26. Buffers Flashcards
What is the definition of a buffer?
An acid–base buffer solution
resists a change of pH
when an acid or base is added to it.
It consists of a weak acid and its conjugate base (salt).
The general equation for a buffer system is:
HA (undissociated acid)
↔ H+ (hydrogen ion) + A¯ (conjugate base)
> Le Chatelier’s principle states that if H+ ions are added to the solution, the equilibrium shifts to the left, and the H+ ions are ‘neutralised’ by the conjugate base, minimising an increase in free [H+]
and maintaining a constant pH.
If a base is added,
H+ and OH− react to form water,
but more HA dissociates to
maintain the [H+] constant,
therefore the equation
shifts to the right.
> By applying the Law of Mass Action,
(Ka = [H+] [A¯ ]/[HA]),
the Henderson–Hasselbach equation can
be derived and the (pKa) for a
buffer system can be calculated:
pH = pKa + log [conjugate base]/[acid]
pH = pKa + log [A–]/[HA]
(where Ka is the dissociation constant of a buffer and
pKa is the pH at which
50% of the buffer’s acid is dissociated).
What is a buffer–titration curve?
> A titration curve is a plot of
pH vs. the amount of acid or base added to a
buffer solution (titration).
> It is useful for determining the
pKa of weak acids or bases.
> The pH is plotted on the
y axis and the buffer composition on the x axis.
> For the bicarbonate and
carbonic acid buffer system,
on the left side of the plot,
most of the buffer is in the form of carbon dioxide or carbonic acid and on the right side of the plot, most of the buffer is in the form of bicarbonate ion.
> The curve is sigmoid in shape,
with greater pH changes occurring at the
extremes of buffer compositions.
> If acid is added,
the pH decreases,
and the buffer shifts towards a greater
H2CO3 and CO2 concentration.
> Conversely, as base is added,
the pH increases and the buffer shifts
towards a greater HCO3− concentration.
> The flatter part of the slope represents the area of greatest buffering capacity where a
shift in the relative concentrations
of bicarbonate and carbon dioxide
produces only a small change in the pH of the solution.
> The steeper part of the slope
represents the area of least buffering capacity,
where even a small shift in relative concentrations of acid and base produces a large change in the pH.
> A t the central (equivalence) point,
both acid and base are present in equal
proportions, and the pH is equal to the pKa
(6.1) for the buffer.
> A t the physiological blood pH of
7.4 (outside of the zone of greatest
buffering capacity), small changes in the relative compositions cause alarge pH change.
> Therefore, in order to maintain a constant pH, the body relies additionally
on other buffer and organ systems.
What are the characteristics of an ideal buffer?
A good buffer solution must maintain
a nearly constant pH when
either acid or base is added.
Two features render this possible:
> Range of buffer (defined as pH = pKa ± 1).
• The buffer functions most effectively when its pKa is within one unit of the desired pH of the solution.
> Buffering capacity • This is defined by the ratio of the concentrations of weak acid to conjugate base, which must remain fairly constant, such that the addition of acid or base will not cause a change of pH.
What are the physiological buffer systems in the body?
1
The bicarbonate/carbonic acid buffer system:
• This is the most important system.
• Despite it having a low pKa (6.1)
relative to blood pH,
it is effective due to the ready excretion
of carbonic acid in the form
of CO2 by the lungs,
and the continuous regeneration of
bicarbonate by the kidneys.
• It is more efficient at buffering
acids since its efficiency
increases as the pH falls.
• It is the main buffer system in the blood due to the abundance of plasma bicarbonate. The production of carbonic acid is catalysed by the enzyme carbonic anhydrase,
which is present in red blood cells,
but not in plasma.
The reaction for this buffer system is:
CO2 + H2O ↔ H2CO3 ↔ HCO3− + H+
The Henderson–Hasselbalch equation
for this buffer system is:
pH = 6.1 + log [HCO3−] / [H2CO3]
and since H2CO3 is proportional to PaCO2:
pH = 6.1 + log [HCO3−] / 0.225 × PaCO2
(0.225 is the solubility coefficient for PaCO2 in kPa).
Haemoglobin:
• Acts as a blood buffer due to the
imidazole groups of its
histidine residues
(each molecule has 38 histidine residues).
Imidazole side chains
are anionic and accept H+.
• Deoxygenated haemoglobin (pKa 8.2) dissociates more readily than oxygenated haemoglobin (pKa 6.6), making it a better buffer and weaker acid (Haldane effect).
The advantage at capillary level is that
after O2 has been offloaded,
oxyhaemoglobin is reduced to deoxyhaemoglobin, which has a better buffering capacity, explaining why venous pH is only slightly more acidic than arterial pH.
• It has six times the
buffering capacity of plasma proteins.
Plasma and proteins:
• These are effective buffers because both their carboxyl (COOH) and free amino (NH2) groups dissociate.
Intracellular proteins are equally
important.
Phosphate:
• Plays a small role in the
extracellular fluid,
but is an important intracellular buffer
due to its
abundance and
dissociation from
phosphoric acid
to dihydrogen phosphate
and then to mono-hydrogen phosphate:
H3PO4 ↔ H2PO4– + H+ ↔ HPO4(2)– + H+
Urinary buffering:
• Occurs in the proximal (PCT) and
distal (DCT) tubules and collecting ducts.
• In the PCT,
H+ is secreted in exchange for Na+
and combines with filtered HCO3 −
to form carbonic acid.
This in turn dissociates into H2O and CO2,
which move freely into the tubular cell.
There, the reaction is reversed and the HCO3 −
formed enters the interstitium
and later the plasma.
Thus, for every H+ secreted,
one HCO3 − is reabsorbed.
• Some phosphate buffering takes place
in the PCT, but most of it occurs
in the DCT and collecting ducts.
• H+ combines with secreted NH3 to form NH4+,
which is excreted in the urine.
• Ammonia buffering takes place
mainly in the PCT and DCT.
• Buffering by bicarbonate results in
bicarbonate reabsorption,
whereas buffering with phosphate
and ammonia results in bicarbonate
regeneration.
What is the difference between ‘open’ and ‘closed’ buffer systems?
> Closed buffers:
the total concentration of buffer
within the cell is fixed,
e.g. phosphate and haemoglobin.
The addition of a strong acid or base
results in maintenance of the pH,
by shifting of the equation to the left or
right, respectively.
The buffering capacity is maximal
when the pH = pKa
of the buffer system and
is significantly reduced when the pH varies
by
more than 1 from the buffer’s pKa.
> Open buffers:
> Open buffers:
the total concentration of buffer
within a compartment is not fixed,
e.g. bicarbonate/carbonic acid system.
One of the components (H2CO3) is
fixed while the other (HCO3−) varies
inversely with the [H+].
This is because CO2 is highly permeable,
therefore its intra- and extracellular
concentrations are equal.
As it is in equilibrium with H2CO3,
it follows that the intracellular concentration
of H2CO3 is fixed.
The buffering capacity increases
as the concentration of the
non-fixed component (HCO3 −) increases.
Therefore intracellular pH increases,
despite cell pH moving
further away from the buffer’s pKa.
How much acid is produced by
the body per day
The body produces
metabolic and respiratory acids.
Metabolic acids are produced from
metabolism of amino acids,
phosphoproteins and
phospholipids,
and amount to 70 μmol/min or 0.1 mol/day.
Respiratory acids are formed from
CO2 production and amount to
200 mL/min or 8 mMol/min
which equals 12 mol/day.
Other sources of acid production include:
• Lactic acid (strenuous exercise)
• Ketoacids (diabetes, alcohol, starvation)
• Failure of H+ secretion by diseased kidneys (renal failure)
• Ingestion of acidifying salts (NH4Cl and CaCl2).
What effects does chronic renal
failure have on acid–base
balance?
> More acid is produced by
metabolism than is excreted.
This depletes extracellular buffers
and reduces plasma bicarbonate levels.
> Reduced total number of functioning nephrons → reduced production and secretion of ammonia → reduced buffering of urinary H+ → reduced tubular secretion of H+.
> Excess K+ causes intracellular alkalosis,
which inhibits H+ secretion.
> Bicarbonate reabsorption and
regeneration are reduced.
> Excess acid may be buffered by
calcium carbonate in bone, so
contributing to renal osteodystrophy.
> Haemoglobin levels are reduced due to depressed production of new red blood cells from a diminished erythropoietin secretion.
> Plasma proteins may be
diminished in the presence of
increased glomerular permeability
in certain conditions (glomerulonephritis/nephrotic syndrome).
