Download Protein Buffer Systems

Acid-Base balance and buffer
systems in the human body
Department of General Chemistry
Water and pH relationship
In solution water shows a very low
dissociation (probability of H+ in water is 1.8 x 10-9)
H2O
H+ + OH-
H+ is actually associated with a cluster of
water molecules and exists in solution as
H3O+ or H5O2+ or H7O3+
Ionization of Water
Ionization of Water
H20 + H20
H20
Keq= [H+] [OH-]
[H2O]
H3O+ + OH-
H+ + OHKeq=1.8 X 10-16M
[H2O] = 55.5 M
[H2O] Keq = [H+] [OH-]
(1.8 X 10-16M)(55.5 M ) = [H+] [OH-]
1.0 X 10-14 M2 = [H+] [OH-] = Kw
pH Scale
Devised by Sorenson (1902) [H+] can range
from 1M and 1 X 10-14M using a log scale
simplifies notation pH = -log [H+]
Water and pH relationship
• For Dissociation of water,


[H ][OH ]
K
[H 2 O]
• Where [ ] denotes concentration and K is
the dissociation constant
Water and pH relationship
1 mole of water = 18g
1 L of water contains
1000 ÷ 18 = 55.56 mol (ie pure water = 55.56M)
Molar concentration of H+ (or OH-) ions can be calculated
[H+] = 1.8x10-9 x 55.56 = 1.0x10-7
In order to avoide using –ve numbers,
the [H+] is expressed as pH which is –ve log (base 10) of [H+]
pH = - log10 [H+]
pH of pure water is 7
Acidic solutions have pH < 7
while basic solutions have pH > 7
Water and pH relationship
[H+]=1x10-6 what is the pH?
• pH= - log10 [H+]
• [H+]=0.24x10-4 what is the
pH?
6
4.6
• [H+]=3.4x10-3 what is the pH? 2.5
Calculating pH
Effect of pH Shifts in Biological Systems
• Biochemical processes
& reactions are pHdependent.
• Cells & organisms
maintain constant
specific pH value that
is optimal for
function.
• Changes in charge can
alter molecular
conformation and
activity.
• pH sensitive molecules
include enzymes,
receptors, ligands, ion
channels,
transporters and
structural proteins.
Weak Acids and Bases Equilibria
•Strong acids / bases – disassociate completely
•Weak acids / bases – disassociate only partially
•Enzyme activity sensitive to pH
• weak acid/bases play important role in
protein structure/function
Shape of titration curve is same for all weak
acids
Acid/conjugate base pairs
HA + H2O
HA
A- + H3O+
A- + H+
HA = acid ( donates H+)(Bronstad Acid)
A- = Conjugate base (accepts H+)(Bronstad Base)
Ka = [H+][A-]
[HA]
pKa = - log Ka
Ka & pKa value describe tendency to
loose H+
large Ka = stronger acid
small Ka = weaker acid
Henderson-Hasselbach Equation
Consider the dissociation of a general acid HA
HA
H + + A-
We can define a dissociation constant (K) where
[H  ][A  ]
K
[HA]
Rearranging gives
K[HA]
[H ] 
[A  ]

Taking logarithms on both sides and multiplying by -1 gives:
-log[H+] = -logK – log [HA]/[A-] or
pH = pK + log [A-]/[HA]
Henderson-Hasselbach Equation
1) Ka =
[H+][A-]
[HA]
HA = weak acid
A- = Conjugate base
2) [H+] = Ka [HA]
[A-]
3) -log[H+] = -log Ka -log [HA]
[A-]
4) -log[H+] = -log Ka +log [A-]
[HA]
5) pH = pKa +log [A-]
[HA]
* H-H equation
describes
the relationship
between
pH, pKa and buffer
concentration
Henderson-Hasselbalch Equation
• This equation can be used to determine the pH
if the pK and ratio of the ionised and unionised
forms is known.
• The pKa (a for acid) is the –ve log of the
dissociation constant of the acid. It is the pH at
which the ratio of the ionised and unionised
species is equal to 1. ie the molar concentration
of the ionised and unionsed species is the same.
• Similarly pKb is –ve log of the dissociation
constant of the base
Regulation of H+ concentration
Concentration of hydrogen ions is regulated sequentially
by:
• Chemical buffer systems –act within seconds
• The respiratory center in the brain stem –acts within
1-3 min
• Renal mechanisms –require hours to days to effect pH
changes
Sources of hydrogen ions anaerobic and aerobic
respiration of glucose incomplete oxidation of fatty
acids oxidation of sulfur-containing amino acids
hydrolysis of phosphoproteins and nucleic acids
Buffers
• Definition: A weak acid plus its conjugate base
that cause a solution to resist changes in pH when
an acid or base are added
Effectiveness of a buffer is determined by:
1) the pH of the solution, buffers work best
within 1 pH unit of their pKa
2) the concentration of the buffer; the more
present, the greater the buffering capacity
Buffer capacity
• The buffer capacity of a system is
already defined as the amount of strong
acid or base added to one litre (l) of the
system in order to change the pH one
unit
The Major Body Buffer Systems
The Major Body Buffer Systems
Site
Buffer System
Comment
ISF
Bicarbonate
For metabolic acids
Phosphate
Not important because concentration too low
Protein
Not important because concentration too low
Bicarbonate
Important for metabolic acids
Haemoglobin
Important for carbon dioxide
Plasma protein
Minor buffer
Phosphate
Concentration too low
Proteins
Important buffer
Phosphates
Important buffer
Phosphate
Responsible for most of 'Titratable Acidity'
Ammonia
Important - formation of NH4+
Ca carbonate
In prolonged metabolic acidosis
Blood
ICF
Urine
Bone
• Carbonic
Acid –
Bicarbonate
Buffer
System
~ Most
important in
the ECF
Bicarbonate buffer system
• Present in intra-and extracellular fluid
– Bicarbonate ion acts as weak base, carbonic
acid acts as a weak acid
– Bicarbonate ions combine with excess
hydrogen ions to form carbonic acid
– Carbonic acid dissociates to release
bicarbonate ions and hydrogen ions
H+ + HCO3-  H2CO3  H+ + HCO3-
The blood buffering system, simplified
Phosphate Buffer
system
~ Important in ICF & urine
Phosphate buffer system
• Important in intracellular fluid and urine pH
regulation
• Consists of two phosphate ions
– Monohydrogenphosphate ions act as a weak base
and combine with hydrogen ions to form
dihydrogenphosphate
– Dihydrogenphosphate dissociates to release
hydrogen ions
H+ + HPO4-2  H2PO4-  H+ + HPO4-2
Phosphate has three ionizable H+ and
three pKas
Protein Buffer
Systems
~ Important in ECF and ICF
~ Interact with other buffer systems
Protein buffer system
• Consists of Plasma Proteins (albumin,
hemoglobin)
• Remember proteins are just chains of AAThe
exposed amine group of the AA (NH2) accepts
H+ ions when conditions are acidic
• The exposed carboxyl group of AA can release
H+ ions when conditions are basic Proteins can
act as Acids or Bases
Hemoglobin is an important blood buffer
particularly for buffering CO2
• Protein buffers in blood include haemoglobin (150g/l) and plasma
proteins (70g/l). Buffering is by the imidazole group of the
histidine residues which has a pKa of about 6.8. This is suitable
for effective buffering at physiological pH.
• Haemoglobin is quantitatively about 6 times more important then
the plasma proteins as it is present in about twice the
concentration and contains about three times the number of
histidine residues per molecule. For example if blood pH changed
from 7.5 to 6.5, haemoglobin would buffer 27.5 mmol/l of H+
and total plasma protein buffering would account for only 4.2
mmol/l of H+.
The acid-base buffering systems of the body.
The two buffer systems are in
dynamic equilibrium with the
same hydrogen ion concentration
(pH), so that a change induced in
the concentration of any one
factor in either buffer system
rapidly affects the other
system and a new hydrogen ion
concentration in the blood is
established.
1. The lungs assist in maintaining a
constant blood pH by removing
CO2,
2. while the kidney excretes acid in
the form of H2PO4- and NH4 and
alkali in the form of HCO3-.
Respiratory Buffer Systems
• The respiratory system regulation of acidbase balance is a physiological buffering
system
• There is a reversible equilibrium between:
– Dissolved carbon dioxide and water
– Carbonic acid and the hydrogen and bicarbonate
ions
CO2+ H2O ↔H2CO3↔H++ HCO3¯
Respiratory Buffer Systems
• CO2 is produced by
cellular respiration.
• CO2 is converted to
bicarbonate by
carbonic anhydrase.
• results in LOWER pH in
respiring tissues.
• CO2 is exhaled in lungs.
Haemoglobin binds both CO2 and H+ and so is a powerful buffer. Deoxygenated haemoglobin has the strongest
affinity for both CO2 and H+; thus, its buffering effect is strongest in the tissues. Little CO2 is produced in red
cells and so the CO2 produced by the tissues passes easily into the cell down a concentration gradient. Carbon
dioxide then either combines directly with haemoglobin or combines with water to form carbonic acid. The CO2 that
binds directly with haemoglobin combines reversibly with terminal amine groups on the haemoglobin molecule to form
carbaminohaemoglobin. In the lungs the CO2 is released and passes down its concentration gradient into the alveoli.
Respiratory Acidosis and
Alkalosis
• Result from failure of the respiratory system
to balance pH
• PCO2is the single most important indicator of
respiratory inadequacy PCO2 levels
• Normal PCO2 fluctuates between 35 and 45
mm Hg
• Values above 45 mm Hg signal respiratory
acidosi
• Values below 35 mm Hg indicate respiratory
alkalosis
Respiratory Acidosis
•
•
Respiratory acidosis is the most common cause of acid-base imbalance
Occurs when a person breathes shallowly, or gas exchange is hampered
by diseases such as pneumonia, cystic fibrosis, or emphysema
Respiratory alkalosis
• A common result of hyperventilation
• Excessive loss of CO2 & subsequent loss
of carbonic acid
• Caused by hyperventillation:too much
CO2 lost (↓carbonic acid and H+ ions)
• Anxiety, high altitudes (low O2 levels),
musicians,
• Symptoms:lightheadedness, agitation,
dizziness,
Metabolic Acidosis
• All pH imbalances except those caused by
abnormal blood carbon dioxide levels
• Metabolic acidosis is the second most common
cause of acid-base imbalance
• Typical causes are ingestion of too much
alcohol and excessive loss of bicarbonate ions
• Other causes include accumulation of lactic
acid, shock, ketosis in diabetic crisis,
starvation, vomiting, and kidney failure
Metabolic Alkalosis
•
•
Rising blood pH and bicarbonate levels indicate metabolic alkalosis
Typical causes are:
– Vomiting of the acid contents of the stomach
– Intake of excess base (e.g., from antacids)
– Constipation, in which excessive bicarbonate is reabsorbed
Clinical Application
Acid-Base Imbalances
If the pH of arterial blood drops to 6.8 or rises to 8.0 for
more than a few hours, the person usually cannot survive
acidosis versus alkalosis
factors that lead to
respiratory acidosis
Clinical Application
Metabolic acidosis
Respiratory alkalosis
Metabolic alkalosis
Renal Mechanisms of Acid-Base
Balance
Chemical buffers can tie up excess acids or
bases, but they cannot eliminate them from
the body
– The lungs can eliminate carbonic acid by eliminating
carbon dioxide
– Only the kidneys can rid the body of metabolic
acids (phosphoric, uric, and lactic acids and
ketones) and prevent metabolic acidosis
– The ultimate acid-base regulatory organs are the
kidneys
Kidney buffering power
Whatever the nature of the
disturbance, the response of the
kidney leads to the formation and
extraction from the plasma of a fluid
with an excess or a deficit of acid.
The primary result is a return of the
H ion concentration of the blood
toward the normal level.
Renal Mechanisms of Acid-Base Balance
•
The most important renal mechanisms for regulating acid-base balance
are:
– Conserving (reabsorbing) or generating new bicarbonate ions: decreases
acidity of ECF
– Excreting bicarbonate ions: increases acidity of ECF
– Excreting excess H+
Respiratory and Renal Compensations
• Acid-base imbalance due to inadequacy
of a one system is compensated for by
the other system
• The respiratory system will attempt to
correct metabolic acid-base imbalances
• The kidneys will work to correct
imbalances caused by respiratory
disease
Respiratory Compensation
• In metabolic acidosis:
– The rate and depth of breathing are elevated
Blood pH is below 7.35 and bicarbonate level is low
– As carbon dioxide is eliminated by the respiratory
system, PCO2falls below normal
• In respiratory acidosis, the respiratory rate
is often depressedand is the immediate cause
of the acidosis
• In metabolic alkalosis:
– Compensation exhibits slow, shallow breathing,
allowing carbon dioxide to accumulate in the blood
– Correction is revealed by:High pH (over 7.45) and
elevated bicarbonate ion levelsRising PCO2
Clinical Applications
• Homeostatic mechanisms slow down with
age
– Elders may be unresponsive to thirst clues
and are at risk of dehydration
– The very young and the very old are the
most frequent victims of fluid, acid-base,
and electrolyte imbalances
Buffering power of saliva
1.The pH of the mouth must be maintained near neutral for normal tooth
maintenance.
2.Oral pH is buffered to a small extent by saliva proteins and
phosphate.The major influence on saliva pH is bicarbonate ion which is a
by-product of cell metabolism
3.Bicarbonate concentration increases in saliva as the flow rate rises and is
due to the increased metabolic rate. This, in turn, raises the pH (more
alkaline) of saliva.
4.Bicarbonate ions diffuse into dental plaque and neutralise acid produced
by plaque bacteria when carbohydrate is fermented.
5.This reaction is driven by a unique type of carbonic anhydrase which is
secreted into saliva by serous acinar cells of the parotid and
submandibular glands.
6.Bicarbonate ions maintain the pH of saliva above 6.3.
Role of Bone Buffering
•
Bone consists of matrix within which specialised cells are dispersed. The matrix
is composed of organic [collagen and other proteins in ground substance] and
inorganic [hydroxyapatite crystals: general formula Ca10(PO4)6(OH)2]
components.
•
The hydroxyapatite crystals make up two-thirds of the total bone volume but
they are extremely small and consequently have a huge total surface area. The
crystals contain a large amount of carbonate (CO3-2) as this anion can be
substituted for both phosphate and hydroxyl in the apatite crystals. Bone is the
major CO2 reservoir in the body and contains carbonate and bicarbonate
equivalent to 5 moles of CO2 out of a total body CO2 store of 6 moles. (Compare
this with the basal daily CO2 production of 12 moles/day)
•
CO2 in bone is in two forms: bicarbonate (HCO3-) and carbonate (CO3-2). The
bicarbonate makes up a readily exchangeable pool because it is present in the
bone water which makes up the ‘hydration shell’ around each of the
hydroxyapatite crystals. The carbonate is present in the crystals and its release
requires dissolution of the crystals. This is a much slower process but the
amounts of buffer involved are much larger.
Role of Bone Buffering
• Bone is an important source of buffer in chronic
metabolic acidosis (ie renal tubular acidosis & uraemic
acidosis)
• Bone is probably involved in providing some buffering
(mostly by ionic exchange) in most acute acid-base
disorders but this has been little studied.
• Release of calcium carbonate from bone is the most
important buffering mechanism involved in chronic
metabolic acidosis.
• Loss of bone crystal in uraemic acidosis is
multifactorial and acidosis is only a minor factor
• BOTH the acidosis and the vitamin D3 changes are
responsible for the osteomalacia that occurs with
renal tubular acidosis.
Ammonia is produced in renal tubular
cells by the action of the enzyme
glutaminase on the amino acid glutamine.
This enzyme functions optimally at a
lower (more acidic) than normal pH.
Therefore, more ammonia is produced
during acidosis improving the buffering
capacity of the urine. Ammonia is
unionised and so rapidly crosses into the
renal tubule down its concentration
gradient. The ammonia combines with H+
to form the ammonium ion, which being
ionised does not pass back into the
tubular cell. The ammonium ion is
therefore lost in the urine, along with
the hydrogen ion it contains