or by the salt/weak base ratio (basic buffer).
At the different buffers the pH depends on the pKa or pKb values of the acid or base.
The buffer works efficiently if the salt/weak acid ratio is in the following range:
0,1< salt/weak acid <10
Based on this fact we can prepare a good buffer from a weak acid and a salt pair in the range: pH = pKs +/-1
The features of the buffer solutions II
Buffer capacity
Definition: The amount of strong acid (e.g. HCl) or strong base (e.g. NaOH) in moles that can be added to one litre of buffer in order to change the pH of the buffer by one unit.
Acid capacity of a buffer- base capacity of a buffer
The features of the buffer solutions III
The effect of the dilution on the pH of the buffer (pH is not influenced by dilution)
and on the buffer capacity (the buffer capacity depends on the dilution).
The function of the buffers I
Calculation
1. Calculate the pH value of the buffer made by mixing 0.5 L 2 molar acetic acid and 0.5 L 2 molar sodium acetate solutions.
(Ka=1.8x10-5)
After mixing the solutions [CH3COOH] = [CH3COO-] = 1 mol/L
pH = pK
s+ log
] [
] [
HA A
pH = 4.74 + log
] 1 [
] 1
[ = 4.74
The function of the buffers II
Calculation
2. Calculate the change in pH if 0.4 g NaOH (strong base) is dissolved in one litre of the buffer solution previously mentioned.
The NaOH neutralizes a certain amount (mole) of acid producing the same amount (mole) of salt.
pH = pK + lg
x c
x c
acid salt
= 4.74 + lg
01 . 0 1
01 . 0 1
= 4.74 + lg
99 . 0
01 .
1 =4.74 + 0.0087 = 4.75 1.02
The change in pH is: 0.01 unit.
The function of the buffers III
Calculation
3. By what factor does the pH of 1 L of distilled water change, if 0.4 g NaOH is dissolved in it.
pH1 = 7, 0.4 g NaOH = 0.01 mole
[OH-] = 0.01 M, pOH = 2, pH = 12, pH2 = 12 The change in pH is: 5 units.
Changes in buffers on treatment with
acids or bases
Physiological buffers I
The buffer systems of the organism
Hydrogen carbonate /carbonic acid buffer (blood) Phosphate buffer (having an intracellular role)
The protons of the acids can originate from: carbonic acid (volatile acid)
Bound acids: phosphoric acid and sulphuric acid
Organic acids: lactic acid and acids of ketone bodies.
Hydrogencarbonate/carbondioxide buffer I
Its role
It maintains the pH of the blood plasma at a constant value of 7.35 – 7.45.
Deviations
At pH>7.6 it is not possible to transfer the CO2 from the cells to the blood.
pH < 7.3 enhances the exchange of gases in the lungs.
Acidosis (at pH<7 coma) Alcalosis
Hydrogencarbonate/carbondioxide buffer II
CO2(g) + H2O(l) CO2(aq) K3 =
)]
( [
)]
( [
2 2
g CO
aq CO
K2 =
)]
( [
] [
2 3 2
aq CO
CO
CO2(aq) + H2O(l) H2CO3(aq) H
K1 =
] [
] ][
[
3 2
3
CO H
HCO H
H2CO3(aq) + H2O(l) H3O+ + HCO3-
The equilibria in the buffer system of the blood Dissolution of carbondioxide
Reaction of the dissolved CO2 with The first dissociation step of the carbonic acid
Hydrogencarbonate/carbondioxide buffer III
K1 =
)]
( [
] ][
[
2 2
3
aq CO
K
HCO H
K1xK2xK3 =
)]
( [
] ][
[
2
3
g CO
HCO H
K =
2 2
] ][
[ 3
CO CO P
HCO H
The [H2CO3] was substituted into the K1 expression from K2.
Into the new equation the [CO2 (aq)] - obtained form the K3 equation- was introduced.
Applying a new constant for the three constants and using Henry’s law for the solubility of gases (it depends on the partial pressure and the absorption coefficient):
Hydrogencarbonate/carbondioxide buffer IV
[H+] = K
]
[ 3
2 2
HCO
PCO
CO pH = pK + log
2 2
]
[ 3
CO CO P HCO
pH = 6.1 + log
53 0226 .
0
24
= 6.1 + log 20 = 7.4 The [H+] and the pH is expressed.
The concrete values are substituted into the equation (pK=6.1, pCO2=53 mbar, CO2=0.0226 mmol/mbar at 37 ºC and [HCO3-]=24 mmol/dm3).
This equation gives the physiological arterial pH value.
Hydrogencarbonate/carbondioxide buffer
IV
Phosphate buffer
[H+] = Ka
] [
] [
2 4
4 2
HPO PO H
pH = pKa + log
] [
] [
4 2
2 4
PO H
HPO pH = 7.21 + log
] [
] [
4 2
2 4
PO H
HPO
The second dissociation step is involved here.
The pH is expressed.
pH = pKa = 7.21
Physiological importance of pH: the enzymes function within a quite small pH range optimally. It influences the distribution of the ions, too.
] [
] [
4 2
2 4
PO H
If HPO =1,
10-8 Ka = 6.2
conjugate base weak acid
H+ + HPO42- H2PO4-
Acid –base indicators I
Definition
They are dyes applied to distinguish between acidic and basic solutions by means of colour change.
HInd + H2O H3O+ + Ind-
acid colour alkaline colour
Kind =
] [
] ][
[
HInd H Ind
The law of mass action is applied.
Acid –base indicators II
[H+] = Kind
] [
] [
Ind
HInd
If the [Ind-] /[Hind] ratio is equal to = 1, pH = -logKind
= pKind , pKind or pI = indicator exponent, it characterizes the indicator. Its colour-change interval is pKind+/- 1 pH unit.
The colour of the indicator is influenced by the [Ind-] /[Hind] ratio.
Acid –base indicators III
The reason for the difference in colours
The two colours can be seen in the presence of each other, if the minor one is minimum 10% present.
Both of them can be observed, if
0,1 <
Ind
- < 10HInd
Acid –base indicators IV
The colour-change interval and colour of some indicators
+ H+
Acid –base indicators V
Methyl red
N N COOH
N
CH
3CH
3NH
N COOH
N
CH3
CH3
+
- H+
colour-change interval: 4.4 – 6.2
The structure of some indicators
yellow
red
Phenolphthalein
O
O O
H
OH
O O O
H
O
-
+ H+ - H+
Acid –base indicators VI
colour-change interval: 8.3 – 10.0
The importance of the acid-base indicators:
1. Determination of pH
2. They are used in titrations to indicate the end point.
colourless violet
Determination of pH I
1. Using indicator dyes (accuracy 1 pH unit) Or mixture of indicator dyes (universal indicator)
Determination of pH II
2. Using potentiometric method (more accurate) measuring the electrode potential (pH dependency) (see electrochemistry)
Titrations
Acid –base titration Definition
It is a quantitative analytical method used for the determination of the concentration of an acid or base.
A standard solution of known concentration is given to the unknown solution. The end point or equivalence point of the titration should be indicated (by visual indicators or potentiometric method).
Titration curves II
Titration of a strong acid with a strong base The equivalence point is at pH = 7.
Application of an indicator: the colour-change interval has to be at about pH = 7, e.g. between pH = 4 – 10.
Titration curves III
CH3COOH(aq) + Na+ + OH- CH3COO- + Na+ + H2O(l)
[CH3COOH] = [CH3COO-] pH = pKa = 4.74
Titration of a weak acid with a strong base
The end-point is in the alkaline range (above pH 7) At the beginning of the curve there is a buffer region, where the inflection point corresponds to the pKa value.
