Acid-Base Physiology

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pH Review

• pH = - log [H+]

• H+ is really a proton!!

• Range is from 0 – 14

• If [H+] is high, the solution is acidic; pH < 7

• If [H+] is low, the solution is basic or alkaline ; pH > 7

How Can You Actually Determine the pH of a Solution?

• Use a pH meter.

• Litmus paper – acidic or alkaline.

• Use pH paper (color patterns indicate pH).

• Titrate the solution with precise amounts of base or acid in conjunction with a soluble dye, like Phenolphthalein, whose color changes when a specific pH is reached.

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4

pH scale – to express hydrogen ion concentration.

pH = - log10 [H+] or

pH = 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 nEq/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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7

10-2

10-3

10-5

10-4

10-8

10-7

10-6

[H+] M

10-10

10-9

10-11

10-12

10-13

10-14

10-1

100 A strong acid

A strong base

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

• Nearly60-80mEq of fixed acids/day• 1 mol of glucose 2 moles of lactate• 3g Sulphuric acid and 3g Phosphoric acid /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 HENDERSON K A HASSELBALCH

Simplified form……

• pH = pKa + log ([A-]

[HA])

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

• pH = pKa + log(Proton acceptorProton donor )

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.

Relationship of pH with K +

• When H+ increases in extracellular fluid it is exchanged with K+

• Metabolic acidosis Hyperkalemia• Metabolic alkalosis Hypokalemia

• RENAL TUBULAR ACIDOSIS FAILURE TO EXCRETE H+ K+ IS LOST IN URINE HYPOKALEMIA

• Rem :- SUDDEN HYPOKALEMIA MAY DEVELOP IN CORRECTION OF ACIDOSIS AS IN DKA WITH INSULIN THERAPY

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The body produces more acids than bases• Acids taken with foods…

• Lime juice , Most fruit juices, Colas….

• Acids produced by metabolism of lipids and proteins.

• Cellular metabolism produces CO2.CO2 + H2O ↔ H2CO3 ↔ H+ + HCO3

-

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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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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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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Three major mechanisms1. 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.

2. 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).

3. Kidney excretion of H ion. The slowest mechanism, but the only way to eliminate acids other than carbonic acid, is through their excretion in urine.

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Rates of correction

• Buffers function almost instantaneously

• Respiratory mechanisms take several minutes to hours

• Renal mechanisms may take several hours to days

Buffers • Defn:- Solutions which can resist changes in pH when

acid or alkali is added.• COMPOSITION OF A BUFFER :-

• A) Mixture of weak acids with their salt with a strong base• Mixtures of weak bases with their salt with a strong acid.eg

• H2CO3/NaHCO3 ( BICARBONATE BUFFER)• CH3COOH/CH3COONa (ACETATE BUFFER)• NaHPO4/NaH2PO4 ( PHOSPHATE BUFFER)

•

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BUFFER SYSTEMS IN THE BODY

• FIRST LINE OF DEFENSE.

• THEY ARE EFFECTIVE AS LONG AS THE ACID LOAD IS NOT VERY HIGH .

• THE BODY’S ALKALI RESERVE SHOULD NOT BE EXHAUSTED- THIS HAS TO BE REPLENISHED ONCE EXHAUSTED.

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Buffering of hydrogen Ions in the body fluids

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

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

Buffer + H+ ↔ Hbuffer

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}

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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Relationship between (H+) and the members of a buffer pair is expresses using-Henderson-Hasselbalch Equation

pH = pKa + log[HCO3-] / s*[PCO2 ]

pKa = 6.1(dissociation constant)

What Is The Central Message Of Henderson-Hasselbalch?

pH = pKa + log(HCO3- / s.PCO2)

Plasma pH is a simple function of the HCO3- : PCO2

ratioHCO3

- : PCO2 ↑ = pH ↑ (ALKALOSIS) : Could occur due to either HCO3

- ↑(Metabolic alkalosis) or PCO2 ↓ (respiratory alkalosis)HCO3

- : PCO2 ↓ = pH ↓( ACIDOSIS) : Could occur either HCO3

- ↓(metabolic acidosis)

or PCO2 ↑ (respiratory acidosis)

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

pH 7.2, HCO3- 15 mM and PCO2 40 mm Hg ?

metabolic acidosis

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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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-

INTRACELLULAR BUFFERS ARE VERY IMPORTANT

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6

40

1 1 BUFFERS

TISSUE CELLSRBCBICARBONATE BUFFERPHOSPHATE BUFFER(EXTRACELLULAR )PROTEIN (EXTRACEL-LULAR)

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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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• In THE LUNGS :-HHb + O2 HbO2 + H+

HCO3 - + H+ H2CO3

H2CO3 H2O + CO2

THE ACTIVITY OF CARBONIC ANHYDRASE ACTIVITY ALSO INCREASES IN ACIDOSIS AND DECREASES WITH DECREASED H+.

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2. 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

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.

↑

↑ ↓

↑

↑

↑

↑

↑

↓

↑

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

Central Chemoreceptors: Effect of PCO2 IN REGULATING VENTILATION

CO2 + H2 0 H2 CO3 H+ + HCO-

3

capillary Ventricle

central chemoreceptors

respiratory centers in the medulla

↑CO2

↑ H+

↓pH

• As carbon dioxide increases, so does the number of hydrogen ions, which in turn lowers the pH. The central chemoreceptors actually respond to this pH change caused by the blood PCO2.

Blood brain Barrier

Breath holding

Cellular Respiration Produces CO2 And “Metabolic Acids”

CellsFood

ECF

Lung

CO2

CO2

CO2

H+ + HC03-

Buffering metabolic acid consumes ECF HC0-

3

• Rate of respiration is controlled by chemoreceptors in the respiratory centre– sensitive to pH changes in blood.

• FALL in pH of plasma•

HYPERVENTILATION

• MORE CO2 ELIMINATED

• H2CO3 REMOVED• pH increased)

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

Increased Hydrogen Ion Concentration Stimulates Alveolar Ventilation

The Renal System Is The 3rd Line Of Defense. Changes Are Slow But Powerful

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

RENAL REGULATION

• Can eliminate large amounts of acid.• Can also excrete base .• Can conserve and produce bicarb ions

• MOST EFFECTIVE REGULATOR OF pH• If kidneys fail, pH balance fails

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Normal Urine(freshly passed) has a pH around 6,i.e lower than plasma ; ACIDIFICATION OF URINE

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MAJOR MECHANISMS OF RENAL REGULATION1. SECRETION OF H+

2. RECOVERY OF HCO3- BY REABSORPTION

3. BICARBONATE IONS ARE “TITRATED” AGAINST HYDROGEN IONS

4. COMBINATION OF EXCESS HYDROGEN IONS WITH PHOSPHATE AND AMMONIA BUFFERS IN THE TUBULE—A MECHANISM FOR GENERATING “NEW” BICARBONATE IONS

5. PRIMARY ACTIVE SECRETION OF HYDROGEN IONS IN THE INTERCALATED CELLS OF LATE DISTAL AND COLLECTING TUBULES

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SECRETION OF H+ IN PROXIMAL CONVOLUTED TUBULE AND RECOVERY OF HCO3

- BY REABSORPTION

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H2O + CO2

H2CO3

HCO3- + H+

BLOOD PCT –CELL TUBULAR LUMEN

H+

Na+Na+Na+

Alkali is recovered

HCO3-

CARBONIC ANHYDRASE

Bicarbonate Ions Are “Titrated” Against Hydrogen Ions in the Tubules

55H2O + CO2

H2CO3

HCO3- + H+

BLOOD TUBULAR CELL TUBULAR LUMEN

H+

Na+Na+Na+

Alkali is recovered

HCO3-

CARBONIC ANHYDRASE

NaHCO3

HCO3-

H2CO3

CARBONIC ANHYDRASE

H2O + CO2

Acid Secretion In The Proximal Tubule Recovers Filtered HCO3

-

Lumen Blood

HCO3-

3Na+

2K+

Na+

NHE3

H+

NBCNa+

HCO3-

CO2

H2O +

H2CO3

CA = carbonic anhydrase

H2CO3

CO2

CAH2O + CA

filtration

CO2

VERY LITTLE ACID EXCRETION OCCURS.

• 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 ammonia and eventually excreted as salts.

Excretion Of “Titratable Acid” Also Generates New HC03

-

Lumen Blood

3Na+

2K+

Na+

NHE3

NBCNa+

HCO3-

CO2

H2O +

H2CO3

CA

H+HPO42-

filtration

H2PO4-

Proximal tubule cell

Phosphate Buffer System

59H2O + CO2

H2CO3

HCO3- + H+

BLOOD (DT)TUBULAR CELL TUBULAR LUMEN

H+

Na+Na+Na+

Alkali is recovered

HCO3-

CARBONIC ANHYDRASE

Na2HPO4

NaHPO4-

NaH2PO4

EXCRETED

pH5.4

pH7.4

H+ EXCRETED

• Phosphate Buffer System Carries Excess Hydrogen Ions into the Urine and Generates New Bicarbonate

• Excretion of Excess Hydrogen Ions and Generation of New Bicarbonate by the Ammonia Buffer System

Summary Of Renal Acid Base Handling• Functions of the renal system in acid base balance• Mechanisms for acid excretion, bicarbonate reabsorption and new

bicarbonate generation.• Renal responses to acid base disorders• Interactions between volume and potassium balance and acid-base

balance

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

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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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Acidosis• Principal effect of acidosis is depression of the CNS through ↓

in synaptic transmission.

• Generalized weakness

• Deranged CNS function the greatest threat

• Severe acidosis causes • Disorientation• coma • death

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Alkalosis• Alkalosis causes over excitability of the central and peripheral nervous

systems.

• Numbness

• Lightheadedness

• It can cause :• Nervousness• muscle spasms , cramps • Convulsions • Loss of consciousness• Coma• Death

Primary Changes and Compensations in Simple Acid-Base Disorders

Primary Disturbance

pH HCO3− Pco2 Prediction of Compensation

Metabolic acidosis

< 7.35 Primary decrease Compensatory decrease

1.2 mm Hg decrease in Pco2 for every 1 mmol/L decrease in HCO3

−

Metabolic alkalosis

> 7.45 Primary increase Compensatory increase

0.6–0.75 mm Hg increase in Pco2 for every 1 mmol/L increase in HCO3

− (Pco2 should not rise above 55 mm Hg in compensation)

Respiratory acidosis

< 7.35 Compensatory increase

Primary increase

Acute: 1–2 mmol/L increase in HCO3− for

every 10 mm Hg increase in Pco2

Chronic: 3–4 mmol/L increase in HCO3−

for every 10 mm Hg increase in Pco2

Respiratory alkalosis

> 7.45 Compensatory decrease

Primary decrease

Acute: 1–2 mmol/L decrease in HCO3− for

every 10 mm Hg decrease in Pco2 Chronic: 4–5 mmol/L decrease in HCO3

− for every 10 mm Hg decrease in Pco2

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http://animalsbeingdicks.com/page/6

Clinical Definitions and Diagnostic Aids

• Respiratory acidosis = PaCO2 > 45 mmHg• Respiratory alkalosis = PaCO2 < 35 mmHg• Metabolic acidosis = HCO3

- < 22 mmHg or Base Deficit of < -2

• Metabolic alkalosis = HCO3- > 28 mmHg or Base

Excess of > +2

Acid - Base Diagnosis

PaCO2< 35 or >45?

No VentilatoryComponent

PaCO2< 35?

HCO3-<21 or >28?

No

PaCO2>45?

VentilatoryAlkalosis

VentilatoryAcidosis

Yes

Yes

Yes

No

No MetabolicComponent

HCO3->28?

HCO3-<21?

No

Yes

MetabolicAlkalosis

MetabolicAcidosis

Yes

No

Yes

pH<7.35?

Acidemia Yes

pH>7.45?

No

Alkalemia YesNormal pH

NoDiagram source unknown

Case #2• 36 year old heroin addict found unresponsive with

needle in arm• P = 102, BP = 110/80, T = 35.2 C• ABG: PaO2 = 70, PaCO2 = 80,• pH = 7.00, HCO3

- = 23

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Respiratory Acidosis

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

• Hypercapnia – high levels of CO2 in blood

Causes

DECREASED FUNCTIONING OF LUNGS

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

• Motor neuron disease

• DEPRESSION OF THE RESPIRATORY CENTRE

• Head Injury• Anaesthetics, sedatives (morphine )

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Signs and Symptoms of Respiratory Acidosis

• Breathlessness• Restlessness• Lethargy and disorientation• Tremors, convulsions, coma.• Respiratory rate rapid, then gradually depressed.• Skin warm and flushed due to vasodilation caused by

excess CO2

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Compensation for Respiratory Acidosis

• Kidneys eliminate hydrogen ion and retain bicarbonate ion

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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

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Respiratory Alkalosis

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

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

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

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Metabolic Acidosis• Bicarbonate deficit - blood concentrations of bicarb drop

below 22mEq/L• Causes:

• Loss of bicarbonate through diarrhea or renal dysfunction• Accumulation of acids (lactic acid or ketones)• Failure of kidneys to excrete H+

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Symptoms of Metabolic Acidosis

• Headache, lethargy• Nausea, vomiting, diarrhea• Coma• Death

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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 constiute the anion gap .( 12± 5 mmol/L)

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Anion Gap In Metabolic Acidosis• Anion gap:

[Na+] - ([Cl-] + [HCO3-]) = 8-16 mmol/L

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

Explanation• Say , 5mmol/L Lactic Acid has entered the circulation

Lactate + H+

• Buffered by HCO3-

• 5mmol/L Lactate + 5mmol/L H2CO3

• H2CO3 H2O + CO2 (LUNGS)• Finally what has happened

• HCO3- LOWERED +5 mmol/L of UNMEASURED ANION

(LACTATE )• NO CHANGES IN Na+/K+

• ELEVATED ANION GAP

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So what does this mean?

• Lactic Acid + HCO3 ↔ lactate- + H2O + CO2

So increasing Lactic acid leads to lactate replacing HCO3

If anion gap is unchanged in metabolic acidosis suggest other reason for acidosis (eg diarrhoea – loss of HCO3 but gain in Cl-

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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

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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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I.V NaHCO3 is kept as a last reserve !

• Never give in Cl- losing situations e.g Vomiting• Never in congestive cardiac failure and renal insufficiency • Can cause hypernatremia especially dangerous in children• Celllulitis around the site of infusion

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 PaO2 = 90, PaCO2 = 44, • pH = 7.62, HCO3

- = 36,

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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• Certain diuretics• Endocrine disorders:Hyperaldosteronism• Heavy ingestion of antacids• Severe dehydration

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Symptoms of Metabolic Alkalosis

• Respiration slow and shallow• Hyperactive reflexes ; tetany• Often related to depletion of electrolytes• Atrial tachycardia• Dysrhythmias

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Compensation for Metabolic Alkalosis

• RENAL COMPENSATION – decreased reabsorption of HCO3-• Kidneys conserve H+ ions

• Sometimes , Alkalosis occurs with renal dysfunction can’t count on kidneys

• Respiratory compensation difficult – hypoventilation limited by hypoxia

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Treatment of Metabolic Alkalosis

• Electrolytes to replace those lost• IV chloride containing solution• Treat underlying disorder

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Diagnosis of Acid-Base Imbalances1. 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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acid base condition

pH 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 alkalosis

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

Pa.CO2 = 23 mmHg (normal 35-45) respiratory alkalosis

4. See if appropriate compensation has occurred:

compensation for respiratory alkalosis is HCO3- excretion

HCO3- = 18 mmHg (normal 22-30)

partially compensated respiratory alkalosis

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Case F2: acid base conditionpH 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 acidosis

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

Pa.CO2 = 26 mmHg (normal 35-45) not resp. acidosis

4. See if appropriate compensation has occurred:

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

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.)

124Siggard Andersen Normogram

125

• Recall HH – compensation aims to normalize pH by restoring [HCO3]:PCO2 ratio towards normal.

• The “Primary” disturbance is the one that is consistent with the pH

Comments On Compensation…

Mixed Acid-Base Disorders

• Most common acid-base disorders• Multiple disorders• Usually one acidosis and one alkalosis• pH usually partially or completely corrected

Key Points

• Acid-base disorders are common and important clinical concerns• Accurate diagnosis is essential to proper treatment• Primary disorders are complicated by secondary disorders occurring

at a different time course