CO2 Transport

CO2 is a waste product of aerobic metabolism. The body must be able to eliminate this waste product to maintain normal function. CO2 dissolves out of the tissue cells into the blood stream where it is carried to the lungs for elimination through ventilation.

 

co2 to Hco3.jpg

Fig. 7-1.  How CO2 is converted to HCO3 at the tissue sites.  Most of the CO2 that is produced at the tissue cells is carried to the lungs in the form of HCO3.

bicarb to co2.jpg

Fig. 7-2.  How HCO3 is transformed back into CO2 and eliminated in the alveoli.

Dissolved in solution

CO2 dissolves into the plasma and the intracellular fluid of the erythrocyte. The partial pressure exerted by the CO2 in solution is what drives the rest of the reactions, so even though the dissolved component is only responsible for approximately 5% of the CO2 that is released to the lungs, it is still an important transport role.

For every mm Hg PCO2 pressure, 0.03 mEq of CO2 are physically dissolved in one liter of plasma. The normal arterial PCO2 is 40 mm Hg, therefore the amount of CO2 dissolved in the plasma can be calculated as follows:

Be careful not to confuse the factor for quantifying dissolved CO2 (0.03) with the factor for quantifying the dissolved O2 (0.003)!

Combined with proteins

CO2 combines with proteins in the plasma and forms carbamino compounds, and combines with hemoglobin in the RBC to form carbaminohemoglobin.

Converted to Bicarbonate

In the plasma at the tissue / systemic capillary level, the hydrolysis (the combining of CO2 with H2O) is a very slow reaction, so only a small amount forms carbonic acid which rapidly dissociates into H+ and HCO3- ions. However, in the RBC there is a catalyst called carbonic anhydrase that accelerates the hydrolysis (13,000 times more rapidly), so that the majority of the CO2 in the RBC is converted quickly to carbonic acid which then dissociates into the hydrogen and bicarbonate ions. Bicarbonate is exchanged for the chloride ion in the plasma, and the reduced hemoglobin binds with the hydrogen ion. At the alveoli / pulmonary capillary level the reversal of the pressure gradients causes all of these processes to reverse so that CO2 diffuses out into the lungs.

table 7-1.jpg

 

 

fig7-3.jpg

Fig. 7-3.  Carbon dioxide dissociation curve.

fig7-4.jpg

Fig. 7-4.  Carbon dioxide dissociation curve. An increase in the PCO2 from 40 mm Hg to 46 mm Hg raise the CO2 content by about 5 vol.%.  PCO2 changes have a greater effect on CO2 content levels than PO2 changes on O2 levels.

fig7-5.jpg

Fig. 7-5.  Carbon dioxide dissociation curve at two different oxygen/hemoglobin saturation levels (SaO2 of 97% and 75%).  When the saturation of O2 increases in the blood, the CO2 content decreases at any given PCO2.  This is known as the Haldane effect.

The low SaO2 at the tissue increases the bloods capacity to hold CO2 and facilitates the loading of CO2 into the blood at the tissues; the high SaO2 at the lungs decreases the bloods capacity to hold CO2 and this facilitates its unloading at the lungs.

fig7-6.jpg

Fig. 7-6.  Comparison of the oxygen and carbon dioxide dissociation curves in terms of partial pressure, content, and shape.

 

Acid - Base Balance

As stated in Egan's, the body's metabolic processes continually generate hydrogen ions . . . even small hydrogen ion concentration changes can cause metabolic processes to fail. Acid-base balance relies on the mechanisms that keep hydrogen ion concentration of body fluids compatible with life by maintaining the pH between 7.35 and 7.45.

Three major organ systems devote much of their activity to regulating acid and alkali levels: the lungs, the blood, and the kidneys. Buffering of acids is largely carried out by the blood and kidneys; excretion of acids is handled by the lungs and kidneys. Of the 12,000mEq pf H+ produced per day, 99% of this is excreted by the lungs. (Respiratory Anatomy and Physiology, Martin & Youtsey)

Everything we do as Respiratory Care Practitioners has this goal - even oxygen administration, by preventing anaerobic metabolism and the generation of lactic acid, helps to maintain proper pH.

 

Hydrogen Ion Regulation in Body Fluids

Hydrogen ions come from volatile or fixed acids. Most H+ is the result of the breakdown of proteins, anaerobic metabolism, or the metabolism of body fats. These are called fixed acids. H+ that are the result of CO2 hydrolysis (H2CO3) is called a volatile acid as it is an acid in equilibrium with a dissolved gas.

Fixed acids are not in equilibrium with a dissolved gas and come from the catabolism of proteins:

Sulfuric and phosphoric acids

Lactic acid (from anaerobic glycolysis)

Ketoacidosis (from fatty acid metabolism)

H+ of fixed acids combines with HCO3- to form H2CO3 and causes an equilibrium imbalance that drives the reaction to the left to form CO2 and H2O

Two major mechanisms that stabilize pH in response to CO2 generated by metabolism are:

Isohydric buffering: H+ is carried by deoxygenated hemoglobin and is released as CO2 in the lungs

Ventilation: elimination of carbonic acid = production

Strong Acids

Acidity of a solution reflects only the free hydrogen ions, not the hydrogen ions still combined with anions. 

Thus, strong acids, which dissociate completely (i.e., they liberate all the H+) and irreversibly in water, dramatically change the pH of the solution.

For example, if 100 hydrochloric (HCl) acid molecules were placed in 1 mL of water, the hydrochloric acid would dissociate into 100 H+ and 100 Cl ions. There would be no undissociated hydrochloric acid molecules in the solution.

Weak Acids

Do not dissociate completely in a solution and have a much smaller effect on pH.

Although weak acids have a relatively small effect on changing pH levels, they have a very important role in resisting sudden pH changes.

Examples of weak acids are carbonic acid (H2CO3) and acetic acid (HC2H3O2

If 100 acetic acid molecules were placed in 1 mL of water, the following reaction would occur:

weak acid.jpg

Because undissociated acids do not alter the pH, the acidic solution will not be as acidic as the HCl solution discussed above

The dissociation of acetic acid can be written as follows:

acetic acid.jpg

Using this equation, it can be seen that:

When H+ (released by a strong acid) is added to the acetic acid solution, equilibrium moves to the left as some of the additional H+ bonds with C2H3O2 to form HC2H3O2

On the other hand, when a strong base is added to the solution (adding additional OH and causing the pH to increase), the equilibrium shifts to the right. This occurs because the additional OH consumes the H+, causing more HC2H3O2 molecules to dissociate and replenish the H+ .

Weak acids play a very important role in the chemical buffer systems of the human body

Strong Bases

 Remember, bases are proton acceptors 

Strong bases dissociate easily in water and quickly tie up H+

Hydroxides are strong bases

Weak Bases

In contrast, weak bases:

Sodium bicarbonate or baking soda is a weak base

Dissociate incompletely and reversibly and are slower to accept protons 

Because sodium bicarbonate accepts a relatively small amount of protons

Its released bicarbonate ion is described as a weak base

 

pH: Acid-Base Concentration

As the concentration of hydrogen ions in a solution increase, the more acidic the solution becomes. 

As the level of hydroxide ions increases the more basic, or alkaline, the solution becomes.

Clinically, the concentration of hydrogen ions in the body is measured in units called pH units, which is a way to quantify the amount of [H+] in the form of a positive number.

pH scale runs from 0 to 14 and is logarithmic

Each successive unit change in pH represents a tenfold change in hydrogen ion concentration

pH of a solution, therefore, is defined as the negative logarithm, to the base 10, of the hydrogen ion concentration [H+] in moles per liter

pH.jpg

When the pH is 7 ([H+] = 10-7 mol/liter):

Number of hydrogen ions precisely equals the number of hydroxide ions (OH)

And the solution is neutral - Neither acidic or basic

Pure water has a neutral pH of 7, or 10-7 mol/liter (0.0000001 mol/liter) of hydrogen ions. 

A solution with a pH below 7, is acidic

There are more hydrogen ions than hydroxide ions

For example, a solution with a pH of 6 has 10 times more hydrogen ions than a solution with a pH of 7

A solution with a pH greater than 7, is alkaline

Hydroxide ions outnumber the hydrogen ions

For example, a solution with a pH of 8 has 10 times more hydroxide ions than a solution with a pH of 7

Thus, as the hydrogen ion concentration increases hydroxide ion concentration falls, and vice versa.

 

pH Values of Representative Substances

 

fig7-7.jpg

Figure 7-7  The pH values of representative substances.  The pH scale represents the number of hydrogen ions in a substance. The concentration of hydrogen ions (H+) and the corresponding hydroxyl concentration (OH–) for each representative substance is also provided. Note that when the pH is 7.0, the amount of H+ and OH– are equal and the solution is neutral.

 

Acid-base balance is what keeps [H+] in normal range.

 Tissue metabolism produces massive amounts of CO2, which is hydrolyzed  into the volatile acid H2CO3.

CO2 + H2O → H2CO3 → H+ + HCO3

 

Reaction is catalyzed in RBCs by carbonic anhydrase. As the lungs eliminate CO2, the falling CO2 reverses the reaction.

 

Ventilation

  ↑

  CO2 + H2O ← H2CO3 ← HCO3 +  H+

Buffer solution characteristics

A solution that resists changes in pH when an acid or a base is added, in other words they minimize the effects of [H+] or [OH-] changes

Composed of a weak acid and its conjugate base

i.e., carbonic acid/bicarbonate, which in blood exists in reversible combination as NaHCO3 and H2CO3

Add strong acid HCl + NaHCO3 → NaCl + H2CO3, and it is buffered with only a small acidic pH change.

Add base NaOH + H2CO3 → NaHCO3 + H2O, and it is buffered with only a slight alkaline pH change.

Can be open or closed buffer systems.

Bicarbonate and nonbicarbonate buffer systems

Bicarbonate: composed of HCO3 and H2CO3

Open system as H2CO3 is hydrolyzed to CO2. Ventilation continuously removes CO2 preventing equilibration, driving reaction to the right

   HCO3 + H+ → H2CO3 → H2O + CO2

 Removes vast amounts of acid from body per day

Can buffer both strong acid and base and thus prevent major changes in pH of the solution

Able to buffer nonvolatile fixed acids, however cannot buffer volatile acids as it is in equilibrium with the volatile acid.

Has the greatest buffering capacity

Nonbicarbonate: composed of phosphate and proteins (including hemoglobin)

Consists of a weak acid and its base or salt:

HHB (acid hemoglobin) KHB(potassium hemoglobin)

KH2PO4 (potassium acid phosphate) K2HPO4 ( potassium alkaline phosphate)

H Protein (acid porteinate) NA protein (sodium proteinate)

NaH2PO4 (sodium acid phosphate) NaHPO4 (sodium alkaline phosphate)

Closed system as no gas is available to remove acid by ventilation. All components remain in the system - when equilibrium is reached no further buffering can occur

Hbuf/buf– represents acid and conjugate base.

   H+ + buf– ↔ Hbuf reaches equilibrium, buffering stops

Buffers volatile acids.

Both systems are important so that the buffering of both fixed and volatile acids occurs.

 

Chemical Buffer Systems and Acid-Base Balance

Chemical buffers resist pH changes and are the body's first line of defense. 

Ability of an acid-base mixture to resist sudden changes in pH is called its buffer action. 

Tissue cells and vital organs of the body are extremely sensitive to even the slightest change in the pH environment 

In high concentrations, both acids and bases can be extremely damaging to living cells

Essentially every biological process within the body is disrupted

Buffers work against sudden and large changes in the pH of body fluids by

1. Releasing hydrogen ions (acting as acids) when the pH increases, and

2. Binding hydrogen ions (acting as bases) when the pH decreases.

Three major chemical buffer systems in the body are the:

Carbonic acid-bicarbonate buffer system

Phosphate buffer system

Protein buffer system

Carbonic Acid-Bicarbonate Buffer System and Acid-Base Balance

The carbonic acid-bicarbonate buffer system plays an extremely important role in maintaining pH homeostasis of the blood

Carbonic acid (H2CO3) dissociates reversibly and releases bicarbonate ions (HCO3) and protons (H+) as follows:

Response to an increase in pH - H+ proton donor

  H2CO3 -> HCO3  +  H+

Response to a decrease in pH  - H+ proton acceptor 

H2CO3 <- HCO3  +  H+

Under normal conditions, the ratio between the HCO3 and H2CO3 in the blood is 20:1

co2 to Hco3.jpg

Chemical equilibrium between carbonic acid (weak acid) and bicarbonate ion (weak base) works to resist sudden changes in blood pH.

For example, when the blood pH increases (i.e., becomes more alkaline from the addition of a strong base), the equilibrium shifts to the right.

A right shift forces more carbonic acid to dissociate, which in turn causes the pH to decrease.

In contrast, when the blood pH decreases (i.e., becomes more acidic from the addition of a strong acid), the equilibrium moves to the left. 

A left shift forces more bicarbonate to bind with protons.

Carbonic acid-bicarbonate buffer system converts:

1. Strong bases to a weak base (bicarbonate ion), and

2. Strong acids to a weak acid (carbonic acid) 

Blood pH changes are much less than they would be if this buffering system did not exist.

 

The Henderson-Hasselbalch Equation (H-H)

H-H equation mathematically illustrates how the pH of a solution is influenced by the HCO3 to H2CO3 ratio (the bicarbonate buffer system); the base to acid ratio

H-H equation is written as follows:

hh equa.jpg

pK is derived from the dissociation constant of the acid portion of the buffer combination

pK is 6:1 and, under normal conditions, the HCO3 to H2CO3 ratio is 20:1

Clinically, the dissolved CO2 (PCO2 x 0.03) can be used for the denominator of the H-H equations, instead of the H2CO3  

This is possible since the dissolved carbon dioxide is in equilibrium with, and directly proportional to, the blood [H2CO3], the PaCO2 is easily measured via blood gas analysis and can easily be converted to mmol/L (same as mEq/L).

Thus, the H-H equation can be written as follows:

hh2.jpg

H-H Equation Applied During Normal Conditions

When the HCO3 is 24 mEq/L, and the PaCO2 is 40 mm Hg, the base to acid ratio is 20:1 and the pH is 7.4 (normal). 

H-H equation confirms the 20:1 ratio and pH of 7.4 as follows:

app HH.jpg

The ratio is the important factor, not the individual concentrations.

A HCO3- of 48 and a PCO2 of 80 would still give a ratio of 20/1

H-H Equation Applied During Abnormal Conditions

 When the HCO3 is 29 mEq/L, and the PaCO2 is 80 mm Hg, the base to acid ratio decreases to 12:1 and the pH is 7.18 (acidic) 

H-H equation confirms the 12:1 ratio and the pH of 7.18 as follows:

 

abn hh.jpg

In contrast, when the HCO3 is 20 mEq/L, and the PaCO2 is 20 mm Hg, the base to acid ratio increases to 33:1 and the pH is 7.62 (alkalotic) 

H-H equation confirms the 33:1 ratio and the pH of 7.62 as follows:

abn hh alk.jpg

 

The Respiratory System and Acid-Base Balance

Respiratory system does not respond as fast as the chemical buffer systems.

However, it has up to two times the buffering power of all of the chemical buffer systems combined. 

CO2 produced by the tissue cells enters the red blood cells and is converted to HCO3 ions as follows:

dissoc of carbonic acid.jpg

The first set of double arrows illustrates a reversible equilibrium between the dissolved carbon dioxide and the water on the left and carbonic acid on the right.

The second set of arrows shows a reversible equilibrium between carbonic acid on the left and hydrogen and bicarbonate ions on the right

Because of this relationship, an increase in any of these chemicals causes a shift in the opposite direction 

Note also that the right side of this equation is the same as that for the carbonic acid-bicarbonate buffer system

Under normal conditions, the volume of CO2 eliminated at the lungs is equal to the amount of CO2 produced at the tissues.

When the CO2 is unloaded at the lungs, the preceding equation flows to the left, and causes the H+ generated from the carbonic acid to transform back to water.

bicarb to co2.jpg

Because of the protein buffer system, the H+ generated by the CO2 transport system is not permitted to increase

Therefore, it has little or no effect on blood pH

The Renal System and Acid-Base Balance

kidney.jpg

Kidneys

Physically remove H+ from body

Excrete <100 mEq fixed acid per day

Also control excretion or retention of HCO3

If blood is acidic, then more H+ are excreted and all the HCO3 is retained, vice versa

While lungs can alter [CO2] in seconds, the kidneys require hours to days change HCO3  and affect pH.

Role of urinary buffers in excretion of excess H+

Once H+ has reacted with all the available HCO3, the excess reacts with phosphate and ammonia.

If all urinary buffers are consumed, further H+ filtration ends when pH falls to 4.5.

Activation of ammonia buffer system enhances Cl loss and HCO3 gain.

 

Even though the chemical buffer systems can inactivate excess acids and bases momentarily, they are unable to eliminate them from the body. 

Similarly, although the respiratory system can expel the volatile carbonic acid by eliminating CO2, it cannot expel other acids generated by cellular metabolism

Only the renal system can rid the body of acids such as phosphoric acids, uric acids, lactic acids, and ketone acids (also called fixed acids).

Only the renal system can regulate alkaline substances in the blood and restore chemical buffers that are used in managing H+ levels in extracellular fluids

Some HCO3, which helps to adjust H+ concentrations, is lost from the body when CO2 is expelled from the lungs.

When the extracellular fluids become acidic, the renal system retains HCO3 and excretes H+ ions into the urine: This causes the blood pH to increase. 

When the extracellular fluids become alkaline, the renal system retains H+ and excretes basic substances primarily HCO3 into the urine: This causes the blood pH to decrease

open closed system.jpg

 

Acid-Base Balance Disturbances

Normal bicarbonate (HCO3) to carbonic acid (H2CO3) ratio in the blood plasma is 20:1. 

In other words, for every H2CO3 produced in blood plasma, 20 HCO3 ions must be formed to maintain a 20:1 ratio (normal pH). 

Or, for every H2CO3 loss in the blood plasma, 20 HCO3 ions must be eliminated to maintain a normal pH. 

In other words, the H2CO3 is 20 times more powerful than the HCO3 ion in changing the blood pH.

Under normal conditions, the 20:1 acid-base balance in the body is automatically regulated by the:

Chemical buffer systems

Respiratory system

Renal system

However, these normal acid-base regulating systems have their limits. 

The bottom line is this:

The body's normal acid-base watchdog systems cannot adequately respond to sudden changes in H+ and HCO3 concentrations regardless of the cause.

For example:

Hypoventilation causes the partial pressure of the alveolar carbon dioxide (PACO2) to increase, which in turn causes the plasma PCO2, HCO3, and H2CO3 to all increase.

  1. This causes HCO3 to H2CO3 ratio to decrease, and the pH to fall.

 

fig 7-8.jpg

Fig. 7-8.  Alveolar hypoventilation causes the PACO2 and the plasma PCO2, H2CO3, and HCO3 to increase. This action decreases the HCO3/H2CO3 ratio, which in turn decreases the blood pH.

Or, when the PACO2 decreases, as a result of alveolar hyperventilation, the plasma PCO2, HCO3 and H2CO3 all decrease which in turn causes:

HCO3 to H2CO3 ratio to increase, and the pH to rise

 

fig 7-9.jpg

Fig. 7-9.  Alveolar hyperventilation causes the PACO2 and the plasma PCO2, H2CO3, and HCO3– to decrease.  This action increases the HCO3–/H2CO3 ratio, which in turn increases the blood pH.

Acid Excretion

Buffers are a temporary measure; if acids were not excreted, life-threatening acidosis would follow.

Lungs

Excrete CO2, which is in equilibrium with H2CO3

Crucial as body produces huge amounts of CO2 during aerobic metabolism (CO2 + H2O → H2CO3)

In addition, through HCO3 eliminate fixed acids indirectly as the byproducts are CO2 and H2O

Lungs remove ~24,000 mmol/L CO2 daily

The is inversely proportional to the PaCO2. As minute ventilation increases PaCO2 decreases; if minute ventilation decreases PaCO2 increases. Since minute ventilation is tidal volume multiplied by the respiratory rate (or frequency) then we can alter the PaCO2 by changing either or both of these components.

Kidneys

Physically remove H+ from body

Excrete <100 mEq fixed acid per day

Also control excretion or retention of HCO3

If blood is acidic, then more H+ are excreted and all the HCO3 is retained, vice versa

While lungs can alter [CO2] in seconds, the kidneys require hours to days change HCO3  and affect pH.

Role of urinary buffers in excretion of excess H+

Once H+ has reacted with all the available HCO3, the excess reacts with phosphate and ammonia.

If all urinary buffers are consumed, further H+ filtration ends when pH falls to 4.5.

Activation of ammonia buffer system enhances Cl loss and HCO3 gain.

 

Relationship between acute PCO2 changes, and the resultant pH and HCO3– changes that occur is graphically illustrated in the PCO2/HCO3–/pH nomogram

fig 7-10.jpg

 

PCO2/HCO3–/pH nomogram is an excellent clinical tool that can be used to identify a specific acid-base disturbance

Anion Gap

In the body electrical neutrality is maintained. Cations = Anions. The anion gap is a way to differentiate between a metabolic acidosis caused by HCO3- loss, or an increase in fixed acids. By looking at the difference between the measurable cations and ions we can evaluate the cause of a metabolic acidois.

The formula for calculating the anion gap is;

[Na+] - ([Cl-] + [HCO3-] = anion gap

Normal anion gap is 9 - 14 mEq/L

A gap of > 14 mEq/L is considered a metabolic acidosis due to the presence of fixed acids in the body

The following is taken from Mosby's online course for Egan's:

The anion gap clarifies whether the metabolic acidosis is caused by the accumulation of fixed acids or loss of HCO3.

When fixed acids accumulate, the anion gap increases, resulting in the following chain of events:

With HCO3 loss, the anion gap does not change. As a result:

 

Common Acid-Base Disturbance Classifications

Respiratory Acid-Base Disturbances

Primary respiratory disturbances

As PaCO2 is controlled by the lung, changes in pH caused by PaCO2 are considered respiratory disturbances

Hyperventilation lowers PaCO2, which raises pH, so is referred to as respiratory alkalosis.

Hypoventilation raises PaCO2 which decreases the pH, so is called respiratory acidosis.

Disorders are identified as

Acute ventilatory failure (respiratory acidosis)

Acute ventilatory failure with partial renal compensation

Chronic ventilatory failure with complete renal compensation

Acute alveolar hyperventilation (respiratory alkalosis)

Acute alveolar hyperventilation with partial renal compensation

Chronic alveolar hyperventilation with complete renal compensation

Metabolic Acid-Base Disturbances

Primary metabolic disturbances

These disturbances involve a gain or loss of fixed acids or HCO3.

Both will appear as changes in HCO3 as changes in fixed acids will alter the amount of HCO3 used in buffering.

A decrease in HCO3 results in a metabolic acidosis.

An increase in HCO3- results in metabolic alkalosis.

Disorders are identified as:

Metabolic acidosis with partial respiratory compensation

Metabolic acidosis with complete respiratory compensation

Metabolic alkalosis with partial respiratory compensation

Metabolic alkalosis with  complete respiratory compensation

Combined Disorders

Both metabolic and respiratory acidosis

Both metabolic and respiratory alkalosis

 

 

Base Excess / Base Deficit

Calculation of the base excess or deficit is a way of quantifying HCO3-.

Base excess is the quantity of base (HCO3-, in mEq/L) that is above or below the normal range of buffer base in the body (22 -28 mEq/L). This cannot be calculated from PCO2 and pH as the hemoglobin also contributes to the buffer base. One can use the Siggaard-Andersen nomogram to estimate base excess or deficit. Another way is to estimate using the following calculations:

For example: if the pH is 7.24, PaCO2 is 28, and the Hb is 15 gm/dl

a) Find the difference between PaCO2 and 40 mm Hg

40 mm Hg - 28 mm Hg = 12 mm Hg

b) Move the decimal two places to the left

12 becomes 0.12

c1) If the PaCO2 is > 40 mm Hg, subtract half the difference from 7.40 to estimate a predicted respiratory pH

c2) If the PaCO2 is < 40 mm Hg, add the entire difference to 7.40.

7.40 + 0.12 = 7.52

d) Find the difference between the measured and predicted respiratory pH

7.24 - 7.52 = - 0.28

e) Move the decimal two places to the right

-0.28 becomes -28

f) Multiply by 2/3 to calculate base excess

(-28 x 2)/3 = -18 mEq/L

Base excess/deficit of +/- 2 mEq/L is normal.

Severe metabolic acidosis is associated with a base deficit of -10 mEq/L

A positive number is called a base excess and indicates a metabolic alkalosis.

A negative number is called a base deficit and indicates a metabolic acidosis.

 

Compensation: Restoring pH to normal

•Any primary disturbance immediately triggers a compensatory response.

Any respiratory disorder will be compensated for by the kidneys (process takes hours to days - 1 to 3 days).

Any metabolic disorder will be compensated for by the lungs (rapid process, occurs within minutes - 1 to 3 minutes).

Respiratory acidosis (hypoventilation)

Compensatory Response: Renal retention of HCO3 raises pH toward normal

Respiratory alkalosis

Compensatory Response: Renal elimination of HCO3  lowers pH toward normal

Metabolic acidosis

Compensatory Response: Hyperventilation ↓CO2, raising pH toward normal

Metabolic alkalosis

Compensatory Response: Hypoventilation ↑CO2, lowering pH toward normal

The CO2 hydration reaction's effect on [HCO3–]

A large portion of CO2 is transported as HCO3. So any thing that affects CO2 will also affect HCO3 levels, yet this is not part of the renal compensatory response - only a result of the hydrolysis reaction. As such, slight changes in the HCO3 in the blood are an expected consequence of elevated or decreased PCO2 and should not be interpreted as compensatory.

As CO2 increases, it also increases HCO3.

In general, the effect is an acute increase of ~1 mEq/L HCO3 for every 10 mm Hg acute increase in PaCO2 .

An increase in CO2 of 30 would increase HCO3 by ~3 mEq/L . This should not be seen as compensation.