This document discusses biological buffers and acid-base balance in the human body. It covers: 1) Definitions of acids, bases, and pH. Buffers resist changes in pH when acids or bases are added. 2) The major buffer systems in the body are bicarbonate-carbonate and phosphate. Bicarbonate buffering is the most important for maintaining pH between 7.38-7.42. 3) The body regulates pH through buffers, respiration, and renal excretion of acids and bases. Respiration controls carbonic acid levels while the kidneys regulate bicarbonate.

1. Dental Biochemistry 2 – (Lec. 9)
Biological Buffer
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2. ACIDS AND BASES
• Definition
• According to the definition proposed by Bronsted,
acids are substances that are capable of donating
protons and bases are those that accept protons.
• For example: HCI ↔ H++CI-
HCO-3 + H+ ↔ H2C03
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3. • Acidity of a solution and pH:
• The acidity of a solution is measured by
noting the hydrogen ion concentration in the
solution.
• pH = log 1/[H+]
• Thus, the pH value is inversely proportional
to the acidity. Lower the pH, higher the
acidity or hydrogen ion concentration while
higher the pH, the acidity is lower. The pH 7
indicates the neutral pH.
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4. BUFFERS
1. Definition
• Buffers are solutions which can resist changes in pH when
acid or alkali is added.
2. Composition of a buffer
Buffers are 2 types:
a, Mixtures of weak acids with their salt with a strong base or
b, Mixtures of weak bases with their salt with a strong acid.
• A few examples are given below:
i. Bicarbonate buffer
ii. Acetate buffer
iii. Phosphate buffer.
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5. • Buffer Capacity
The buffering capacity of a buffer is
defined as the ability of the buffer to
resist changes in pH when an acid or base
is added.
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6. How do Buffers Act?
• i. Buffer solutions consist of mixtures of a weak
acid or base and its salt.
• ii. To take an example, when hydrochloric acid is
added to the acetate buffer, the salt reacts with
the acid forming the weak acid, acetic acid and its
salt. Similarly when a base is added, the acid
reacts with it forming salt and water. Thus
changes in the pH are minimized.
• CH3-COOH + NaOH ↔ CH3-COONa + H20
• CH3-COONa + HCI ↔ CH3-COOH + NaCl
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7. Acid-Base balance
• Normal pH:
• The pH of plasma is 7.4 in normal life, the variation
of plasma pH is very small. The pH of plasma is
maintained within a narrow range of 7.38 to 7.42.
The pH of the interstitial fluid is generally 0.5 units
blow that of the plasma.
• Acidosis:
• If the pH is blow 7.38, it is called acidosis. Life is
threatened when pH is lowered below 7.25. Death
occurs when pH is below 7.
• Alkalosis:
• When the pH is more than 7.42, it is alkalosis. It is
very dangerous if pH is increased above 7.55. Death
occurs when the pH is above 7.6.
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8. Volatile and Fixed Acids
• i. During the normal metabolism, the acids
produced may be volatile acid like carbonic acid
or nonvolatile (fixed) acids like lactate, keto acids,
sulfuric acid and phosphoric acid.
• ii. The carbonic acid, being volatile, is eliminated
as CO2 by the lungs. The fixed acids are buffered
and later on the H+ are excreted by the kidney.
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9. Mechanisms of regulation of pH
• Buffers of body fluids.
• Respiratory system.
• Renal excretion.
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10. 1- Buffers of the body fluids:
• Buffers are the first line of defense
against acid load. The buffers are
effective as long as the acid load is not
excessive, and the alkali reserve is not
exhausted.
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11. A- Bicarbonate buffer system:
• The most important buffer system in the
plasma is the bicarbonate-carbonic acid system
(NaHCO3/H2CO3). It accounts for 65% of
buffering capacity in plasma and 40% of
buffering action in the whole body.
• The base constituent, bicarbonate (HCO3-), is
regulated by the kidney (metabolic
component).
• While the acid part, carbonic acid (H2CO3), is
under respiratory regulation (respiratory
component).
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12. • The normal bicarbonate level of plasma is 24
mmol/liter.
• The ratio of HCO3- to H2CO3 at pH 7.4 is 20
under normal conditions.
• The bicarbonate carbonic acid buffer system is
the most important for the following reasons:
• Presence of bicarbonate in relatively high
concentrations.
• The components are under physiological
control, CO2 by lung and bicarbonate by
kidneys.
•
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13. B- Alkali reserve:
• Bicarbonate represents the alkali reserve
and it has to be sufficiently high to meet
the acid load. If it was too low, all the
HCO3- would have been exhausted within a
very short time; and buffering will not be
effective. So, under physiological
circumstances, the ratio of 20 (a high alkali
reserve) ensure high buffering efficiency
against acids.
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14. C- Phosphate buffer system:
• It is mainly intracellular buffer. Its
concentration in plasma is very low.
• The phosphate buffer system is found to
be effective at a wide pH range, because it
has more than one ionizable group.
• In the body, Na2HPO4/NaH2PO4 is an
effective buffer system.
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15. D- Buffers act quickly, but not
permanently:
• Buffers can respond immediately to addition
of acid or base, but they do not serve to
eliminate the acid from the body.
• They are also unable to replenish the alkali
reserve of the body.
• For the final elimination of acids, the
respiratory and renal regulations are very
essential.
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16. 2- Respiratory regulation of pH
• It is the second line of defense
• This is achieved by changing the pCO2.
• The CO2 diffuses from the cells into the
extracellular fluid and reaches the lungs through
the blood.
• When there is a fall in pH of plasma (acidosis), the
respiratory rate is stimulated resulting in
hyperventilation.
• This would eliminate more CO2, thus lowering
H2CO3 level.
• However, this cannot continue for long. The
respiratory system responds to any change in pH
immediately, but cannot proceed to completion.
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17. 3- Renal regulation of pH
• Kidneys excrete urine (pH around 6) with a pH
lower than that of extracellular fluid (pH= 7.4). This
is called acidification of urine. The pH of the urine
may vary from as low as 4.5 to as high as 9.8,
depending on the amount of acid excreted.
The major kidney mechanisms for regulation of pH
are:
• Excretion of H+.
• Reabsorption of bicarbonate.
• Excretion of titratable acid (net acid excretion).
• Excretion of NH4+ (ammonium ion).
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18. Classification of Acid-Base Disturbances
1. Acidosis (fall in pH)
Where acids accumulate or base is lost, it is acidosis.
a. Respiratory acidosis: Primary excess of carbonic acid
b. Metabolic acidosis: Primary deficiency of bicarbonate.
2. Alkalosis (rise in pH)
A loss of acid or accumulation of base is alkalosis
a. Respiratory alkalosis: Primary deficiency of carbonic
acid.
b. Metabolic alkalosis: Primary excess of bicarbonate
3. Mixed Responses
i. In mixed disturbances, both HCO3 and H2C03 levels are
altered. 18

19. ELECTROLYTE AND WATER BALANCE
• The maintenance of extracellular fluid volume
and pH are closely interrelated.
• Body is composed of about 60-70% water.
• Osmolality of the intra-and extracellular fluid
is the same, but there is marked difference in
the solute content.
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20. Intake and Output of Water
• During oxidation of food stuffs, 1 g carbohydrate
produces 0.6 ml of water, 1 g protein releases 0.4
ml water and 1 g fat generates 1.1 ml of water.
Intake of 1000 kcal produces 125 ml water
• The major factors controlling the intake are thirst
and the rate of metabolism.
• The renal function is the major factor controlling
the rate of output.
• The rate of loss through skin is influenced by the
weather.
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