Acid-Base Biochemistry Prof. Dr. M. Sait KELEŞ Department of Biochemistry

Acid-Base BiochemistryProf. Dr. M. Sait KELEŞ

Department of Biochemistry

Acid-Base Biochemistry

WHAT THE SOURCES OF ACID IN THE BODY?


Sources of acidMetabolism of foodMetabolism of drugs Inborn errors of metabolism


Acid production from metabolism of foodSulphuric acid from metabolism of sulphur-

containing amino acids of proteinsLactic acid from sugarsKetoacids from fats


Acid production from metabolism of drugs Direct metabolism of drug to more acidic

compound eg. salicylates, urates etc.Induction of enzymes which metabolise other

compounds (endogenous or exogenous) to acids


Inborn errors of metabolism Organic acid disordersLactic acidosis


Greatest potential source of acid Carbon dioxide

(1) CO2 + H2O <=> H2CO3

(2) H2CO3 <=> H+ + HCO3-

Potentially 15,000 mmol/24 hours


Hydrogen ion homeostasis1. buffering2. excretion


Buffering of hydrogen ionsIn health as hydrogen ions are produced they are buffered – limiting the rise in [H+]


Buffer solutions consist of a weak acid and its conjugate baseAs hydrogen ions are added some will combine with the conjugate base and convert it to undissociated acid

Buffer Know-how

Buffers are important in biochemical processes. Whether they occur naturally in plasma or in the cytosol of cells, buffers assure biological reactions occur under conditions of optimal pH. They do this by controlling the hydrogen ion concentration of solutions. The word “buffer” is so common in biochemistry its replaces the word “water” in experimental protocols. For example, typical of the statement seen in publications is “the pellet was dissolved in pH 7.5 buffer”. All this should alert you to the importance of thoroughly understanding buffers and buffering agents. Words such as pH, pKa, conjugate acid, conjugate base, Henderson-Hasselbalch equation are used frequently in biochemical language and every publication that describes an experiment performed “in vitro” (Lat., in glass), must include a clear description of the buffer that was used. In this tutorial we will revisit buffers and attempt to understand their make up and mechanism of action. We will also give some insights into how to solve buffer problems.

Rules Governing Buffer Reactions

To help you understand buffer action, consider the equation that describes a buffer reaction. The HA and A- represent the two components of any buffer: the conjugate acid and the conjugate base, respectively. Note that the right side and left side of the equation are the same. Rule1 gives the meaning of the reaction .

Rules

HA + A- + H2O HA + A- + H2O

1. Buffer reactions never go to completion.

2. H+ can only react with A-, OH- can only react with HA.

3. The conjugate acid and conjugate base must change proportionately and in opposite directions.

Adding NaOH or HCl to this buffer would shift the reaction toward the right, but with different results. This is because of Rule 2. Both components in the buffer must change any time acid or base is added to the buffer. This is because of Rule 3.

Finally, we take into account the HA and A- components, expecting to see a decline in the overall buffer. Such is not found because of Rule 4 . Hence, Rule 5 summarizes an important principle that we must know .

4. The sum of the concentrations of base and acid components stays constant, i.e, HA + A- before = HA + A- after.

5. Only the ratio, never the total, of HA to A- changes as a result of buffer action.

Rules

Putting the Rules to Work

Let’s now see how the 5 rules apply to a buffer reaction. Lets start with the buffer. Suppose we add OH- to the buffer . Rule 2 tells us the following would occur

HA + A- + H2O

OH-

HA + A- + H2O

We see that HA decreases and A- increases, both to the same increment. This illustrates Rule 3. H2O would also increase, but we can ignore H2O because the increase would be insignificant compared to the amount of H2O present .

2 moles

5 moles 5 moles

Finally, if we calculate HA and A- before and after the addition of OH-, we see that both add up to 10 moles in either case . This verifies Rule 4. Thus, we started with a ratio of A- to HA = 5/5 = 1.0 before the reaction and after the reaction the ratio of A to HA- = 7/3= 2.3 to 1.0, but the total did not change.

10 moles, ratio = 1.0

10 moles, ratio A-/HA = 2.43/1.0

3 moles 7 moles

If we applied numbers to the concentrations of OH-, HA and A we can test the other rules . Here we see that we are adding 2 moles of OH- to 5 moles each of HA and A-. The reaction of OH- with HA lowers the HA from 5 to 3 and raises the A- from 5 to 7 . The decrease in HA was matched by an equal increase in A-, which is Rule 3.

Principle Behind Buffer Action

Buffers are composed of weak acids and their salts. A salt is the acid minus its proton. Weak acids and their salts have two properties that are important for buffering action. First, weak acids are a reserve of the protons that neutralize OH- and prevent the solution from becoming alkaline. Salts of weak acids are strong bases and prevent the solution from becoming acidic. Both components are needed and both are interchangeable through the loss (or gain) of a single proton.

HA

HA

HAHA HA HAHA

HA HA

HAHA

Reserve acid11 moles

Reserve salt11 moles

A-A-

A-

A-

A-A-

A-

A-

A-

A-

A-

OH-

A-

A-

A-

A-

A-

6 moles 16 moles

OH

OH-

Neutralized

OH-

Neutralized

A buffer’s power lies in its reserves. A buffer is at optimal strength when there is an equal amount of HA and A- in solution as shown. This will only occur when the pH of the solution equals the pKa of the acid’s group. Adding OH- causes the buffer to respond by calling on the reserve pool of HA. A- is formed at the expense of HA. This continues until all the excess OH- is neutralized. At the end the salt pool has increased (and the acid pool has decreased) by the same number of moles of base that were added.

Focus on the Ratio of [A-]/[HA]

In the previous illustration you saw the importance of knowing the ratio of HA and A-. Now you will see that it is the ratio that determines the pH of the solution, and vice versa, the pH allows you to determine the ratio. It all begins with an equilibrium expression (click 1). If we take the log of all components we derive a logarimic expression of the same equation (click 1),

[H+] = Keq [HA] [A-] Log [H+] = Log Keq + Log

[HA] [A-]

Mutiplying components on both sides of the equation by -1 gives (click 1)

–Log [H+] = –Log Keq – Log [HA] [A-]

Substituting pH and pK for the appropriate terms in the equation and making the log of HA/A- positive by reversing numerator and denominator gives (click 1)

pH = pK + Log [A-] [HA]

Note, in the equation, pK is a constant and A/HA is the only variable (click 1). This is the Henderson-Hasselbalch equation. Click to go on.

constant variable

Putting Henderson-Hasselbalch to use

Knowing the ratio of [A-]/[HA] allows you to calculate pH. Always treat the ratio as a whole number, i.e., do not separate numerator from denominator. As an example, assume 2 moles of NaOH are added to 10 moles of a pH 5.2, pK = 4.8 buffer (HA + A-). You want to know the moles of HA after neutralization and the new pH. Follow these steps to the solutionFirst determine the moles of HA at the start

pH = pK + Log [A-] [HA]

Solving for Log [A-] [HA]

[A-] [HA]Log = pH – pK

= 5.2 – 4.8

= 0.4 [A-] [HA]= 2.5 / 1.0

The ratio of A- to HA is 2.5 parts to 1 part. This means the 10 moles are represented by 3.5 parts. If 2.5 parts of the 10 are moles of A- and 1.0 part is HA, then before OH- was added there were 7.1 moles of A- and 2.9 moles of HA. Together the two add up to 10 and their ratio is 2.5:1.0 . When OH- is added, 2.0 moles of NaOH react with 2.9 moles of HA. As a consequence, HA goes from 2.9 to 0.9 and A- goes up from 7.1 to 9.1 moles. The new pH is determined from the ratio 9.1 to 0.9 or 10.1. This computes to pH = 5.8

The common ion effect is the shift in equilibrium caused by the addition of a compound having an ion in common with the dissolved substance.

The presence of a common ion suppresses the ionization of a weak acid or a weak base.

Consider mixture of CH3COONa (strong electrolyte) and CH3COOH (weak acid).

CH3COONa (s) Na+ (aq) + CH3COO- (aq)

CH3COOH (aq) H+ (aq) + CH3COO- (aq)

common ion

16.2

Consider mixture of salt NaA and weak acid HA.

HA (aq) H+ (aq) + A- (aq)

NaA (s) Na+ (aq) + A- (aq)Ka =

[H+][A-][HA]

[H+] =Ka [HA]

[A-]

-log [H+] = -log Ka - log[HA]

[A-]

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

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

pKa = -log Ka

Henderson-Hasselbalch equation

16.2

pH = pKa + log[conjugate base]

[acid]

A buffer solution is a solution of:

1. A weak acid or a weak base and

2. The salt of the weak acid or weak base

Both must be present!

A buffer solution has the ability to resist changes in pH upon the addition of small amounts of either acid or base.

16.3

Add strong acid

H+ (aq) + CH3COO- (aq) CH3COOH (aq)

Add strong base

OH- (aq) + CH3COOH (aq) CH3COO- (aq) + H2O (l)

Consider an equal molar mixture of CH3COOH and CH3COONa

HCl H+ + Cl-

HCl + CH3COO- CH3COOH + Cl-

16.3

Which of the following are buffer systems? (a) KF/HF (b) KBr/HBr, (c) Na2CO3/NaHCO3

(a) KF is a weak acid and F- is its conjugate base

buffer solution

(b) HBr is a strong acid

not a buffer solution

(c) CO32- is a weak base and HCO3

- is it conjugate acid

buffer solution

16.3

Chemistry In Action: Maintaining the pH of Blood

16.3

TitrationsIn a titration a solution of accurately known concentration is added gradually added to another solution of unknown concentration until the chemical reaction between the two solutions is complete.

Equivalence point – the point at which the reaction is complete

Indicator – substance that changes color at (or near) the equivalence point

Slowly add baseto unknown acid

UNTIL

The indicatorchanges color

(pink)

4.7

Strong Acid-Strong Base Titrations

NaOH (aq) + HCl (aq) H2O (l) + NaCl (aq)

OH- (aq) + H+ (aq) H2O (l)

16.4

Weak Acid-Strong Base Titrations

CH3COOH (aq) + NaOH (aq) CH3COONa (aq) + H2O (l)

CH3COOH (aq) + OH- (aq) CH3COO- (aq) + H2O (l)

CH3COO- (aq) + H2O (l) OH- (aq) + CH3COOH (aq)

At equivalence point (pH > 7):

16.4

Strong Acid-Weak Base Titrations

HCl (aq) + NH3 (aq) NH4Cl (aq)

NH4+ (aq) + H2O (l) NH3 (aq) + H+ (aq)

At equivalence point (pH < 7):

16.4

H+ (aq) + NH3 (aq) NH4Cl (aq)

Exactly 100 mL of 0.10 M HNO2 are titrated with a 0.10 M NaOH solution. What is the pH at the equivalence point ?

HNO2 (aq) + OH- (aq) NO2- (aq) + H2O (l)

start (moles)

end (moles)

0.01 0.01

0.0 0.0 0.01

NO2- (aq) + H2O (l) OH- (aq) + HNO2 (aq)

Initial (M)

Change (M)

Equilibrium (M)

0.05 0.00

-x +x

0.05 - x

0.00

+x

x x

[NO2-] =

0.01

0.200= 0.05 MFinal volume = 200 mL

Kb =[OH-][HNO2]

[NO2-]

=x2

0.05-x= 2.2 x 10-11

0.05 – x 0.05 x 1.05 x 10-6

= [OH-]

pOH = 5.98

pH = 14 – pOH = 8.02

Acid-Base Indicators

HIn (aq) H+ (aq) + In- (aq)

10[HIn]

[In-]Color of acid (HIn) predominates

10[HIn]

[In-]Color of conjugate base (In-) predominates

16.5

The titration curve of a strong acid with a strong base.

16.5

Which indicator(s) would you use for a titration of HNO2 with KOH ?

Weak acid titrated with strong base.

At equivalence point, will have conjugate base of weak acid.

At equivalence point, pH > 7

Use cresol red or phenolphthalein

16.5

pH and Solubility

• The presence of a common ion decreases the solubility.

• Insoluble bases dissolve in acidic solutions• Insoluble acids dissolve in basic solutions

Mg(OH)2 (s) Mg2+ (aq) + 2OH- (aq)

Ksp = [Mg2+][OH-]2 = 1.2 x 10-11

Ksp = (s)(2s)2 = 4s3

4s3 = 1.2 x 10-11

s = 1.4 x 10-4 M

[OH-] = 2s = 2.8 x 10-4 M

pOH = 3.55 pH = 10.45

At pH less than 10.45

Lower [OH-]

OH- (aq) + H+ (aq) H2O (l)

remove

Increase solubility of Mg(OH)2

At pH greater than 10.45

Raise [OH-]

add

Decrease solubility of Mg(OH)216.9

Flame Test for Cations

lithium sodium potassium copper

16.11


Bicarbonate – carbonic acid buffer systemH+ + HCO3

- <=> H2CO3

Addition of H+ drives reaction to the rightConversely Fall in H+ drives reaction to the left as carbonic

acid dissociates producing more H+


Buffering systems in bloodBicarbonate ions-most importantProteins including intracellular proteinsHaemoglobin

Buffer Systems in Body Fluids

Figure 27–7

3 Major Buffer Systems1. Protein buffer systems:

help regulate pH in ECF and ICF interact extensively with other buffer

systems2. Carbonic acid–bicarbonate buffer system:

most important in ECF3. Phosphate buffer system:

buffers pH of ICF and urine


Buffer solutions operate most efficiently at [H+] that result in approximately equal concentration of undissociated acid and conjugate base

But at normal extracellular fluid pH [H2CO3] 1.2 mmol

whereas [HCO3-] is twenty times greater


The bicarbonate system is enhanced by the fact that carbonic acid can be formed from CO2 or disposed of by conversion to CO2

CO2 + H2O <=> H2CO3


For every hydrogen ion buffered by bicarbonate – a bicarbonate ion is consumed.

To maintain the capacity of the buffer system, the bicarbonate must be regenerated

However, when bicarbonate is formed from carbonic acid (CO2 and H2O) equimolar amounts of [H+] are formed


Bicarbonate formation can only continue if these hydrogen ions are removed

This process occurs in the cells of the renal tubules where hydrogen ions are secreted into the urine and where bicarbonate is generated and retained in the body


2 different processesBicarbonate regeneration (incorrectly

reabsorption)Hydrogen ion excretion


Importance of Renal Bicarbonate Regeneration Bicarbonate is freely filtered through the

glomerulus so plasma and glomerular filtrate have the same bicarbonate concentration

At normal GFR approx 4300 mmol of bicarbonate would be filtered in 24 hr

Without re-generation of bicarbonate the buffering capacity of the body would be depleted causing acidotic state

In health virtually all the filtered bicarbonate is recovered


Renal Bicarbonate Regeneration involves the enzyme carbonate dehydratase (carbonic anhydrase)

Luminal side of the renal tubular cells impermeable to bicarbonate ions

Carbonate dehydratase catalyses the formation of CO2 and H2O from carbonic acid (H2CO3) in the renal tubular lumen

CO2 diffuses across the luminal membrane into the tubular cells


within the renal tubular cells carbonate dehydratase catalyses the formation of carbonic acid (H2CO3) from CO2 and H2O

Carbonic acid then dissociates into H+ and HCO3-

The bicarbonate ions pass into the extracellular fluid and the hydrogen ions are secreted back into the lumen in exchange for sodium ions which pass into the extracellular fluid

Exchange of sodium and hydrogen ions an active process involving Na+/K+/H+ ATP pump

K+ important in electrolyte disturbances of acid-base


Regeneration of bicarbonate does not involve net excretion of hydrogen ions

Hydrogen ion excretion requires the same reactions occurring in the renal tubular cells but also requires a suitable buffer in urine

Principal buffer system in urine is phosphate80% of phosphate in glomerular filtrate is in the

form of the divalent anion HPO42-

This combines with hydrogen ionsHPO4

2- + H+ ↔ H2PO4-


Hydrogen ion excretion capacityThe minimum urine pH that can be generated

is 4.6 ( 25µmol/L)Normal urine output is 1.5LWithout the phosphate buffer system the free

excretion of Hydrogen ions is less than 1/1000 of the acid produced by normal metabolism


The phosphate buffer system increases hydrogen ion excretion capacity to 30-40 mmol/24 hours

In times of chronic overproduction of acid another urine buffer system

Ammonia


Ammonia produced by deamination of glutamine in renal tubular cells

Catalysed by glutaminase which is induced by chronic acidosis

Allows increased ammonia production and hence increased hydrogen ion excretion via ammonium ions

NH3 + H+ ↔ NH4+


At normal intracellular pH most ammonia is present as ammonium ions which can’t diffuse out of the cell

Diffusion of ammonia out of the cell disturbs the equilibrium between ammonia and ammonium ions causing more ammonia to be formed

Hydrogen ions formed at the same time!These are used up by the deamination of glutamine

to glutamate during gluconeogenesis


Carbon dioxide transportCarbon dioxide produced by aerobic respiration

diffuses out of cells and into the ECFA small amount combines with water to form

carbonic acid decreasing the pH of ECFIn red blood cells metabolism is anaerobic and

very little CO2 is produced hence it diffuses into red cells down a concentration gradient to form carbonic acid (carbonate dehydratase) buffered by haemoglobin .


Haemoglobin has greatest buffering capacity when it is dexoygenated hence the buffering capacity increases as oxygen is lost to the tissues

Net effect is that carbon dioxide is converted to bicarbonate in red cells

Bicarbonate diffuses out of red cells down concentration gradient and chloride ions diffuse in to maintain electrochemical neutrality (chloride shift)

•Acid-Base Biochemistry

In the lungs this process is reversedHaemoglobin is oxygenated reducing its

buffering capacity and generating hydrogen ions

These combine with bicarbonate to form CO2 which diffuses into the alveoli

Bicarbonate diffuses into the cells from the plasma


Most of the carbon dioxide in the blood is present as bicarbonate

Carbon dioxide, carbonic acid and carbamino compounds less than 1/10 th of the total

Bicarbonate /total CO2 used interchangeably though not strictly the same

Most analytical methods actually measure total CO2 as bicarbonate difficult to measure


The hydrogen ion concentration of plasma is directly proportional to the PCO2 and inversely proportional to bicarbonate

[H+] = k pCO2/[HCO3-]

[H+] in nmoles/L, [HCO3-] in mmoles/L

pCO2 in kPa k = 180

pCO2 in mm Hg k= 24


Derived bicarbonatePossible to use the equation to calculate the

bicarbonate concentration from the pCO2 and pH (blood gas analysers)

?how valid in non-ideal solutionsAuto analysers – measured bicarbonate


The relationship between [H+], pCO2 and bicarbonate fundamental to understanding pathophysiology of hydrogen ion homeostasis


4 Components to acid-base disordersGenerationBufferingCompensationCorrectionOccurring concurrently


Classification of acid-base disordersAcidosis [H+] above normal, pH below normalAlkalosis[H+] below normal, pH above normal


Further classified asRespiratoryNon-respiratory (metabolic)Mixed – difficult to distinguish between primary

mixed condition and compensated disorder


Respiratory disorders involve a change in pCO2

Metabolic disorders involve change in production or excretion of hydrogen ions or both


Non-respiratory acidosisIncreased production/reduced excretion of acid?causes


Non-respiratory acidosisOverproduction of acid

Keto acidosis (diabetes, starvation, alcohol)Lactic acidosis (inherited metabolic defect or

drugs)Inherited organic acidosesPoisoning (salicylate, ethylene glycol, alcohol)Excessive parenteral amino acids


Non-respiratory acidosisReduced excretion of acid

Generalised renal failure Renal tubular acidosesCarbonate dehydratase inhibitors


Non-respiratory acidosisLoss of Bicarbonate

DiarrhoeaPancreatic, intestinal, biliary fistula or drainage


Compensation of non-respiratory acidosisExcess hydrogen ions are buffered by bicarbonate forming carbonic acid which dissociates to carbon dioxide to be lost in expired air

The buffering limits the rise in [H+] at the expense of reduction in bicarbonate


Compensation of non-respiratory acidosisHyperventilation increases removal of CO2

lowering pCO2

PCO2 / [HCO3-] ratio falls reducing [H+]

Hyperventilation is the direct result of increased [H+] stimulating the respiratory centre (Kussmaul respiration)Respiratory compensation of non-respiratory acidosis


Compensation of non-respiratory acidosisLimitationsRespiratory compensation cannot

completely normalise the [H+] because the hyperventilation is stimulated by the increase in [H+] and as this falls the drive on the respiratory centre is reduced

Increased work of respiratory muscles during hyperventilation produces CO2 limiting the degree to which PCO2 can be lowered


The degree of compensation may be limited further if respiratory function is compromised

If it is not possible to correct the cause of the acidosis may get a new steady state of chronic acidosis

[H+] [HCO3-] and ↓PCO2


In the absence of acidosis - hyperventilation would normally generate a respiratory alkalosis

Compensatory mechanisms usually involve generation of a second opposing disturbance

In non-respiratory acidosis the hyperventilation limits the severity of the acidosis but is not great enough to cause alkalosis in the patient


Non-respiratory compensation of non-respiratory acidosis

If renal function is normal excess [H+] can be excreted by the kidneys

But renal function is often impaired even if not the primary cause of the acidosis


Correction of acidosis Complete correction requires reversal or

removal of the underlying causeEthylene glycol poisoning – slow the rate of

metabolism with ethanol Diabetes – rehydration and insulin


Summary of non-respiratory acidosispH [H+] PCO2

[HCO3-]


Management of non-respiratory acidosis1. Removal of cause2. Administration of Bicarbonate – only in

severe cases pH <7.0 and where 1 is not possible

Must be given in small quantities with constant monitoring of pH


Respiratory acidosisPrimarily an increase in PCO2

Number of different causes


Retention of CO2

Production of carbonic acidFor every hydrogen ion produced a

bicarbonate ion is generatedMost of the [H+] is buffered by intracellular

buffers (haemoglobin)Development of renal compensation if renal

function is normal


Acute respiratory acidosis For every KPa increase in PCO2

increase in bicarbonate < 1 mmoleIncrease in [H+] 5.5 nmol/L

ChronicFor every KPa increase in PCO2

increase in bicarbonate 2-3 mmoleIncrease in [H+] 2.5 nmol/L


Compensation of respiratory acidosisIncreased renal excretion of hydrogen ions


Management of respiratory acidosisWith reduced ventilation it is usually the

hypoxaemia that is life threatening 4 mins if ventilation ceases

Improve alveolar ventilation bronchodilators and antibiotics

Artificial ventilation close monitoring required to avoid over correction esp in chronic acidosis


Summary of respiratory acidosisSummary of respiratory acidosis

AcuteAcute ChronicChronic

pHpH Slight Slight or low or low normalnormal

[H[H++]] Slight Slight or or high normalhigh normal

PCOPCO22

[HCO[HCO33--]] Slight Slight


Non respiratory alkalosisLoss of un-buffered hydrogen ions

Gastrointestinal - vomiting with pyloric stenosis - diarrhoea

- nasogastric aspiration


Causes of non respiratory alkalosisRenal Mineralo-corticoid excess

Conn’s syndromeCushings syndrome

Drugs with mineralocorticoid activityDiuretic therapy (not K+ sparing)


Causes of non respiratory alkalosisAdministration of alkali

Over-treatment of acidosis

Chronic alkali ingestion (antacids)


Non respiratory alkalosisCharacterised by primary increase in ECF

bicarbonateConsequent reduction in [H+]Normally increase in bicarbonate causes

reduction in renal bicarbonate regeneration and increased urinary excretion of bicarbonate


non respiratory alkalosisMaintenance requires inappropriate renal

bicarbonate reabsorption/regeneration- decrease in ECF volume (hypovolaemia)- mineralocorticoid excess- potassium depletion


non respiratory alkalosisHypovolaemia

Increased stimulus to sodium reabsorption Dependant on adequate anionsIf chloride deficient (GI losses) electrochemical

neutrality during Na+ absorption maintained by increased bicarbonate absorption and by H+ and K+ excretion


non respiratory alkalosisMineralocorticoid excess

Alkalosis perpetuated by increased hydrogen ion excretion secondary to increased sodium reabsorption

Potassium depletionPotassium and hydrogen ion excretion compete

for exchange with sodium so depletion of potassium causes increased H+ excretion


non respiratory alkalosisCompensationLow H+ inhibits the respiratory centre causing

hypoventilation and increase in PCO2

Self- limiting as increase in PCO2 increases drive on respiratory centre

In chronic state development of reduced sensitivity to PCO2 – more significant compensation BUT

Hypoventilation causing hypoxaemia will provide stimulation of RC and prevent further compensation


non respiratory alkalosisManagementDependent on severity and cause- severe hypovolaemia /hypochloraemia

correct with saline infusion- potassium supplements/removal of diuretics


Summary of non respiratory alkalosis[H+] pH PCO2 [HCO3

-]


Respiratory alkalosisCauses Hypoxia

High altitudeSevere anaemiaPulmonary disease


Respiratory alkalosisCausesIncreased respiratory drive

Stimulants eg salicylatesCerebral – trauma, infection, tumoursHepatic failure


Respiratory alkalosis Causes

Pulmonary disease- Pulmonary oedema- Pulmonary embolism

Mechanical over-ventilation


Respiratory alkalosisCharacterised by reduction in PCO2

Reduces the PCO2/ [HCO3-] ratio

For every KPa decrease in PCO2

decrease in [H+] 5.5 nmol/LSmall decrease in bicarbonate


Respiratory alkalosisCompensation

-reduction in renal hydrogen ion excretion Develops slowly maximal in 36-72 hours


Respiratory alkalosis managementMainly removal of underlying causeIncreasing inspired PCO2 by rebreathing of

expired air for temporary measure - Prolonged – risk of hypoxia


Summary of respiratory alkalosis Acute ChronicpH Slight or low normal[H+] Slight or high normalPCO2

[HCO3-] Slight

Interpretation of resultsReference rangespH 7.35 – 7.46[H+] 35-45 nmol/LpCO2 4.5-6.0 kPa (35-46 mm Hg)pO2 11-15 kPa (85-105 mm Hg) Total Bicarbonate (CO2) 22-30 mmol/L

Figure 27–15 (1 of 2)

Acid-Base BiochemistryRECOMMENDED READINGAnalytical/methods

Tietz Textbook of Clinical Chemistry by Carl A. Burtis (Author), Edward R. Ashwood (Author)

ClinicalClinical Biochemistry by William J. Marshall

and Stephen Bangert

Documents

Acid-Base Biochemistry Prof. Dr. M. Sait KELEŞ Department of Biochemistry