This document discusses acids, bases, pH, buffers, and the regulation of pH in the body. It defines acids and bases, describes pH and how it is measured. It explains the carbonic acid-bicarbonate buffer system, which is one of the most important buffer systems in the body. It also discusses how respiratory regulation and kidney regulation help maintain pH levels through increasing or decreasing ventilation and excreting acid and bases in the urine. The kidney regulates pH through reabsorbing bicarbonate and secreting hydrogen ions into the tubule, where they react with phosphate and ammonia to generate buffers without lowering urine pH.

1. ACID, BASES, pH,
BUFFER and regulation of
pH

2. ACIDS
• An acid is any substance that dissociates (ionizes) in solution and
releases hydrogen ions (H+).
• Acids have the following characteristics:
• Taste sour
• Turns litmus indicators red
• React with bases to form salts
• Cause some metals to liberate hydrogen
• Examples of acids in body includes:
Hydrochloric acid, lactic, phosphoric, carbonic, citric and carboxylic
acids

3. BASES
• A base is any substance that picks up or accepts H+ to form hydroxide
ions (OH-) in water solutions.
• Base or alkaline solutions have the following characteristics:
• Bitter taste
• Slippery to the touch
• Turns litmus indicators blue
• React with acids to form water
• Examples of bases in the body include sodium and calcium hydroxide
and aqueous solutions of ammonia that form ammonium hydroxide.

4. Litmus Paper
• To test whether the solution is acidic or alkaline.
Test with acid Test with alkali
Red litmus paper No changes Turns blue
Blue litmus paper Turns red No changes

5. pH
• pH refers to a solution’s concentration of protons or H+ .
• Solutions with relatively more OH- than H+ have a pH above 7.0 and
are called basic or alkaline.
• Conversely , solutions with more H+ than OH- have a pH below 7.0
and are termed as acidic.
• Chemically pure (distilled) water has a pH of 7.0 (neutral) with equal
amounts of H+ and OH-
• The pH scale, devised in 1909 by Danish chemist Soren Sorenson,
ranges from +1.0 to +14.0

6. pH Paper
pH Scale

7. • An inverse relation exists between pH and the H+ concentration
([H+]), as pH refers to negative logarithm of the [H+]
• The relation between [H+] and pH can be expressed as:
• pH= log10 1/[H+]
• pH= -log10 [H+]
• It means a one unit change in pH corresponds to a tenfold change
in [H+]
• For example, Lemon juice and gastric juice (pH=2.0) have 1000
times greater [H+] than black coffee (pH=5.0), whereas
hydrochloric acid (pH=1.0) has approximately 1000000 times the
[H+] of blood (pH= 7.4)

8. • The pH of body fluids ranges from a low of 1.0 for the digestive acid
hydrochloric acid to a slightly basic pH between 7.35 and 7.45 for
arterial and venous blood (and most other body fluids)
• pH 7.4 +/- 0.05; alkaline

10. Alkalosis and Acidosis
• The term ALKALOSIS refers to an increase in pH above the normal
average of 7.4; this results directly from of a decrease in [H+] (increase
in pH)
• Conversely, ACIDOSIS refers to an increase in [H+] (decrease in pH)
• The highly specific acid-base quality of various body fluid remains
regulated within narrow limits because of the high sensitivity of
metabolism to the [H+] of the reacting medium

11. Result of Acidosis
H2O + CO2 ← H2CO3 H+ + HCO3
-
• An increase in plasma carbon dioxide or H concentration immediately
stimulates ventilation to eliminate “excess” carbon dioxide.
• Conversely, a decrease in plasma H concentration inhibits the
ventilatory drive and retains carbon dioxide that then combines with
water to increase acidity (carbonic acid) and normalize pH.
Result of Alkalosis
H2O + CO2 H2CO3 H+ + HCO3
-

13. Mechanism of Regulation of pH
• First line of defence- Blood buffers or Chemical buffers
• Second line of defence- Respiratory regulation or Ventilatory
buffer
• Third line of defence- Kidney regulation or Renal buffer

14. Buffer System
• A buffer is a substance that has the ability to bind or release
H+ in solution.
• A buffer in a solution consists of a weak acid and its
conjugate base, thus keeping the pH of a solution relatively
constant despite the addition of considerable quantities of
acid or base.
• Buffering is the primary means by which large changes in
[H+] are minimised

15. Henderson- Hasselbalch equation
• The general equation for a buffer system is
HA H+ + A-
• Where A- represents any anion and HA the undissociated
acid.
• If an acid stronger than HA is added to a solution containing
this buffer system, the equilibrium is shifted to the left.
• H+ are ‘tied up’ in the formation of more undissociated HA,
so the increase in H+ concentration is much less than it would
otherwise be.

16. • Conversely, if a base is added to the solution, H+ and OH-
react to form H2O; but more HA dissociates, limiting the
decrease in H+ concentration.
• By the Law of mass action, the product of the concentrations
of the products in a chemical reaction divided by the product
of the reactants at equilibrium is a constant :
[H+] [A-]/ [HA]= K

17. • If this equation is solved for H+ and put in pH notation (pH is
the negative log of [H+]), the resulting equation is that
originally derived by Henderson and Hasselbalch to
describe the pH changes resulting from addition of H+ or OH-
to any buffer system (Henderson- Hasselbalch equation):
pH= pK + log [A-]/[HA]

18. • It is apparent from these equations that the buffering capacity of a
system is greatest when the amount of free anion is equal to the
amount of undissociated HA i.e. when [A-]/[HA] =1, so that
log [A-]/[HA]= 0 and pH=pK.
• This is why the most effective buffers in the body would be expected
to be those with pKs close to the pH in which they operate.
• The pH of the blood is normally 7.4; that of the cells is probably about
7.2; and that of urine varies from 4.5 to 8.0.
• The pK of the bicarbonate system is 6.1.
• The pK of the dibasic system is 6.8
• The pK of Ammonia system is 9.0

19. Buffer Systems in Body
1. Haemoglobin system
2. Protein system
3. Carbonic acid- Bicarbonate system
4. Phosphate system

20. Haemoglobin Buffer
• The buffering action of haemoglobin is due mainly to the imidazole
groups of the histidine residues.
• Haemoglobin in blood has “six” times the buffering capacity of the
plasma proteins.
• In addition, imidazole groups of deoxy-haemoglobin i.e. reduced
haemoglobin (Hb- ) dissociate less than those of oxyhaemoglobin
(HbO2); therefore, Hb- produces less H+ at a given pH than does
oxyhaemoglobin (HbO2), making it a weaker acid and thus Hb-
becomes a more effective buffer when CO2 (and hence H+) are added
from the tissues.

21. Protein System
• Plasma proteins are effective buffers, because both their carboxyl
(COOH) and their free amino (NH3+) groups dissociate, therefore,

22. Carbonic Acid – Bicarbonate System
• This system consists of H2CO3 (weak acid) and HCO3
-
• Carbonic acid (H2CO3) is only partially dissociated into H+ and HCO3
-.
Therefore,
• If H+ is added to a solution of , the equilibrium shifts to the left and
most of the added H+ is removed from solution, and
• If OH- is added, H+ and OH- combine, taking H+ out of solution.
• However, the decrease is countered by more dissociation of H2CO3,
and the decline in H+ concentration is minimised.

23. • Carbonic acid-bicarbonate system is one of the most effective buffer
systems in the body, because the H2CO3 level in plasma is in
equilibrium with the dissolved CO2 (H2CO3 CO2 + H2O)and the
amount of dissolved CO2 is controlled by respiration.
• In addition, the plasma concentration of HCO3
- is regulated by
kidneys.
• The reaction CO2 + H2O H2CO3 proceeds slowly in either
direction unless the enzyme carbonic anhydrase (CA) is present.
• There is no “CA” in plasma, but there is an abundant supply in RBC,
gastric acid secreting cells and renal tubular cells.

24. Phosphate System
• The system has a H2PO4
- H+ + H2PO4
2- PK of 6.8.
• In plasma, the phosphate concentration is too low for this system to be
a quantitatively important buffer, but it is important intracellularly, and
it frequently plays a significant role in the urine.

25. Respiratory Regulation of pH
• When the quantity of free H+ in extracellular fluid and plasma
increases, it directly stimulates the respiratory center to
immediately increase alveolar ventilation.
• This rapid adjustment reduces alveolar PCO2 and causes
carbon dioxide to be “blown off” from the blood.
• Reduced plasma carbon dioxide levels accelerate the
recombination of H+ and HCO3
-, lowering free H+
concentration in plasma.

26. • For example, doubling alveolar ventilation by
hyperventilation at rest increases blood alkalinity and pH by
0.23 units from 7.40 to 7.63.
• Conversely, reducing normal alveolar ventilation
(hypoventilation) by half increases blood acidity by
approximately 0.23 pH units. The potential magnitude of
ventilatory buffering equals twice the combined effect of all
the body’s chemical buffers.

28. Renal Regulation of Acid-base Balance
INTRODUCTION
• The kidneys are responsible for clearing the body of
metabolically produced non-carbonic acids.
• Therefore, normal urine reaction is acidic in nature.
• In all body fluids, electrical neutrality must be maintained i.e.
anions like phosphate or sulphate must be ‘covered’ by an
equal amount of cations.
• This can only be achieved by daily excretion of the excess of
anion in the urine.

29. • If these anions which are excreted in the urine are fully
‘covered’ by an equivalent amount of cations (mainly Na+)
this will produce serious consequences.
• However, this situation is prevented due to manufacture of
two important cations, H+ and NH4
+ by the kidneys which
‘cover’ the excreted anions.

30. Role of the kidneys
1. The kidneys perform two major functions:
• The kidneys stabilize HCO3
- pool by obligatory reabsorption
(mainly by PCT) and by controlled reabsorption of filtered
HCO3
- (by the DCT and CT)
• The kidneys excrete metabolically produced non-carbonic
acid, which represents a H+ excretion.

31. 2. The major sites of urine acidification are DCT and CT
• Essentially, all of the H+ within the tubular lumen is from the tubular
secretion of H+ generated by metabolism.
• There is no significant contribution of H+ from the glomerular filtrate.
• The amount of acid secreted depends upon the subsequent events in the
tubular urine.
• The maximal H+ gradient against which the transport mechanism can
secrete H+ (in humans) corresponds to a urine pH of about 4.4 i.e. a
[H+] of 40 x 10-6 Eq/L. Since the plasma [H+] is 40 x 10-9 Eq/L, the
kidney can cause a 1000 fold [H+] gradient between the plasma and
urine. Because the lowest pH attainable in urine is 4.4, which is thus
called the limiting pH of urine

33. • If there were no buffers that ‘tied up’ H+ in the urine, this pH
would be reached rapidly and H+ secretion would stop.
• However, three important reactions (buffer systems) in the
tubular fluid remove free H+, permitting more acid to be
secreted
• The major buffer systems present in kidneys are:
• Bicarbonate system
• Dibasic phosphate system, and
• Ammonia system

34. Buffer systems in the kidney
General
• At normal blood pH of 7.4 (approx. that of the glomerular
filtrate), since pK = 6.1, the ratio of [HCO3
-]/[H2CO3] is 20/1
for pH=pK+ log [HCO3
-]/[H2CO3]
i.e. 7.4= 6.1 + 1.3
• 7.4= 6.1+ log 20 (as log 20=1.3)
• Thus, at a urinary pH of 6.1,
• Log [HCO3
-]/[H2CO3] is zero, and hence [HCO3
-] = [H2CO3]

35. Bicarbonate System
Reaction with bicarbonate ions (HCO3
-)
• The majority of secreted H+ in the PCT reacts with HCO3
- to
form H2CO3 and is used to bring about HCO3
- reabsorption.
• The CO2 formed in the lumen from secreted H+ returns to no
net H+ secretion occurs.
• Since most of the secreted H+ is removed from the tubule
fluid is changed very little; as a result it does not contribute to
the urinary excretion of acid.
• The H+ secreted in excess of those required for HCO3
-
reabsorption are buffered in the tubular fluid by the remaining
other two buffer systems.

36. Dibasic Phosphate System
Reactions with dibasic phosphate ions (HPO4
2-)
• Besides HCO3
- , HPO4
2- represents a major filtered conjugate
base.
• Approx. 75% of the filtered HPO4
2- is reabsorbed by the PCT,
therefore, only 25% of the filtered HPO4
2- is available for
buffering in DCT and CT (because it is here that the
phosphate which escapes reabsorption in the PCT gets
concentrated by the reabsorption of water).
• The H+ secreted into the tubules, therefore, can react with
filtered HPO4
2- rather than the filtered HCO3
-

37. • The exchange of H+ for Na+ converts dibasic sodium
phosphate (Na2HPO4) in the glomerular filtrate into acidic
sodium dihydrogen phosphate (NaH2PO4) and is excreted in
the urine as Titratable Acid.
• The PCT is the major nephron site where titratable acid is
formed.
• Additional titratable acid is generated along the collecting
duct by a H+ ATPase pump.

38. Phosphate buffer system carries excess H+
into the urine and generates new HCO3
-

39. Ammonia System
Reaction with Ammonia (NH3)
A) Secretion of ammonia
• Ammonia (NH3) enters the tubular lumen NOT BY
FILTERATION but by tubular synthesis and secretion, which
normally is confined to the DCT and CT
• Most of the ammonium ion (NH4
+ ) excreted in urine is
produced in the PCT cells from amino acids, primarily
glutamine.

40. B) Formation of Ammonium ion (NH4
+ )
• NH3 passively diffuses out of tubular cells easily (because it
is lipid soluble) along its concentration gradient, called non-
ionic diffusion of NH3.
• In the lumen NH3 combines with H+ to form ammonium ion
(NH4
+ ).
• NH4
+ is relatively lipid insoluble, therefore, stays in the
lumen.

41. C) Excretion of Ammonium ion (NH4
+ )
• The important physiologic characteristic of the ammonia
system (NH3 / NH4
+ ) is that, as H+ is combined with intra-
luminal buffer (NH3), H+ is excreted in urine as NH4
+ , a
substance that does not cause the pH of urine to fall.
• Ammonia system has a very high pK (about 9), which means
that, at the usual urine pH of 6.0, practically all of the
nonpolar NH3 that enters the DCT lumen immediately
combines with H+ to form NH4
+

43. • Therefore, when urine pH is more than 6.0, NH3 being lipid
soluble and in gaseous form, diffuses to the other side of the
membrane i.e. absorbed into the peritubular capillary.
• Thus, ammonium content of the urine is negligible until pH
falls below 6.0.
• Ammonium excretion then increases linearly as the urinary
pH falls below this value.
• The renal excretion of NH4
+ causes the net addition of HCO3
-
to the plasma.

44. • The amount of NH4
+ formed depends upon the pH of the
tubular fluid and the rate of NH3 production. At any given
rate of NH3 production, the amount of NH4
+ formed is
proportionate to the amount of H+ available and, therefore, to
the rate of H+ secretion.
• Thus, the NH4
+ content of an alkaline urine (pH more than
6.0) is nil whereas that of a maximally acid urine is high.
• The Ammonium system assists the conservation of Na+ and
HCO3
- by the body.

46. Effects of Intense Exercise
• Increased H+ concentration from carbon dioxide production
and lactate formation during strenuous exercise makes pH
regulation progressively more difficult.
• Acid–base regulation becomes exceedingly difficult during
repeated, brief bouts of all-out exercise that elevate blood
lactate values to 30 mM (270 mg of lactate per dL of blood)
or higher.

47. • Top. General relationship between blood pH and blood lactate concentration during rest and increasing
intensities of short-duration exercise up to maximum.
• Bottom. Blood pH and blood lactate concentration related to exercise intensity expressed as a percentage of
the maximum. Decreases in blood pH accompany increases in blood lactate concentration.

48. • There occurs inverse linear relationship between blood lactate
concentration and blood pH. Blood lactate concentration
varied between a pH of 7.43 at rest and 6.80 during
exhaustive exercise.
• This response indicates that humans temporarily tolerate
pronounced disturbances in acid–base balance during
maximal exercise, at least to an overall blood pH as low as
6.80.
• A plasma pH below 7.00 does not occur without
consequences; this level of acidosis produces nausea,
headache, and dizziness in addition to discomfort and pain
that ranges from mild to severe within active muscles.

49. Thankyou