Acid - base equilibrium

2

pH concept

pH = - log [H+]

pH [H+]pH [H+]

1 100 mmol/L1 100 mmol/L

= 90 mmol/L= 90 mmol/L

2 10 mmol/L2 10 mmol/L

= 9 mmol/L= 9 mmol/L

3 1 mmol/L3 1 mmol/L

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pH = - log [H+]

4

pH

pH of capillary blood norm: 7,35 – 7,45 Sorensen unit

44,7 nmol/L – 35,5 nmol/L

pathology: 6,9 – 7,7 Sorensen unit

126 nmol/L – 20 nmol/L

pH = - log [H+]

pH [H+]

7,4 – 7,1 (0,3 unit) 40 – 80 nmol/L ( = 40 nmol/L) 7,4 – 7,7 (0,3 unit) 40 – 20 nmol/L ( = 20 nmol/L)

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Buffer solutionsMixtures:

weak acid and its salt with strong base

CH3COOH+CH3COO-

weak base and its salt with strong acid

NH3 aqueous

+NH4+

two salts of polyprotic acids

-H2PO4-

+-HPO42-

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Henderson – Hasselbalch equationAH A- + H+

[A-] [H+] Ka = [AH]

[A-] [H+] - log Ka = - log

[AH]because - log Ka = pKa

[A-] [H+] pKa = - log [AH]

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Henderson - Haselbalch equation

[A-] [H+]

pKa = - log [AH]

AH A- + H+

pKa = - log [A-] – log [H+] + log [AH]-log [H+] = pKa + log[A-] – log [AH]

-log [H+] = pH

[A-] pH = pKa + log [AH]

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Henderson – Hasselbalch equation

[A-]

pH = pKa + log [AH]

pH of the buffer mixture depends on:

kind of the acid ratio of the concentrations of buffer components

pH of the buffer mixtures does not change with dilution of the solution!

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Buffer solutions – mechanism of action

acetate buffer: CH3COOH i CH3COO-

at equilibrium: CH3COOH + H2O CH3COO- + H3O+

CH3COOH CH3COO- + H+

acetic acid is a weak acid, acetate ions come entirely from salt (acetate)

add some amount of strong acid: +H+

CH3COOH CH3COO- + H+

strong acid removes weak acid from its salt

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Buffer capacity

c =

pH – buffer capacity

c – amount of strong acid or strong base added to the buffer solution mol/L

pH – observed change of pH

Buffer capacity depends on concentration of components:

increases with their increment decreases with dilution of the buffer

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Buffer capacity

is the highest when pH = pK

Excess of acid in the buffer – better buffering of bases

Excess of a salt – better buffering of acids.

During addition of acid or base buffer capacity decreases is equal to zero when whole salt in the buffer is converted to weak

acid or whole weak acid is converted to salt.

c =

pH

[A-]

pH = pKa + log

[AH]

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Intracellular pH

cytoplasm 6,0

mitochondria endoplasmatic reticulum 7,0 – 7,4 nucleus

average pH of intracellular fluid:

6,95 (112 nmol/L)

differences in pH of extracellular fluid between cells from different organs:

erytrocytes 7,20 renal tubular epithelial cells 7,32 cells of skeletal muscles. 6,9

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Mechanisms of pH regulation in organism

Organ’s regulation regulation by kidney regulation by lungs regulation by bones

Buffer regulation proteinate buffer

H-proteins Proteinates

phosphate buffer-H2PO4-

-HPO42-

bicarbonate buffer

HCO3- H2CO3

Hemoglobin buffer

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Hemoglobin buffer is the most important proteinate buffer

1. Hemoglobin makes about ¾ of all blood’s protein

2. Hemoglobin has an acidic character because of the presence of majority of acidic groups of hem over basic groups of globins. Therefore, hemoglobin (Hb) has a large capacity for base binding.

3. Acidity of hemoglobin changes significantly with degree of oxidation.

Hemoglobin transports not only oxygen from lungs to cells but also carbon dioxide from cells to lungs.

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Proteinate buffer-hemoglobin buffer

Hemoglobine buffer is an intracellular buffer.

The Hb and oxygen connection is reversible and the Hb molecule changes

its conformation when it binds O2.

HHb (Hb without oxygen) and HHbO2 (Hb connected with oxygen, oxy-

hemoglobin) differ in their ability to donate or accept H+ ions.

HbO2 + H+ HHb + O2 In tissues

stronger acid weaker acid

HHb + O2 + HCO3- HbO2 + H2O + CO2 In Lungs

Buffering properties depend on the equilibrium between oxy-hemoglobinand hemoglobin.Hemoglobin is less acidic than oxy-hemoglobin.With increasing CO2 pressure and H+ concentration, the amount of HHb increasesand buffering capacity increases as well.

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Proteinate buffer

In acidic environment: carboxyl groups of amino acids do not dissociate basic groups (amine, imidasole) are proton acceptors

In basic environment: carboxyl groups are proton donors and they neutralize hydroxide ions in weak basic environment - pH 7,4, proteins are in form of anions.

concentration of proteins in the blood is 16 mEq/l

Buffer capacity of proteinate buffer is:5 mmols/ l/one pH unit

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Proteinate Buffer – hemoglobin buffer

I system II system

HHbO2 HHb KHbO2 KHb

hemoglobin is an important buffer in blood hemoglobin makes ¾ of whole proteins in the blood. hemoglobin is acidic because of excess of acidic groups of

heme over basic groups from globin. acidity of hemoglobin can change depending on oxygenation It is indispensable for full effectiveness of carbonate buffer

in an open system

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Phosphate buffer – important urinary buffer

phosphate buffer in blood :

H2PO4- HPO4

2- + H+

KH2PO4- = 6,2 x 10-8, pK2 = 7,21

HPO42- / H2PO4- = 4/1

phosphate buffer in urine:

HPO42- / H2PO4

- = 1/4

HPO42- + H+ H2PO4-

main intracellular buffer optimal pK for this buffer is

6,8. in phosphate buffer in urine

(pH about 6,0) ratio of hydrogen phosphate to dihydrogen phosphate is 1/4

Difference between ratios of phosphates in blood and urine is because hydrogen phosphate of urine binds protons secreted by distal tubules of kidneys and is converted to dihydrogen phosphate.

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Bicarbonate buffer – an extracellular buffer; in equilibrium with atmospheric air.

The most important buffer sytem in the blood is:

HCO3-/H2CO3

organism removes by lungs a product of dehydration of carbonic acid - carbon dioxide.

This buffer acts in an open system.

H2CO3 i CO2 dissolved in water phase remain at equilibrium with CO2 which is in gasous phase.

CO2 in blood circulating in lungs remains at equilibrium with CO2 in lung vesicles.

buffer in an open system has several times higher capacity in comparison to a closed system

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Bicarbonate buffer is responsible for the physiological pH of blood, 7.35-7.45

CO2 + H2O H2CO3 H+ + HCO3-

99% 1%

gas liquid liquid liquid liquid

% - expresses amounts in equilibrium state

[A] pH = pKH2CO3 + log [ AH]

[ HCO3- ]

pH = pKH2CO3 + log

[CO2 ]

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Gasometric parameters of blood (physiological levels):

pKH2CO3 = 6,11

[HCO3- ] = 24 mmol/L

[CO2 ] = x p

– coefficient of solubility of CO2 in plasma

= 0,225 mmol/L

p – partial pressure of CO2 in lung vessels

pCO2 = 5,32 kPa

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Bicarbonate buffer

[ HCO3- ] pH = pKH2CO3 + log

[CO2 ]

pKH2CO3 = 6,11[HCO3

- ] = 24 mmol/L[CO2 ] = x p = 0,225 mmol/LpCO2 = 5,32 kPa

pH of bicarbonate buffer in blood

24 pH = 6,11 + log = 7,4 0,225 x 5,32

CO2 + H2O H2CO3 H+ + HCO3-

99% 1%

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Bicarbonate buffer

CO2 + H2O H2CO3 H+ + HCO3-

+H+

Bicarbonate buffer – mechanism of action

CO2 + H2O H2CO3 H+ + HCO3-

99% 1%

The kidneys regulate HCO-3 ion concentration in blood plasma and protect the

organism against metabolic acidosis.The role of kidneys is:• reabsorption of HCO-

3 ions in kidney tubules

• excretion of HCO-3 ions when their level changes in extracellular fluid

• regeneration of lost HCO-3 ions or in reactions with H+ derived from water

(in proximal tubular cells and collecting tubule cells)

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Bicarbonate buffer

most impotant buffer in blood, in acid-base balance acts in open system in normal conditions HCO3/CO2 is 20:1

metabolic component

[HCO3- ]

pH = pKH2CO3

+ log pCO2 x

respiratory component

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Disturbances of acid-base balance

[HCO3- ] metabolic component

pH = pKH2CO3

+ log pCO2 x respiratory component

Too much acid in the body resulting from accumulation of acid or depletion of alkaline reserves leads to acidosis – abnormally low pH.

May be caused by : diabetic ketoacidosis, lung disease, kidney disease.

The condition opposite to acidosis is – alkalosis – pH is to high due toexcess base or insufficient acid in the body.

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Disorder pH HCO3- pCO2

Metabolic acidosis

Respiratory acidosis

Metabolic alkalosis

Respiratory alkalosis

primary changes

secondary changes

[HCO3- ]

pH = pKH2CO3 + log

pCO2 x

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Bicarbonate bufferTo 1 liter of plasma 10mmols of strong acid were added . Calculate pH change

when the system is:• open • closed. pKH2CO3 = 6,11; [HCO3- ] = 24 mmol/L; = 0,225 mmol/l /kPapCO2 = 5,32 kPa; pH=7,4

Open system

24 HCO3- + 10 H+ -------> 10 CO2 + 10 H2O + 14 HCO3-

24 – 10 14

pH = 6,11 + log = 6,11 + log --------- = 7,17 0,225 x 5,32 1,2

Closed system

24 HCO3- + 10 H+ -------> X 10 CO2 + 10 H2O + 14 HCO3

-

24 – 10 14

pH = 6,11 + log = 6,11 + log --------- = 6,2 1,2 + 10 11,2

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Red blood cells’ role in acid-base regulation of blood . Reaction in capillary tissues

Erythrocyte

HbO2- + H+ H-Hb + O2

O2

To tissues

CO2 + H2O H2CO3 H+ + HCO3-

HCO3-

Cl-

Cl-Offset chloride

carbonic anhydrase

CO2

From tissues

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Red blood cells role in blood’s acid-base regulation; reaction in lung capillaries

Erythrocyte

H-Hb + O2 HbO2- + H+

O2 from lungs

HCO3 + H+ H2CO3 H2O + CO2

HCO3-

Cl-

Cl-

carbonic anhydrase

CO2

Exhaled

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THE END