Acid-base balance

From Iusmphysiology

• started here on 03/30/11.

Contents 1 Acid-base balance 1.1 Objectives 1.2 Normal pH in arterial blood 1.3 Threats to pH of extracellular fluid 1.4 Chemical buffers 1.5 Isohydric principle 1.6 Negative feedback controls on endogenous acid production 1.7 The importance of the HCO3 / CO2 buffer system 1.7.1 HCO3 / CO2 equilibrium 1.8 The lungs as an opening to the environment 1.9 4 acid-base distrubances 1.9.1 Respiratory acidosis 1.9.2 Respiratory alkalosis 1.9.3 Metabolic acidosis 1.9.3.1 Anion gap 1.9.4 Metabolic alkalosis 1.10 Davenport diagram: A magic, visual explanation 1.11 Examples of acid-base values

[edit] Acid-base balance

[edit] Objectives

• Define the following: acid, base, buffer, pH. Give the normal range of arterial blood pH and the limits compatible with life. Explain why constancy of pH is important.
• State the isohydric principle. List the important chemical buffers present in extracellular fluid, intracellular fluid, and bone.
• Write the Henderson-Hasselbalch equation for the bicarbonate/CO2 system. Write Henderson's equation for calculating [HCO3-] from [H+] and PCO2 measurements.
• Explain why the bicarbonate/CO2 system is so important.
• List the four simple acid-base disturbances. Describe for each: 1) the primary defect, 2) changes in arterial blood chemistry (pH, PCO2, and plasma [HCO3-]), 3) some common causes, 4) chemical buffering processes, and 5) respiratory and renal compensations.
• Given plasma electrolyte concentrations, calculate and interpret the anion gap.
• Given values for arterial blood pH, plasma [HCO3-], and PCO2 (or any two of the three), be able to identify the type of acid-base disturbance present.

[edit] Normal pH in arterial blood

• A normal pH is 7.4 (7.38-7.42), at which point the concentration of H+ is 40 (38 - 42) nmol / liter.
  • Survivable pH is 7.0-7.6 which is considered acidosis and alkalosis, respectively.
  • Note that this is a four fold change in H+ concentration.
• Death by change in pH occurs as a result of the changes in intracellular proteins upon change in pH.

[edit] Threats to pH of extracellular fluid

• There are two major forces that affect extracellular pH: oxidative phosphorylation and protein metabolism.
  • Both of these processes produce sources of acid: oxidative phosphorylation produces CO2 and protein metabolism produces H2SO4 and HCl.

• Oxidative phosphorylation as a source of extracellular acid:
  • Recall that oxphos dumps electrons onto oxygen and secretes this waste as CO2.
  • Recall that CO2--via carbonic anhydrase--affects the extracellular levels of H+ and HCO3-.
  • That is, as CO2 rises, more H2CO3 and H+ are found in the blood; the pH decrease.
  • We say CO2 is a volitile acid because it indirectly affects H+ levels.
  • 13-20 moles of CO2 are produced each day by oxphos.

• Protein metabolism as a source of extracellular acid:
  • Recall that proteins are digested for energy and as a source of amino acids.
  • Metabolism requires the removal and storage of the many hydrogens on the proteins and thus can generate H+.
    • Met or Cys metabolism generates H2SO4.
    • Lys or Arg metabolism generates HCl.
  • We say protein metabolism generates fixed acids because it directly affects H+ levels.
  • 40-60 moles of fixed acid are produced each day by protein metabolism.

[edit] Chemical buffers

• Recall that a chemical buffer is an acid (HA) and it's conjugate base (A-).
  • There is an equilibrium among these two components as the hydrogen dissociates and reassociates: HA <-> H+ + A-
  • The K_a describes the balance between the two entities (A- and HA): K_a = [H+][A-] / [HA]
• Recall the Henderson-Hassalbach equation which describes the pH in terms of the K_a and the amount of conjugate base (A-) present in a sample: pH = pK_a + log([A-] / [HA])
  • Recall that pK_{a = log(Ka) = log([H+][A-] / [HA])}
  • So the Henderson-Hasselbach equation takes the (log of the) expected balance ([H+][A-] / [HA]) and modifies it by adding (the log of) the ratio of conjugate base to conjugate acid.

• Chemical buffers of the body come exist in three compartments: ECF, ICF, and bone.
  • Each of the buffers in these three compartments can add or remove H+ from the system in order to buffer the pH.
  • The ECF contains HCO3- / CO2, plasma proteins, and inorganic phosphates.
  • The ICF contains proteins, organic phosphates, some HCO3- / CO2.
  • The bone contains pohosphate and carbonate salts.

[edit] Isohydric principle

• The isohydric principle says that at any given concentration of H+, all conjugate acid / base pairs are in equilibrium.
• This means that the pH can be determined from any pair of conjugate acid / base pairs using their pKa.
  • pH = pKa + log([A-] / [HA])
    • pH = 6.1 + log([HCO3-] / 0.03*PCO2)
    • pH = 6.8 + log([HPO4] / [H2PO4])
    • pH = pK_HPr + log([Pr-] / [HPr])

[edit] Negative feedback controls on endogenous acid production

• There is a negative feedback system that helps resist changes in systemic pH when endogenous acid production is high.
  • This system may have an effect during vigorous activity or diabetic ketoacidosis.
• This negative feedback system is based on cellular enzymes that produce endogenous acid (think back to oxphos and protein metabolism as endogenous sources of acid).
  • Named examples of acid-producing processes include ketoacidosis and lactic acidosis.
• Negative feedback of pH change is probably mediated by the inherent effect that a change in pH has on these enzymes that make endogenous acid: as as pH goes up, endogenous acid decreases (because the enzymes don't function as well at high pH) and as pH drops endogenous acid increases (because the enzymes work better at higher pH).

[edit] The importance of the HCO3 / CO2 buffer system

• The HCO3 / CO2 buffer has two characteristics that make it very effective: abundance of buffer pair molecules, an "open" system.

• Bicarb / CO2 is a good buffer because of it's abundance:
  • There is 24 mM of HCO3 in the blood.
  • There is 1.2 mM of CO2 in the blood and 13-20 moles produced each day.
  • Having lots of molecules means that many molecules of H+ can be bound or unbound as needed.

• Bicarb / CO2 is a good buffer because it is an "open" system:
  • Recall that CO2 can be breathed off at the lungs and HCO3 can be secreted by the kidneys.
  • Because of these physiological connections to organs that "open" up to the external environment, the HCO3 / CO2 system is considered "open".
  • Note that the lungs can respond quickly (within minutes) and the kidneys respond slowly (within days) to balance HCO3 / CO2.

[edit] HCO3 / CO2 equilibrium

• Recall the carbonic anhydrase driving equation: CO2 (aq) + H20 <-> H2CO3 <-> H+ + HCO3
• Recall that the amount of aqueous CO2 is always 0.03 * PCO2.
  • This is useful because PCO2 is measurable in the clinical setting.
  • When PCO2 is 40 mmHg (pretty normal), CO2 (aq) is 1.2 mM.
• So when we calculate pH based on this:
  • pH = pKa + log([A-] / [HA])
  • pH = 6.1 + log([HCO3 / [CO2])
  • pH = 6.1 + log([HCO3 / 0.03 * PCO2)
  • pH = 6.1 + log(24 / 1.2)
    • The 20:1 ratio of HCO3 : aqueous CO2 is key to having a normal pH.
  • pH = 6.1 + log(20)
  • pH = 6.1 + 1.3
  • pH = 7.4

• But there is a simpler way: the clinical equation.
  • In clinical settings, logs are not really practical so we use an alternative equation.
  • The clinical equation uses two values easily procured by labs (PCO2 and [HCO3]) to calculate the concentration of H+ ([H+]) which is compared to an expected value (40 nM) to determine if the pt is acidotic ([H+] higher than 40 nM) or alkalotic ([H+] lower than 40 nM).
  • [H+] = 24 * PCO2 / [HCO3]
    • Recall that PCO2 and [HCO3] are usually known so we are calculating to find [H+].
  • We have an expectation as to what a normal [H+] should be (40 nM) and this equation tells us if [H+] is high (low pH, acidosis) or low (high pH, alkalosis).

[edit] The lungs as an opening to the environment

• Recall that the lungs can breath off CO2 and thus decrease the pH of the blood (shifts equation away from H+ + HCO3-).
• As ventilation increases, more CO2 is breathed off.
• Ventilation is stimulated by increased arterial PCO2 or decreased pH.
• Note that neither of these forces can cause at pt to ventilate to their full voluntary ventilation capacity; thus, one can hyperventilate themselves into passing out--because one can throw off acid-base balance by ventilating too rapidly so the body arrests auto-hyperventilation by blacking out.

[edit] 4 acid-base distrubances

• There are four acid-base distrubances: acidosis and alkalosis arising from respiratory or metabolic origins.
• Recall that there are three compensation mechanisms that are very fast, fast, and slow: chemical buffering of the blood by HCO3 / CO2 (very fast, seconds), respiratory compensation by breathing off CO2 (fast, minutes), and renal compensation by excreting H+ / HCO3 (slow, days).
• Note that in respiratory acid-base imbalances, the primary variable that has changed is CO2 (because the lungs can only increase or decrease this one variable).
  • Therefore, compensation mechanisms in pulmonary malfunction will always attempt to change HCO3 in the same direction as CO2.
• Note that in metabolic acid-base imbalances, the primary variable that has changed can be H+ or HCO3- (because metabolic processes can cause a change in either acid or base production).
  • Therefore, compensation mechanisms in metabolic malfunction can attempt to change the opposite factor (base in an acid-malfunction or acid in a base-malfunction) in the same direction or can attempt to counteract the changes to CO2 levels.

[edit] Respiratory acidosis

• Respiratory acidosis is defined as "any abnormal pulmonary function that results in CO2 accumulation".
• As CO2 accumulates the equation shifts to the right, toward H+ and HCO3.
  • However, HCO3 production does not keep up with the shift so H+ > HCO3-.

• Common cause: hypoventilation

• Compensation:
  • Recall that compensation mechanisms attempt to normalize the HCO3 / CO2 ratio by moving the opposite factor in the same direction; therefore as CO2 increases, the compensation mechanisms will attempt to increase HCO3, also.
  • Chemical buffering: mostly achieved by proteins binding H+ in cells (think Hb).
  • Pulmonary compensation: none, lungs are the problem
  • Renal compensation: raise the HCO3 content of the blood; attempt to restore a HCO3 / CO2 ratio of 20:1.
    • This is an attempt to modify the HCO3 / CO2 ratio by increasing HCO3.
    • Acute respiratory acidosis results in a renal compensation of about 1 mEq / L of HCO3 production for every 10 mmHg increase in PCO2.
    • Chronic respirartory acidosis results in a renal compensation of about 4 mEq / L of HCO3 production for every 10 mmHg increase in PCO2.

[edit] Respiratory alkalosis

• Definition: "any abnormal pulmonary function that results in CO2 deficiency".
• As CO2 is released, the equation shifts to the left, decreasing H+ and leaving HCO3 to cause alkalosis.

• Common causes: hyperventilation

• Compensation:
  • Recall that compensation mechanisms attempt to normalize the HCO3 / CO2 ratio (to 20:1) by changing the non-effected value in the same direction.
  • In respiratory alkalosis, the CO2 level is too low so the body attempts to compensate by decreasing the amount of HCO3-.
  • Chemical buffering: mostly achieved by release of H+ by intracellular proteins.
  • Pulmonary compensation: none, lungs are the problem.
  • Renal compensation: decrease the plasma HCO3- levels
    • decrease HCO3- production, increase HCO3- secretion

[edit] Metabolic acidosis

• Definition: "any abnormal function that results in a gain of acid or loss of base (excepting the gain of H2CO3)".
  • Recall that while pulmonary acid-base imbalances result from changes in CO2,] metabolic acid-base imbalances can result from acid or base changes.
  • Adding an acid makes the pt acidotic and pushes the HCO3 / CO2 reaction toward CO2.
  • Removing a base makes the pt acidotic and pulls the HCO3 / CO2 reaction toward H+ / HCO3-.

• Common causes:
  • renal failure,
  • excessive intake of nonvolatile acids,
  • excessive production of nonvolatile acids (ketoacidosis, lactic acidosis, ingestion of acidosis, ingestion of acidifying agents),
  • poisons (salicylate, methanol, ethylene glycol),
  • loss of bicarbonate (excessive urinary excretion of bicarbonate [renal tubular acidosis], diarrhea)

• Compensation:
  • Chemical buffering: half of the buffering occurs in the cells and bone, HCO3- is the main ECF buffering base.
  • Pulmonary compensation: prompt hyperventilation to lower the PCO2; cannot completely compensate, however.
    • Note that in the case of added acid, this would breathe off the elevated CO2, decreasing acid.
    • Note that in the case of removed base, this would provide counter force along the equation axis to keep some of the acid in CO2 form.
  • Renal compensation: increased H+ secretion, increased ammonia synthesis, increased bicarb reabsorption / production.

[edit] Anion gap

• Recall that metabolic acidosis can result from a depletion of HCO3.
• The chief purpose of calculating the anion gap is to determine the source of HCO3 depletion; that is, to determine the etiology of metabolic acidosis.

• Recall from general chemistry that charges in a solution like to be in equilibrium; the body maintains a certain equilibrium of positively charged ions (cations) to negatively charged ions (anions).
  • In general, the anions and the cations are in equilibrium (as with all solutions).
• We can simplify this to a short equation if we only take into account the major ions:
  • [Na+] = [Cl-] + [HCO3-] + [unmeasured anions]
• The anion gap is the difference in the cations and the anions.
  • Note that we drop the unmeasured anions term because...well...they are unmeasured so we don't know the value.
    • Unmeasured anions include lactate, ketones, proteins, phosphate, citrate, and sulfate to name a few.
    • These will become causes of acidosis when they are aberrently elevated.
  • Anion gap = [Na+] - [Cl-] - [HCO3-]
  • Normal anion gap = 140 - 105 - 24 = 11 mEq / L
  • This makes sense because Na is the highest concentration of cations, Cl is the major important anion, and HCO3- is the variable.

• So what does an increased anion gap mean?
  • Recall that a normal anion gap = 140(Na+) - 105(Cl-) - 24(HCO3-) = 11 mEq / L.
  • An increase in the anion gap means that there is less HCO3 to subtract from the Na or there is less Na from which to subtract the HCO3.
    • Metabolic acidosis generates an increased anion gap because of a deficiency of HCO3 secondary to over production of metabolic acids like lactic acid, ketone body acids, or toxins.

• So what does a normal anion gap mean?
  • It could mean that there is no acid-base problem. :)
  • It could mean that the metabolic acidosis is caused by renal tubular acidosis, here's why:
    • When HCO3- is lost (loss of a base leads to acidosis) at the renal tubule (perhaps because of diarrhea in which high volume requires lots of HCO3- secretion to maintain filtrate pH; perhaps because of poor HCO3 reabsorption at the PCT) it can be exchanged for Cl-.

Is this exchange an active exchange like Na for K in water balance or is it indirect like "well because HCO3 wasn't reabsorbed and yet the filtrate must be ionically balanced, Cl gets reabsorbed"? 

• • • When HCO3- is exchanged for CL- at the renal tubule, the anion gap doesn't change, yet HCO3- is being depleted. That is, when metabolic acidosis occurs without an anion gap, one knows that the renal tubule is the source of HCO3 depletion.
    • Metabolic acidosis generates a normal anion gap because of a deficiency of HCO3 secondary to over excretion of HCO3 at the renal tubule.

• Common causes of anion-gap acidosis:
  • MULEPaKS
  • Methanol
  • Uremia (renal failure, urine in the blood)
  • Lactic acid
  • Ethylene glycol
  • pAldehyde
  • Ketone body acids
  • Salicylates

[edit] Metabolic alkalosis

• Definition: "any abnormal function that results in a gain of base or loss of acid (especially the gain of bicarbonate but not the loss of H2CO2)".
  • Recall that gain of bicarb or another base will push the HCO3 / CO2 reaction toward CO2.
  • Recall that loss of acid will drag the HCO3 / CO2 reaction toward H+ / HCO3.

• Common causes:
  • Acid loss: vomiting (stomach juices have lots of H+), hyperaldosteronism (excessive H+ loss at the kidney), hypokalemia (excessive H+ loss at the kidney via the H-K exchanger that will reabsorb K by sacrificing H+ to the filtrate)
  • Base gain: excessive alkali intake

• Compensation:
  • Chemical buffering: 1/3 in the cell compartment
  • Pulmonary compensation: hypoventilation to increase PCO2 (but cannot fully compensate)
    • Recall that HCO3- cannot be generated as quickly as acid when the HCO3 / CO2 equation is shifted toward H+ / HCO3.
    • So increasing PCO2 will compensate for excessive acid loss by pushing the equation toward H+ / HCO3 and thus regenerating some of the acid.
    • So increasing PCO2 will compensate for base gain by pushing the equation toward H+ / HCO3 and regenerating acid (faster than it regnerates HCO3) to counter the excess in base.
  • Renal compensation: reabsorb less PCO3 and therefore lower the plasma HCO3.
    • Note that this will compensate for increased base by decreasing another base: HCO3.
    • Note that this will compensate for decreased acid by equalizing the ratio of acid to base.

[edit] Davenport diagram: A magic, visual explanation

[edit] Examples of acid-base values

pH	PCO2 (mmHg)	[HCO3-] (mEq / L)	State
7.08	49	14	Respiratory acidosis + Metabolic acidosis (pulmonary disease -> low O -> high lactic acid)
7.32	28	14	Metabolic acidosis	Respiratory compensation
7.40	40	24	Normal	NA
7.51	49	38	Metabolic alkalosis	Respiratory compensation
7.53	20	16	Respiratory alkalosis	Renal compensation
7.62	20	20	Respiratory alkalosis	Respiratory compensation

• stopped here on 03/30/11.