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Acid - base
equilibrium
pH concept
pH = - log [H+]
pH
1
[H+]
100 mmol/L
D = 90 mmol/L
2
10 mmol/L
D = 9 mmol/L
3
1 mmol/L
2
pH = - log [H+]
3
pH = - log [H+]
pH
pH of capillary blood
norm:
7,35
–
7,45 Sorensen unit
–
7,7
44,7 nmol/L – 35,5 nmol/L
pathology:
6,9
126 nmol/L
pH
7,4 – 7,1
7,4 – 7,7
(0,3 unit)
(0,3 unit)
–
Sorensen unit
20 nmol/L
[H +]
40 – 80 nmol/L (D = 40 nmol/L)
40 – 20 nmol/L (D = 20 nmol/L)
4
Buffer solutions
Mixtures:

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 equation
AH
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
AH
A- + H+
[A-] [H+]
pKa = - log
[AH]
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
CH3COOH
CH3COO- + H3O+
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
Dc
b = 
DpH
b–
buffer capacity
Dc – amount of strong acid or strong base added to the buffer
solution mol/L
DpH – 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
Dc
b = 
DpH
[A-]
pH = pKa + log
[AH]

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.
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Intracellular pH
 cytoplasm 6,0
 mitochondria
 endoplasmatic reticulum
 nucleus
7,0 – 7,4
 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
cells of skeletal muscles. 6,9
7,32
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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
HCO3H2CO3

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, oxyhemoglobin) differ in their ability to donate or accept H+ ions.
HbO2 + H+
stronger acid
HHb + O2
In tissues
weaker acid
HHb + O2 + HCO3-
HbO2 + H2O + CO2
In Lungs
Buffering properties depend on the equilibrium between oxy-hemoglobin
and hemoglobin.
Hemoglobin is less acidic than oxy-hemoglobin.
With increasing CO2 pressure and H+ concentration, the amount of HHb increases
and buffering capacity increases as well.
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Proteinate buffer-hemoglobin buffer
is based on binding of H + ions within capillary vessels of all
cells and its release in capillary vessels in lungs.
 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
HHbO2
KHbO2
II system
HHb
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 :


H2PO4K
H2PO4-
= 6,2 x
10-8,
HPO42- + H+
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 CO 2
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
99%
gas
liquid
H+ + HCO3-
H2CO3
1%
liquid
liquid
liquid
% - expresses amounts in equilibrium state
[A]
pH = pKH CO + log
2
3
[ AH]
[ HCO3- ]
pH = pKH2CO3 + log
[CO2 ]
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Gasometric parameters of blood
(physiological levels):
pKH2CO3 = 6,11
[HCO3- ] = 24 mmol/L
[CO2 ] = a
x
p
a – coefficient of solubility of CO 2 in plasma
a = 0,225 mmol/L
p – partial pressure of CO2 in lung vessels
pCO2 = 5,32 kPa
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Bicarbonate buffer
CO2 + H2O
99%
H+ + HCO3-
H2CO3
1%
[ HCO3- ]
pH = pKH2CO3 + log
[CO2 ]
pKH2 CO3 = 6,11
[HCO3 - ] = 24 mmol/L
[CO2 ] = a x p
a = 0,225 mmol/L
pCO2 = 5,32 kPa
pH of bicarbonate buffer in blood
24
= 7,4
pH = 6,11 + log
0,225 x 5,32
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Bicarbonate buffer
CO2 + H2O
99%
H2CO3
1%
H+ + HCO3-
Bicarbonate buffer – mechanism of action
+H+
CO2 + H2O
H2CO3
H+ + HCO3-
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 blod, in acid-base balance
 acts in open system
 in normal conditions HCO3/CO2 is 20:1
metabolic component
[HCO3- ]
pH = pKH2CO3 + log
pCO2 x a
respiratory component
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Disturbances of acid-base balance
pH = pKH2CO3 + log
[HCO3- ]
pCO2 x a
metabolic component
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 to
excess 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
pH = pKH2CO3 + log
[HCO3- ]
pCO2 x a
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Bicarbonate buffer
To 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; a = 0,225 mmol/l /kPa
pCO2 = 5,32 kPa; pH=7,4
Open system
24 HCO3- + 10 H+ -------> 10 CO2 + 10 H2O + 14 HCO3-
pH = 6,11 + log
24 – 10
0,225 x 5,32
14
= 6,11 + log --------- = 7,17
1,2
Closed system
24 HCO3- + 10 H+ -------> X 10 CO2 + 10 H2O + 14 HCO324 – 10
pH = 6,11 + log
1,2 + 10
14
= 6,11 + log --------- = 6,2
11,2
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Red blood cells’ role in acid-base regulation
of blood . Reaction in capillary tissues
O2
Erythrocyte
HbO2- + H+
CO2 + H2O
To
tissues
H-Hb + O2
H2CO3
H+ + HCO3-
carbonic anhydrase
Cl-
HCO3Cl-
CO2
From tissues
Offfset chloride
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Red blood cells role in blood’s acid-base
regulation; reaction in lung capillaries
O2 from lungs
Erythrocyte
H-Hb + O2
HCO3 + H+
Cl-
HbO2- + H+
H2CO3
H2O + CO2
carbonic anhydrase
HCO3- Cl-
CO2
Exhaled
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THE END
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