The document discusses pH, buffers, and acid-base balance in the body. It provides information on: - The definition of pH and how it relates to hydrogen ion concentration. pH ranges from 0-14 with lower values being more acidic and higher more basic. - Common buffers in the body including the bicarbonate buffer system, hemoglobin buffer system, and phosphate buffer system which help regulate pH. - How the body maintains acid-base balance through buffering, respiratory compensation, and renal regulation of bicarbonate and acid excretion. - The four main types of acid-base imbalances: metabolic acidosis, metabolic alkalosis, respiratory acidosis, and respiratory alk

1. pH and Buffers
DR HAJA RAMATULAI WURIE

2. References
Harper's Illustrated Biochemistry

3. pH
It is the negative log of the hydrogen ion
concentration pH = -log [H+ ]
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
The pH scale corresponds to the
concentration of hydrogen ions.

5. Acid – base balance
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
Normal pH : 7.35-7.45
Acidosis: Physiological state resulting from abnormally low plasma Ph
Acidemia: plasma pH < 7.35
Alkalosis Physiological state resulting from abnormally high plasma
Alkalemia: plasma pH > 7.45
When an acid reacts with a base, they undergo neutralization:

6. Some Common Acids and Bases and their
Household Uses

7. Strong and Weak Acids
A strong acid dissociates completely into ions in water:
HA+ H2O(l) → H3O+(aq) + A-(aq)
A dilute solution of a strong acid contains no HA molecules.
A weak acid dissociates slightly to form ions in water:
HA(aq) + H2O(l) H3O+(aq) + A-(aq)
In a dilute solution of a weak acid, most HA molecules are
undissociated.

8. Strong acid: HA + H2O(l) → H3O+(aq) + A-(aq)
The extent of dissociation for strong acids.
There are no HA molecules in solution.

9. The extent of dissociation for weak acids.
Weak acid: HA(aq) + H2O(l) H3O+(aq) + A-(aq)
Most HA molecules are undissociated.

10. The Acid Dissociation Constant, Ka
The value of Ka is an indication of acid strength.
Stronger acid larger Ka
higher [H3O+]
Weaker acid smaller Ka
lower % dissociation of HA
HA(aq) + H2O(l) H3O+(aq) + A-(aq)

11. Functional Groups That Are Weak Acids Have
Great Physiologic Significance
Many biochemicals possess functional groups that are weak acids or bases
Carboxyl groups, amino groups, and phosphate esters, whose second dissociation falls within
the physiologic range, are present in proteins and nucleic acids, most coenzymes, and most
intermediary metabolites.
Knowledge of the dissociation of weak acids and bases thus is basic to understanding the
influence of intracellular pH on structure and biologic activity.

12. MEASUREMENT OF PH
The pH can be measured by: pH strips, pH indicators pH meter

13. Buffers
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.

14. 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)
How do buffers work?
Equilibrium between acid and base.
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
The greater the buffer capacity the less the pH changes upon addition of H+ or OH-

15. Buffer systems in body fluids

16. 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
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

17. RESPONSE TO ACID BASE CHALLENGE
Buffering
Compensation
BUFFERS
First line of defence
Two most common chemical buffer groups
Bicarbonate
Non bicarbonate (Hb,protein,phosphate)
Blood buffer systems act instantaneously
Regulate pH by binding or releasing H⁺

18. 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
CARBONIC ACID–BICARBONATE BUFFER
SYSTEM
 Consists of carbonic acid, a weak acid, and
the bicarbonate anion, its conjugate base
 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

19. CARBONIC ACID–BICARBONATE BUFFER SYSTEM
If a strong acid is introduced in the system, which is
equivalent to having an increased concentration of hydronium
ions, it will react with the bicarbonate anion and form
carbonic acid, a weak acid.
The equilibrium will thus shift to the left.
The fact that a strong acid is converted to a weak one will
prevent to acidity of the solution to increase significantly.
Likewise, if a strong base is introduced, it will react with the
carbonic acid to form the bicarbonate anion, thus reducing the
potential increase in pH.
The equilibrium will shift right.
This buffer is actually used by the body to regulate blood
acidity.

20. THE HEMOGLOBIN BUFFER SYSTEM

21. THE HEMOGLOBIN BUFFER SYSTEM

22. Reading assignment
Buffers in the blood during exercise

23. 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 the phosphate buffer system
Provide only temporary solution to acid–base imbalance
Do not eliminate H+ ions
Supply of buffer molecules is limited

24. 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

25. RENAL ACID-BASE CONTROL MECHANISMS
The kidneys are the third line of defense 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
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

26. Proteins as buffers

27. ACID–BASE BALANCE DISTURBANCES
FOUR BASIC TYPES OF IMBALANCE
Metabolic Acidosis
Metabolic Alkalosis
Respiratory Acidosis
Respiratory Alkalosis

28. Reading assignment
Write short notes on the causes and compensations for the four types of imbalances