Ph and buffer

PH AND BUFFER
Acid Base Balance
Prakash Pokhrel

PH
It is the negative log of the hydrogen ion concentration.
pH = -log [H+]
ACID BASE BALANCE

 pH is a unit of measure which describes the degree of
acidity or alkalinity (basic) of a solution.
 It is measured on a scale of 0 to 14.
 Low pH values correspond to high concentrations of H+
and high pH values correspond to low concentrations of
H+.

PH VALUE
 The pH value of a substance is directly related to the
ratio of the hydrogen ion and hydroxyl ion
concentrations.
 If the H+ concentration is higher than OH- the material
is acidic.
 If the OH- concentration is higher than H+ the material
is basic.
 7 is neutral, < is acidic, >7 is basic

THE PH SCALE
 The pH scale corresponds to the concentration of
hydrogen ions.
 For example pure water H+ ion concentration is 1 x 10^-
7 M, therefore the pH would then be 7.

Acid
Any compound which forms H⁺ ions in solution
(proton donors)
eg: Carbonic acid releases H⁺ ions
Base
Any compound which combines with H⁺ ions in
solution (proton acceptors)
eg:Bicarbonate(HCO3⁻) accepts H+ ions

ACID–BASE BALANCE
 Normal pH : 7.35-7.45
 Acidosis
Physiological state resulting from abnormally low plasma
 Alkalosis
Physiological state resulting from abnormally high plasma
 Acidemia: plasma pH < 7.35
 Alkalemia: plasma pH > 7.45

MEASUREMENT OF PH
The pH can be measured by:
ph strips
Ph indicators
Ph meter

SOME IMPORTANT INDICATORS USED IN A CLINICAL
BIOCHEMISTRY LABORATORY ARE LISTED BELOW:
sr,.
No.
INDICATOR Ph range Colour in
acidic ph
Colour in
basic ph
1 Phenophthalein 9.3-10.5 colourless pink
2 Methyl orange 3.1-4.6 red yellow
3 Bromophenol blue 3.0-4.6 yellow blue
4 Methyl red 4.4-6.2 Red yellow
5 Phenol red 6.8 – 8.4 yellow red
6 Litmus 4.5-8.3 red Blue

PH METER
 The pH meter is a laboratory equipment which used to measure
acidity or alkalinity of a solution
 The pH meter measures the concentration of hydrogen ions [H+]
using an ion-sensitive electrode.
 It is the most reliable and convenient method for measuring ph.

BUFFER
A buffer solution is a solution which resists changes in pH when
a small amount of acid or base is added.
Typically a mixture of a weak acid and a salt of its conjugate
base or weak base and a salt of its conjugate acid.

TYPES OF BUFFERS
Two types :
 ACIDIC BUFFERS –
Solution of a mixture of a weak acid and a salt of this weak
acid with a strong base.
E.g. CH3COOH + CH3COONa
( weak acid ) ( Salt )
 BASIC BUFFERS –
Solution of a mixture of a weak base and a salt of this weak
base with a strong acid.
e.g. NH4OH + NH4Cl
(Weak base) ( Salt)

HOW BUFFERS WORK
 Equilibrium between acid and base.
 Example: ACETATE BUFFER
 CH3COOH  CH3COO- + H+
 If more H+ is added to this solution, it simply shifts the
equilibrium to the left, absorbing H+, so the [H+]
remains unchanged.
 If H+ is removed (e.g. by adding OH-) then the
equilibrium shifts to the right, releasing H+ to keep the
pH constant

•HANDERSON HASSELBALCH
EQUATION
 Lawrence Joseph Henderson wrote an equation, in 1908,
describing the use of carbonic acid as a buffer solution.
 Karl Albert Hasselbalch later re-expressed that formula
in logarithmic terms, resulting in the
Henderson–Hasselbalch equation.

Ka =
[H+] [A-]
[HA]
take the -log on both sides
The Henderson-Hasselbalch Equation derivation
-log Ka = -log [H+] -log
[A-]
[HA]
pH = pKa + log
[A-]
[HA]
= pKa + log
[Proton acceptor]
[Proton donor]
HA H+ + A-
pKa = pH -log [A-]
[HA]
apply p(x) = -log(x)
and finally solve for pH…

 - The greater the buffer capacity the less the pH
changes upon addition of H+ or OH-
Choose a buffer whose pKa is closest to the desired
pH.
pH should be within pKa ± 1

ACIDS
 VOLATILE ACIDS:
 Produced by oxidative metabolism of CHO,Fat,Protein
 Average 15000-20000 mmol of CO₂ per day
 Excreted through LUNGS as CO₂ gas
• FIXED ACIDS (1 mEq/kg/day)
 Acids that do not leave solution ,once produced they
remain in body fluids Until eliminated by KIDNEYS
Eg: Sulfuric acid ,phosphoric acid , Organic acids
Are most important fixed acids in the body
Are generated during catabolism of:
amino acids(oxidation of sulfhydryl gps of cystine,methionine)
Phospholipids(hydrolysis)
nucleic acids

RESPONSE TO ACID BASE CHALLENGE
1. Buffering
2. Compensation

BUFFERS
 First line of defence (> 50 – 100 mEq/day)
 Two most common chemical buffer groups
 Bicarbonate
 Non bicarbonate (Hb,protein,phosphate)
 Blood buffer systems act instantaneously
 Regulate pH by binding or releasing H⁺

CARBONIC ACID–BICARBONATE BUFFER SYSTEM
Carbon Dioxide
 Most body cells constantly generate carbon dioxide
 Most carbon dioxide is converted to carbonic acid, which dissociates into
H+ and a bicarbonate ion
Prevents changes in pH caused by organic acids and fixed acids in ECF
 Cannot protect ECF from changes in pH that result from elevated
or depressed levels of CO2
 Functions only when respiratory system and respiratory control
centers are working normally
 Ability to buffer acids is limited by availability of bicarbonate ions

ACID–BASE BALANCE
The Carbonic Acid–Bicarbonate Buffer System

THE CARBONIC ACID HYDROGENCARBONATE
BUFFER SYSTEM
• The carbonic acid-hydrogen Bicarbonate ion buffer is
the most important buffer system.
• Carbonic acid, H2CO3, acts as the weak acid
• Hydrogen carbonate, HCO3
-, acts as the conjugate base
• Increase in H+(aq) ions is removed by HCO3
-(aq)
• The equilibrium shifts to the left and most of the H+(aq)
ions are removed

 The small concentration of H+(aq) ions reacts with the
OH-(aq) ions
 H2CO3 dissociates, shifting the equilibrium to the right,
restoring most of the H+(aq) ions
 Any increase in OH-(aq) ions is removed by H2CO3

THE HEMOGLOBIN BUFFER SYSTEM
CO2 diffuses across RBC membrane
 No transport mechanism required
As carbonic acid dissociates
Bicarbonate ions diffuse into plasma
In exchange for chloride ions (chloride shift)
 Hydrogen ions are buffered by hemoglobin
molecules
 Is the only intracellular buffer system with an immediate effect
on ECF pH
 Helps prevent major changes in pH when plasma PCO
2
is rising
or falling

PHOSPHATE BUFFER SYSTEM
 Consists of anion H2PO4
- (a weak acid)(pKa-6.8)
 Works like the carbonic acid–bicarbonate buffer system
 Is important in buffering pH of ICF
Limitations of Buffer Systems
 Provide only temporary solution to acid–base
imbalance
 Do not eliminate H+ ions
 Supply of buffer molecules is limited

RESPIRATORY ACID-BASE CONTROL
MECHANISMS
 When chemical buffers alone cannot prevent
changes in blood pH, the respiratory system is the
second line of defense against changes.
 Eliminate or Retain CO₂
 Change in pH are RAPID
 Occuring within minutes
 PCO₂ ∞ VCO₂/VA

29
PHOSPHATE BUFFER SYSTEM
 The phosphate buffer system (HPO4
2-/H2PO4
-)
plays a role in plasma and erythrocytes.
 H2PO4
- + H2O ↔ H3O+ + HPO4
2-
 Any acid reacts with monohydrogen phosphate
to form dihydrogen phosphate
dihydrogen phosphate monohydrogen phosphate
 H2PO4
- + H2O ← HPO4
2- + H3O+
 The base is neutralized by dihydrogen phosphate
dihydrogen phosphate monohydrogen phosphate
 H2PO4
- + OH- → HPO4
2- + H3O+

RENAL ACID-BASE CONTROL
MECHANISMS
 The kidneys are the third line of defence against
wide changes in body fluid pH.
 movement of bicarbonate
 Retention/Excretion of acids
 Generating additional buffers
 Long term regulator of ACID – BASE balance
 May take hours to days for correction

RENAL REGULATION OF ACID BASE BALANCE
 Role of kidneys is preservation of body’s bicarbonate
stores.
 Accomplished by:
 Reabsorption of 99.9% of filtered bicarbonate
 Regeneration of titrated bicarbonate by excretion of:
Titratable acidity (mainly phosphate)
Ammonium salts

32
PROTEINS AS A BUFFER
 Proteins contain – COO- groups, which, like acetate ions
(CH3COO-), can act as proton acceptors.
 Proteins also contain – NH3
+ groups, which, like
ammonium ions (NH4
+), can donate protons.
 If acid comes into blood, hydronium ions can be
neutralized by the – COO- groups
 - COO- + H3O+ → - COOH + H2O
 If base is added, it can be neutralized by the – NH3
+
groups
- NH3
+ + OH- → - NH2 + H2O

TITRATABLE ACIDITY
 Occurs when
secreted H+
encounter & titrate
phosphate in tubular
fluid
 Refers to amount of
strong base needed
to titrate urine back to
pH 7.4
 40% (15-30 mEq) of
daily fixed acid load
 Relatively constant
(not highly adaptable)

RENAL REABSORPTION OF
BICARBONATE
Proximal tubule:
70-90%
Loop of Henle:
10-20%
Distal tubule and
collecting ducts:
4-7%

FACTORS AFFECTING RENAL
BICARBONATE REABSORPTION
 Filtered load of
bicarbonate
 Prolonged changes in
pCO2
 Extracellular fluid
volume
 Plasma chloride
concentration
 Plasma potassium
concentration
 Hormones (e.g.,
mineralocorticoids,
glucocorticoids)

If secreted H+ ions combine with filtered
bicarbonate, bicarbonate is reabsorbed
If secreted H+ ions combine with
phosphate or ammonia, net acid excretion
and generation of new bicarbonate occur

NET ACID EXCRETION
 Hydrogen Ions
Are secreted into tubular fluid along
 Proximal convoluted tubule (PCT)
 Distal convoluted tubule (DCT)
 Collecting system

AMMONIUM EXCRETION
 Occurs when
secreted H+
combine with NH3
and are trapped as
NH4
+ salts in
tubular fluid
 60% (25-50 mEq)
of daily fixed acid
load
 Very adaptable (via
glutaminase
induction)

AMMONIUM EXCRETION
 Large amounts
of H+ can be
excreted
without
extremely low
urine pH
because pKa
of NH3/NH4
+
system is very
high (9.2)

ACID–BASE BALANCE DISTURBANCES
Interactions among the Carbonic Acid–Bicarbonate Buffer
System and Compensatory Mechanisms in the Regulation of
Plasma pH.

FOUR BASIC TYPES OF IMBALANCE
Metabolic Acidosis
Metabolic Alkalosis
Respiratory Acidosis
Respiratory Alkalosis

ACID BASE DISORDERS
Disorder pH [H+] Primary
disturbance
Secondary
response
Metabolic acidosis    [HCO3
-]  pCO2
Metabolic alkalosis    [HCO3
-]  pCO2
Respiratory acidosis    pCO2  [HCO3
-]
Respiratory alkalosis    pCO2  [HCO3
-]

METABOLIC ACIDOSIS
 Primary AB disorder
 ↓HCO₃⁻ → ↓ pH
 Gain of strong acid
 Loss of base(HCO₃⁻)

CAUSES OF METABOLIC ACIDOSIS
 LACTIC ACIDOSIS
 KETOACIDOSIS
 Diabetic
 Alcoholic
 Starvation
 RENAL FAILURE (acute
and chronic)
 TOXINS
 Ethylene glycol
 Methanol
 Salicylates
 Propylene glycol

.
Responses to Metabolic Acidosis

METABOLIC ACIDOSIS
 Symptoms are specific and a result of the underlying pathology
 Respiratory effects:
 Hyperventilation
 CVS:
 ↓ myocardial contractility
 Sympathetic over activity
 Resistant to catecholamines
 CNS:
 Lethargy, disorientation,stupor,muscle twitching, COMA,
CN palsies
 Others : hyperkalemia

METABOLIC ALKALOSIS
↑ pH due to ↑HCO₃⁻ or ↓acid
 Initiation process :
 ↑in serum HCO₃⁻
 Excessive secretion of net daily production of fixed
acids
 Maintenance:
 ↓HCO₃⁻ excretion or ↑ HCO₃⁻ reclamation
 Chloride depletion
 Pottasium depletion
 ECF volume depletion
 Magnesium depletion

CAUSES OF METABOLIC ALKALOSIS
I. Exogenous HCO3 − loads
A. Acute alkali administration
B. Milk-alkali syndrome
II. Gastrointestinal origin
1. Vomiting
2. Gastric aspiration
III. Renal origin
1. Diuretics
2. Posthypercapnic state
3. Hypercalcemia/hypoparathyroidism
4. Recovery from lactic acidosis or ketoacidosis
5. Nonreabsorbable anions including penicillin, carbenicillin
6. Mg2+ deficiency
7. K+ depletion

COMPENSATION FOR METABOLIC ALKALOSIS
 Respiratory compensation: HYPOVENTILATION
 ↑PCO₂=0.6 mm  pCO2 per 1.0 mEq/L ↑HCO3
-
 Maximal compensation: PCO₂ 55 – 60 mmHg
 Hypoventilation not always found due to
 Hyperventilation
 due to pain
 due to pulmonary congestion
 due to hypoxemia(PO₂ < 50mmHg)

.
Metabolic Alkalosis

METABOLIC ALKALOSIS
 Decreased myocardial contractility
 Arrythmias
 ↓ cerebral blood flow
 Confusion
 Mental obtundation
 Neuromuscular excitability

RESPIRATORY ACIDOSIS
 ↑ PCO₂ → ↓pH
 Acute(< 24 hours)
 Chronic(>24 hours)

COMPENSATION IN RESPIRATORY ACIDOSIS
Acute resp.acidosis:
 Mainly due to intracellular buffering(Hb,Pr,PO₄)
 HCO₃⁻ ↑ = 1mmol for every 10 mmHg ↑ PCO₂
 Minimal increase in HCO₃⁻
 pH change = 0.008 x (40 - PaCO₂)
Chronic resp.acidosis
 Renal compensation (acidification of urine & bicarbonate retention)
comes into action
 HCO₃⁻ ↑= 3.5 mmol for every 10 mm Hg ↑PCO₂
 pH change = 0.003 x (40 - PaCO₂)
 Maximal response : 3 - 4 days

Respiratory Acid–Base Regulation.

Ph and buffer

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