Bichemistry

1. PH AND BUFFER
Acid Base Balance
Prakash Pokhrel

2. PH
It is the negative log of the hydrogen ion concentration.
pH = -log [H+]
ACID BASE BALANCE

3.  pH is a unit of measure which describes the degree of
acidity or alkalinity (basic) of a solution.
 It is measured on a scale of 0 to 14.
 Low pH values correspond to high concentrations of H+
and high pH values correspond to low concentrations of
H+.

4. PH VALUE
 The pH value of a substance is directly related to the
ratio of the hydrogen ion and hydroxyl ion
concentrations.
 If the H+ concentration is higher than OH- the material
is acidic.
 If the OH- concentration is higher than H+ the material
is basic.
 7 is neutral, < is acidic, >7 is basic

5. THE PH SCALE
 The pH scale corresponds to the concentration of
hydrogen ions.
 For example pure water H+ ion concentration is 1 x 10^-
7 M, therefore the pH would then be 7.

6. Acid
Any compound which forms H⁺ ions in solution
(proton donors)
eg: Carbonic acid releases H⁺ ions
Base
Any compound which combines with H⁺ ions in
solution (proton acceptors)
eg:Bicarbonate(HCO3⁻) accepts H+ ions

7. ACID–BASE BALANCE
 Normal pH : 7.35-7.45
 Acidosis
Physiological state resulting from abnormally low plasma
 Alkalosis
Physiological state resulting from abnormally high plasma
 Acidemia: plasma pH < 7.35
 Alkalemia: plasma pH > 7.45

8. MEASUREMENT OF PH
The pH can be measured by:
ph strips
Ph indicators
Ph meter

9. SOME IMPORTANT INDICATORS USED IN A CLINICAL
BIOCHEMISTRY LABORATORY ARE LISTED BELOW:
sr,.
No.
INDICATOR Ph range Colour in
acidic ph
Colour in
basic ph
1 Phenophthalein 9.3-10.5 colourless pink
2 Methyl orange 3.1-4.6 red yellow
3 Bromophenol blue 3.0-4.6 yellow blue
4 Methyl red 4.4-6.2 Red yellow
5 Phenol red 6.8 – 8.4 yellow red
6 Litmus 4.5-8.3 red Blue

10. PH METER
 The pH meter is a laboratory equipment which used to measure
acidity or alkalinity of a solution
 The pH meter measures the concentration of hydrogen ions [H+]
using an ion-sensitive electrode.
 It is the most reliable and convenient method for measuring ph.

12. BUFFER
A buffer solution is a solution which resists changes in pH when
a small amount of acid or base is added.
Typically a mixture of a weak acid and a salt of its conjugate
base or weak base and a salt of its conjugate acid.

13. TYPES OF BUFFERS
Two types :
 ACIDIC BUFFERS –
Solution of a mixture of a weak acid and a salt of this weak
acid with a strong base.
E.g. CH3COOH + CH3COONa
( weak acid ) ( Salt )
 BASIC BUFFERS –
Solution of a mixture of a weak base and a salt of this weak
base with a strong acid.
e.g. NH4OH + NH4Cl
(Weak base) ( Salt)

14. HOW BUFFERS WORK
 Equilibrium between acid and base.
 Example: ACETATE BUFFER
 CH3COOH  CH3COO- + H+
 If more H+ is added to this solution, it simply shifts the
equilibrium to the left, absorbing H+, so the [H+]
remains unchanged.
 If H+ is removed (e.g. by adding OH-) then the
equilibrium shifts to the right, releasing H+ to keep the
pH constant

15. •HANDERSON HASSELBALCH
EQUATION
 Lawrence Joseph Henderson wrote an equation, in 1908,
describing the use of carbonic acid as a buffer solution.
 Karl Albert Hasselbalch later re-expressed that formula
in logarithmic terms, resulting in the
Henderson–Hasselbalch equation.

16. Ka =
[H+] [A-]
[HA]
take the -log on both sides
The Henderson-Hasselbalch Equation derivation
-log Ka = -log [H+] -log
[A-]
[HA]
pH = pKa + log
[A-]
[HA]
= pKa + log
[Proton acceptor]
[Proton donor]
HA H+ + A-
pKa = pH -log [A-]
[HA]
apply p(x) = -log(x)
and finally solve for pH…

17.  - The greater the buffer capacity the less the pH
changes upon addition of H+ or OH-
Choose a buffer whose pKa is closest to the desired
pH.
pH should be within pKa ± 1

18. BUFFER
SYSTEM IN
BODY FLUIDS

19. ACIDS
 VOLATILE ACIDS:
 Produced by oxidative metabolism of CHO,Fat,Protein
 Average 15000-20000 mmol of CO₂ per day
 Excreted through LUNGS as CO₂ gas
• FIXED ACIDS (1 mEq/kg/day)
 Acids that do not leave solution ,once produced they
remain in body fluids Until eliminated by KIDNEYS
Eg: Sulfuric acid ,phosphoric acid , Organic acids
Are most important fixed acids in the body
Are generated during catabolism of:
amino acids(oxidation of sulfhydryl gps of cystine,methionine)
Phospholipids(hydrolysis)
nucleic acids

20. RESPONSE TO ACID BASE CHALLENGE
1. Buffering
2. Compensation

21. BUFFERS
 First line of defence (> 50 – 100 mEq/day)
 Two most common chemical buffer groups
 Bicarbonate
 Non bicarbonate (Hb,protein,phosphate)
 Blood buffer systems act instantaneously
 Regulate pH by binding or releasing H⁺

22. CARBONIC ACID–BICARBONATE BUFFER SYSTEM
Carbon Dioxide
 Most body cells constantly generate carbon dioxide
 Most carbon dioxide is converted to carbonic acid, which dissociates into
H+ and a bicarbonate ion
Prevents changes in pH caused by organic acids and fixed acids in ECF
 Cannot protect ECF from changes in pH that result from elevated
or depressed levels of CO2
 Functions only when respiratory system and respiratory control
centers are working normally
 Ability to buffer acids is limited by availability of bicarbonate ions

23. ACID–BASE BALANCE
The Carbonic Acid–Bicarbonate Buffer System

24. THE CARBONIC ACID HYDROGENCARBONATE
BUFFER SYSTEM
• The carbonic acid-hydrogen Bicarbonate ion buffer is
the most important buffer system.
• Carbonic acid, H2CO3, acts as the weak acid
• Hydrogen carbonate, HCO3
-, acts as the conjugate base
• Increase in H+(aq) ions is removed by HCO3
-(aq)
• The equilibrium shifts to the left and most of the H+(aq)
ions are removed

25.  The small concentration of H+(aq) ions reacts with the
OH-(aq) ions
 H2CO3 dissociates, shifting the equilibrium to the right,
restoring most of the H+(aq) ions
 Any increase in OH-(aq) ions is removed by H2CO3

26. THE HEMOGLOBIN BUFFER SYSTEM
CO2 diffuses across RBC membrane
 No transport mechanism required
As carbonic acid dissociates
Bicarbonate ions diffuse into plasma
In exchange for chloride ions (chloride shift)
 Hydrogen ions are buffered by hemoglobin
molecules
 Is the only intracellular buffer system with an immediate effect
on ECF pH
 Helps prevent major changes in pH when plasma PCO
2
is rising
or falling

27. PHOSPHATE BUFFER SYSTEM
 Consists of anion H2PO4
- (a weak acid)(pKa-6.8)
 Works like the carbonic acid–bicarbonate buffer system
 Is important in buffering pH of ICF
Limitations of Buffer Systems
 Provide only temporary solution to acid–base
imbalance
 Do not eliminate H+ ions
 Supply of buffer molecules is limited

28. RESPIRATORY ACID-BASE CONTROL
MECHANISMS
 When chemical buffers alone cannot prevent
changes in blood pH, the respiratory system is the
second line of defense against changes.
 Eliminate or Retain CO₂
 Change in pH are RAPID
 Occuring within minutes
 PCO₂ ∞ VCO₂/VA

29. 29
PHOSPHATE BUFFER SYSTEM
 The phosphate buffer system (HPO4
2-/H2PO4
-)
plays a role in plasma and erythrocytes.
 H2PO4
- + H2O ↔ H3O+ + HPO4
2-
 Any acid reacts with monohydrogen phosphate
to form dihydrogen phosphate
dihydrogen phosphate monohydrogen phosphate
 H2PO4
- + H2O ← HPO4
2- + H3O+
 The base is neutralized by dihydrogen phosphate
dihydrogen phosphate monohydrogen phosphate
 H2PO4
- + OH- → HPO4
2- + H3O+

30. RENAL ACID-BASE CONTROL
MECHANISMS
 The kidneys are the third line of defence against
wide changes in body fluid pH.
 movement of bicarbonate
 Retention/Excretion of acids
 Generating additional buffers
 Long term regulator of ACID – BASE balance
 May take hours to days for correction

31. RENAL REGULATION OF ACID BASE BALANCE
 Role of kidneys is preservation of body’s bicarbonate
stores.
 Accomplished by:
 Reabsorption of 99.9% of filtered bicarbonate
 Regeneration of titrated bicarbonate by excretion of:
Titratable acidity (mainly phosphate)
Ammonium salts

32. 32
PROTEINS AS A BUFFER
 Proteins contain – COO- groups, which, like acetate ions
(CH3COO-), can act as proton acceptors.
 Proteins also contain – NH3
+ groups, which, like
ammonium ions (NH4
+), can donate protons.
 If acid comes into blood, hydronium ions can be
neutralized by the – COO- groups
 - COO- + H3O+ → - COOH + H2O
 If base is added, it can be neutralized by the – NH3
+
groups
- NH3
+ + OH- → - NH2 + H2O

33. TITRATABLE ACIDITY
 Occurs when
secreted H+
encounter & titrate
phosphate in tubular
fluid
 Refers to amount of
strong base needed
to titrate urine back to
pH 7.4
 40% (15-30 mEq) of
daily fixed acid load
 Relatively constant
(not highly adaptable)

34. RENAL REABSORPTION OF
BICARBONATE
Proximal tubule:
70-90%
Loop of Henle:
10-20%
Distal tubule and
collecting ducts:
4-7%

35. FACTORS AFFECTING RENAL
BICARBONATE REABSORPTION
 Filtered load of
bicarbonate
 Prolonged changes in
pCO2
 Extracellular fluid
volume
 Plasma chloride
concentration
 Plasma potassium
concentration
 Hormones (e.g.,
mineralocorticoids,
glucocorticoids)

36. If secreted H+ ions combine with filtered
bicarbonate, bicarbonate is reabsorbed
If secreted H+ ions combine with
phosphate or ammonia, net acid excretion
and generation of new bicarbonate occur

37. NET ACID EXCRETION
 Hydrogen Ions
Are secreted into tubular fluid along
 Proximal convoluted tubule (PCT)
 Distal convoluted tubule (DCT)
 Collecting system

38. AMMONIUM EXCRETION
 Occurs when
secreted H+
combine with NH3
and are trapped as
NH4
+ salts in
tubular fluid
 60% (25-50 mEq)
of daily fixed acid
load
 Very adaptable (via
glutaminase
induction)

39. AMMONIUM EXCRETION
 Large amounts
of H+ can be
excreted
without
extremely low
urine pH
because pKa
of NH3/NH4
+
system is very
high (9.2)

40. ACID–BASE BALANCE DISTURBANCES
Interactions among the Carbonic Acid–Bicarbonate Buffer
System and Compensatory Mechanisms in the Regulation of
Plasma pH.

41. FOUR BASIC TYPES OF IMBALANCE
Metabolic Acidosis
Metabolic Alkalosis
Respiratory Acidosis
Respiratory Alkalosis

42. ACID BASE DISORDERS
Disorder pH [H+] Primary
disturbance
Secondary
response
Metabolic acidosis    [HCO3
-]  pCO2
Metabolic alkalosis    [HCO3
-]  pCO2
Respiratory acidosis    pCO2  [HCO3
-]
Respiratory alkalosis    pCO2  [HCO3
-]

43. METABOLIC ACIDOSIS
 Primary AB disorder
 ↓HCO₃⁻ → ↓ pH
 Gain of strong acid
 Loss of base(HCO₃⁻)

44. CAUSES OF METABOLIC ACIDOSIS
 LACTIC ACIDOSIS
 KETOACIDOSIS
 Diabetic
 Alcoholic
 Starvation
 RENAL FAILURE (acute
and chronic)
 TOXINS
 Ethylene glycol
 Methanol
 Salicylates
 Propylene glycol

45. ACID–BASE BALANCE DISTURBANCES
.
Responses to Metabolic Acidosis

46. METABOLIC ACIDOSIS
 Symptoms are specific and a result of the underlying pathology
 Respiratory effects:
 Hyperventilation
 CVS:
 ↓ myocardial contractility
 Sympathetic over activity
 Resistant to catecholamines
 CNS:
 Lethargy, disorientation,stupor,muscle twitching, COMA,
CN palsies
 Others : hyperkalemia

47. METABOLIC ALKALOSIS
↑ pH due to ↑HCO₃⁻ or ↓acid
 Initiation process :
 ↑in serum HCO₃⁻
 Excessive secretion of net daily production of fixed
acids
 Maintenance:
 ↓HCO₃⁻ excretion or ↑ HCO₃⁻ reclamation
 Chloride depletion
 Pottasium depletion
 ECF volume depletion
 Magnesium depletion

48. CAUSES OF METABOLIC ALKALOSIS
I. Exogenous HCO3 − loads
A. Acute alkali administration
B. Milk-alkali syndrome
II. Gastrointestinal origin
1. Vomiting
2. Gastric aspiration
III. Renal origin
1. Diuretics
2. Posthypercapnic state
3. Hypercalcemia/hypoparathyroidism
4. Recovery from lactic acidosis or ketoacidosis
5. Nonreabsorbable anions including penicillin, carbenicillin
6. Mg2+ deficiency
7. K+ depletion

49. COMPENSATION FOR METABOLIC ALKALOSIS
 Respiratory compensation: HYPOVENTILATION
 ↑PCO₂=0.6 mm  pCO2 per 1.0 mEq/L ↑HCO3
-
 Maximal compensation: PCO₂ 55 – 60 mmHg
 Hypoventilation not always found due to
 Hyperventilation
 due to pain
 due to pulmonary congestion
 due to hypoxemia(PO₂ < 50mmHg)

50. ACID–BASE BALANCE DISTURBANCES
.
Metabolic Alkalosis

51. METABOLIC ALKALOSIS
 Decreased myocardial contractility
 Arrythmias
 ↓ cerebral blood flow
 Confusion
 Mental obtundation
 Neuromuscular excitability

52. RESPIRATORY ACIDOSIS
 ↑ PCO₂ → ↓pH
 Acute(< 24 hours)
 Chronic(>24 hours)

53. COMPENSATION IN RESPIRATORY ACIDOSIS
Acute resp.acidosis:
 Mainly due to intracellular buffering(Hb,Pr,PO₄)
 HCO₃⁻ ↑ = 1mmol for every 10 mmHg ↑ PCO₂
 Minimal increase in HCO₃⁻
 pH change = 0.008 x (40 - PaCO₂)
Chronic resp.acidosis
 Renal compensation (acidification of urine & bicarbonate retention)
comes into action
 HCO₃⁻ ↑= 3.5 mmol for every 10 mm Hg ↑PCO₂
 pH change = 0.003 x (40 - PaCO₂)
 Maximal response : 3 - 4 days

54. ACID–BASE BALANCE DISTURBANCES
Respiratory Acid–Base Regulation.

55. ACID–BASE BALANCE DISTURBANCES
Respiratory Acid–Base Regulation.