Biochemistry
Exercise 1
Student’s name: ____________________________
Date: __________________
Assistant’s signature:
Points:
Index numer: _________________________
/3
EXERCISE 1
ACID – BASE BALANCE IN ORGANISM.
BUFFERS.
Buffers can be defined as weak acids or weak bases in the presence of their salts. Also equimolar mixture of
two salts is used to prepare buffer, the one that has more hydrogen atoms, is considered to be the acid, whereas
the salt with less hydrogen atoms is considered to be the salt (see Table 1.). Buffers can resist changes in pH
when small amounts of strong acids or bases are added.
Table 1. Examples of buffers.
Name of a buffer
Components of a buffer
acetate buffer
a weak acid / a salt of weak acid
CH3COOH / CH3COONa
(acetic acid / sodium acetate)
bicarbonate buffer
a weak acid / a salt of weak acid
H2CO3 / NaHCO3
(carbonic acid / sodium bicarbonate)
ammonia buffer
a weak base / a salt of weak base
NH3 / NH4Cl
(ammonia / ammonium chloride)
phosphate buffer
two salts: one salt / second salt
KH2PO4 / K2HPO4
(potassium dihydrogen phosphate / potassium hydrogen phosphate)
Blood has three powerful buffer systems that protect against the introduction of acid or base. These are:
1) proteins in the plasma and haemoglobin in the red cells,
2) the phosphate buffer,
3) bicarbonate buffer in the plasma.
Department of Biochemistry
Second Faculty of Medicine with the English Division and the Physiotherapy Division
1
Biochemistry
Exercise 1
Ad.1.
Proteins are effective, because the blood contains a large amount of them. Proteins act as a buffer because of
their ability to neutralize either acid or base. Protein molecules consist of amino acids with amino groups, to
neutralize acids, and carboxyl groups, which neutralize bases.
Ad.2.
Bicarbonate acts as a buffer against acidity – reacting with a strong acid it produces carbonic acid (H2CO3),
which is a weak acid. For example, if strong acid HCl is added to a solution containing sodium (potassium)
bicarbonate, the following change occurs:
HCO3- + H3O+ ↔ H2CO3 + H2O
↓
H2O + CO2
In this reaction the strong acid (HCl) is replaced by H2CO3, which acidity is so weak, that it causes very little
change in the pH of the solution. At the same moment the amount of H2CO3 is increasing, as much as amount
of HCO3- was diminished.
The carbonic acid thus formed acts as a buffer against strong bases, neutralizing them and producing water and
the poorly ionized compound, sodium bicarbonate, which slightly changes pH of the solution:
H2CO3 + OH- ↔ HCO3- + H2O
Ad.3.
A mixture of H2PO4- and HPO42- acts as a buffer as well. By analogy:
H2PO4- + OH- ↔ HPO42- + H2O
HPO42- + H3O+ ↔ H2PO4- + H2O
And let us consider following reactions, that illustrate mechanisms of action of acetate and ammonia buffer,
respectively:
CH3COO- + H3O+ ↔ CH3COOH + H2O
CH3COOH + OH- ↔ CH3COO- + H2O
NH3aq + H3O+ ↔ NH4+ + H2O
NH4+ + OH- ↔ NH3aq + H2O
Buffers play a very important part in living organism. Without them the human body could never withstand the
acids produced in normal metabolism, the excesses of acids and bases that are sometimes encountered as a
result of accidental intake of extra acid or base, or the abnormal amounts of acids or bases resulting from an
unbalanced diet or disease.
Department of Biochemistry
Second Faculty of Medicine with the English Division and the Physiotherapy Division
2
Biochemistry
Exercise 1
pH of any buffer is defined by Henderson – Hasselbach equation:
𝑝𝐻 = 𝑝𝐾𝑎 + log
𝐶𝑀𝑠𝑎𝑙𝑡
𝐶𝑀𝑎𝑐𝑖𝑑
where:
𝑝𝐾𝑎 = − log 𝐾𝑎
𝐾𝑎 – acid dissociation constant
𝐶𝑀𝑠𝑎𝑙𝑡 – molar concentration (molarity) of a salt
𝐶𝑀𝑎𝑐𝑖𝑑 – molar concentration (molarity) of an acid
Expressions for pH of selected buffers:
 Carbonate buffer
𝑝𝐻 = 𝑝𝐾𝑎 + log
[𝐻𝐶𝑂3− ]
[𝐻2 𝐶𝑂3 ]
 Phosphate buffer
𝑝𝐻 = 𝑝𝐾𝐻2 𝑃𝑂4− + log
[𝐻𝑃𝑂42− ]
[𝐻2 𝑃𝑂4− ]
 Ammonia buffer
𝑝𝐻 = 14 − 𝑝𝑂𝐻 = 14 − (𝑝𝐾𝑏 + log
[𝑁𝐻4 𝐶𝑙]
)
[𝑁𝐻3 ∙ 𝐻2 𝑂]
where:
𝑝𝐾𝑏 = − log 𝐾𝑏
𝐾𝑏 – base dissociation constant
or:
𝑝𝐻 = 𝑝𝐾𝑎 + log
[𝑁𝐻3 ∙ 𝐻2 𝑂]
[𝑁𝐻4 𝐶𝑙]
where:
𝑝𝐾𝑎 = − log 𝐾𝑎
𝐾𝑎 – acid dissociation constant
Buffer capacity is a measure of the ability of solution to resist pH change. Buffer capacity is the number of
moles of strong acid or strong base needed to change the pH of 1 litre of buffer solution by 1 pH unit.
The more concentrated the components of a buffer, the greater the buffer capacity. Since the concentrations
ratio of the buffer components determines the pH, the less the ratio changes. After addition of a certain volume
of acid or base, the ratio changes less when buffer components concentrations are similar than when they are
much different. It follows that a buffer has the highest capacity when its components are present at equal
concentration, that is, when
𝐶𝑀𝑠𝑎𝑙𝑡
𝐶𝑀𝑎𝑐𝑖𝑑
= 1, which gives 𝑝𝐻 = 𝑝𝐾𝑎 .
Buffer capacity (A) is a ratio of acid or base added (to 1 litre of a buffer) to change its pH:
𝐴=
∆𝑛
∆𝑝𝐻
where:
𝐴 – buffer capacity
∆𝑛 – number of moles of added acid or base
∆𝑝𝐻 – pH change
Department of Biochemistry
Second Faculty of Medicine with the English Division and the Physiotherapy Division
3
Biochemistry
Exercise 1
Experiment 1
Preparation of buffers of a known pH
Buffer solutions of define hydrogen ion concentration may be made up from a series of stock solutions in
define proportions.
Prepare buffer solution using stock solutions of
1
M
15
Na2HPO4 and
1
M
15
KH2PO4 or 0.2M CH3COONa and
0.2M CH3COOH (see Table 2. and Table 3.) according to assistant’s instruction.
Table 2. Phosphate buffer.
A–
B–
1
M Na2 HPO4
15
1
M KH2PO4
15
Table 3. Acetate buffer.
A – 0.2M CH3COONa
B – 0.2M CH3COOH
No.
A [ml]
B [ml]
No.
A [ml]
B [ml]
1.
3.70
16.30
1.
3.60
16.40
2.
5.30
14.70
2.
5.30
14.70
3.
7.50
12.50
3.
7.40
12.60
4.
10.00
10.00
4.
9.80
10.20
5.
12.22
7.78
5.
12.00
8.00
6.
14.30
5.70
6.
14.10
5.80
7.
16.08
3.92
7.
15.80
4.20
Determination of pH of buffer solution using indicators
A. The first step in this procedure is to determine the approximate pH of the buffer solution using indicate
paper:
Treat a small piece of indicator paper with 1 drop of buffer solution and compare the colour obtained with
the colour scale. Such paper is convenient and sufficiently accurate for many purposes. It is possible to
obtain a precision of ±1 pH units in this method.
Conclusion:
Department of Biochemistry
Second Faculty of Medicine with the English Division and the Physiotherapy Division
4
Biochemistry
Exercise 1
B. For obtaining a precision of 0.5-0.1 pH and determine pH of buffer solution use indicator solutions (see
Table 4.).
Using 12-well dish (“porcelain plate”) treat a small portion (a few drops) of the buffer solution with 1 drop
of an indicator solution. Compare the colour obtained with the alkaline colour and acid colour of the
indicator. If the colour obtained is intermediate between the acid and alkaline colour of the indicator, the
pH of the solution lies within the effective range of this indicator. If, on the other hand, the solution shows
either the full acid or full alkaline colour with the indicator selected, it is unsuitable and another indicator
must be tried in a similar manner, until the pH will be determined with 1.0-0.5 pH unit precision.
Table 4. Colours and pH range of indicators.
Indicator
Colour at lower
pH
Range of colour
change (pH)
Colour at higher
pH
Thymol Blue
pink-red
1.2–2.8
yellow
Töpfer’s reagent
red
2.8–4.6
yellow
Methyl Orange
red
3.1–4.4
yellow
Bromcresol Green
yellow
3.8–5.4
dark blue
Methyl Red
red
4.2–6.3
yellow
Litmus
red
5.0–8.0
blue
Bromthymol Blue
yellow
6.0–7.6
dark blue
Cresol Red
orange
7.2–8.8
purple
Neutral Red
purple-red
6.8–8.0
orange-brown
Phenol Red
yellow-orange
6.8–8.4
red-purple
Thymol Blue
yellow
8.0–9.6
purple-blue
Phenolphthalein
colourless
8.3–10.0
pink
Conclusion:
Department of Biochemistry
Second Faculty of Medicine with the English Division and the Physiotherapy Division
5
Biochemistry
Exercise 1
C. Using the Henderson-Hasselbach equation calculate the pH of your buffer solution.
The values of Ka and pKa are:
KCH3COOH = 1.8×10-5 (pKa = 4.74)
KH2PO4- = 6.2×10-8 (pKa = 7.21)
Calculations:
Determination of pH of buffer solution using pH-meter
Place the buffer solution into a small beaker and introduce (carefully) an electrode of pH-meter. Wash the
electrode (carefully) with distilled water before and after each measurement. Write down the result.
Conclusion:
Department of Biochemistry
Second Faculty of Medicine with the English Division and the Physiotherapy Division
6
Biochemistry
Exercise 1
Effect of dilution on buffer capacity
Procedure:
1. Measure precisely (into 4 clean test tubes) 5.0 ml, 2.5 ml, 1.0 ml and 0.5 ml of buffer solution and dilute all
to 5.0 ml with distilled water.
2. Add to all test tubes 1-2 drops of appropriate indicator (consult it with an assistant). pH of your buffer
solution should be near to effective pH range of indicator used, but not be within this range!).
3. Titrate all 4 solution with HCl or NaOH (depending on effective pH range). Write down the volume of
titrant used.
4. Calculate buffer capacity for all 4 dilutions (according to Table 5.) and draw a conclusion about the effect of
dilution on buffer capacity and on pH of a solution.
Table 5. Sample presentation of data.
No.
ml of
buffer
ml of
water
dilution
pH of
solution
________________________
(name of indicator used)
colour of solution
(by pH-meter)
before
titration
after
titration
ΔpH
ml of used
acid/base
per 1 dm3 of
solution
Δn
A
1.
2.
3.
4.
where:
𝐴 – buffer capacity
∆𝑛 – number of moles of added acid or base (per 1 litre of a buffer)
∆𝑝𝐻 – pH change
Conclusion:
Department of Biochemistry
Second Faculty of Medicine with the English Division and the Physiotherapy Division
7