WaterWater, , ElectrolyteElectrolyte andandAcidAcid--BaseBase BalanceBalance

IuriiIurii L. L. KuchynKuchyn, M.D., Ph.D., M.D., Ph.D.

Water Balance

Water is the main component of blood and cells. It fillsmost of the spaces around cells. To functionnormally, the body must keep the amount of waterin balance and relatively constant.

too little water (dehydration) ortoo much water (overhydration)The water in the body contains dissolved minerals

called electrolytes. They include sodium, potassium, calcium etc. The body must also keeplevels of electrolytes in balance and relativelyconstant. The balance of electrolytes is closely tiedto the balance of water in the body: if one changesthe other usually also changes.

Water Balance

Many disorders, especialy those that causefever, vomiting or diarrhea, can result inproblems with water and electrolytebalance.

These disorders may be acute (ex. pneumonia) or chronic (ex. kidney failure). Many drugs, especially diuretics, can alsocause problems.

Water Balance

34 %40 %Intracellular water

< 2 - 3 %< 2 - 3 %Transcellularwater

15 %15 %Extravascular,extracellular

water

5 %5 %Intravascularwater

20%20%Extracellular water:50%60%Total Body Water

FemaleMale

The Composition of FluidCompartments

Exchange between FluidCompartments

• Diffusion is the random movement of moleculesdue to their kinetic energy and is responsible forthe majority of fluid and solute exchangebetween compartments. The rate of diffusion ofa substance across a membrane depends on (1) the permeability of that substance through thatmembrane, (2) the concentration difference forthat substance between the two sides, (3) thepressure difference between either side becausepressure imparts greater kinetic energy, and (4) the electrical potential across the membrane forcharged substances.

Diffusion is the random movement of molecules due to theirkinetic energy and is responsible for the majority of fluidand solute exchange between compartments. The rate ofdiffusion of a substance across a membrane depends on(1) the permeability of that substance through thatmembrane, (2) the concentration difference for thatsubstance between the two sides, (3) the pressuredifference between either side because pressure impartsgreater kinetic energy, and (4) the electrical potentialacross the membrane for charged substances.Diffusion between interstitial fluid and intracellular fluid maytake place by one of several mechanisms: (1) directlythrough the lipid bilayer of the cell membrane, (2) throughprotein channels within the membrane, or (3) by reversiblebinding to a carrier protein that can traverse the membrane(facilitated diffusion). Oxygen, CO2, water, and lipid-solublemolecules penetrate the cell membrane directly. Cationssuch as Na+, K+, and Ca2+ penetrate the membranepoorly because of the cell transmembrane voltage potential(which is positive to the outside) created by the Na+–K+ pump. Therefore, these cations can diffuse only throughspecific protein channels. Passage through these channelsis dependent on membrane voltage and the binding ofligands (such as acetylcholine) to the membrane receptors. Glucose and amino acids diffuse with the help ofmembrane-bound carrier proteins.Fluid exchange between the intracellular and interstitialspaces is governed by the osmotic forces created bydifferences in nondiffusible solute concentrations. Relativechanges in osmolality between the intracellular andinterstitial compartments result in a net water movementfrom the hypoosmolar to the hyperosmolar compartment.

Normal Water Balance• The normal adult daily water intake averages 2500 mL,

which includes approximately 300 mL as a by-product ofthe metabolism of energy substrates. Daily water lossnecessarily averages 2500 mL and can roughly beaccounted for by 1500 mL in urine, 400 mL in respiratorytract evaporation, 400 mL in skin evaporation, 100 mL insweat, and 100 mL in feces. The evaporative losses arevery important in thermoregulation because theynormally account for 20–25% of heat loss.

• Both ICF and ECF osmolalities are closely regulated insuch a way as to maintain a normal water content intissues. Changes in water content and cell volume caninduce serious impairment of function, particularly in thebrain

DehydrationDehydrationDehydration is not having enough water in thebody

The levels of many electrolytes tend to becomeabnormal.

The body tries to keep blood pressure from fallingby moving water from cell and the spaces around thecell into blood vessels.

Then tissues dry out.The kidney try to conserve water by concentratingurine more or by not making any urine.

Conditions that make dehydrationmore likely include the following:

Hot weather, because sweating is increased. Fever, because sweating is increased and breathing

becomes more rapid (causing more water to be lostin the air that is breathed out).

Diarrhea, because water is lost in the stool. Vomiting, because water is lost in vomit. Diabetes that is poorly controlled, because the body

produces more urine. Kidney disorders, because the kidneys are less able

to concentrate urine as needed. Problem with walking, because getting water is

difficult. Dementia, because the sense of thirst is reduced and

the ability to get water when needed is impaired. Use of diuretics, because these drugs increased the

amount of water (and salt) excreted by kidneys.

Isotonic dehydration Na loss in the isoosmolalic solution form. Redundand body fluid loss by digestive track

(vomiting, diarrhea). Hyperexcretion of sodium and water by kidneys. Blood loss. Intestinal obstruction.MainMain symptom:symptom: hypovolemia, circulatory disturbances,

collapse, pressure fall, kidney function disturbances(prerenal failure) with oliguria and uremia.

Laboratory tests:Ht , Hb, Total protein , RBC Na - normal

Hypertonic dehydrationFree water loss; more water is losed than sodium. High loss of body fluids by: kidney (renal diabetes

insipidus, osmotic diuresis in diabetes), lungs(hiperventilation), skin.

Low water intake: young children, senseless people, oldpeople.

Laboratory tests:Ht (0), Hb , Osm , Na , RBC

Hypotonic dehydration

Sodium loss syndromLoss of sodium by kidneys: chronic renal

failure, Addison disease, organic lesion ofcentral nervous system.Laboratory tests:

Ht , Hb , Osm , RBC Na ,

Overhydration

• Overhydration is having too much water inthe body.

• When more fluid is consumed than can be excreted, overhydration occures.

• The blood vessels overfill and fluid movesfrom the blood vessels into the spacesaround cells, causing swelling (edema).

OverhydrationOverhydration has many causes:Heart failure.Kidney disorders.When the body produces too much

antidiuretic hormone (may be coused by pneumonia, stroke and by drugs likecarbamazepine and sertraline).

Too rapid intravenous fluids infusions orblood transfusions.

OverhydrationBlood test may be done to measure levels of

electrolytes or other substances that indicatehow well the kidney are functioning

A chest x-ray can show the back of fluid in thelungs.

Test may be needed to determine whetherheartfailure is present.

For people who are overhydrated, treatmentinvolves helping the body excrete the excesswater. Diuretic are drugs that help thekidneys do just that.

Isotonic overhydrationExcessive retention of body water and Na.Occure in hepatic, cardiac and hunger edema.Exciting cause:Chronic circulatory insufficiency,Cirrhosis,Chronic glomerulitis.Laboratory tests:Na – 0/ Ht , Hb , RBC

Hypertonic overhydration

Sodium excessPrimary hyperaldosteronism (Conn’s syndrom)Excessive supply of hypertonic fluids:Parenteral (ex.patients with renal failure)Oral ( ex. see water drinking by castaways).Laboratory tests:Na Ht , Hb , RBC , Osm

Hypotonic Overhydration„Water poisoning”Excessive administration of hypotonic

fluid (glomerular failure patients).Increase of antidiuretic hormone

secretion (after surgery, encephalitis)I phase – hyperdiuresisII phase – oliguria, anuriaLaboratory test:Na , Ht /0, Hb , RBC

Comparison dehydration withoverhydration

Hypertonic overhydration0 Hypotonic overhydration

0Isotonic overhydration0 Hypertonic dehydration

Hypotonic dehydration0Isotonic dehydration

HbRBCHtNa

OsmolalityOsmolality is a count of the total number of

osmotically active particles in a solution and is equal to the sum of the molalities of all the solutes present in that solution.

The osmolality of plasma is closely regulated by anti-diuretic hormone (ADH).

The osmolality of a solution can be measured using an osmometer. The most commonly used instrument in modern laboratories is a freezing point depression osmometer.

Osmolality Plasma osmolality can also be calculated from the

measured components. While there are many equations, a simple one is as follows:

Serum oSerum osmolalitysmolality [[mOsmmOsm/kg H/kg H22O]O] = 2 x Na + Glucose+ = 2 x Na + Glucose+ urea (all measurements in urea (all measurements in mmolmmol/L)/L) OrOr

Serum Serum osmolalityosmolality [[mOsmmOsm/kg H/kg H22O]O] = 2 x = 2 x Na[mmolNa[mmol/L] + /L] + Glucose[mg/dGlucose[mg/dL]/18+ + L]/18+ + urea[mg/durea[mg/dL]/6L]/6

Reference range (plasma,serum): 275 – 300 mOsm/L Reference range (urine): 50 – 1400 mOsm/L The difference between the measured and calculated

plasma osmolality is known as the the osmolarosmolar gapgap and normally is between 0 and 10 0 and 10 mOsmmOsm/kg./kg.

Clinical Uses Serum Serum osmolalityosmolality is used in two main

circumstances: investigation of hyponatraemia and identification of an osmolar gap.

Urine Urine osmolalityosmolality is an important test:- of renal concentrating ability,- for identifying disorders of the ADH mechanism, - identifying causes of hyper-or hyponatraemia.

Faecal Faecal osmolalityosmolality can be used to assist with diagnosis of the cause of diarrhoea.

Serum Osmolality Plasma osmolality is a useful preliminary investigation for

identifying the cause of hyponatraemia.

If a patient with significant hyponatraemia (serum sodium < 130 mmol/L) has a normal plasma osmolality, the patient may have pseudohyponatraemia due to excess lipids or proteins, or the sample may have been collected from a drip arm containing dextrose.

If the patient has an increased osmolality it is likely the patient has reactive hyponatraemia due to an excess of solute pulling water out of cells. Examples of this include glucose in diabetes mellitus. The finding of a hypo-osmolarhyponatraemia ("true hyponatraemia") then leads to further investigation of the cause.

Serum Osmolar Gap The osmolar gap is determined by subtracting the calculated

osmolality from the measured osmolality. While there are many formulae for the calculated osmolality, the most commonly used is:

Calculated Calculated osmolalityosmolality = 2 x serum sodium + serum = 2 x serum sodium + serum glucose + serum urea (all in glucose + serum urea (all in mmolmmol/L)./L).

The normal osmolarosmolar gapgap is up to 10 10 mmolmmol/L/L((mOsmmOsm/L)/L) and values in excess of this usually indicate the presence of an exogenous agent.

The most common by far is: ethanol, but methanol, ethylene glycol, acetone and isopropyl alcohol can occasionally be present in sufficient quantities to produce an increased osmolar gap.

Na and Water

Major Causes of Hypernatremia.

Treatment of Hypernatremia• The treatment of hypernatremia is aimed at restoring

plasma osmolality to normal as well as correcting theunderlying problem. Water deficits should generally becorrected over 48 h with a hypotonic solution such as 5% dextrose in water. Abnormalities in extracellular volumemust also be corrected. Hypernatremic patients withdecreased total body sodium should be given isotonicfluids to restore plasma volume to normal prior totreatment with a hypotonic solution. Hypernatremicpatients with increased total body sodium should betreated with a loop diuretic along with intravenous 5% dextrose in water. The treatment of diabetes insipidus isdiscussed above.

Example

Classification ofHypoosmolalHyponatremia

Hyponatremia & Low Total BodySodium

• Progressive losses of both sodium and water eventually lead toextracellular volume depletion. As the intravascular volume deficitreaches 5–10%, nonosmotic ADH secretion is activated (seeabove). With further volume depletion, the stimuli for nonosmoticADH release overcome any hyponatremia-induced suppression ofADH. Preservation of circulatory volume takes place at the expenseof plasma osmolality.

• Fluid losses resulting in hyponatremia may be renal or extrarenal inorigin. Renal losses are most commonly related to thiazide diureticsand result in a urinary [Na+] greater than 20 mEq/L. Extrarenallosses are typically gastrointestinal and usually produce a urine[Na+] of less than 10 mEq/L. A major exception to the latter ishyponatremia due to vomiting, which can result in a urinary [Na+] greater than 20 mEq/L. In those instances, bicarbonaturia from theassociated metabolic alkalosis obligates concomitant excretion ofNa+ with HCO3 to maintain electrical neutrality in the urine; urinarychloride concentration, however, is usually less than 10 mEq/L.

Hyponatremia & Increased TotalBody Sodium

• Edematous disorders are characterized by an increasein both total body sodium and TBW. When the increasein water exceeds that in sodium, hyponatremia occurs. Edematous disorders include congestive heart failure, cirrhosis, renal failure, and nephrotic syndrome. Hyponatremia in these settings results from progressiveimpairment of renal free water excretion and generallyparallels underlying disease severity. Pathophysiologicalmechanisms include nonosmotic ADH release anddecreased delivery of fluid to the distal diluting segmentin nephrons. The "effective" circulating blood volume isreduced.

Hyponatremia with Normal TotalBody Sodium

• Hyponatremia in the absence of edema or hypovolemia may beseen with glucocorticoid insufficiency, hypothyroidism, drug therapy(chlorpropamide and cyclophosphamide), and the syndrome ofinappropriate antidiuretic hormone secretion (SIADH). Thehyponatremia associated with adrenal hypofunction may be due tocosecretion of ADH with corticotropin-releasing factor (CRF). Patients with AIDS often exhibit hyponatremia, which may due toadrenal infection by cytomegalovirus or mycobacteria. Diagnosis ofSIADH requires exclusion of other causes of hyponatremia and theabsence of hypovolemia, edema, and adrenal, renal, or thyroiddisease. A variety of malignant tumors, pulmonary diseases, andcentral nervous system disorders are commonly associated withSIADH. In most such instances, plasma ADH concentration is notelevated but is inadequately suppressed relative to the degree ofhypoosmolality in plasma; urine osmolality is usually > 100 mOsm/kg and urine sodium concentration is > 40 mEq/L.

Clinical Manifestations ofHyponatremia

• Symptoms of hyponatremia are primarily neurological and result from anincrease in intracellular water. Their severity is generally related to therapidity with which extracellular hypoosmolality develops. Patients with mildto moderate hyponatremia ([Na+] > 125 mEq/L) are frequentlyasymptomatic. Early symptoms are typically nonspecific and may includeanorexia, nausea, and weakness. Progressive cerebral edema, however, results in lethargy, confusion, seizures, coma, and finally death. Seriousmanifestations of hyponatremia are generally associated with plasmasodium concentrations < 120 mEq/L. Compared with men, premenopausalwomen appear to be at greater risk of neurological impairment and damagefrom hyponatremia.

• Patients with slowly developing or chronic hyponatremia are generally lesssymptomatic. A gradual compensatory loss of intracellular solutes (primarilyNa+, K+, and amino acids) appears to restore cell volume to normal. Neurological symptoms in patients with chronic hyponatremia may berelated more closely to changes in cell membrane potential (due to a lowextracellular [Na+]) than to changes in cell volume.

Treatment of Hyponatremia

Treatment of Hyponatremia• As with hypernatremia, the treatment of hyponatremia is directed at correcting

both the underlying disorder as well as the plasma [Na+]. Isotonic saline isgenerally the treatment of choice for hyponatremic patients withdecreased total body sodium content. Once the extracellular fluid deficit iscorrected, spontaneous water diuresis returns plasma [Na+] to normal.

• Acute symptomatic hyponatremia requires prompt treatment. In suchinstances, correction of plasma [Na+] to > 125 mEq/L is usually sufficient toalleviate symptoms. The amount of NaCl necessary to raise plasma [Na+] to the desired value, the Na+ deficit, can be estimated by the followingformula:

• Very rapid correction of hyponatremia has been associated withdemyelinating lesions in the pons (central pontine myelinolysis), resulting in serious permanent neurological sequelae. The rapidity withwhich hyponatremia is corrected should be tailored to the severity ofsymptoms. The following correction rates have been suggested: for mildsymptoms, 0.5 mEq/L/h or less; for moderate symptoms, 1 mEq/L/h orless; and for severe symptoms, 1.5 mEq/L/h or less.

Example

Disorders of Potassium Balance• Potassium plays a major role in the electrophysiology of cell

membranes as well as carbohydrate and protein synthesis (seebelow). The resting cell membrane potential is normally dependenton the ratio of intracellular to extracellular potassium concentrations. Intracellular potassium concentration is estimated to be 140 mEq/L, whereas extracellular potassium concentration is normally about 4 mEq/L. Although the regulation of intracellular [K+] is poorlyunderstood, extracellular [K+] generally reflects the balancebetween potassium intake and excretion.

• Under some conditions (see below), a redistribution of K+ betweenthe ECF and ICF compartments can result in marked changes inextracellular [K+] without a change in total body potassium content.

Normal Potassium Balance• Dietary potassium intake averages 80 mEq/d in adults (range, 40–

140 mEq/d). About 70 mEq of that amount is normally excreted inurine, whereas the remaining 10 mEq is lost through thegastrointestinal tract.

• Renal excretion of potassium can vary from as little as 5 mEq/L toover 100 mEq/L. Nearly all the potassium filtered in glomeruli isnormally reabsorbed in the proximal tubule and the loop of Henle. The potassium excreted in urine is the result of distal tubularsecretion. Potassium secretion in the distal tubules is coupled toaldosterone-mediated reabsorption of sodium

Regulation of Extracellular Potassium Concentration• Extracellular potassium concentration is closely regulated by cell

membrane Na+–K+ ATPase activity as well as plasma [K+]. Theformer regulates the distribution of potassium between cells andECF, whereas the latter is the major determinant of urinarypotassium excretion.

Hypokalemia• Hypokalemia is defined as plasma [K+] less than

3.5 mEq/L and can occur as a result of (1) anintercompartmental shift of K+, (2) increasedpotassium loss, or (3) an inadequate potassiumintake. Plasma potassium concentration typicallycorrelates poorly with the total potassium deficit. A decrease in plasma [K+] from 4 mEq/L to 3 mEq/L usually represents a 100- to 200-mEq deficit, whereas a plasma [K+] below 3 mEq/L can represent a deficit anywhere between 200 mEq and 400 mEq.

Major Causes ofHypokalemia

Clinical Manifestations ofHypokalemia

• Hypokalemia can produce widespread organ dysfunction. Mostpatients are asymptomatic until plasma [K+] falls below 3 mEq/L. Cardiovascular effects are most prominent and include an abnormalelectrocardiogram (ECG), arrhythmias, decreased cardiaccontractility, and a labile arterial blood pressure due to autonomicdysfunction. Chronic hypokalemia has also been reported to causemyocardial fibrosis. ECG manifestations are primarily due todelayed ventricular repolarization and include T-wave flatteningand inversion, an increasingly prominent U wave, ST-segmentdepression, increased P-wave amplitude, and prolongation ofthe P–R interval. Increased myocardial cell automaticity anddelayed repolarization promote both atrial and ventriculararrhythmias.

Effects of Hypokalemia

• Electrocardiographic effects of hypokalemia. Note progressive flattening ofthe T wave, an increasingly prominent U wave, increased amplitude of the P wave, prolongation of the P–R interval, and ST-segment depression.

Treatment of Hypokalemia• The treatment of hypokalemia depends on the presence and severity of any

associated organ dysfunction. Significant ECG changes such as ST-segmentchanges or arrhythmias mandate continuous ECG monitoring, particularly duringintravenous K+ replacement. Digoxin therapy—as well as the hypokalemiaitself—sensitizes the heart to changes in potassium ion concentration. Musclestrength should also be periodically assessed in patients with weakness.

• Oral replacement with potassium chloride solutions is generally safest (60–80 mEq/d). Replacement of the potassium deficit usually requires several days. Intravenous replacement of potassium chloride should usually be reserved forpatients with or at risk for serious cardiac manifestations or muscle weakness. The goal of intravenous therapy is to remove the patient from immediate dangerand not necessarily to correct the entire potassium deficit. Peripheralintravenous replacement should not exceed 8 mEq/h because of the irritativeeffect of potassium on peripheral veins. Dextrose-containing solutions shouldgenerally be avoided because the resulting hyperglycemia and secondaryinsulin secretion may actually lower plasma [K+] even further. Fasterintravenous replacement (10–20 mEq/h) requires a central venous catheter andclose monitoring of the ECG. Higher replacement rates may be safest through a femoral catheter, because very high localized K+ concentrations may occurwithin the heart with standard central venous catheters. Intravenousreplacement should generally not exceed 240 mEq/d.

• Potassium chloride is the preferred potassium salt when a metabolic alkalosis isalso present because it also corrects the chloride deficit discussed above. Potassium bicarbonate or equivalent (K+ acetate or K+ citrate) is preferable forpatients with metabolic acidosis. Potassium phosphate is a suitable alternativewith concomitant hypophosphatemia (diabetic ketoacidosis).

Hyperkalemia• Hyperkalemia exists when plasma [K+] exceeds 5.5 mEq/L. Hyperkalemia

rarely occurs in normal individuals because of the kidney's tremendouscapacity to excrete potassium. When potassium intake is increased slowly, the kidneys can excrete as much as 500 mEq of K+ per day. Thesympathetic system and insulin secretion also appear to play importantroles in preventing acute increases in plasma [K+] following potassiumloads.

• Hyperkalemia can result from (1) an intercompartmental shift of potassiumions, (2) decreased urinary excretion of potassium, or, rarely, (3) anincreased potassium intake. Measurements of plasma potassiumconcentration can be spuriously elevated if red cells hemolyze in a bloodspecimen (most commonly due to prolonged application of a tourniquetwhile obtaining a venous sample). In vitro release of potassium from whitecells in a blood specimen can also falsely indicate increased levels in themeasured plasma [K+] when the leukocyte count exceeds 70,000 x 109/L. A similar release of potassium from platelets occurs when the platelet countexceeds 1,000,000 x 109/L.

Causes ofHyperkalemia

Clinical Manifestations ofHyperkalemia

• The most important effects of hyperkalemia are on skeletal andcardiac muscle. Skeletal muscle weakness is generally not seenuntil plasma [K+] is greater than 8 mEq/L. The weakness is due tosustained spontaneous depolarization and inactivation of Na+channels of muscle membrane (similar to succinylcholine), eventually resulting in ascending paralysis. Cardiac manifestationsare primarily due to delayed depolarization and consistently presentwhen plasma [K+] is greater than 7 mEq/L. ECG changescharacteristically progress (in order) from symmetrically peaked T waves (often with a shortened QT interval) widening of the QRS complex prolongation of the P–R interval loss of the P wave lossof R-wave amplitude ST-segment depression (occasionallyelevation) an ECG that resembles a sine wave—beforeprogression to ventricular fibrillation and asystole. Contractilityappears to be relatively well preserved. Hypocalcemia, hyponatremia, and acidosis accentuate the cardiac effects ofhyperkalemia.

• Electrocardiographic effects of hyperkalemia. Electrocardiographic changescharacteristically progress from symmetrically peaked T waves, often with a shortened QT interval, to widening of the QRS complex, prolongation of theP–R interval, loss of the P wave, loss of R-wave amplitude, and ST-segment depression (occasionally elevation)—to an ECG that resembles a sine wave—before final progression into ventricular fibrillation or asystole.

Treatment of Hyperkalemia• Because of its lethal potential, hyperkalemia exceeding 6

mEq/L should always be treated. Treatment is directedat reversing cardiac manifestations, and skeletal muscleweakness, and restoring of plasma [K+] to normal. Thenumber of treatment modalities employed (see below) depends on the severity of manifestations as well as thecause of hyperkalemia. Hyperkalemia associated withhypoaldosteronism can be treated with mineralocorticoidreplacement. Drugs contributing to hyperkalemia shouldbe discontinued and sources of increased potassium

Treatment of Hyperkalemia• intake reduced or stopped.• Calcium (5–10 mL of 10% calcium gluconate or 3–5 mL of 10%

calcium chloride) partially antagonizes the cardiac effects ofhyperkalemia and is useful in patients with marked hyperkalemia. Itseffects are rapid but unfortunately short lived. Care must beexercised in patients taking digoxin, as calcium potentiates digoxintoxicity.

• When metabolic acidosis is present, intravenous sodiumbicarbonate (usually 45 mEq) will promote cellular uptake ofpotassium and can decrease plasma [K+] within 15 min. -Agonistspromote cellular uptake of potassium and may be useful in acutehyperkalemia associated with massive transfusions; low doses ofepinephrine (0.5–2 mg/min) often rapidly decrease plasma [K+] andprovide inotropic support in this setting. An intravenous infusion ofglucose and insulin (30–50 g of glucose with 10 units of insulin) isalso effective in promoting cellular uptake of potassium and loweringplasma [K+], but often takes up to 1 h for peak effect.

Treatment of Hyperkalemia

• For patients with some renal function, furosemide is a useful adjunct in increasingurinary excretion of potassium. In the absence ofrenal function, elimination of excess potassiumcan be accomplished only with nonabsorbablecation-exchange resins such as oral or rectalsodium polystyrene sulfonate (Kayexalate). Each gram of resin binds up to 1 mEq of K+ andreleases 1.5 mEq of Na+; the oral dose is 20 g in 100 mL of 20% sorbitol.

Treatment of Hyperkalemia

• Dialysis is indicated in symptomaticpatients with severe or refractoryhyperkalemia. Hemodialysis is faster andmore effective than peritoneal dialysis indecreasing plasma [K+]. Maximalpotassium removal with hemodialysisapproaches 50 mEq/h, compared with 10–15 mEq/h for peritoneal dialysis.

Electrolites - Magnesium• Magnesium (Mg) is the fourth most plentiful cation in the

body. A 70-kg adult has roughly 2000 mEq of Mg. About 50% is sequestered in bone and is not readilyexchangeable with other compartments.

• Normal plasma Mg concentration ranges from 1.4 to 2.1 mEq/L (0.70 to 1.05 mmol/L or 1.71 – 2.44mg/dL).

• The maintenance of plasma Mg concentration is largely a function of dietary intake and extremely effective renal and intestinal conservation.

Magnesium – reference range

0,7 – 1,0 mmol/LAdults

0,6 – 1,2 mmol/LChildren 6 – 14 yr

0,4 – 1,2 mmol/LNewborns

Magnesium - hypomagnesemia• Plasma magnesium concentration below 1.4 mEq/L

(0.70 mmol/L).• The disorders associated with Mg deficiency are

complex and usually accompanied by multiple metabolic and nutritional disturbances.

• Clinically significant Mg deficiency most commonly is associated with (1) malabsorption syndromes;(2) protein-calorie malnutrition (eg, kwashiorkor);(3) parathyroid disease, in which hypomagnesemiaoccurs after removal of a parathyroid tumor;(4) chronic alcoholism, renal excretion;(5) chronic diarrhea.

Magnesium - hypermagnesemia

• Plasma magnesium concentration above 2.1 mEq/L (1.05 mmol/L).

• renal failure after ingestion of Mg-containing drugs such as antacids or purgatives

• Deep tendon reflexes disappear as the plasma Mg level approaches 10 mEq/L (5.0 mmol/L); hypotension, respiratory depression, and narcosis develop with increasing hypermagnesemia. Cardiac arrest may occur when blood Mg levels exceed 12 to 15 mEq/L (6.0 to 7.5 mmol/L).

ChlorideChloride is an element in the blood and body fluid that functions like sodium to maintain fluid balance.Disporportionate loss of chloride can lead to the body's environment becoming more acidic.

Chloride - The major anion found in the fluid outside of cells and in blood. An anion is the negatively charged part of certain substances such as table salt (sodium chloride or NaCl) when dissolved in liquid.

The balance of chloride ion (Cl-) is closely regulated by the body. Significant increases or decreases in chloride can have deleterious or even fatal consequences.

Chloride – reference range

Increased chloride (hyperchloremia): Elevations in chloride may be seen in diarrhea, certain kidney diseases, and sometimes in overactivity of the parathyroid glands.

Decreased chloride (hypochloremia): Chloride is normally lost in the urine, sweat, and stomach secretions. Excessive loss can occur from heavy sweating, vomiting, and adrenal gland and kidney disease.

95 – 112 mmol/LChildren96 – 110 mmol/LAdults

Bicarbonate Bicarbonate is a measure of the alkalinity of the

body's tissues. The bicarbonate content of the body is delicately balanced by the kidneys and also by the lungs. It's job is to prevent the body's environment from getting too acidic or too basic (that is, having too little acid or too much alkali).

Bicarbonate: Bicarbonate levels are used to monitor the acidity of the blood and body fluids. The acidity is affected by foods or medications that we ingest and the function of the kidneys and lungs. The chemical notation for bicarbonate on most lab reports is HCO3-

or represented as the concentration of carbon dioxide (CO2).

The normal serum range for bicarbonate is 22-30 mmol/L.

Suggested Reading• Boldt J: Volume replacement in the surgical patient—

does the type of fluid make a difference? Br J Anaesth2000;84:783. [PMID: 10895757]

Paradis OC: Fluids and Electrolytes, 2nd ed. Lippincott, 1999.

Vincent JL: Strategies in body fluid replacement. MinervaAnesthesiol 2000;66:278. [PMID: 10965702]

Palmer BF: Managing hyperkalemia caused by inhibitorsof the rennin-angiotensin-aldosterone system. N Engl J Med 2004;351:585. ACE inhibitors are being used withincreased frequency. This article is a good review of oneof their side effects.

Acid–Base Chemistry

• The concentration of water is omitted from the denominator of thisexpression because it does not vary appreciably and is alreadyincluded in the constant. Therefore, given [H+] or [OH–], theconcentration of the other ion can be readily calculated.

• Example: If [H+] = 10–8 nEq/L, then [OH–] = 10–14 ÷ 10–8 = 10–6 nEq/L.

• Arterial [H+] is normally 40 nEq/L, or 40 x 10–9 mol/L. Hydrogen ionconcentration is more commonly expressed as pH, because dealingwith numbers of this order of magnitude is awkward. The pH of a solution is defined as the negative logarithm (base 10) of [H+].

• Normal arterial pH is therefore –log (40 x 10–9) = 7.40. Hydrogen ion concentrations between 16 and 160 nEq/L (pH6.8–7.8) are compatible with life.

The relationship between pH and [H+]. Note that between a pH of 7.10 and 7.50, the relationship between pH and [H+] is nearly linear.

Acids & Bases• An acid is usually defined as a chemical species that can act as a

proton (H+) donor, whereas a base is a species that can act as a proton acceptor (Brönsted–Lowry definitions). In physiologicalsolutions, it is probably better to use Arrhenius' definitions: An acidis a compound that contains hydrogen and reacts with water toform hydrogen ions. A base is a compound that produceshydroxide ions in water. Using these definitions, the SID becomesimportant as other ions in solutions (cations and anions) will affectthe dissociation constant for water and, therefore, the hydrogenion concentration. A strong acid is a substance that readily andalmost irreversibly gives up an H+ and increases [H+], whereas a strong base avidly binds H+ and decreases [H+]. In contrast, weakacids reversibly donate H+, whereas weak bases reversibly bindH+; both tend to have less of an effect on [H+]. Biologicalcompounds are either weak acids or weak bases.

• From this equation, it is apparent that the pH of this solution is related to theratio of the dissociated anion to the undissociated acid.

• The problem with this approach is that it is phenomenological—measure thepH and bicarbonate, and then other variables can be manipulatedmathematically. This approach works well with pure water—theconcentration of [H+] must equal [OH–]. But physiological solutions, although aqueous, are far more complex. Even in such a complex solution, the [H+] can be predicted using three variables: the SID, the PCO2, and thetotal weak acid concentration [ATOT].

• The SID is the sum of all the strong, completely or almost completelydissociated, cations (Na+, K+, Ca2+, Mg2+) minus the strong anions (Cl–, lactate–, etc). Although we can calculate an SID, because the laws ofelectroneutrality must be observed, if there is an SID, other unmeasuredions must be present. PCO2 is an independent variable assumingventilation is ongoing. The conjugate base of HA is A– and is composedmostly of phosphates and proteins that do not change independent of theother two variables. A– plus AH is an independent variable because itsvalue is not determined by any other variable. Note that [H+] is not a strongion (water does not completely dissociate), but it can, does, and mustchange in response to any change in SID, PCO2, or ATOT to comply withthe laws of electroneutrality and conservation of mass. Strong ions cannotbe made to achieve electroneutrality but hydrogen ions, H+, are created orconsumed based on changes in the dissociation of water.

Conjugate Pairs & Buffers• As discussed above, when the weak acid HA is in solution, HA can

act as an acid by donating an H+ and A– can act as a base bytaking up H+. A– is therefore often referred to as the conjugate baseof HA. A similar concept can be applied for weak bases. Considerthe weak base B, where

A buffer is a solution that contains a weak acid and its conjugate base ora weak base and its conjugate acid (conjugate pairs). Buffers minimizeany change in [H+] by readily accepting or giving up hydrogen ions. It isreadily apparent that buffers are most efficient in minimizing changes inthe [H+] of a solution (ie, [A–] = [HA]) when pH = pK. Moreover, theconjugate pair must be present in significant quantities in solution to actas an effective buffer.

Clinical Disorders• A clear understanding of acid–base disorders and

compensatory physiological responses requires preciseterminology. The suffix "-osis" is used here to denote anypathological process that alters arterial pH. Thus, anydisorder that tends to lower pH is an acidosis, whereasone tending to increase pH is termed an alkalosis. If thedisorder primarily affects [HCO3–], it is termedmetabolic. If the disorder primarily affects PaCO2, it istermed respiratory. Secondary compensatory responses(see below) should be referred to as just that and not asan "-osis." One might therefore refer to a metabolicacidosis with respiratory compensation.

• When only one pathological process occurs by itself, the acid–basedisorder is considered to be simple. The presence of two or moreprimary processes indicates a mixed acid–base disorder.

• The suffix "-emia" is used to denote the net effect of all primaryprocesses and compensatory physiological responses (see below) on arterial blood pH. Because arterial blood pH is normally 7.35–7.45 in adults, the term "acidemia" signifies a pH < 7.35 whilealkalemia signifies a pH > 7.45.

Compensatory Mechanisms

• Physiological responses to changes in[H+] are characterized by three phases: (1) immediate chemical buffering, (2) respiratory compensation (wheneverpossible), and (3) a slower but moreeffective renal compensatory responsethat may nearly normalize arterial pH evenif the pathological process is still present.

Body Buffers• Physiologically important buffers in humans include bicarbonate

(H2CO3/HCO3–), hemoglobin (HbH/Hb–), other intracellular proteins(PrH/Pr–), phosphates (H2PO4–/HPO42–), and ammonia (NH3/NH4+). Theeffectiveness of these buffers in the various fluid compartments is related totheir concentration. Bicarbonate is the most important buffer in theextracellular fluid compartment. Hemoglobin, though restricted inside redblood cells, also functions as an important buffer in blood. Other proteinsprobably play a major role in buffering the intracellular fluid compartment. Phosphate and ammonium ions are important urinary buffers.

• Buffering of the extracellular compartment can also be accomplished by theexchange of extracellular H+ for Na+ and Ca2+ ions from bone and by theexchange of extracellular H+ for intracellular K+. Acid loads can alsodemineralize bone and release alkaline compounds (CaCO3 and CaHPO4). Alkaline loads (NaHCO3) increase the deposition of carbonate in bone.

• Buffering by plasma bicarbonate is almost immediate whereas that due tointerstitial bicarbonate requires 15–20 min. In contrast, buffering byintracellular proteins and bone is slower (2–4 h). Up to 50–60% of acidloads may ultimately be buffered by bone and intracellular buffers

The Bicarbonate Buffer

• Note that its pK' is not close to the normal arterial pH of7.40, which means that bicarbonate would not beexpected to be an efficient extracellular buffer (seeabove). The bicarbonate system is, however, importantfor two reasons: (1) bicarbonate (HCO3–) is present inrelatively high concentrations in extracellular fluid, and(2) more importantly—PaCO2 and plasma [HCO3–] areclosely regulated by the lungs and the kidneys, respectively. The ability of these two organs to alter the[HCO3–]/PaCO2 ratio allows them to exert importantinfluences on arterial pH.

• This equation is very useful clinicallybecause pH can be readily convertedto [H+]. Note that below 7.40, [H+] increases 1.25 nEq/L for each 0.01 decrease in pH; above 7.40, [H+] decreases 0.8 nEq/L for each 0.01 increase in pH.

Hemoglobin as a Buffer• Hemoglobin is rich in histidine, which is an effective buffer

from pH 5.7 to 7.7 (pKa 6.8). Hemoglobin is the mostimportant noncarbonic buffer in extracellular fluid. Simplistically, hemoglobin may be thought of as existing inred blood cells in equilibrium as a weak acid (HHb) and a potassium salt (KHb). In contrast to the bicarbonate buffer, hemoglobin is capable of buffering both carbonic (CO2) andnoncarbonic (nonvolatile) acids:

Pulmonary Compensation• Changes in alveolar ventilation responsible for

pulmonary compensation of PaCO2 are mediated bychemoreceptors within the brain stem. These receptorsrespond to changes in cerebrospinal spinal fluid pH. Minute ventilation increases 1–4 L/min for every 1 mmHg increase in PaCO2. In fact, the lungs are responsiblefor eliminating the approximately 15 mEq of carbondioxide produced every day as a byproduct ofcarbohydrate and fat metabolism. Pulmonarycompensatory responses are also important in defendingagainst marked changes in pH during metabolicdisturbances.

Pulmonary Compensation during Metabolic Acidosis• Decreases in arterial blood pH stimulate medullary respiratory centers. The

resulting increase in alveolar ventilation lowers PaCO2 and tends to restorearterial pH toward normal. The pulmonary response to lower PaCO2 occursrapidly but may not reach a predictably steady state until 12–24 h; pH isnever completely restored to normal. PaCO2 normally decreases 1–1.5 mmHg below 40 mm Hg for every 1 mEq/L decrease in plasma [HCO3–].

Pulmonary Compensation during Metabolic Alkalosis• Increases in arterial blood pH depress respiratory centers. The resulting

alveolar hypoventilation tends to elevate PaCO2 and restore arterial pHtoward normal. The pulmonary response to metabolic alkalosis is generallyless predictable than the response to metabolic acidosis. Hypoxemia, as a result of progressive hypoventilation, eventually activates oxygen-sensitivechemoreceptors; the latter stimulates ventilation and limits thecompensatory pulmonary response. Consequently, PaCO2 usually does notrise above 55 mm Hg in response to metabolic alkalosis. As a general rule, PaCO2 can be expected to increase 0.25–1 mm Hg for each 1 mEq/L increase in [HCO3–].

Renal Compensation• The ability of the kidneys to control the amount of

HCO3– reabsorbed from filtered tubular fluid, form newHCO3–, and eliminate H+ in the form of titratable acidsand ammonium ions allows them to exert a majorinfluence on pH during both metabolic and respiratoryacid–base disturbances. In fact, the kidneys areresponsible for eliminating the approximately 1 mEq/kgper day of sulfuric acid, phosphoric acid, andincompletely oxidized organic acids that are normallyproduced by the metabolism of dietary and endogenousproteins, nucleoproteins, and organic phosphates (fromphosphoproteins and phospholipids). Metabolism ofnucleoproteins also produces uric acid. Incompletecombustion of fatty acids and glucose produces ketoacids and lactic acid. Endogenous alkali are producedduring the metabolism of some anionic amino acids(glutamate and aspartate) and other organic compounds(citrate, acetate, and lactate), but the quantity isinsufficient to offset the endogenous acid production.

Renal Compensation duringAcidosis

• The renal response to acidemia is 3-fold: (1) increased reabsorption of the filteredHCO3–, (2) increased excretion oftitratable acids, and (3) increasedproduction of ammonia.

• Although these mechanisms are probablyactivated immediately, their effects aregenerally not appreciable for 12–24 h andmay not be maximal for up to 5 days.

Increased Reabsorption of HCO3–• Bicarbonate reabsorption is shown in Figure 30–3. CO2 within

renal tubular cells combines with water in the presence ofcarbonic anhydrase. The carbonic acid (H2CO3) formed rapidlydissociates into H+ and HCO3–. Bicarbonate ion then entersthe bloodstream while the H+ is secreted into the renal tubule, where it reacts with filtered HCO3– to form H2CO3. Carbonicanhydrase associated with the luminal brush border catalyzesthe dissociation of H2CO3 into CO2 and H2O. The CO2 thusformed can diffuse back into the renal tubular cell to replace theCO2 originally consumed. The proximal tubules normallyreabsorb 80–90% of the filtered bicarbonate load along withsodium, whereas the distal tubules are responsible for theremaining 10–20%. Unlike the proximal H+ pump, the H+ pumpin the distal tubule is not necessarily linked to sodiumreabsorption, and is capable of generating steep H+ gradientsbetween tubular fluid and tubular cells. Urinary pH candecrease to as low as 4.4 (compared with a pH of 7.40 inplasma).

• Reclamation offiltered HCO3– bythe proximal renaltubules

Increased Excretion of TitratableAcids

• After all the HCO3– in tubular fluid isreclaimed, the H+ secreted into thetubular lumen can combine withHPO42– to form H2PO4– ; the latter isnot readily reabsorbed because of itscharge and is eliminated in urine. Thenet result is that H+ is excreted fromthe body as H2PO4–, and the HCO3–that is generated in the process canenter the bloodstream. With a pK of6.8, the H2PO4–/HPO42– pair isnormally an ideal urinary buffer. Whenurinary pH approaches 4.4, however, all the phosphate reaching the distaltubule is in the H2PO4– form; HPO42–ions are no longer available foreliminating H+.

Increased Formation of Ammonia• After complete reabsorption of

HCO3– and consumption of thephosphate buffer, theNH3/NH4+ pair becomes themost important urinary buffer(Figure 30–5). Deamination ofglutamine within themitochondria of proximaltubular cells is the principalsource of NH3 production in thekidneys. Acidemia markedlyincreases renal NH3 production. The ammoniaformed is then able to passivelycross the cell's luminalmembrane, enter the tubularfluid, and react with H+ to formNH4+. Unlike NH3, NH4+ doesnot readily penetrate theluminal membrane and istherefore trapped within thetubules. Thus, excretion ofNH4+ in urine effectivelyeliminates H+.

Renal Compensation duringAlkalosis

• The tremendous amount of HCO3– normally filtered andsubsequently reabsorbed allows the kidneys to rapidly excrete largeamounts of bicarbonate if necessary. As a result, the kidneys arehighly effective in protecting against metabolic alkalosis, whichtherefore generally occurs only in association with concomitantsodium deficiency or mineralocorticoid excess. Sodium depletiondecreases extracellular fluid volume and enhances Na+reabsorption in the proximal tubule. To maintain neutrality, the Na+ion is brought across with a Cl– ion. As Cl– ions decrease in number(< 10 mEq/L of urine), HCO3– must be reabsorbed. In addition, increased H+ secretion in exchange for augmented Na+reabsorption favors continued HCO3– formation with metabolicalkalosis. Similarly, increased mineralocorticoid activity augmentsaldosterone-mediated Na+ reabsorption in exchange for H+ secretion in the distal tubules. The resulting increase in HCO3–formation can initiate or propagate metabolic alkalosis. Metabolicalkalosis is commonly associated with increased mineralocorticoidactivity even in the absence of sodium and chloride depletion.

Base Excess• Base excess is the amount of acid or base that must be

added for blood pH to return to 7.40 and PaCO2 toreturn to 40 mm Hg at full O2 saturation and 37°C. Moreover, it adjusts for noncarbonic buffering in theblood. Simplistically, base excess represents themetabolic component of an acid–base disturbance. A positive value indicates metabolic alkalosis, whereas a negative value reveals metabolic acidosis. Base excessis usually derived graphically or electronically from a nomogram originally developed by Siggaard-Andersenand requires measurement of hemoglobin concentration

Diagnosis of simple acid–basedisorders.

AcidosisPhysiological Effects of Acidemia

• [H+] is strictly regulated in the nanomole/liter (36–43 nmol/L) range as H+ ions have high charge densities and "large" electric fields that can affect thestrength of hydrogen bonds that are present on most physiologicalbiochemicals. Biochemical reactions are very sensitive to changes in [H+]. The overall effects of acidemia seen in patients represent the balancebetween its direct effects and sympathoadrenal activation. With worseningacidosis (pH < 7.20), direct depressant effects predominate. Directmyocardial and smooth muscle depression reduces cardiac contractility andperipheral vascular resistance, resulting in progressive hypotension. Severeacidosis can lead to tissue hypoxia despite a rightward shift in hemoglobinaffinity for oxygen. Both cardiac and vascular smooth muscle become lessresponsive to endogenous and exogenous catecholamines, and thethreshold for ventricular fibrillation is decreased. Progressive hyperkalemiaas a result of the movement of K+ out of cells in exchange for extracellularH+ is also potentially lethal. Plasma [K+] increases approximately 0.6 mEq/L for each 0.10 decrease in pH.

• Central nervous system depression is more prominent with respiratoryacidosis than with metabolic acidosis. This effect, often termed CO2 narcosis, may be the result of intracranial hypertension secondary toincreased cerebral blood flow and of severe intracellular acidosis. UnlikeCO2, H+ ions do not readily penetrate the blood–brain barrier.

Respiratory Acidosis

RespiratoryAcidosis

• Carbon dioxide production is a byproduct of fat and carbohydratemetabolism. Muscle activity, bodytemperature, and thyroidhormone activity can all havemajor influences on CO2 production. Because CO2 production does not appreciablyvary under most circumstances, respiratory acidosis is usually theresult of alveolar hypoventilation. In patients with a limited capacityto increase alveolar ventilation, however, increased CO2 production can precipitaterespiratory acidosis.

Acute Respiratory Acidosis• The compensatory response to acute (6–12 h)

elevations in PaCO2 is limited. Buffering is primarilyprovided by hemoglobin and the exchange ofextracellular H+ for Na+ and K+ from bone and theintracellular fluid compartment (see above). The renalresponse to retain more bicarbonate is very limitedacutely. As a result, plasma [HCO3–] increases onlyabout 1 mEq/L for each 10 mm Hg increase in PaCO2 above 40 mm Hg.

Chronic Respiratory Acidosis• "Full" renal compensation characterizes chronic

respiratory acidosis. Renal compensation is appreciableonly after 12–24 h and may not peak until 3–5 days. During that time, the sustained increase in PaCO2 hasbeen present long enough to permit maximal renalcompensation. During chronic respiratory acidosis, plasma [HCO3–] increases approximately 4 mEq/L foreach 10 mm Hg increase in PaCO2 above 40 mm Hg.

Treatment of Respiratory Acidosis• Respiratory acidosis is treated by reversing the imbalance between CO2 production and alveolar

ventilation. In most instances, this is accomplished by increasing alveolar ventilation. Measures aimed atreducing CO2 production are useful only in specific instances (eg, dantrolene for malignant hyperthermia, muscle paralysis for tetanus, antithyroid medication for thyroid storm, and reduced caloric intake). Temporizing measures aimed at improving alveolar ventilation include bronchodilation, reversal ofnarcosis, administration of a respiratory stimulant (doxapram), or improving lung compliance (diuresis). Moderate to severe acidosis (pH < 7.20), CO2 narcosis, and impending respiratory muscle fatigue areindications for mechanical ventilation. An increased inspired oxygen concentration is also usuallynecessary, as coexistent hypoxemia is common. Intravenous NaHCO3 is rarely necessary unless pH is < 7.10 and HCO3– is < 15 mEq/L. Sodium bicarbonate therapy will transiently increase PaCO2:

• Buffers that do not produce CO2, such as carbicarb or tromethamine (THAM), have been proposed asalternatives but are not of proven benefit (below). Carbicarb is a mixture of 0.3 M sodium bicarbonate and 0.3 M sodium carbonate; buffering by this mixture mainly produces sodium bicarbonate instead of CO2. Tromethamine has the added advantage of lacking sodium and may be a more effective intracellular buffer.

• Patients with a baseline chronic respiratory acidosis require special consideration. When such patientsdevelop acute ventilatory failure, the goal of therapy should be to return PaCO2 to the patient's "normal" baseline. Normalizing the patient's PaCO2 to 40 mm Hg will result in respiratory alkalosis (see below). Oxygen therapy must also be carefully controlled, because the respiratory drive in these patients may bedependent on hypoxemia, not PaCO2; or may increase physiological dead space. "Normalization" of PaCO2 or relative hyperoxia can precipitate severe hypoventilation.

Metabolic Acidosis

• Metabolic acidosis is defined as a primary decrease in [HCO3–]. Pathological processes can initiatemetabolic acidosis by one of threemechanisms: (1) consumption of HCO3–by a strong nonvolatile acid, (2) renal orgastrointestinal wasting of bicarbonate, or (3) rapid dilution of the extracellularfluid compartment with a bicarbonate-free fluid.

• A fall in plasma [HCO3–] without a proportionate reduction in PaCO2 decreases arterial pH. The pulmonarycompensatory response in a simplemetabolic acidosis (see above) characteristically does not reducePaCO2 to a level that completelynormalizes pH but can produce markedhyperventilation (Kussmaul's respiration).

• Note that differential diagnosis ofmetabolic acidosis may be facilitated bycalculation of the anion gap.

The Anion Gap• The anion gap in plasma is most commonly

defined as the difference between the majormeasured cations and the major measuredanions:

• "Unmeasured cations" include K+, Ca2+, and Mg2+, whereas"unmeasured anions" include all organic anions (including plasmaproteins), phosphates, and sulfates. Plasma albumin normallyaccounts for the largest fraction of the anion gap (about 11 mEq/L). The anion gap decreases by 2.5 mEq/L for every 1 g/dL reduction inplasma albumin concentration. Any process that increases"unmeasured anions" or decreases "unmeasured cations" willincrease the anion gap. Conversely, any process that decreases"unmeasured anions" or increases "unmeasured cations" willdecrease the anion gap.

• Mild elevations of plasma anion gap up to 20 mEq/L may not behelpful diagnostically during acidosis, but values > 30 mEq/L usuallyindicate the presence of a high anion gap acidosis (below). Metabolic alkalosis can also produce a high anion gap because ofextracellular volume depletion, an increased charge on albumin, anda compensatory increase in lactate production. A low plasma aniongap may be encountered with hypoalbuminemia, bromide or lithiumintoxication, and multiple myeloma.

High Anion Gap Metabolic Acidosis

• Metabolic acidosis with a high anion gap is characterizedby an increase in relatively strong nonvolatile acids. These acids dissociate into H+ and their respectiveanions; the H+ consumes HCO3– to produce CO2, whereas their anions (conjugate bases) accumulate andtake the place of HCO3– in extracellular fluid (hence theanion gap increases). Nonvolatile acids can beendogenously produced or ingested.

• Failure to Excrete Endogenous Nonvolatile Acids• Endogenously produced organic acids are normally

eliminated by the kidneys in urine (above). Glomerularfiltration rates below 20 mL/min (renal failure) typicallyresult in progressive metabolic acidosis from theaccumulation of these acids.

Increased Endogenous NonvolatileAcid Production

• Severe tissue hypoxia following hypoxemia, hypoperfusion (ischemia), or inability toutilize oxygen (cyanide poisoning) can result in lactic acidosis. Lactic acid is the endproduct of the anaerobic metabolism of glucose (glycolysis) and can rapidlyaccumulate under these conditions. Decreased utilization of lactate by the liver, andto a lesser extent by the kidneys, is less commonly responsible for lactic acidosis; causes include hypoperfusion, alcoholism, and liver disease. Lactate levels can bereadily measured and are normally 0.3–1.3 mEq/L. Acidosis resulting from D-lacticacid, which is not recognized by -lactate dehydrogenase (and not measured byroutine assays), may be encountered in patients with short bowel syndromes; D-lacticacid is formed by colonic bacteria from dietary glucose and starch and is absorbedsystemically.

• An absolute or relative lack of insulin can result in hyperglycemia and progressiveketoacidosis from accumulation of -hydroxybutyric and acetoacetic acids. Ketoacidosis may also be seen following starvation and alcoholic binges. Thepathophysiology of the acidosis often associated with severe alcoholic intoxicationand nonketotic hyperosmolar coma is complex and may represent a build-up of lactic, keto, or other unknown acids.

• Some inborn errors of metabolism, such as maple syrup urine disease, methylmalonic aciduria, propionic acidemia, and isovaleric acidemia, produce a highanion gap metabolic acidosis as a result of accumulation of abnormal amino acids.

Ingestion of Exogenous NonvolatileAcids

• Ingestion of large amounts of salicylates frequently results inmetabolic acidosis. Salicylic acid as well as other acid intermediatesrapidly accumulate and produce a high anion gap acidosis. Becausesalicylates also produce direct respiratory stimulation, most adultsdevelop mixed metabolic acidosis with superimposed respiratoryalkalosis. Ingestion of methanol (methyl alcohol) frequently producesacidosis and visual disturbances (retinitis). Symptoms are typicallydelayed until the slow oxidation of methanol by alcoholdehydrogenase produces formic acid, which is highly toxic to theretina. The high anion gap represents the accumulation of manyorganic acids, including acetic acid. The toxicity of ethylene glycol isalso the result of the action of alcohol dehydrogenase to produceglycolic acid. Glycolic acid, the principal cause of the acidosis, isfurther metabolized to form oxalic acid, which can be deposited inthe renal tubules and result in renal failure.

Normal Anion Gap MetabolicAcidosis

• Metabolic acidosis associated with a normal anion gap is typicallycharacterized by hyperchloremia. Plasma [Cl–] increases to take theplace of the HCO3– ions that are lost. Hyperchloremic metabolicacidosis most commonly results from abnormal gastrointestinal orrenal losses of HCO3–.

• Calculation of the anion gap in urine can be helpful in diagnosing a normal anion gap acidosis.

•• The urine anion gap is normally positive or close to zero. The

principal unmeasured urinary cation is normally NH4+, which shouldincrease (along with Cl–) during a metabolic acidosis; the latterresults in a negative urinary anion gap. Impairment of H+ or NH4+ secretion, as occurs in renal failure or renal tubular acidosis (below), results in a positive urine anion gap in spite of systemic acidosis.

Increased Gastrointestinal Loss ofHCO3–

• Diarrhea is the most common cause of hyperchloremicmetabolic acidosis. Diarrheal fluid contains 20–50 mEq/L of HCO3–. Small bowel, biliary, and pancreatic fluids areall rich in HCO3–. Loss of large volumes of these fluidscan lead to hyperchloremic metabolic acidosis. Patientswith ureterosigmoidostomies and those with ileal loopsthat are too long or that become partially obstructedfrequently develop hyperchloremic metabolic acidosis. The ingestion of chloride-containing anion-exchangeresins (cholestyramine) or large amounts of calcium ormagnesium chloride can result in increased absorptionof chloride and loss of bicarbonate ions. Thesenonabsorbable resins bind bicarbonate ions, whereascalcium and magnesium combine with bicarbonate toform insoluble salts within the intestines.

Increased Renal Loss of HCO3–• Renal wasting of HCO3– can occur as a result of failure to reabsorb filtered HCO3–

or to secrete adequate amounts of H+ in the form of titratable acid or ammonium ion. These defects are encountered in patients taking carbonic anhydrase inhibitors suchas acetazolamide and in those with renal tubular acidosis.

• Renal tubular acidosis comprises a group of nonazotemic defects of H+ secretionby the renal tubules, resulting in a urinary pH that is too high for the systemicacidemia. These defects may be a result of a primary renal defect or may besecondary to a systemic disorder. The site of the H+-secreting defect may be in thedistal (type 1) or proximal (type 2) renal tubule. Hyporeninemic hypoaldosteronism iscommonly referred to as type 4 renal tubular acidosis. With distal renal tubularacidosis, the defect occurs at a site after most of the filtered HCO3– has beenreclaimed. As a result, there is a failure to acidify the urine, so that net acid excretionis less than daily net acid production. This disorder is frequently associated withhypokalemia, demineralization of bone, nephrolithiasis, and nephrocalcinosis. Alkali(NaHCO3) therapy (1–3 mEq/kg/d) is usually sufficient to reverse those side effects. With the less common proximal renal tubular acidosis, defective H+ secretion in theproximal tubule results in massive wasting of HCO3–. Concomitant defects in tubularreabsorption of other substances such as glucose, amino acids, or phosphates arecommon. The hyperchloremic acidosis results in volume depletion and hypokalemia. Treatment involves giving alkali (as much as 10–25 mEq/kg per day) and potassiumsupplements.

Other Causes of HyperchloremicAcidosis

• A dilutional hyperchloremic acidosis can occur whenextracellular volume is rapidly expanded with a bicarbonate-free fluid such as normal saline. The plasmaHCO3– decreases in proportion to the amount of fluidinfused as extracellular HCO3– is diluted. Amino acidinfusions (parenteral hyperalimentation) contain organiccations in excess of organic anions and can producehyperchloremic metabolic acidosis because chloride iscommonly used as the anion for the cationic aminoacids. Lastly, the administration of excessive quantitiesof chloride-containing acids such as ammonium chlorideor arginine hydrochloride (usually given to treat a metabolic alkalosis) can cause hyperchloremic metabolicacidosis.

Treatment of Metabolic Acidosis• Several general measures can be undertaken to control the severity

of acidemia until the underlying processes are corrected. Anyrespiratory component of the acidemia should be corrected. Respiration should be controlled if necessary; a PaCO2 in the low30s may be desirable to partially return pH to normal. If arterialblood pH remains below 7.20, alkali therapy, usually in the form ofNaHCO3 (usually a 7.5% solution), may be necessary. PaCO2 maytransiently rise as HCO3– is consumed by acids (emphasizing theneed to control ventilation in severe acidemia). The amount ofNaHCO3 given is decided empirically as a fixed dose (1 mEq/kg) oris derived from the base excess and the calculated bicarbonatespace (see below). In either case, serial blood gas measurementsare mandatory to avoid complications (eg, overshoot alkalosis andsodium overload) and to guide further therapy. Raising arterial pH to> 7.25 is usually sufficient to overcome the adverse physiologicaleffects of the acidemia. Profound or refractory acidemia may requireacute hemodialysis with a bicarbonate dialysate.

Treatment of Metabolic Acidosis• The routine use of large amounts of NaHCO3 in treating cardiac arrest and

low flow states is no longer recommended. Paradoxical intracellular acidosismay occur, particularly when CO2 elimination is impaired, because the CO2 formed readily enters cells but bicarbonate ion does not. Alternate buffersthat do not produce CO2 may be theoretically preferable, but are unprovenclinically.

• Specific therapy for diabetic ketoacidosis includes replacement of theexisting fluid deficit (as a result of a hyperglycemic osmotic diuresis) first aswell as insulin, potassium, phosphate, and magnesium. The treatment oflactic acidosis should be directed first at restoring adequate oxygenationand tissue perfusion. Alkalinization of the urine with NaHCO3 to a pHgreater than 7.0 increases elimination of salicylate following salicylatepoisoning. Ethanol infusions (an intravenous loading dose of 8–10 mL/kg ofa 10% ethanol in D5 solution over 30 min with the concomitantadministration of a continuous infusion at 0.15 mL/kg/h to achieve a bloodethanol level of 100–130 mg/dL) are indicated following methanol orethylene glycol intoxication. Ethanol competes for alcohol dehydrogenaseand slows down the formation of formic acid from methanol and glycolic andoxalic acids from ethylene glycol, respectively.

Bicarbonate Space• The bicarbonate space is defined as the volume to which HCO3–

will distribute when it is given intravenously. Although thistheoretically should equal the extracellular fluid space(approximately 25% of body weight), in reality it ranges anywherebetween 25% and 60% of body weight depending on the severityand duration of the acidosis. This variation is at least partly relatedto the amount of intracellular and bone buffering that has takenplace.

AlkalosisPhysiological Effects of Alkalosis

• Alkalosis increases the affinity of hemoglobin for oxygen and shiftsthe oxygen dissociation curve to the left, making it more difficult forhemoglobin to give up oxygen to tissues. Movement of H+ out ofcells in exchange for the movement of extracellular K+ into cells canproduce hypokalemia. Alkalosis increases the number of anionicbinding sites for Ca2+ on plasma proteins and can thereforedecrease ionized plasma [Ca2+], leading to circulatory depressionand neuromuscular irritability. Respiratory alkalosis reduces cerebralblood flow, increases systemic vascular resistance, and mayprecipitate coronary vasospasm. In the lungs, respiratory alkalosisincreases bronchial smooth muscle tone (bronchoconstriction) butdecreases pulmonary vascular resistance.

RespiratoryAlkalosis

• Respiratory alkalosis is defined asa primary decrease in PaCO2. The mechanism is usually aninappropriate increase in alveolarventilation relative to CO2 production. Plasma [HCO3–] usually decreases 2 mEq/L foreach 10 mm Hg acute decrease inPaCO2 below 40 mm Hg. Thedistinction between acute andchronic respiratory alkalosis isnot always made, because thecompensatory response tochronic respiratory alkalosis isquite variable: plasma [HCO3–] decreases 2–5 mEq/L for each 10 mm Hg decrease in PaCO2 below40 mm Hg.

Treatment of Respiratory Alkalosis

• Correction of the underlying process is theonly treatment for respiratory alkalosis. Forsevere alkalemia (arterial pH > 7.60), intravenous hydrochloric acid, argininechloride, or ammonium chloride may beindicated

Metabolic Alkalosis

• Metabolic alkalosis isdefined as a primaryincrease in plasma[HCO3–]. Most cases ofmetabolic alkalosis canbe divided into (1) thoseassociated with NaCldeficiency andextracellular fluiddepletion, oftendescribed as chloridesensitive, and (2) thoseassociated withenhancedmineralocorticoid activity, commonly referred to aschloride resistant

Chloride-Sensitive MetabolicAlkalosis

• Depletion of extracellular fluid causes the renal tubules to avidlyreabsorb Na+. Because not enough Cl– is available to accompanyall the Na+ ions reabsorbed, increased H+ secretion must take placeto maintain electroneutrality. In effect, HCO3– ions that mightotherwise have been excreted are reabsorbed, resulting in metabolicalkalosis. Physiologically, maintenance of extracellular fluid volumeis therefore given priority over acid–base balance. Becausesecretion of K+ ion can also maintain electroneutrality, potassiumsecretion is also enhanced. Moreover, hypokalemia augments H+ secretion (and HCO3– reabsorption) and will also propagatemetabolic alkalosis. Indeed, severe hypokalemia alone can causealkalosis. Urinary chloride concentrations during a chloride-sensitivemetabolic alkalosis are characteristically low (< 10 mEq/L).

Chloride-Sensitive MetabolicAlkalosis

• Diuretic therapy is the most common cause of chloride-sensitivemetabolic alkalosis. Diuretics such as furosemide, ethacrynic acid, and thiazides increase Na+, Cl–, and K+ excretion, resulting in NaCldepletion, hypokalemia, and usually mild metabolic alkalosis. Lossof gastric fluid is also a common cause of chloride-sensitivemetabolic alkalosis. Gastric secretions contain 25–100 mEq/L of H+, 40–160 mEq/L of Na+, about 15 mEq/L of K+, and approximately200 mEq/L of Cl–. Vomiting or continuous loss of gastric fluid bygastric drainage (nasogastric suctioning) can result in markedmetabolic alkalosis, extracellular volume depletion, andhypokalemia. Rapid normalization of PaCO2 after plasma [HCO3–] has risen in chronic respiratory acidosis results in metabolicalkalosis (posthypercapnic alkalosis; see above). Infants being fedformulas containing Na+ without chloride readily develop metabolicalkalosis because of the increased H+ (or K+) secretion that mustaccompany sodium absorption.

Chloride-Resistant MetabolicAlkalosis

• Increased mineralocorticoid activity commonlyresults in metabolic alkalosis even when it is notassociated with extracellular volume depletion. Inappropriate (unregulated) increases inmineralocorticoid activity cause sodium retentionand expansion of extracellular fluid volume. Increased H+ and K+ secretion takes place tobalance enhanced mineralocorticoid-mediatedsodium reabsorption, resulting in metabolicalkalosis and hypokalemia. Urinary chlorideconcentrations are typically greater than 20 mEq/L in such cases.

Other Causes of MetabolicAlkalosis

• Metabolic alkalosis is rarely encountered in patientsgiven even large doses of NaHCO3 unless renalexcretion of HCO3– is impaired. The administration oflarge amounts of blood products and some plasmaprotein-containing colloid solution frequently results inmetabolic alkalosis. The citrate, lactate, and acetatecontained in these fluids are converted by the liver intoHCO3–. Patients receiving high doses of sodiumpenicillin (particularly carbenicillin) can developmetabolic alkalosis. Because penicillins act asnonabsorbable anions in the renal tubules, increased H+ (or K+) secretion must accompany sodium absorption. For reasons that are not clear, hypercalcemia thatresults from nonparathyroid causes (milk-alkalisyndrome and bone metastases) is also often associatedwith metabolic alkalosis. The pathophysiology ofalkalosis following refeeding is also unknown.

Treatment of Metabolic Alkalosis

• As with other acid–base disorders, correction ofmetabolic alkalosis is never complete until theunderlying disorder is treated. When ventilation iscontrolled, any respiratory component contributing toalkalemia should be corrected by decreasing minuteventilation to normalize PaCO2. The treatment of choicefor chloride-sensitive metabolic alkalosis isadministration of intravenous saline (NaCl) andpotassium (KCl). H2-blocker therapy is useful whenexcessive loss of gastric fluid is a factor. Acetazolamidemay also be useful in edematous patients. Alkalosisassociated with primary increases in mineralocorticoidactivity readily responds to aldosterone antagonists(spironolactone). When arterial blood pH is greater than7.60, treatment with intravenous hydrochloric acid (0.1 mol/L), ammonium chloride (0.1 mol/L), argininehydrochloride, or hemodialysis should be considered.

Diagnosis of Acid–Base Disorders• Interpretation of acid–base status from analysis of blood gases

requires a systematic approach. A recommended approach follows:– 1. Examine arterial pH: Is acidemia or alkalemia present? – 2. Examine PaCO2: Is the change in PaCO2 consistent with a

respiratory component? – 3. If the change in PaCO2 does not explain the change in arterial pH,

does the change in [HCO3–] indicate a metabolic component? – 4. Make a tentative diagnosis. – 5. Compare the change in [HCO3–] with the change in PaCO2. Does a

compensatory response exist? Because arterial pH is related to the ratioof PaCO2 to [HCO3–], both pulmonary and renal compensatorymechanisms are always such that PaCO2 and [HCO3–] change in thesame direction. A change in opposite directions implies a mixed acid–base disorder.

– 6. If the compensatory response is more or less than expected, bydefinition a mixed acid–base disorder exists.

– 7. Calculate the plasma anion gap in the case of metabolic acidosis. – 8. Measure urinary chloride concentration in the case of metabolic

alkalosis.

Normal Compensatory Responsesin Acid–Base Disturbances

• An alternative approach that is rapid butperhaps less precise is to correlate changesin pH with changes in CO2 or HCO3. For a respiratory disturbance, every 10 mm Hgchange in CO2 should change arterial pH byapproximately 0.08 U in the oppositedirection. During metabolic disturbances, every 6 mEq change in HCO3 also changesarterial pH by 0.1 in the same direction. If thechange in pH exceeds or is less thanpredicated, a mixed acid–base disorder islikely to be present.

Measurement of Blood GasTensions & pH

• Values obtained by routine blood gas measurement include oxygen and carbondioxide tensions (PO2 and PCO2), pH, [HCO3–], base excess, hemoglobin, and thepercentage oxygen saturation of hemoglobin. As a rule, only PO2, PCO2, and pH aremeasured. Hemoglobin and percentage oxygen saturation are measured with a cooximeter. [HCO3–] is derived using the Henderson–Hasselbalch equation and baseexcess from the Siggaard-Andersen nomogram.

Sample Source & Collection• Arterial blood samples are most commonly utilized clinically, though capillary or

venous blood can be used if the limitations of such samples are recognized. Oxygentension in venous blood (normally 40 mm Hg) reflects tissue extraction, notpulmonary function. Venous PCO2 is usually 4–6 mm Hg higher than PaCO2. Consequently, venous blood pH is usually 0.05 U lower than arterial blood pH. Despite these limitations, venous blood is often useful in determining acid–basestatus. Capillary blood represents a mixture of arterial and venous blood, and thevalues obtained reflect this fact. Samples are usually collected in heparin-coatedsyringes and should be analyzed as soon as possible. Air bubbles should beeliminated, and the sample should be capped and placed on ice to prevent significantuptake of gas from blood cells or loss of gases to the atmosphere. Although heparinis highly acidic, excessive amounts of heparin in the sample syringe usually lower pHonly minimally but decrease PCO2 in direct proportion to percentage dilution, andhave a variable effect on PO2.

pH Measurement• When a metal is placed in solution with its salt, the

tendency of the metal to ionize into the solution leavesthe metal with a negative charge. If two different metals(electrodes) and their salts are separated by a porouspartition (allowing transfer of charge), the tendency forone metal to go into solution more than the other resultsin an electromotive force between the two electrodes. For pH measurements, a silver/silver chloride electrodeand a mercury/mercurous chloride (calomel) electrodeare most commonly used. The silver electrode is incontact with the test solution through pH-sensitive glass. The calomel electrode interfaces with the test solutionthrough a potassium chloride solution and a porous plug. The electromotive force developed between the twoelectrodes is proportionate to [H+].

Suggested Reading• Kellum JA: Metabolic acidosis in the critically ill: lessons

from physical chemistry [clinical management of acuterenal failure in the ICU]. Kidney Int 1998;53:S-81. Anoutstanding review of the changes in understanding ofacid–base balance using the work of Stewart and others.

Kraut JA, Madias NE: Approach to patients with acid–base disorders. Respir Care 2001;46:392. [PMID: 11262558]

Longenecker JC: High-Yield Acid-Base. Williams & Wilkins, 1998.

Pestana C: Fluids and Electrolytes in the SurgicalPatient, 5th ed. Lippincott, Williams & Wilkins, 2000.

Blood gases

Why Blood Gases are obtained ?

Assess the oxygenation capacity of thecardiopulmonary system,

Assessment of oxygen pressure to guidetherapy,

Assessment of adequacy of ventilation,Assessment of acid-base status,Assessment of Hb concentration,electrolytes

and glucose

Blood Gases

Arterial - peripheral arterial bloodVenous - peripheral venous bloodMixed venous - pulmonary artery bloodCapillary - scalp stick, finger stick

Acid-Base Balance

The blood's acid-base balance is precisely controlled, because even a

minor deviation from the normal range can severely affect many organs. The body uses different mechanisms to

control the blood's acid-base balance.

Blood Gas Analysis

PO2 measuredpH measuredPCO2 measuredHCO3- CalculatedHb measured indirectlySat O2 measured (functional, fractional)Electrolytes Na+, K+, Cl-, Ca++,Mg++Other Lactate, Glucose

Acid-Base Balance

• The blood hydrogen ion (H+) concentration is maintained within narrow limits.

• The pH (negative logarithm of H+

concentration) is widely used in clinical medicine. The normal arterial blood pH ranges from 7.37 to 7.43

Acid-Base Balance

Both pulmonary and renal function act to compensate for disturbances in acid-base balance to maintain blood pH within normal ranges.

Acid-Base Balance - pH• Wide fluctuations in H+ concentration are also

prevented by the presence of several pH buffers.

• major pH buffer in the blood, and that which is most relevant to clinical acid-base disturbances, is the bicarbonate/carbonic acid system,

• Other buffers: proteins including hemoglobin,phosphates; ammonia; and bone

pHpH kidneykidney regulationregulation

H2CO3

lumen lumen cellcell bloodblood

HH++

ATPATPazaaza

NHE 3NHE 3

HCO3 + H+

H2CO3

AW

H2O + CO2 CO2+H2O

AW

H HCO3

Na

PROXIMAL

DISTALH+ + HCO3

H2CO3

AW

H2O + CO2

HPO4- -

H2PO4-

NH3

NH4+

pHpH kidneykidney regulationregulation

lumen lumen cellcell bloodblood

What is pCO2 ?

The partial preassure of CO2 dissolved inplasma

The enzyme carbonic anhydrase quickly converts carbonic acid in the blood to CO2and water. The partial pressure of CO2 gas (pCO2) is readily measured in blood samples and is directly proportional to blood CO2 content; therefore, PCO2 is used to represent the concentration of acid in the system.

Acid-Base Balance

the Henderson-Hasselbalch equation:pH = pH = pKapKa + log+ log HCOHCO33

0,03 x pCO0,03 x pCO22

Where: pKa is negative logarithm of the aciddissociation constant for carbonic acid is 6,1

0,03 – relates pCO2 to the amount of CO2

dissolved in plasma

Clinical disturbances of acid-base metabolism

• Clinical disturbances of acid-base metabolism classically are defined in terms of the HCO3

-/CO2buffer system.

• Rises or falls in HCO3- are termed metabolic

alkalosis or acidosis, respectively.• Rises or falls in PCO2 are termed respiratory

acidosis or alkalosis, respectively. • Simple acid-base disturbances include both the

primary alteration and an expected compensation. For example, in metabolic acidosis there is a primary fall in plasma HCO3

- concentration and a secondary fall in the PCO2 due to respiratory compensation.

BE – base excess• BE of normal patient (blood – pH 7.4;

pCO2 – 40 mm Hg; Hb – 15 mg/dL; temp. 37ºC) is zero.

BE=BB – NBB

BB – bufer baseNBB – normal bufore base

Oxygen saturation – SO2

• Oxygen saturation of hemoglobin.• Assess the effectiveness of oxygen

therapy.• Oxygen saturation is calculated from the

measured values of pH and pO2 .

Blood gases –reference range

25 – 4080 - 100pO2 (mm Hg)40 - 7096 - 97SO2 (%)

(- 2,5) –(+2,5)

(- 2,5) – (+2,5)BE (mmol/L)25 – 2923 - 27T CO2 (mmol/L)41 - 5135 – 45pCO2 (mm Hg)24 – 2822 - 26HCO3 (mmol/L)

7,32 –7,427,35 – 7,45pHVenous bloodArterial blood

Disturbances of Acid-BaseBalance

Above ?Is the pH above or below 7,35 – 7,45?

Below ?

Alkalosis pH = 7,5

Acidosis pH = 7,2

Acidosis• Acidosis is excessive blood acidity

caused by an overabundance of acid in the blood or a loss of bicarbonate from the blood (metabolic acidosis), or by a buildup of carbon dioxide in the blood that results from poor lung function or slow breathing (respiratory acidosis).

• If an increase in acid overwhelms the body's pH buffering systems, the blood will become acidic.

Respiratory Acidosis• Respiratory acidosis develops when the lungs do not

expel carbon dioxide adequately, a problem that can occur in diseases that severely affect the lungs (such as emphysema, chronic bronchitis, severe pneumonia, pulmonary edema, and asthma). Respiratory acidosis can also develop when diseases of the nerves or muscles of the chest impair breathing. In addition, a person can develop respiratory acidosis if overly sedated from opioids (narcotics) and strong sleeping medications that slow respiration

Metabolic Acidosis• Metabolic acidosis develops when the amount of

acid in the body is increased through ingestion of a substance that is, or can be metabolized to, an acid--such as wood alcohol (methanol), antifreeze (ethylene glycol), or large doses of aspirin (acetylsalicylic acid). Metabolic acidosis can also occur as a result of abnormal metabolism (diabetes mellitus).

• Even the production of normal amounts of acid may lead to acidosis when the kidneys are not functioning normally and are therefore not able to excrete sufficient amounts of acid in the urine.

Anion gapCalculation of the anion gap is often helpful in the differential diagnosis of metabolic acidosis.The anion gap is estimated by subtracting the sum of the Cl and HCO3

- concentrations from the plasma Na concentration. Negatively charged plasma proteins account for most of the anion gap; the charges of other plasma cations (K, Ca, and Mg) and anions (PO4, sulfate, and organic anions) tend to balance out. The accepted range for the normal anion gap is 7 – 16 mmol/L.

AG = (Na+ + K+) – (Cl- +HCO3-)

AlkalosisAlkalosis is excessive blood alkalinity caused by an overabundance of bicarbonate in the blood or a loss of acid from the blood (metabolic alkalosis), or by a low level of carbon dioxide in the blood that results from rapid or deep breathing (respiratory alkalosis).

Respiratory AlkalosisRespiratory alkalosis develops when rapid, deep breathing (hyperventilation) causes too much carbon dioxide to be expelled from the bloodstream. The most common cause of hyperventilation, and thus respiratory alkalosis, is anxiety. Other causes of hyperventilation and consequent respiratory alkalosis include pain, cirrhosis, low levels of oxygen in the blood, fever, and aspirin overdose (which can also cause metabolic acidosis)

Metabolic AlkalosisMetabolic alkalosis develops when the body loses too much acid or gains too much base. For example, stomach acid is lost during periods of prolonged vomiting or when stomach acids are suctioned with a stomach tube (as is sometimes done in hospitals). In rare cases, metabolic alkalosis develops in a person who has ingested too much base from substances such as baking soda (bicarbonate of soda). In addition, metabolic alkalosis can develop when excessive loss of sodium or potassium affects the kidneys' ability to control the blood's acid-base balance. For instance, loss of potassium sufficient to cause metabolic alkalosis may result from the use of diuretics or corticosteroids

Respiratory effects on pH

Respiratory Alkalosis

CO2 excessive loss

Hyperventilation

CO2 excessiveretained

Hypoventilation

Respiratory Acidosis

Metabolic effects on pH

Metabolic Alkalosis

Loss of H+VomitingExcessive AntacidIntake

Metabolic Acidosis

Excess of H+Diabetic

KetoacidosisRenal failure

Examples

95-100% (on air)92%95%95%98%SaO2(air)

+/-2 mmol/L-9-14+10+17BE22-26 mmol/L19123348SB11.3-14 kPa9.014.08.511.6pO2

4.7-6.0 kPa4.95.08.66.0pCO2

7.35-7.457.267.157.297.48pH

Normal ValuesPatient Results

Examples

-9

15.2

12.1

7.2

6.95

+/-2 mmol/L+8-9+10-9BE22-26 mmol/L34.915.23224.1SB11.3-14 kPa8.312.113.513.8pO2

4.7-6.0 kPa9.127.25.82.95pCO2

7.35-7.457.317.157.57.5pH

Normal ValuesPatient Results