2. Introduction:
Molecules are formed by interaction of
positively and negatively charged ions
through the formation of ionic bonds.
when certain compounds are dissociated
in water they contribute either H+ which
referred to an acid or OH- which referred
to a base in a solution, and that
compound which contribute neither are
called salts.
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3. About 50 to 100 mmol of hydrogen ions are
released from cells into extra cellular fluid
each day either from metabolism of amino
acids or incomplete metabolism of organic
compounds. The body has a wide array of
mechanisms to maintain homeostasis in the
blood and extracellular fluid.
The body yield H+ ions more than OH- ions
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4. Definitions:
(1) Acid : is a substance which donate H+ ions .
(2) Base : is a substance which accept H+ ions or
donate OH- ions.
(3) Neutral solution: have equal concentration of H+
ions & OH- ions.
(4) pH : is an expression of the H+ ions concentration
in a solution. pH is a measure of the acidity or
basicity of a solution.
pH is an abbreviation for "power of hydrogen"
where "p" is short for the German word for power,
potenz and H is the element symbol for hydrogen.
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5. It is typically defined as the negative logarithm of
the hydrogen ion concentration, & the pH scale
uses the no. 0 – 14 to describe the acidity or
alkalinity of a solution.
0……Acid……7……alkaline………14
The equation used to determine the pH of a
solution is stated as
pH = - log [H+]
where [H+] is the concentration of hydrogen in mol/l.
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6. Examples
1.What is the pH of a solution with [H+] = 1 x 10-6 M.
pH = - log [H+]
pH = - log (1 x 10-6)
pH = -(-6)
pH = 6
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7. Example
2. when [H+] = 4.0 x 10¯ 8 mol/l
then pH = (-log 4.0) + (-log 10¯8)
= -0.6 + 8
= 7.4
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8. Definitions cont:
The normal concentration of H+ in the extracellular
body fluid ranges from 36–44 nmol/L ( pH lies between 7.35 to
7.45)
(*) The most important way that the pH of the blood is kept
relatively constant is by buffers dissolved in the blood.
(5)A buffer is a solution containing substances which have the
ability to minimize changes in pH when an acid or base is
added to it.
A buffer typically consists of a solution which contains a
weak acid mixed with the salt of the a strong base
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9. The Major Body Buffer Systems
(1) Proteins: more important buffering
capacity intracellular because it found in
high concentration.
(2) Hemoglobin: also acts as a pH buffer in
the blood that hemoglobin protein can
reversibly bind either H+ (to the protein)
or O2 (to the Fe of the heme group) , but
that when one of these substances is
bound, the other is released.
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10. The Major Body Buffer Systems
cont
(3) Phosphate buffers: The phosphate buffer
consists of phosphoric acid (H3PO4) in
equilibrium with dihydrogen phosphate
ion (H2PO4
_) and H+. The concentration of
phosphate in the blood is so low that it is
quantitatively unimportant. Phosphates are
important buffers intracellularly and in
urine where their concentration is higher.
(4) Ammonia: one of the most important
urinary buffers.
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11. The Major Body Buffer Systems
cont
(5)The Bicarbonate Buffer System: This is responsible for
about 80% of extracellular buffering.
Co2 + H2O →H2CO3 → H+ + HCO3-
The Henderson-Hasselbalch equation expresses the
relation between pH and buffer pair
pH = pK + log base
Acid
pH = pKa + log ([A-]/[HA])
.
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12. cont.
Where,
pK is the disassociation constant
[A-] = molar concentration of a conjugate base
[HA] = molar concentration of a undissociated weak acid (M)
The equation can be rewritten to solve for pOH:
pOH = pKb + log ([HB+]/[ B ])
[HB+] = molar concentration of the conjugate base (M)
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13. cont.
where A is proton acceptor, or base (e.g., HCO3 HA is proton donor, or
weak acid (e.g., H2CO3), and pK’ is pH at which there is an equal
concentration of protonated and unprotonated species.
Knowing any of the three variables allows for the calculation of the
fourth.
In plasma and at body temperature (37°C), the pK of the bicarbonate
buffering system is 6.1. The equilibrium between H2CO3 and CO2 in
plasma is approximately 1:800.
In health, when the kidneys and lungs are functioning properly, a 20:1
ratio of HCO3 to H2CO3 will be maintained (resulting in a pH of 7.40).
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14. Disturbance of hydrogen ion
homeostasis
Acidosis is associated with pH less than the lower limit of the
reference range (7.35).Which may be due to :
Metabolic acidosis : reduction in bicarbonate (HCO3
_)
Respiratory acidosis: rise in PCO2
Alkalosis occurs when pH exceeds the upper limit of the
reference range (7.45), which may be due to:
(a) Metabolic Alkalosis : rise in bicarbonate (HCO3_)
(b) Respiratory Alkalosis: fall in Pco2
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15. Measurement of pH :
- pH meter.
- blood gas analyzer.
- calculation from
Henderson-Hasselbalch
equations.
- pH indicator.
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16. Example Problem Applying the Henderson-
Hasselbalch Equation
Calculate the pH of a buffer
solution made from 0.20 M
HC2H3O2 and 0.50 M C2H3O2
-
that has an acid dissociation
constant for HC2H3O2 of 1.8 x
10-5.
Answer: Use equation:
pH = pKa + log ([A]/[HA])
pH = pKa + log
([C2H3O2
-] /
[HC2H3O2])
pH = -log (1.8 x 10-5) + log
(0.50 M / 0.20 M)
pH = -log (1.8 x 10-5) + log
(2.5)
pH = 4.7 + 0.40
pH = 5.1
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