acid-base physiology. 2 ph scale – to express hydrogen ion concentration. ph = - log 10 [h+] or ph...

Acid-Base Physiology

2

pH scale – to express hydrogen ion concentration.

pH = - log10 [H+] orpH = log 1 / [H+] log to the base 10 of the reciprocal

of hydrogen-ion concentration.

1) Because [H+] is in the denominator,A high [H+] low pH andA low [H+] high pH.

2) pH unit change of 1 = 10X change in [H+]

The [H+] of ECF is very low (0.00004 mEq/L = 40 nmoles/L). Normal variations are are markably small 3-5 mEq/L. It is customary to express these very small numbers using thelogarithmic pH scale.

The Conceptual Problem with pH

• Because it’s a logarithmic scale, it doesn’t make “sense” to our brains.

• EASY TO REMEMBER FACTS :- – Every factor of 10 difference in [H+] represents 1.0

pH units, – Every factor of 2 difference in [H+] represents 0.3

pH units.

• Therefore, even numerically small differences in pH, can have profound biological effects…

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ACIDS

• Acids are H+ donors.

• Acids can be:– Strong – dissociate completely in solution

–HCl– Weak – dissociate only partially in solution

–Lactic acid, carbonic acid

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Volatile and Fixed Acids

• VOLATILE ACIDS :- carbonic acid– Nearly 20,000 mEq of carbonic acid /day

• FIXED ACID :- lactate , keto acids, sulphuric acid, phosphoric acid

-- Nearly 60-80mEq of fixed acids/day

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BASES

• Bases are acceptors of H+(protons) or give up OH- in solution

• Bases can be:-– Strong – dissociate completely in solution

-NaOH– Weak – dissociate only partially in solution

– NaHCO3

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Weak acids thus are in equilibrium with their ionized species:

HA H+ + A-

Ka = , pKa = -log Ka[H+][A-] [HA]

Governed by the Law of Mass Action, and characterized by an equilibrium constant:

Derivation of the Henderson-Hasselbalch equation

• Ka = [H+] [A-] [HA]• so [H+] = Ka [HA] [A-]• TAKING THE NEGATIVE LOG OF BOTH SIDES• As pH = - log [ H+],

• pH = -log Ka [HA]

[A-])• pH = -log(Ka)-log([HA]

[A-])• pH = pKa + log([A-]/[HA])

The Henderson Hasselbalch Equation

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pH = pKa + log [A-] [HA]

L J HENDERSONK A HASSELBALCH

Simplified form……

• pH = pKa + log ([A-] [HA])

• pH = pKa + log(Conjugate base Conjugate acid)

• pH = pKa + log(Proton acceptorProton donor )

The Body and pH

• Homeostasis of pH is tightly controlled• Extracellular fluid = 7.4• Blood = 7.35 – 7.45• < 6.8 or > 8.0 death occurs• Acidosis (acidemia) below 7.35• Alkalosis (alkalemia) above 7.45

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Importance Of Maintenance Of pH Between 7.35 - 7.45(7.4)

Acidosis pH<7.35 and AlkalosispH>7.45.Death occurs if pH falls outside the range of 6.8 to 8.0

• Altered [H+] results in changes in protein structure (Enzymes, Receptors and ligands, Ion channels,Transporters,Structural proteins)

• Function of excitable tissues– Acidosis: hypoexcitability, CNS depression– Alkalosis: hyperexcitability, tetany

• Affects K+ levels in the body.

Continuous addition of H+ ions to the body fluids and 3 Lines Of Defense Against pH Changes due to this:

• Buffering• Changes in ventilation• Changes in renal handling of H+ and HCO3

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Mechanisms of Regulation of pH

• FIRST LINE OF DEFENSE : BLOOD BUFFERS

• SECOND LINE OF DEFENSE :- RESPIRATORY REGULATION

• THIRD LINE OF DEFENSE :RENAL REGULATION

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FIRST LINE OF DEFENSE : BLOOD BUFFERS

• Buffer systems. Buffers act quickly to temporarily bind H+ removing the highly reactive, excess H+ from solution. Buffers thus raise pH of body fluids but do not remove H+ from the body.

• Buffers function almost instantaneously

Buffers Are The1st Line Of Defense. They Minimize (But Do Not Prevent) Changes In pH.

Buffer + H+ ↔ Hbuffer

Buffering of hydrogen Ions in the body fluids

• Bicarbonate buffer system• Intracellular protein• Hemoglobin Buffer system.• Phosphate buffer system

Bicarbonate Buffer

• The most important buffer in plasma.• 65% of buffering capacity.• BASE CONSTITUENT :- (HCO3

-) Renal Regulation

• ACID CONSTITUENT :- (H2CO3) Respiratory Regulation

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Bicarbonate buffer

• Sodium Bicarbonate (NaHCO3) and carbonic acid (H2CO3)

• Maintain a 20:1 ratio : HCO3- : H2CO3

HCl + NaHCO3 ↔ H2CO3 + NaCl ; {excess H2CO3 , excess CO2}

NaOH + H2CO3 ↔ NaHCO3 + H2O; { decre H2CO3 ,dec CO2}

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CO2 + H2O ↔ H2CO3 ↔ H+ + HCO3-

• Normal bicarbonate level of plasma is 24mmol/L

• The normal pCO2 is 40mm Hg

• The normal carbonic acid concentration is 1.2 mmol/L

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Remember these values!!

• pKa for carbonic acid is 6.1• So, applying Henderson –Hasselbalch’s equation

pH= pKa + log [HCO3- ]

[H2CO3]

= 6.1 + log 24 1.2

= 6.1 + log 20 = 6.1 +1.3= 7.4

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What Is The Central Message Of Henderson-Hasselbalch?

pH = pKa + log(HCO3- / H2CO3)

Plasma pH is a simple function of the HCO3

- :H2CO3 ratioHCO3

- : H2CO3 ↑ = pH ↑ (ALKALOSIS) : Could occur due to either HCO3

- ↑(Metabolic alkalosis) or PCO2 ↓ (respiratory alkalosis)

HCO3- : H2CO3 ↓ = pH ↓( ACIDOSIS) :

Could occur either HCO3- ↓(metabolic acidosis)

or PCO2 ↑ (respiratory acidosis)

Phosphate buffer:

• Major intracellular buffer• The main elements of the phosphate buffer

system are H2PO4– and HPO4=.

• H+ + HPO42- ↔ H2PO4-

• OH- + H2PO4- ↔ H2O + H2PO4

2-

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Protein Buffers

• Buffering capacity of protein dependson the pKa value of the ionizable side chains.

• Includes hemoglobin• In general ,

– Carboxyl group gives up H+ – Amino Group accepts H+– Side chains that can buffer H+ are present on

amino acids.

Protein Buffer System

• The free carboxyl group at one end of a protein acts like an acid by releasing H+ when pH rises; it dissociates as follows:

ACTION OF HEMOGLOBIN

• GENERATES BICARBONATE BY CARBONIC ANHYDRASE

• In tissues :-

CO2 + H2O Carbonic Anhydrase H2CO3

H2CO3 HCO3- + H+

H+ + Hb- HHb

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SECOND LINE OF DEFENSE :- RESPIRATORY REGULATION

• Exhalation of carbon dioxide. By increasing the rate and depth of breathing, more carbon dioxide can be exhaled. Within minutes this reduces the level of carbonic acid in blood, which raises the blood pH (reduces blood H+ level).

• Respiratory mechanisms take several minutes to hours

Respiratory mechanisms

• 2nd Line of Defence • Exhalation of carbon dioxide• Powerful, but only works with volatile acids• Doesn’t affect fixed acids like lactic acid• CO2 + H2O ↔ H2CO3 ↔ H+ + HCO3

-

• Body pH can be adjusted by changing rate and depth of breathing

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Volatile and Non Volatile acid secretion

The peripheral chemoreceptors also respond to pH changes caused by PCO2 changes, however they directly monitor changes in the arterial blood, not the cerebrospinal fluid as the central chemoreceptors do.

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↑ ↓

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↑

↑

↑

↑

↓

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The peripheral chemoreceptors also respond to acids such as lactic acid, which is produced during strenuous exercise

Respiratory System is the Second Line of Defense

Increased Hydrogen Ion Concentration Stimulates Alveolar Ventilation

Increasing Alveolar Ventilation Decreases Extracellular Fluid Hydrogen Ion Concentration and Raises pH

THIRD LINE OF DEFENSE :RENAL REGULATION

• Kidney excretion of H ion. The slowest mechanism, but the only way to eliminate non volatile acids, is through their excretion in urine.

• Renal mechanisms may take several hours to 2-3 days.

• Changes are slow but powerful.

The Renal System mechanism

1. Regulation of plasma HCO3-

2. Excretion of fixed (metabolic) acid load

…..Most of the time the urine is acidic to balance metabolic acid production

MAJOR MECHANISMS OF RENAL REGULATION

1. Reabsorption of filtered HCO3-

2. Excretion of fixed H+

Reabsorption of filtered HCO3-

• The process results in net reabsorption of filtered HCO3

-. However, it does not result in net secretion of H+.

Regulation of reabsorption of filtered HCO3-

Filtered load PCO2

ECF volume Angiotensin II

2. Excretion of fixed H+

• Fixed H+ produced in the body is excreted by two mechanisms,

• Simultaneously excreting urinary buffers (titratable acid )

• Attaching H+ ion to ammonia (NH3) and excrete it as (NH4

+)

a. Excretion of H+ as titratable acid (H2PO4-)

b. Excretion of H+ as NH4+

• In Alkalosis• there is an excess of

HCO3– over H+ in the tubular filterate, the excess HCO3– cannot be reabsorbed; therefore, the excess HCO3– is left in the tubules and eventually excreted into the urine, which helps correct the metabolic alkalosis.

• In Acidosis• there is excess H+

relative to HCO3–, causing complete reabsorption of the bicarbonate; the excess H+ passes into the urine.

• The excess H+ is buffered in the tubules by phosphate and exc

• And excreted as ammonium ion

PLEASE REMEMBER !!!

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Normal Values

pH 7.35 – 7.45

Bicarbonate 22-26mmol/L

Chloride 96-106mmol/L

Potassium 3.5-5mmol/L

Sodium 136-145mmol/L

pO2 95(85-100) mmHg

pCO2 40(35-45) mmHg

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COMA

CRAMPS

Acid-Base Imbalances

• pH< 7.35 acidosis• pH > 7.45 alkalosis• The body response to acid-base imbalance is

called compensation• May be complete if brought back within

normal limits• Partial compensation if range is still outside

norms.

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Case #2

• 36 year old heroin addict found unresponsive with needle in arm

• Pulse = 102, BP = 110/80, Temp = 35.2 C• ABG(Arterial Blood Gas): PaO2 = 70, PaCO2 = 80,

• pH = 7.00, HCO3- = 23mEq/L

Respiratory Acidosis

• Carbonic acid excess caused by blood levels of CO2 above 45 mm Hg.

• Hypercapnia – high levels of CO2 in blood

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Causes

DEPRESSION OF THE RESPIRATORY CENTRE

• Head Injury• Anaesthetics, sedatives

(morphine )

DECREASED FUNCTIONING OF LUNGS

• Pneumonia • Bronchitis • Asthma • Pneumothorax• COPD (Emphysema)• ARDS- Adult Respiratory Distress

Syndrome

• Motor neuron disease

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Treatment of Respiratory Acidosis

• MOST IMP - Restore ventilation• IV lactate solution• Treat underlying dysfunction or disease

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Question :- Why is lactate used ??

Case #3

• 16 year old with closed head injury after a fall from 15 feet

• P = 132, BP = 115/90, • T = 37.2 C• ABG: PaO2 = 110, PaCO2 = 26,

• pH = 7.52, HCO3- = 22

Respiratory Alkalosis

• Carbonic acid deficit• pCO2 less than 35 mm Hg (hypocapnea)• Most common acid-base imbalance

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Causes

• Hyperventilation(most common )– Anxiety, Hysteria etc

• Conditions that stimulate respiratory center:– Oxygen deficiency at high altitudes– Pulmonary disease and Congestive heart failure – caused by

hypoxia – Acute anxiety– Fever, anemia– Meningitis– Cirrhosis– Gram-negative sepsis

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Compensation of Respiratory Alkalosis

• Kidneys conserve hydrogen ion

• Excrete more bicarbonate ion( i.e less is resorbed)

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Treatment of Respiratory Alkalosis

• Treat underlying cause

• Breathe into a paper bag

• IV Chloride containing solution – Cl- ions replace lost bicarbonate ions

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Case #4

• 22 year old diabetic found unresponsive• P = 102, BP = 110/80, • T = 36.2 C• ABG: PaO2 = 90, PaCO2 = 36,

• pH = 7.12, HCO3- = 8

Metabolic Acidosis

• Bicarbonate deficit - blood concentrations of bicarb drop below 22mEq/L

• Causes:– Loss of bicarbonate through diarrhea or renal

dysfunction(Type 2 RTA)– Accumulation of acids (lactic acid or ketones)– Failure of kidneys to excrete H+ (Type 1 and Type

4 RTA)

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Anion gap; Difference b/w measured cations and measured anions .

• Actually the sum of CATIONS and ANIONS in ECF is always equal.

• There is no gap whatsoever .

• The unmeasured anions constitute the anion gap .( 10± 2mmol/L)

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Anion Gap In Metabolic Acidosis

• Anion gap: [Na+] - ([Cl-] + [HCO3

-]) = 8-12 mmol/L

• If > 18, there are unmeasured anions, such as:– lactate– ketones– salicylate– ethanol– ethylene glycol (anti-freeze)

MUDPILES

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High Anion-Gap Acidosis1. Ketoacidosis

•Diabetic ketoacidosis •Starvation ketoacidosis

2. Lactic Acidosis

3. Renal Failure- Excretion of H+ and regeneration of HCO3- DEFICIENT

4. Toxins

•Ethylene glycol •Methanol•Salicylates

MUDPILES (methanol, uremia, diabetic ketoacidosis, propylene glycol, isoniazid, lactic acidosis, ethylene glycol, salicylates)

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Normal Anion-Gap Acidosis(Loss of both CATIONS AND ANIONS)

1. Renal Causes

•Renal tubular acidosis•Carbonic anhydrase inhibitors

2. GIT Causes

•Severe diarrhoea •Uretero-enterostomy or Obstructed ileal conduit•Drainage of pancreatic or biliary secretions•Small bowel fistula

3. Other Causes

•Addition of HCl, NH4Cl

Compensation for Metabolic Acidosis

• Increased ventilation- to decrease volatile acid• Increased reapsorption of HCO3- by kidneys• Renal excretion of hydrogen ions if possible• K+ exchanges with excess H+ in ECF• ( H+ into cells, K+ out of cells)

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Case #5

• 6 week old infant is lethargic with history of vomiting increasing for 1 week

• P = 122, BP = 85/60, • T = 37.2 C• ABG analysis: PaO2 = 90, PaCO2 = 44,

• pH = 7.62, HCO3- = 36,

Metabolic Alkalosis

• Bicarbonate excess - concentration in blood is greater than 26 mEq/L

• Causes:– Excess vomiting = loss of stomach acid– Excessive use of alkaline drugs,antacids(NaHCO3)– Excess aldosterone

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Diagnosis of Acid-Base Imbalances

1. Note whether the pH is low (acidosis) or high (alkalosis)

2. Decide which value, pCO2 or HCO3- , is outside the

normal range and could be the cause of the problem.

• If the cause is a change in pCO2, the problem is respiratory.

• If the cause is HCO3- the problem is metabolic.

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3. Look at the value that doesn’t correspond to the observed pH change.

If it is inside the normal range, there is no compensation occurring.

If it is outside the normal range, the body is partially compensating for the problem.

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Example

• A patient is in intensive care because he suffered a severe myocardial infarction 3 days ago. The lab reports the following values from an arterial blood sample:– pH 7.3– HCO3- = 20 mEq / L ( 22 - 26)– pCO2 = 32 mm Hg (35 - 45)

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Diagnosis

• Metabolic acidosis• With compensation

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Q

• pH 7.58;

• Pa.CO2 23 mm Hg;

• [HCO3-] 18 mEq/L

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acid base conditionpH 7.58; Pa.CO2 23 mm Hg; [HCO3

-] 18 mEq/L

1. Look at pH (is it acidosis or alkalosis?)

pH = 7.58 alkalosis

2. Look at HCO3- (is it metabolic alkalosis?)

HCO3- = 18 mEq/L (normal 22-30) not metabolic

alkalosis3. Look at Pa.CO2 (is it respiratory alkalosis?)

Pa.CO2 = 23 mmHg (normal 35-45) respiratory alkalosis4. See if appropriate compensation has occurred:

compensation for respiratory alkalosis is HCO3-

excretionHCO3

- = 18 mmHg (normal 22-30) partially compensated respiratory alkalosis

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Case F2: acid base condition

pH 7.29; Pa.CO2 26 mm Hg; [HCO3-] 12 mEq/L

1. Look at pH (is it acidosis or alkalosis?)

pH = 7.29 acidosis

2. Look at HCO3- (is it metabolic acidosis?)

HCO3- = 12 mEq/L (normal 22-30) metabolic

acidosis3. Look at Pa.CO2 (is it respiratory acidosis?)

Pa.CO2 = 26 mmHg (normal 35-45) not resp. acidosis4. See if appropriate compensation has occurred:

compensation for metabolic acidosis is hyperventilationPa.CO2 = 26 mmHg (normal 35-45); partial compensation

• In the patient described , which of the following laboratory results would be expected, compared with normal?

A) Increased renal excretion of HCO3-

B) Decreased urinary titratable acidC) Increased urine pHD) Increased renal excretion of NH4

+

Mixed disturbances

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Here several problems of acid-base management are colliding at the same time.

It’s definitely not just a matter of the body trying to compensate for one such disorder.

a. An example would be a DIABETIC with KETOACIDOSIS, who also happens tohave C.O.P.D, or develops a bad PNEUMONIA (and as a result develops a respiratory acidosis.)

Davenport diagram showing the relationships among HCO3-, pH, and PCO2. Normal buffer line BAC

Davenport diagram showing the relationships among HCO3, pH, and PCO2. . B shows the changes/compensation occurring in respiratory and metabolic acidosis and alkalosis

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Which point on the graph below would most likely represent the systemic arterial blood of a mountain climber after several weeks at high altitude?

Which arrow on the graph below could represent the change in status of an individual with metabolic acidosis who was then given an intravenous injection of sodium bicarbonate?

114Siggard Andersen Normogram