15.

Clinical Biochemistry

A. Definitions of acid, base, pH and pKa

B. Buffers-Definition
 Henderson Hasselbach equation
 Principal buffer systems in the ECF, ICF and urine
 Bicarbonate and Phosphate buffer systems (pKa value, normal ratio of
base/acid in the plasma)
 Acidosis & Alkalosis (Definition, classification, causes and biochemical
findings).
C. Normal serum levels and condition where they are altered.
 Glucose, protein, Urea, Uric acid, and Creatinine
 Bilirubin, Cholesterol, Serum electrolytes.

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ACID: An acid is any hydrogen-containing substance that is capable of

donating a proton (hydrogen ion) to another substance. (acids are proton
donors). The word acid comes from a Latin word ‘acere’ which means ‘sour’.
Acids are known to turn blue litmus red. Eg. HCl, HNO3, H2SO4, CH3COOH,
H2CO3

Strong acids: Acids which can dissociate (ionizes) completely in solution are
called as strong acids. Eg. Hydrochloric acid (HCl), Nitric acid (HNO3), Sulphuric
acid (H2SO4 ) etc.,

Strong acids are corrosive in nature and cause severe burns when they come in
contact with skin.

Weak acids: Acids which dissociate only partially in solution are called as weak
acids. Ex. Carbonic acid (H2CO3), lactic acid (CH3CHOHCOOH), acetic acid
(present in vinegar) (CH3COOH), citric acid (present in citrus fruits) etc.,

Weak acids are only mildly corrosive and are even present in our food and body.

BASE: A base is a molecule or ion able to accept a hydrogen ion from an acid
(bases are proton acceptors) or a substance that gives hydroxyl ions in solution.
A base that can be dissolved in water is known as an alkali. Bases are known to
turn red litmus blue. Ex. Sodium hydroxide (NaOH), Potassium hydroxide (KOH),
Calcium hydroxide (Ca(OH)2), Barium hydroxide (Ba(OH)2), Ammonium
hydroxide (NH4OH), Calcium carbonate (CaCO3) etc.,
A base is a compound, which reacts with an acid to form salt and water.

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Eg. HCl + NaOH  NaCl + H2O

Sodium chloride (NaCl) is also known as table salt.

Strong base – A base which completely dissociate into its ions when dissolved in
water Or a base that has an ability to remove a proton from a very weak acid is
known as a strong base. Examples are potassium hydroxide (KOH), sodium
hydroxide (NaOH).

Weak base – A base which dissociates incompletely when dissolved in water.

The aqueous solution contains both the weak base as well as its conjugate acid.
Examples are ammonia (NH3), water (H2O), pyridine (C5H5N).

Theories of Acids and Bases

Three different theories have been put forth in order to define acids and
bases. These theories include the Arrhenius theory, the Bronsted-Lowry
theory, and the Lewis theory of acids and bases. Acids and bases can be defined
via three different theories.

 The Arrhenius theory of acids and bases states that “an acid generates
H+ ions in a solution whereas a base produces OH– ion in its solution”.
 The Bronsted-Lowry theory defines “an acid as a proton donor and a
base as a proton acceptor”.
 Finally, the Lewis definition of acids and bases describes “acids as
electron-pair acceptors and bases as electron-pair donors”.
Properties of Acids and Bases
1. Properties of Acids
 Acids are corrosive in nature.
 They are good conductors of electricity.
 Their pH values are always less than 7.
 When reacted with metals, these substances produce hydrogen gas.
 Acids are sour-tasting substances.
 Examples: Sulphuric acid [H2SO4], Hydrochloric acid [HCl], Acetic acid
[CH3COOH].
2. Properties of Bases
 They are found to have a soapy texture when touched.
–
 These substances release hydroxide ions (OH ions) when dissolved in
water.
 In their aqueous solutions, bases act as good conductors of electricity.
 The pH values corresponding to bases are always greater than 7.

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  Bases are bitter-tasting substances which have the ability to turn red
litmus paper blue.
 Examples: Sodium hydroxide [NaOH], milk of magnesia [Mg(OH)2],
calcium hydroxide [Ca(OH)2].
Uses of Acids and Bases
1. Uses of Acids
 Vinegar, a diluted solution of acetic acid, has various household
applications. It is primarily used as a food preservative.
 Citric acid is an integral part of lemon juice and orange juice. It can also
be used in the preservation of food.
 Sulphuric acid is widely used in batteries. The batteries used to start the
engines of automobiles commonly contain this acid.
 The industrial production of explosives, dyes, paints, and fertilizers
involve the use of sulphuric acid and nitric acid.
 Phosphoric acid is a key ingredient in many soft drinks.

2. Uses of Bases

 The manufacturing of soap and paper involve the use of sodium hydroxide.
NaOH is also used in the manufacture of rayon.
 Ca(OH)2, also known as slaked lime or calcium hydroxide, is used to
manufacture bleaching powder.
 Dry mixes used in painting or decoration are made with the help of
calcium hydroxide.
 Magnesium hydroxide, also known as milk of magnesia, is commonly used
as a laxative. It also reduces any excess acidity in the human stomach and
is, therefore, used as an antacid.
 Ammonium hydroxide is a very important reagent used in laboratories.
 Any excess acidity in soils can be neutralized by employing slaked lime.

pH:
pH is defined as the negative logarithm of H+ ion concentration. Hence the
meaning of the name pH is justified as the power of hydrogen.
pH = -log [H+]
We know that all the acids and bases do not react with the same chemical
compound at the same rate. Some react very vigorously, some moderately while
others show no reaction. To determine the strength of acids and bases
quantitatively, we use a universal indicator which shows different colours at

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different concentration of hydrogen ion in solution. Generally, the value of pH
of acids and bases are used to quantitatively determine their strength.

pH of Acids and Bases

The pH of a solution varies from 0 to 14.

 Solutions having a value of pH ranging between 0 to 7 on pH scale are

known as acidic solutions and for the value of pH ranging between 7 to 14
on pH scale are known as basic solutions.
 Solutions having the value of pH equal to 7 on pH scale are known as
neutral solutions.
Solutions having the value of pH equal to 0 are known to be strongly acidic
solutions. Further, the acidity decreases as the value of pH increases from 0 to
7 whereas, solutions with the value of pH equal to 14 are known as strongly
basic solutions.

The basicity decreases as the value of pH decreases from 14 to 7. The strength

of acids and bases depends on the number of H+ and OH– ions produced. Acids
furnishing more number of H+ ions are known to be strong acids and vice versa.

The degree of ionisation of acids and bases differ for different acids and
bases. It helps in the determination of the strength of acids and bases. The
strength of an acid depends on the concentration of hydronium ion (H 3O+) too.
With the help of the comparison between the concentration of hydronium ion
and the hydroxyl ion, we can distinguish between acids and bases.

 For acidic solution: [H3O+] > [OH–]

 For neutral solution: [H3O+] = [OH–]
 For basic solution: [H3O+] < [OH–]

Importance of pH

 Only a narrow range of pH change can be sustained by a living organism,

any further change in pH can make the living difficult. For example: in the
case of acid rain, the pH of water is less than 7. As it flows into a river,
it lowers the pH of river water which makes the survival of aquatic life
difficult.
 We know that our stomach contains hydrochloric acid which helps in the
digestion of food. When the stomach produces too much of hydrochloric
acid during indigestion, we feel a lot of pain and irritation. Hence, we

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 generally use antacids or a mild base which increases the pH of the acidic
stomach and thus decreases the pain.
 Bacteria present in our mouth sometimes lower the pH of our mouth by
producing acids through degradation of the food particle. Hence, we are
instructed to clean our mouths with toothpaste (which are generally
basic) to prevent their decay by maintaining the pH.
 We experience a lot of pain in case of bee-sting as the bee injects the
methanoic acid through its sting. Hence, we are generally advised to
apply baking soda or other mild bases on the surface as it helps in
maintaining the pH of the surface.
pKa: The pKa is the negative log of the acid dissociation constant or Ka value.
A lower pKa value indicates a stronger acid. That is, the lower value indicates
the acid more fully dissociates in water. The pKa value is one method used to
indicate the strength of an acid.

BUFFERS: Solutions which can resist changes in pH when acid or alkali is

added in small quantities are called as buffers. Buffers are a) mixture of weak
acids with their salt with strong base or b) mixture of weak bases with their
salt with strong acid. Examples: H2CO3 / NaHCO3 (Carbonate/bicarbonate),
CH3COOH / CH3COONa (acetic acid/sodium acetate), Na2HPO4 / NaH2PO4
(Phosphate buffer)
Buffer Solutions are used in fermentation, food preservatives, drug delivery,
electroplating, printing, the activity of enzymes, blood oxygen carrying capacity
need specific hydrogen ion concentration (pH).

Uses of Buffer Solutions

 Use of bicarbonate and carbonic acid buffer system in order to regulate

the pH of blood.
 Buffer solutions are also used to maintain an optimum pH for enzyme
activity in the cells.
 The absence of these buffers may lead to the slowing of the enzyme
action, loss in enzyme properties, or even denature of the enzymes. This
denaturation process can even permanently deactivate the catalytic
action of the enzymes.

Henderson-Hasselbalch Equation
The Henderson-Hasselbalch equation provides a relationship between the pH of
acids (in aqueous solutions) and their pKa (acid dissociation constant). The pH of

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a buffer solution can be estimated with the help of this equation when the
concentration of the acid and its conjugate base, or the base and the
corresponding conjugate acid, are known. The Henderson-Hasselbalch equation
can be written as:

pH = pKa + log10 ([A–] / [HA])

Where [A–] denotes the molar concentration of the conjugate base (of the acid)
and [HA] denotes the molar concentration of the weak acid. Therefore, the
Henderson-Hasselbalch equation can also be written as:

An equation that could calculate the pH value of a given buffer solution was
first derived by the American chemist Lawrence Joseph Henderson. This
equation was then re-expressed in logarithmic terms by the Danish chemist Karl
Albert Hasselbalch. The resulting equation was named the Henderson-
Hasselbalch Equation.

Derivation of the Henderson-Hasselbalch Equation

The ionization constants of strong acids and strong bases can be easily
calculated with the help of direct methods. However, the same methods cannot
be used with weak acids and bases since the extent of ionization of these acids
and bases are very low (weak acids and bases hardly ionize). Therefore, in order
to approximate the pH of these types of solutions, the Henderson-Hasselbalch
Equation is used.

Let us take an example of ionization of weak acid HA:

HA + H2O ⇋ H+ + A−
Acid dissociation constant, K a can be given as:

K a = [H+][A−] / [HA]
Taking, negative log of RHS and LHS:

−log Ka = −log [H+][A−] / [HA]

⇒−log Ka = −log [H+] – log [A−] / [HA]
As we know, −log [H+] = pH and −log Ka = pKa,
The equation above can also be written as,

pKa = pH − log [A−] / [HA]

Rearranging the equation,

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⇒pH = pKa + log[A−] / [HA]

The above equation is known as Henderson-Hasselbalch equation, popularly

known as Henderson equation. It is very useful for estimating the pH of a
buffer solution and finding the equilibrium pH in acid-base reactions.

Similarly, for a weak base B:

B + H2O ⇋ OH− + HB+

Base dissociation constant, Kb, of the base can be given as,

Kb = [BH+][OH−] / [B]
Taking negative log of RHS and LHS

−logKb = −log[BH+][OH−] / [B]

⇒−logKb = −log [OH−]−log[BH+] / [B]
As we know, −log [OH−] = pOH and −logKb = pKb,
Above equation can be written as,

pKb = pOH − log[BH+]/[B]

Rearranging the equation,

⇒pOH = pKb+log [BH+] / [B]

Limitations of the Henderson-Hasselbalch Equation

The Henderson-Hasselbalch equation fails to predict accurate values for the

strong acids and strong bases because it assumes that the concentration of the
acid and its conjugate base at chemical equilibrium will remain the same as the
formal concentration (the binding of protons to the base is neglected).

Since the Henderson-Hasselbalch equation does not consider the self-

dissociation undergone by water, it fails to offer accurate pH values for
extremely dilute buffer solutions.

The three major buffer systems in our body: They are carbonic acid
bicarbonate buffer system, phosphate buffer system and protein buffer
system.

Principal buffer system in the Extracellular fluid (ECF):

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Bicarbonate buffer system (NaHCO3/H2CO3) is the most important and
predominant extracellular buffer and is therapeutically the most widely used in
the treatment of acidosis.
 The bicarbonates buffer systems are composed of weak acid i.e, carbonic
acid(H2CO3) and a weak base i.e, bicarbonate(HCO3). When combined,
they function to keep the pH of the blood and other internal fluids in a
particular range. This reduces the concentration of H+ ion and brings the
pH back to its original value.
 Accounts for 65% buffering capacity in plasma & 40 % buffering action in
the whole body.
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 The normal average HCO3 level of plasma is 24 mmol/L
 The normal pCO2 of arterial blood is 40mmHg.
 The normal H2CO3 concentration in blood is 1.2mmol/L (product of pCO2 &
solubility constant of CO2 (0.03)).
 The pka for carbonic acid is 6.1

Substituting these values in H-H’s equation

pH= pka + log [HCO3-] /[H2CO3]

7.4 = 6.1 + log 24 /1.2

= 6.1 + log 20

= 6.1 + 1.3 = 7.4

 Blood pH is 7.4, but the ratio of HCO3- to H2CO3 is 20:1

 HCO3- conc. is 20 times higher than H2CO3
 Hence the ratio of HCO3- to H2CO3 at pH 7.4 is 20 under normal
conditions.
 Cellular respiration produces carbon dioxide as a waste product. This is
immediately converted to bicarbonate ion in the blood. On reaching the
lungs it is again converted to and released as carbon dioxide.
While in the blood, it neutralises acids released due to other metabolic
processes. In the stomach and duodenum it also neutralises gastric acids
and stabilises the intra cellular pH of epithelial cells by the secretions of
bicarbonate ions into the gastric mucosa.

Principal buffer system in the Intracellular fluid (ICF): Phosphate

buffer (Na2HPO4 / NaH2PO4) and protein buffer system are an important
intracellular buffer systems, basically involved in the maintenance of a constant
pH of intracellular fluid.

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Phosphate buffer (Na2HPO4 / NaH2PO4):
 Concentration in plasma is very low.
 The pka value is 6.8
 Effective at wide range of pH as it has more than one ionizable group
with different pka values (1.9 & 12.4).
 Would have been more effective if it was present in higher concentration.
 Phosphate buffer system operates in the internal fluids of all cells. It
consists of dihydrogen phosphate ions as the hydrogen ion donor (acid)
and hydrogen phosphate ion as the ion acceptor (base). If additional
hydroxide ions enter the cellular fluid, they are neutralised by the
dihydrogen phosphate ion. If extra hydrogen ions enter the cellular fluid
then they are neutralised by the hydrogen phosphate ion.

Protein buffer: Plasma proteins & hemoglobin constitute the protein buffer.

 The buffering capacity depends on pka of ionizable groups of amino acids.

 Histidine imidazole group with pka of 6.7 is the most effective

contributor of protein buffers.
 16 histidine residues in albumin & 38 in hemoglobin.
 Hemoglobin of RBC is an important buffer.
 Mainly buffers the fixed acids (in addition to its role in transport of O 2 &
CO2).

Protein buffer system helps to maintain acidity in and around the cells.
Haemoglobin makes an excellent buffer by binding to small amounts of acids in
the blood, before they can alter the pH of the blood. Other proteins containing
amino acid histidine are also good at buffering.
The main purpose of all these buffers is to maintain proper pH (7.35 to 7.45)
within the body system so that all biochemical process can take place.

Phosphate buffer in urine:

 According to the American Association for Clinical Chemistry, the

average value for urine pH is 6.0, but it can range from 4.5 to
8.0. Urine under 5.0 is acidic, and urine higher than 8.0 is alkaline, or
basic.
 Normally phosphate is the only buffer in urine, although carbonic acid/
bicarbonate is also present. The developing urine contains
NaH2PO4/Na2HPO4 in the same concentration as present in blood plasma.

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Disorders of acid-base balance are broadly grouped as an acidosis or an
alkalosis depending on whether there is a general accumulation of hydrogen ions,
an acidosis (low pH), or if there is a loss of hydrogen ions, an alkalosis (high
pH).These are both further subdivided into two more groups depending on the
cause of the disorder. If the pH imbalance is due to metabolic or renal damage,
it is classified as ‘metabolic’ and where the problem is due to lung function, it is
classified as respiratory.
Acidosis: If the pH of blood is below 7.35 it is called acidosis.
Types:
a) Metabolic acidosis – due to decrease in bicarbonate.
b) Respiratory acidosis – due to an increase in carbonic acid.
 Life is threatened when pH is below 7.25
 Acidosis leads to CNS depression & coma.
 When pH is below 7.0 death occurs.
Alkalosis: If the pH of the blood is more than 7.42 it is called alkalosis.
Types:
a) Metabolic alkalosis – due to an increase in bicarbonate.
b) Respiratory alkalosis – due to decrease in carbonic acid.
 Very dangerous if it is increased above 7.55
 Alkalosis induces neuromuscular hyper excitability & tetany.
 When pH is above 7.6, death occurs.
Causes: all the four acid-base disorders are primarily due to alterations in
either bicarbonate or carbonic acid.
 Metabolic acid-base disorders are due to alteration in bicarbonate
concentration.
 Respiratory acid-base disturbances are due to a change in carbonic acid
level (CO2).
Metabolic acidosis
- Diabetes mellitus (ketoacidosis)
- Lactic acidosis, diarrhea
- Renal failure, Poisoning or overdose (methanol poisoning) etc.,
Respiratory acidosis
-Severe asthma
-Pneumonia
-Cardiac arrest

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 -Obstruction in airways
-Chest deformities
-Depression of respiratory center (by drugs e.g. opiates)
Metabolic alkalosis
-Severe vomiting
-Hypokalemia
-Intravenous administration of bicarbonate

Respiratory alkalosis
- Hyperventilation
-Anemia
-High altitude
-Salicylate poisoning
-Pneumonia, Asthama, pulmonary edema
Biochemical findings:
Metabolic acidosis Plasma findings:
 Characterized by primary decrease in plasma HCO3- ( 24mmol/L)
 Decreased pCO2
 pH decreases (uncompensated state), normal in compensated state.
 Cl- is normal, but increases in renal tubular acidosis & acetazolamide
therapy.
Respiratory acidosis Plasma findings:
 Decrease in pH
 Increase in pCO2
 Normal plasma (HCO3-) (uncompensated state).
Metabolic alkalosis Plasma findings:
 Marked increase in pH.
 Increase in HCO3-
 Normal pCO2 (uncompensated state),
 Increase in pCO2 (compensated state).
Respiratory alkalosis Plasma findings:
 Marked increase in pH
 Marked decrease in pCO2
 Normal HCO3- (uncompensated state)
 Decrease HCO3- (compensated state)

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Normal serum levels and conditions where they are altered:

Parameters Serum /Plasma Altered conditions

level
Glucose
Increased in Diabetes mellitus,
Fasting 70-110 mg/dL hyperthyroidism, hyperpituitarism.
(SI Unit: 4-6 mmol/L) Decreased in hyperinsulinism
Random/post prandial 75-140 mg/dL
(<8 mmol/L)
Renal threshold for glucose 180 mg/dL
(10 mmol/L)
Decreases in Cirrhosis of liver,
malnutrition, nephrotic syndrome.
Protein, Total 6-8 g/dL

Increased in
Pre-renal- increased catabolism
(major surgery, diabetic coma,
thyrotoxicosis, fever, wasting
Urea diseases).
20-40 mg/dL
Renal- acute glomerulonephritis,
polycystic kidney.
Post-renal- tumours, renal calculi,
enlargement of prostate obstructing
the urinary tract.
Uric acid
Male 3-7 mg/dL Increased in Gout, leukemia,
Female 3-6 mg/dL polycythaemia, febrile conditions,
Children 2-5.5 mg/dL eclampsia.

Creatinine 0.7-1.5 mg/dL Increased in

Renal disease (nephritis)
Bilirubin,

Total 0.2-1 mg/dL Increased in

Conjugated(direct) 0-0.2 mg/dL Biliary obstruction,
Unconjugated (indirect) 0.2-0.8 mg/dL Hemolytic anemia

Increased in
Cholesterol, Total Diabetes mellitus, hypothyroidism,
150-200 mg/dL
nephritic syndrome, obstructive
jaundice.
Serum electrolytes

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 135-145 mEq/L Decrease in adrenocortical
Sodium insufficiency
3.5-5mEq/L Increased in
Potassium Uremia, acute infections, pneumonia.
96-106 mEq/L Congestive cardiac failure,
Chloride eclampsia, nephritis

QUESTIONS

1. Define a) Acids, b) Bases c) Salts with one example for each.

2. Write Henderson-Hasselbalch equation and mention its uses.
3. Name two blood buffers.
4. What is the normal pH of blood? Name the fluid buffers. Explain
any one buffer in the regulation of blood pH.
5. Define acids and bases with suitable examples.
6. Define buffers with examples. How do you determine the pH of a
buffer using Henderson-Hasselbalch equation?
7. Explain bicarbonate buffer system.
8. SI units for glucose and total protein in serum.
9. What are bases? Give two examples.
10. Define acids. Give two examples.
11. Give the normal serum levels of a) Creatinine b) Uric acid. Mention
one clinical condition for each where their levels are elevated.
12. Name the electrolytes and mention their normal serum levels.
13. Write the normal serum level of a) Bilirubin and b) Cholesterol.
Mention one clinical condition for each where their levels are
altered.
14. Mention the normal serum levels of a) Glucose and b) Total protein.
15. Define buffer and name the principle buffer system in urine.
16. Define acidosis. Mention the types and biochemical findings in each
type.
17. Define alkalosis. Mention the types and biochemical findings in each
type.
18. Define pH and pKa. Mention the normal blood pH.
19. Mention two renal markers.
20. Define diabetes mellitus and mention the cause.
21. What is jaundice? Mention the types.

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22. Mention the causes for Respiratory acidosis and alkalosis.
23. Mention the causes for Metabolic acidosis and alkalosis.
24. Name the clinical condition each where the levels of a) Glucose and
b) Bilirubin are altered.
25. Mention the clinical conditions where the serum electrolytes are
altered.
26. What is hypercholesterolaemia? Mention the normal serum levels
of cholesterol.
27. What is atherosclerosis?
28. Name four acid-base disorders

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