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
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.
11.
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
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
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
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)
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)
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)