Dr. Saidunnisa

Professor of BiochemistryAcids, bases, conjugate acid base

pairs, body buffers.

Learning objectives

• At the end of the session student shall be able to :• List the major sources of acids and bases in the body.• Define an acid ; base ; buffer; conjugate acid base pairs with

suitable examples.• List the various buffer systems.• Revise your concept of pH ; pKa and Henderson – hasselbalch

equation• Study how the bicarbonate phosphate protein and

Hemoglobin buffer system works • Explain how the above are linked with respiration and kidneys.• What is meant by the term Alkali reserve.

Introduction

• Under normal conditions the pH of ECF usually does not vary beyond the range 7.35-7.45 and is maintained approximately at 7.4.

• Maintenance of this pH is one of the prime requisites of life and any variation on either side seriously disturbs the vital processes and may lead to death.

• pH less than 7.3 leads to acidosis causing CNS depression , coma and death .

• pH more than 7.5 leads to alkalosis which induces neuromuscular hyper excitability and tetany and death.

• Large amounts of H+ are continuously contributed to ECF from intracellular metabolic reactions. Hence to maintain a constant pH it is necessary that they are removed from the fluids promptly and effectively.

Productions of acids by the Body

• Carbonic acid (H2CO3):- Chief acid produced in the body during

oxidation of carbon compounds• Sulphuric acid (H2So4) :-Produced during oxidation of ‘S’ –

containing aminoacids• Phosphoric acid(H2Po4) :- Produced during metabolism of

phospho proteins and nucleoproteins• Organic acids:- Pyruvic acids, Lactic acids, and ketone bodies

are the intermediates in the metabolism• Iatrogenic :- Antacids• Diet rich in animal protein

Production of bases by the body

• Formation of basic compounds in the body in

normal circumstances is negligible.• Bicarbonate produced from organic acids.• Diet rich in vegetables.

Acid/Base definitions

• Definition #1: Arrhenius (traditional)

Acids – produce H+ ions (or hydronium ions H3O+)

Bases – produce OH- ions

(problem: some bases don’t have hydroxide ions!)

Acid/Base Definitions

• Definition #2: Brønsted – Lowry

Acids – proton donor

Bases – proton acceptor

A “proton” is really just a hydrogen atom that has lost it’s electron!

ACID-BASE THEORIESACID-BASE THEORIES

The Brønsted definition means NH3 is a BASE in water — and water is itself an ACID

BaseAcidAcidBaseNH4

+ + OH-NH3 + H2OBaseAcidAcidBase

NH4+ + OH-NH3 + H2O

Conjugate Pairs

Learning Check!

Label the acid, base, conjugate acid, and conjugate base in each reaction:

HCl + OH-   Cl- + H2O HCl + OH-   Cl- + H2O

H2O + H2SO4   HSO4- + H3O

+ H2O + H2SO4   HSO4- + H3O

+

Conjugate acid-base• A strong acid and a weak base and vice-versa are

called conjugate acid-base pairs.• Acid and Bases can be :• Strong: dissociate completely• HCL H+ + Cl- (complete)

Strong weak baseAcid • The concentration of H+ is very high in strong acid.

– Weak – dissociate only partially in solution• carbonic acid , Lactic acid,

• H2CO3 H+ + HCO3- (partially)

Weak acid• The number of acid molecules existing may be

only 50%.

• The weak acids exist in two forms: The dissociated form will release H+ ions

and the conjugate base For e.g acetic acid, the dissociation will give

rise to free H+ ions and acetate ions The undissociated form existing as acetic acid.• CH3COOH CH3COO- + H+

pKa- Definition

• The pH at which the rate of forward reaction (dissociation) is in equilibrium with the rate of backward reaction. It is designated as pKa or dissociation constant.

• CH3COOH CH3COO- + H+• Ka = CH3COO- + H+• CH3COOH (undissociated molecules)• Strong acids have low pKa and weak acids have

higher pKa.

pKa-Importance

• The ratio between the dissociated and un dissociated form is equal to 1 at a pH equal to its pKa.

• Example in the case of acetic acid, at a pH of 4.76 (pKa), the amount of dissociated form is equal to the un dissociated form.

The combination of these two forms is buffer.

Acidity of the solution

• Is equal to the ratio of the activities of the acid to the base multiplied by the dissociation constant.

• [H+] = Ka [Acid] or [HA]• [Base] [A-]• Ka = Dissociation constant.

pH- Definition

• Sorensen expressed the H+ concentration as the negative logarithm of hydrogen ion activity and is designated as the pH.

• pH = -log [H+]• pH value is inversely proportional to the

acidity.

Henderson-Hasselbalch equation

• Effects of salt upon the dissociation of acids.• Relationship between pH, pKa, concentration of

acid and base expressed by Henderson-Hasselbalch equation.

• pH = pKa + log [Base] or pH = pKa + log [Salt] • [Acid] [Acid]

• when the concentration of the base and the acid are the same then pH= pKa .

Indicators

• Weak organic acids or bases• They change in colour during pH change.• The change is over a specific range only.• Each indicator exists as conjugate pair• Each member differs sharply in colour.• It follows the law of Handerson –Hasselbalch

equation.

• An average rate of metabolic activity produces roughly 22,000mEq acid per day.

• If all the acids were dissolved at one time in unbuffered body fluids their pH would be less than 1.

• However the pH of blood is maintained between 7.36-7.44.

• How is it possible?

Buffer

• A buffer is defined as a solution which resist the change in pH when an acid or alkali is added.

• HCL + NaHCO3 H2CO3 + NaCl

Strong Buffer weak Neutral Acid Acid salt

• Buffers can resist the change in pH when an acid or alkali is added upto approximately ±pKa 1 .

• Buffers can act quickly but not permanently.• It cannot remove the H+ ions from the body.• It acts as a temporary measure to decrease

the free H+ ions.• Final elimination is through lungs or kidneys.

• The body has developed three mechanisms to regulate the body’s acid –base balance and maintain the ECF at pH 7.4

• BUFFERS• RESPIRATORY MECHANISMS• RENAL MECHANISMS

Body buffers-Distribution

1. Bicarbonate-Carbonic acid buffer

2. Hemoglobin buffer

3. Phosphate buffer

4. Protein buffer

5. Renal cells-Ammonia buffer• Relative distribution of various buffer systems in the body :- • 42% are present in ECF and • 58% in ICF

Buffer systems of the body

Renal cells-Ammonia buffer pKa 9.25

Bicarbonate buffer system

• It is most predominant buffer system of the ECF particularly the plasma

H2CO3 H+ HCO3

-

By the law of mass action pKa= [H+ ] [HCO3-]

[ H2CO3 ]

• pKa = dissociation constant of carbonic acid = 6.1

• The equation may be rewritten as • [H+ ] = pKa [H2CO3]

[HCO3-]

• pH = log 1 [H+] • By taking the reciprocals and logarithms we get pH =pKa +

log [ HCO3-] = [Base]

[H2CO3] [Acid]

This reaction is called as Henderson – Hassel Balch equation which is valid for any buffer pair.

• NaHCO3 = 20 H2CO3 1

• At blood pH 7.4 the ratio of bicarbonate to carbonic

acid is 20:1.• Bicarbonate is much higher i.e 20 times higher than

carbonic acid in blood. • This is referred to as alkali reserve and is responsible

for effective buffering of H+ ions generated in the body.

Neutralization

• Eg:- Neutralization of strong and non-volatile acids such as HCL ,H2SO4 and lactic acid entering the ECF is achieved by bicarbonate buffer .

• Thus lactic acid is buffered as follows:

• Carbonic Acid thus formed is volatile and is eliminated by diffusion of CO2 through alveoli of lungs

• Note: Proper lung functioning is important

hence bicarbonate buffer is linked up with respiration .

• Similarly when alkaline substances e.g NaOH enters the ECF it reacts with the acid component that is H2CO3 of the buffer system.

• Alkali reserve is represented

by the sodium bicarbonate concentration in the blood that has not combined with strong and non-volatile acids.

• Advantages • It is a very good

physiological buffer and acts as a first line of defense because

• It is present in very high concentration than other buffer systems

• Produces carbonic acids which is weak and volatile acid CO2 is exhaled out.

• Disadvantages: • As a chemical buffer it

is rather weak.• pKa is further away

from the physiological pH .

Phosphate Buffer System

• Salt / Acid = ( Na2HPO4/ NaH2PO4)• It is mostly an intracellular buffer and less

important in plasma due to its low concentration with pKa of 6.8 close to blood pH 7.4.

• It would have been more affective had it been present in high concentration.

• Normal ratio in plasma is 4:1• The phosphate buffer system is directly linked up

with the kidneys.

• PH = pKa + log [ salt ] [ acid ] 7.4 = 6.8 + log [ salt ] [ Acid ] 0.6 anti log is 4

Phosphate buffer

• E.g : When strong acid enters the blood it is buffered by base of the phosphate as follows

• When the alkali enters the blood it is buffered by the acid of phosphate as follows

• When the alkali enters the blood it is buffered by the acid of

phosphate as follows• Thus phosphate buffer system works in conjunction with the

kidneys. A normal healthy kidney is necessary for proper functioning

• Advantages:• As a chemical buffer it is very effective as pKa approaches the

physiological pH.• Disadvantage:• As physiological buffer it is less efficient because it is present in

low concentration in the Blood.

Protein Buffer System

• Salt / Acid = Na + Pr- / H+ Pr-. • Buffering capacity of plasma proteins much less than Hb.• Eg: Hb present in one litre if blood can buffer 27.5mEq of H+ ions • Where as plasma proteins can buffer 4.24mEq of H+ ions at pH 7.5 • In acidic medium protein acts as a base amino group takes up H+ ions

from the medium forming NH3+ and proteins become positively charged.

• In alkaline medium protein acts as an acid COOH group dissociates and gives H+ forming COO- so proteins become negatively charged.

Hb as a buffering System

• Buffering capacity of Hb is due to the presence of

imidazole nitrogen group of histidine amionoacid .• 38 histidine residues are present in 1 molecule of hb and

has a pka 6.1.• Oxygenated Hb is Stronger acid than deoxygenated Hb .• On oxygenation; the imidazole N2 group acts as acid and

donates protons in the medium. • Deoxygenated Hb ; the imidazole N2 group acts as a

base and takes up protons from the medium.

• Acidity of the medium favours delivery of oxygen ; alkalinity favours oxygenation of Hb.

• The role of the Hb buffer is considered along

with the respiratory regulation of pH.

Ammonia buffer

• Ammonia is an urinary buffer.• Ammonia is toxic and hence not present in the

blood, but generated in the renal tubules to combine with H+ ions to be excreted as ammonium salts.