The document discusses acid-base balance and pH. It defines acids and bases, noting that acids donate protons while bases accept protons. pH is defined as the negative logarithm of hydrogen ion concentration. Strong acids ionize completely while weak acids only partially ionize. The Henderson-Hasselbalch equation relates pH, pKa, and concentrations of acid and conjugate base. Buffers resist pH changes when acids or bases are added. The body maintains acid-base balance through bicarbonate, phosphate, protein, and hemoglobin buffer systems as well as respiratory and renal regulation.

1. ACID BASE BALANCE AND pH
Dr Neha Rani Verma
Assistant professor
Dept. of Biochemistry

2. Definition
 Acids are substances that are capable of
donating protons and
 Acids are proton donors
 Examples:
HA  H
+
+ A
–
HCl  H
+
+ Cl
–
H2CO3  H
+
+ HCO3
–
 pH <7
 Bases are those that accept protons.
 Bases are proton acceptors
 Examples:
NH3 + H
+
 NH
+
4
HCO3
–
+H
+
 H2CO3
 pH > 7

3.  pH= Negative logarithm of hydrogen ion concentration.
 Normal value 7.4 (range 7.38 -7.42)
 The pH value is inversely proportional to the acidity.
 Lower the pH, higher the acidity or hydrogen ion concentration.
 Strong acids- Acids which ionize completely; e.g. HCl
 HCl = H+ + Cl–- (Complete)
 Weak acids- Acids which ionize incompletely e.g. H2CO3
 H2CO3 = H
+
+ HCO3
–-
(Partial)
 Buffers= Solutions minimize changes in pH
_1_
pH = –log [H+] = log [H+]

4. Dissociation Constant
 Dissociation of an acid is a freely reversible reaction. At equilibrium the ratio between
dissociated and undissociated particle is a constant.
 Dissociation constant (Ka) of an acid is
 Ka = [H+] [A–-]
[HA]
 pKa value =pH at which the acid is half ionized; Salt : Acid = 1:1
 Strong acids will have a low pKa and weak acids have a higher pKa.

5.  The relationship between pH, pKa, concentration of acid and
conjugate base (or salt) is expressed by the
 Henderson-Hasselbalch equation,
[base] [salt]
 pH = pKa + log ––––– or pH = pKa + log –––––
[acid] [acid]
 When [base] = [acid]; then pH = pKa
 Thus, when the acid is half ionized, pH and pKa have the same
values.
Practical
application in
clinical practice in
assessing the acid-
base status, and
predicting the
limits of the
compensation of
body buffers.

6. Buffers
 Buffers are solutions which can resist changes in pH when acid or alkali is added.
 Buffering capacity is the number of grams of strong acid or alkali which is necessary for a
change in pH of one unit of one liter of buffer solution
 Buffers are of two 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.
Examples
 i. H2CO3 / NaHCO3 (Bicarbonate buffer- carbonic acid and sodium bicarbonate)
 ii. CH3COOH / CH3COO Na (Acetate buffer- acetic acid and sodium acetate)
 iii. Na2HPO4 / NaH2PO4 (Phosphate buffer)
Factors affecting pH of a
buffer
1. pK
2. The ratio of salt to acid
concentrations
Factors affecting buffer
capacity
1. Actual concentration of salt
and acid

7. ACID BASE BALANCE
 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 below that of the plasma.
 During metabolism Volatile acids produced are carbonic acid, is eliminated as CO2 AND
 Nonvolatile (fixed) acids produced are lactate, keto acids, sulfuric acid and phosphoric acid, are
buffered and excreted by the kidney.
Acidosis
If the pH is below 7.38
Acidosis leads to CNS depression
and coma.
Death occurs when pH is below
7.0.
Alkalosis
When the pH is more than 7.42
Alkalosis induces neuromuscular
hyperexcitability and tetany.
Death occurs when the pH is above
7.6.

9. Bicarbonate buffer system
 The most important buffer system in the plasma is the bicarbonate-carbonic acid system
(NaHCO3 / H2CO3). Reasons are:
 Present in High concentration
 Bicarbonate (HCO3
–), is regulated by the kidney (metabolic component).
 carbonic acid (/ H2CO3), is under respiratory regulation (respiratory component).
fig. During acidosis reaction
follow red arrow while during
alkalosis reaction follows blue
arrow

10. Phosphate buffer system
 The main elements of the phosphate buffer
system are HPO4 – – and H2PO4 – .
 When a strong acid such as HCl is added to a
phosphate buffer system, the H+ is accepted
by the base HPO4 – – and converted to H2PO4
– and strong acid HCl is replaced by a weak
acid H2PO4 and decrease in pH is minimized.
 When strong base, such as NaOH, is added to
the buffer system, the OH– is buffered by the
H2PO4 – to form HPO4 – – and water. Thus,
strong base NaOH is replaced by weak base
HPO4 – –, causing slight increase in the pH.

11. Protein buffer system
 In the blood, plasma proteins especially albumin act as buffer because:
 Proteins contain a large number of dissociable acidic (COOH) and basic (NH2) groups in
their structure.
 In acid solution they act as a buffer in that, the basic amino group (NH2) takes up excess
H+ ions forming (NH3+).
 In basic solutions the acidic COOH groups give up hydrogen ion forming OH– of alkali to
water.
 Other important buffer groups of proteins, are the imidazole groups of histidine.
Each albumin molecule contains 16 histidine residues.

12. Haemoglobin Buffer
 • Hemoglobin is the major intracellular buffer which is present in erythrocytes.
• It buffers carbonic acid (H2CO3) and its anhydride CO2 from the tissues.
In the tissues the CO2 formed by
metabolic processes diffuses into red
blood cell and is converted to carbonic
acid (H2CO3) by carbonic anhydrase
(CA). The H2CO3 thus formed ionizes
to form H+ and HCO3– and results in
decrease in blood pH.
• The deoxyhemoglobin (KHb) acts as
a buffer and accepts these H+ ions to
form HHb (weak acid). Thus, H+ ions
produced from H2CO3 does not
cause any change in pH

14. Contd.
 In the lungs deoxyhemoglobin (HHb)
carried from tissue is oxygenated to
oxyhemoglobin (HHbO2).
 Since, oxyhemoglobin (HHbO2) is a
stronger acid results in the release of
H+, which is buffered by HCO3– to give
H2CO3.
 This buffering effect reduces the pH
change as a result of the oxygenation of
HHb.
 The carbonic acid formed is converted
quickly in the presence of the carbonic
anhydrase (CA) to carbon dioxide and
water which is eliminated by ventilation.

15. Respiratory Regulation of pH
 Second line of defense against acid-bases
disturbances
 An increase in (H+) or (H2CO3) stimulates the
respiratory center to increase the rate of
respiratory ventilation. When the ventilation
rate increases, more CO2 is released from the
blood and pH increases.
 Similarly, an increase in (OH–) or (HCO3–)
depresses respiratory ventilation. A decrease
in ventilation rate will cause a decrease in
release of CO2 from the blood. The increased
blood CO2 will result in the formation of more
H2CO3. Thus, there will be decrease in pH

16. Renal Regulation of pH
 Third line of defense in acid-base balance
 Long-term acid-base control is exerted by renal mechanisms.
 Kidney participates in the regulation of acid base balance primarily by
conservation of HCO3– (alkali reserve), which occur through four key
mechanisms-

17. Contd.
 1. Exchange of H+ for Na+ of tubular
fluid.
 2. Reabsorption (reclamation) of
bicarbonate from tubular fluid.
 3. Formation of ammonia and excretion
of ammonium ion (NH4+) in the urine.
 4. Excretion of H+ as H2PO4– in urine.

18. 1. Exchange of H+ for Na+ of tubular fluid.
 In renal tubular cells, the carbonic
anhydrase catalyzes(CA) the
formation of carbonic acid (H2CO3)
from CO2 and water. The carbonic
acid, thus formed dissociates to yield
H+ and HCO3–.
 The H+ ions formed in tubular cells
are secreted into the tubular fluid in
exchange for Na+ present in tubular
fluid.
 The bicarbonate anion formed by the
dissociation of H2CO3 in the tubular
cell diffuses into the blood as the
accompanying ion to Na+ and
HCO3– is thus conserved and
increases the 'alkali reserve' of the
body.

19. 2. Reabsorption (reclamation) of bicarbonate
from tubular fluid.
 Some H+ that are secreted into the
tubular fluid in exchange of Na+ react
with HCO3– in the tubular fluid to form
H2CO3, which is dehydrated to CO2
and H2O by an enzyme carbonic
anhydrase.
 The increase in CO2 in tubular fluid
causes carbon dioxide to diffuse into
the tubular cell where it react with
H2O to form H2CO3 and
subsequently, H+ and HCO3–.
 The process of bicarbonate
reabsorption is enhanced in states of
acidosis and decreased in alkalosis.

20. 3. Formation of ammonia and excretion of ammonium
ion (NH4+) in the urine.
 Ammonia (the urinary buffer) is produced by
deamination of glutamine in renal tubular cell.
Glutaminase catalyzes this reaction.
 Ammonia is a gas and diffuses readily across
the cell membrane into the tubular lumen,
where it buffers hydrogen ions to form
ammonium (NH4+) ions.
 The NH4+ ions formed in the tubular lumen
cannot diffuse back into tubular cells and
thus, is trapped in the tubular urine and
excreted with anions, such as phosphate,
chloride or sulphate.
 The removal of hydrogen ions as NH4+
decreases the requirement of bicarbonate to
buffer the urine.

21. 4. Excretion of H+ as H2PO4– in urine.
 The hydrogen ions secreted into the
tubular fluid in exchange of Na+ are
buffered by HPO4– – of phosphate
buffer.
 HPO4– – combines with the secreted
H+ and is converted to H2PO4– and
are excreted in the urine as
NaH2PO4