ACID BASE DISORDERS
BASIC SCIENCES :
• What is an Acid ?
• An acid is a potential proton (H +ion) DONOR
• What is a base ?
• A base is a potential proton ACCEPTOR
• The first notion of acids comes from ancient Greece, where
people noticed that some substances tasted sour.
• This is also where the word “acid” comes from; it is derived
from the Greek word “oxein” which in Latin is “acere,” meaning
“to make sour.”
• It was also noted that acids and bases could color certain
substances.
• The strength of an acid is measured by the extent to which it
dissociates in an aqueous solution.
• HA ⇋ [H+] + [A-]
• For a strong acid, { [H+] + [A-] } will be in greater concentration
than HA in the undissociated form.
• Also, at equilibrium, the product of concentration on one side of
the equation will bear a constant relation to the product of
concentration on the other side
• Therefore,
• HA ⇋ [H+] + [A-] ………………………………………(1)
• 𝐾𝑎 =
𝐻
+
𝐴
𝐻𝐴
−
……………………………………….(2)
• Ka = acid dissociation constant
• Strong acids = High Ka
CONCEPT OF PH
Søren Sørensen (1868–1939)
• A Ph.D. from the University of Copenhagen, Sørensen was the
director of the chemical department of the Carlsberg
Laboratory, which was supported by the : ????
• Beer company of the same name, brewing being one of the
oldest chemical industries.
• The letters pH are an abbreviate for "pondus hydrogenii"
(translated as potential hydrogen) meaning hydrogen power as
acidity is caused by a predominance of hydrogen ions (H+).
•
Dr. Sørensen has been credited as the founder of the modern
pH concept.
• The acidity of an aqueous solution is measured by it’s hydrogen
ion concentration [ H+] or activity.
• The [ H+] activity is expressed as pH.
• pH = Negative Logarithm of H+ ion concentration
= Negative Logarithm of [ H+] .
• 𝑝𝐻 = 𝐿𝑜𝑔 10
1
𝐻+
…………………………(3)
• pH = ─ log10 [ H+]
• pH = Negative exponent of an expression of [H+ ] to
the power of 10
• Thus ,
• pH of 7.4 = [ H+] x 10 -7.4
• The normal concentration of H+ ion IN BLOOD :
• [H+] in blood = 40 x 10 -9 moles / L
• pH of 7.4 = Hydrogen ion concentration of 40 x 10 -9
moles / L
• Going to equation (2) :
• 𝐾𝑎 =
𝐻
+
𝐴
𝐻𝐴
−
• Rearranging :
•
𝐻+ =
𝐾𝑎 𝐻𝐴
−
𝐴
…………………………(4)
• Taking reciprocals :
•
1
𝐻+
=
1
𝐾𝑎
∗
𝐴−
𝐻𝐴
• Taking Logs
• Log
1
𝐻+
= 𝑙𝑜𝑔
1
𝐾𝑎
+ 𝑙𝑜𝑔
𝐴−
𝐻𝐴
…………………(5)
• Thus :
• Using (3) { 𝑝𝐻 = 𝐿𝑜𝑔
1
𝐻+
=
} in (5){ Log
1
𝐻+
= 𝑙𝑜𝑔
1
𝐾𝑎
+ 𝑙𝑜𝑔
𝐴−
𝐻𝐴
HENDERSON – HASSELBACH EQUATION
}
BUFFERS
• What are Buffers ?
• Buffers are agents whose function is to prevent a rapid change
in pH when acids or bases are added to the solution.
• The main buffer systems in the body are :
•
•
•
•
1. Haemoglobin
2. Bicarbonates and inorganic phosphates within the blood
3. Plasma proteins
4. Tissue proteins
• The bicarbonate – Carbonic acid reaction is important in
regulating acid – base balance.
•
It constitutes a weak buffer system which can be used as a
measure or reflection of all acid – base reactions within the
body.
• The CO2 produced by tissue metabolism diffuses into plasma
and the RBC’s.
• The enzyme carbonic anhydrase within the RBC’s catalyzes the
formation of carbonic acid.
• H2O + CO2
→ H2CO3
• Carbonic acid within the RBC’s is dissociated thus :
• H2CO3 →
[H+ ] + [HCO3- ]
• The bicarbonate within the RBC’s diffuses out into the plasma,
while the H+ ions are mopped up by the haemoglobin which
acts as a buffer base.
• The loss of one ion from the cell has to be compensated by the
entry of an equivalent ion.
• Thus, the chloride ions from the plasma enter the blood cells in
place of the bicarbonate ions which have diffused into the
plasma.
• When the blood reaches the lungs, the chloride shift is reverse
and bicarbonate enters the blood cells.
• The bicarbonate within the RBC’s breaks down into H2O an
CO2.
• The CO2 diffuses out through the capillaries into the alveoli and
is washed out into the outside air.
• H+ + HCO3- → H2CO3 → CO2 + H2O
• The carbon dioxide is rapidly washed out via the lungs leading
to a rapid control of [H+] ion production but at the expense of a
fall in bicarbonate.
• The Henderson – Hasselbach equation for the bicarbonate –
carbonic acid system :
−
• pH = pK carbonic acid + log
𝐻𝐶𝑂3
𝐻2𝐶𝑂3
−
• pH = 6.1 + log
𝐻𝐶𝑂3
………………………….(6)
0.03𝑥𝑃𝐶𝑂2
• Where :
0.03 = solubility of CO2 in plasma
PCO2 = partial pressure of CO2 in plasma
6.1 = negative logarithm of dissociation constant of
carbonic acid
To be continued………