Exploring Acid-Base Physiology

“I should refresh myself on acid-base stuff” I naively thought to myself, after being asked to give some impromptu teaching at an ALS course. “Probably just recap the bicarb-CO2 thingy and that’ll cover it”.

Oh, sweet summer child.

After many hours of trying to find the ABC answer to this topic, I’ve discovered no such certainty exists. There is no watertight, nicely presented framework or theoretical underpinning to the problem of acid-base balance. Instead, there is a Wild West of chemists slinging contentious theories at each other, with multiple differing algorithms and approaches. It’s a black hole of never-ending tangents and off-shoots.

Disclaimer: I have selectively included and ignored aspects of this treacherous topic at my own whim. For example, I have not covered the Six Simultaneous Equations used by Stewart, nor other such physiochemical wizardry. See references yourself if you want to get into the weeds.

Why is acid-base important?

Probably a reasonable first question. Generally, bodily processes operate within a tight pH between 7.35 and 7.45. 

• In chemistry, neutral pH is 7. The body tends to be slightly more alkali (around 7.35-7.45). 
• pH affects the O2-Hb dissociation curve. Oxygen offloads from Hb more readily in acidic environments, such as metabolically active tissues.
• Protein structure and function becomes more unstable and potentially denatured in lower pH.
• Acidosis is negatively inotropic, initially compensated by catecholamine drive, but worsening with pH <7.2. Risk of tachycardia and arrhythmia is also higher in acidosis. 

There are two key questions before even looking at a blood gas:

1. Do you need one?
   • A blood gas will tell us many things, but uniquely it can give insight into 1) ventilation/oxygenation and 2) acid-base balance. I will be primarily focusing on acid-base balance in this article.
   • Thankfully the fetish for ABGs in every patient with even mild hypoxia is becoming less prevalent, but always consider how a blood gas will influence management. Patient just had a tonic-clonic seizure in front of you? Don’t get the gas for a lactate (it will be high), but it may be useful to see their sodium. I’ve heard of clinically well patients being kept on wards just because their lactate is slightly above normal on repeat gases. That leads us to…
     
2. How does the patient look?
   • You need to know the clinical context attached to the numbers. A mild acidosis may be no big deal in a diabetic ketoacidosis, but could be extremely worrying in a calcium channel blocker overdose. A study of 6 elite rowers found an average lactate of 26 and pH of 6.85 after 2000 metres rowing; needless to say they didn’t need intubated and wheeled off to critical care. pH on its own does not tell the whole story; the pathology driving the numbers is what counts. 

Bare bones chemistry

.What does acid-base balance actually mean? Essentially, its the interaction between different systems within the body to influence hydrogen ion concentration and, thus, pH.

Hydrogen is an atom.

A hydrogen ion is a hydrogen atom without its orbital electron; functionally, therefore, it exists as a proton. It is depicted as H+ due to only being a hydrogen ion with a single proton, with the two terms often being used interchangeably. 

Hydrogen ions exist in much smaller concentrations in the blood than other ions, and so are measured in nanomoles (thousand-millionths of a mole). For comparison, sodium, potassium and other electrolytes are measured in millimoles (thousandths of a mole).

The absolute number of hydrogen ions may be way smaller than other ions, but we tolerate a higher margin of relative flux in hydrogen ion concentration. For example, a 2.5x increase in [H+] from 40 to 100 nmol (equivalent change of pH 7.4 to 7.0) is survivable, but a 2.5x flux in sodium or potassium would kill us.

Fundamentally, acidosis is a rise in hydrogen ion concentration [H+] whilst alkalosis is a reduction in [H+]. 

Hydrogen ions do not generally exist in the body in freely, dissociated form as protons just floating around. If not buffered by substances such as proteins, Hb, phosphate, or bicarbonate, they are associated with water as H₃O⁺. 

H+ + H2​O → H3​O+

For simplicity, this form of associated hydrogen ion is usually just referred to as ‘free’ hydrogen ions (H+). 

pH is a representation of the concentration of hydrogen ions ([H+]), but is not in itself a concentration. Because of this, pH does not have any units: it is just a number. Hydrogen ion concentration determines the pH and, therefore, the acidity of a solution. 

Henderson-Hasselbalch

The Henderson-Hasselbalch (H-H) equation gave us a shorthand for measuring the acidity or alkalinity of a solution; pH. This provides a measurable and sensible representation of hydrogen ion concentration.

• pH = a measure of the acidity/alkalinity. Expressed as the negative logarithm of [H+] – ie a lower pH means a higher [H+]
• pKa = how strong an acid is, determined by how easily it dissociates to H+ and its conjugate base in solution
• [A-] = concentration of conjugate base, after HA has dissociated
• [HA] = concentration of acid that hasn’t yet dissociated and given up its H+
• Log = to keep measurements of [H+] within a reasonable, understandable range, the logarithmic scale is used. This is a nonlinear scale, where a change of one pH unit reflects a 10-fold change in hydrogen ion concentration [H+]. A normal blood pH of 7.4 reflects a [H+] of 40 nmol/l, whereas at pH 7.0 it is 100 nmol/l. 

So, what actually causes a rise or fall in pH in the body?

Two Broad Schools:

Two broad schools of thought have evolved as a framework to explain what may be going on:

1) Traditional Approach

Building on work by Arrhenius in the 1800’s, Bronsted and Lowry established our traditional clinical understanding of acid-base chemistry in 1923. The key idea they proposed is that acids donate hydrogen ions, whereas bases accept them. 

• An acid donates H+; after donating, it becomes a conjugate base, due to loss of a proton.
• A base accepts H+; after accepting, it becomes a conjugate acid, due to gain of a proton.

For example:

HCl + H₂O → Cl⁻ + H₃O⁺

• HCl is an acid, that donates a proton to H2O, to become a conjugate base (Cl-)
• H2O is a base, that accepts a proton, to become a conjugate acid (H₃O⁺)
• The generation of H₃O⁺ results in a reduction of pH through increased acidity.
• If we were to add HCl to water, this increases acidity (reduces the pH) due to the binding of H+ to neutral water molecules.

So the traditional approach holds that if we add HCl to a solution, pH will reduce and acidity will increase as a direct result of the addition of H+. 

As this theory became ingrained in clinical practice, attention was increasingly focused on PaCO2 and HCO3- as the primary determinants of bodily pH. This was further cemented by incorporating carbonic acid and bicarbonate into the Henderson-Hasselbach equation.

The carbonic acid equilibrium equation

CO₂ + H₂O <-> H₂CO₃ <-> HCO₃– + H+

This chemical equation underpins most of clinical reasoning in acid-base interpretation, placing CO2, H+ and HCO3- at the centre of our clinical understanding of acid-base, acidosis and alkalosis.

Whilst it is accurate and sufficient for respiratory acid-base interpretations, we should be aware that it may only provide a superficial, incomplete explanation about what’s going on metabolically. 

It’s main shortcoming is that it relies on only using CO₂, HCO₃– and H+ to explain all disturbances in acid-base physiology. Unfortunately, there is a lot more going on metabolically than just these three.

2) Stewart Approach

In 1978, Peter Stewart proposed a physicochemical approach to acid–base physiology that differed from conventional proton-centred explanations.

He asserted that pH is determined by 3 laws/constraints:

1. The law of conservation of mass means that matter cannot be created or destroyed by ordinary chemical reactions; substances won’t just disappear from solution, but may recombine or redistribute in different ways.
2. The law of mass action means that the concentration of substances in a reversible chemical reaction (think carbonic acid, for example) determine the direction and outcome of that equilibrium reaction.
3. The law of electroneutrality determines that the electrochemical state of a solution, determined by the sum of positive and negative charges, will always be balanced.

Stewart defined an acid as any ion that shifts the dissociation equilibrium of water to higher concentrations of hydrogen ions. Water acts as a reservoir for hydrogen, and is constantly participating in equilibrium reactions in concert with the dynamic shifts in ions, acids, and carbon dioxide in the body.

H2​O ⇌ H+ + OH−

Hydrogen and bicarbonate ions are dependent variables that can increase or decrease in response to changes in the following three independent variables.

Three independent variables:

1. Carbon dioxide / pCO2
2. Strong ion difference (SID)
3. Total nonvolatile weak acids (Atot) – mainly albumin and phosphate

Simply put, the law of electroneutrality dictates that the electrochemical charge of a solution in the body must be balanced. Whatever is added or removed from the body, the new state must remain electrochemically neutral.

Take the dissociation of hydrogen chloride as an example (remember H₃O⁺ is the same as H+)

HCl + H₂O → Cl⁻ + H₃O⁺

HCl dissociates to Cl⁻ and H+. Whilst Cl⁻ contributes a negative charge to the body, H+ is buffered by different proteins or bicarbonate. This Cl⁻ changes the strong ion difference, consumes HCO3-, and redistributes H+ through the body; the addition of a solution of HCl has altered the electrochemical landscape in the body.

The 3 independent variables that can alter the final electrochemical landscape are:

1. Strong ions; negatively charged Cl⁻ may be partially balanced by other cations such as Na+, K+, Ca2+ and so on.
2. Weak acids; such as albumin and phosphate are less likely to shift dramatically.
3. Carbon dioxide; increased H+ will stimulate chemoreceptors to increase respiration as a compensatory response.

Due to changes in the above, the dependent variables of H+ and HCO3- will shift and alter to the new strong ions of Cl⁻. They are second-order outcomes of shifts in the above three independent variables. The net redistribution of hydrogen ions will ultimately determine the pH.

In the final analysis, the increased concentration of chloride will contribute an increased negative charge (and reduced SID – see below for more on this), less bicarbonate, and more hydrogen ions than before. The net amount of hydrogen ions will be greater, resulting in an acidosis.

Water within the body participates in these electrochemical equilibration reactions by shifting left or right, to result in higher or lower concentrations of hydrogen ions (which are positively charged). These auto-ionisation reactions are a result of changes in the three independent variables.

H2​O ⇌ H+ + OH−

Q: What is the role of OH- in all this from water dissociation? Doesn’t its existence just balance out any H+ added?

Water undergoes continuous autoionisation, producing H⁺ and OH⁻. However, the concentrations of these free ions are very low compared with the concentration of water itself.

Although each autoionisation event produces one H⁺ and one OH⁻, their free concentrations in the body are not equal. At physiological pH (approximately 7.4), the concentration of OH⁻ exceeds that of H⁺. This reflects the ion product of water and the combined acid–base equilibria of the body’s fluids.

Changes in strong ion concentrations, weak acids or PaCO₂ alter these equilibria and therefore the free H⁺ concentration and pH. Throughout these changes, electroneutrality is maintained: the net charge of all ions in solution must balance.

A quick checkpoint:

After that initial foray into acid-base history, its useful to pause and take stock.

For my own sake, here is a brief summary of conclusions we can draw so far.

1. Serum hydrogen ion (H+) concentration changes pH and acid-base status.
2. Stewart proposes that the 1. strong ion difference, 2. carbon dioxide concentration, and 3. non-volatile weak acids, rather than direct addition of H+ or HCO3- per se, are the main drivers of acid-base disturbance.
3. The carbonic acid equilibrium equation is a good explanation and clinically useful for respiratory acid-base disturbance. This holds true across both the Traditional and Stewart approaches.

IMPORTANT CAVEAT:
Stewart was a science man rather than a clinician. This means that although his approach may provide a deeper understanding of what causes acid-base disturbance, trying to wholesale adopt his entire framework will usually be prohibitively unwieldy at the bedside.

This doesn’t stop some clinicians doing their darndest to turn bedside acid-base analysis into an impossibly masochistic clinical tool.

Nonetheless, after gaining an appreciation for this, we can parse away and incorporate that which is helpful.

Respiratory Acid-Base balance

Mercifully, there is broad agreement between the Traditional and Stewart approach when it comes to explaining respiratory acid-base balance.

We don’t need to throw the carbonic acid baby out with the protonated bathwater, and can thankfully still cling onto this much touted pillar of acid-base teaching; at least as far as respiratory perturbations are concerned.

Carbon Dioxide / pCO2

At rest, an adult produces about 200-250mls of CO2 per minute, which is ventilated off via the lungs, maintaining arterial pCO2 within a normal physiologic range. 

CO₂ + H₂O <-> H₂CO₃ <-> HCO₃– + H+

A respiratory acidosis occurs when there is a rise in the partial pressure of CO2 in the blood. This can occur from increased presence of inspired CO2 (not common), increased addition of CO2 in the body (such as malignant hyperthermia and laparoscopic CO2), or decreased alveolar ventilation (nearly all cases in clinical practice are due to this). A change in the partial pressure of CO2 will influence serum pH. If CO2 rises, then we see a rightward shift in the carbonic acid dissociation equation, resulting in the formation of H+ and HCO3-. This increase in H+ concentration causes a reduced pH and increased acidity. 

Well why doesn’t the HCO3- produced just buffer this increased H+ concentration? 

After creating the H+ and HCO3 from CO2 in our body, we can’t just immediately correct an underlying hypercapnia by reversing the reaction back to CO2 and H2O. If a respiratory acidosis persists, then the kidneys will increase reabsorption of HCO3-, but this is something that takes 6-12 hours to occur, reaching maximal effect by around 3-4 days.

So if it’s not HCO3- that buffers H+ from an acute respiratory acidosis, then what does?

In the immediate term of a respiratory acidosis, the excess hydrogen concentration is buffered primarily intracellularly by haemoglobin, phosphates and other intracellular proteins. These buffering process are limited in their capacity, leaving increased ventilation (in the acute phase) and increased bicarbonate creation/reabsorption (as a longer compensatory mechanism) as the primary means to correcting the acidosis. If ventilation doesn’t alter accordingly, hypercapnic respiratory failure and worsening acidosis ensues.

Increased pCO2 has a number of physiologic effects:

• Rightward shift of the O2-Hb curve; causing reduced Hb affinity for O2 to offload for tissues
• Cerebral and peripheral vasodilation
• Pulmonary vascular resistance; hypercapnia increases PVR via vasoconstriction
• Sympathetic stimulation
• Increased ventilation via central (medullary) and peripheral (carotid bodies and aortic arch) chemoreceptors; these detect changes in pH and pCO2. 

The body will respond to a respiratory acidosis rapidly by increasing ventilation; this is because CO2 easily crosses the blood brain barrier to stimulate the respiratory control centre in the medulla. However, at a certain level of CO2, the respiratory drive will become blunted and hypercapnic narcosis may develop; this is what we see with significant respiratory acidosis in, for example, an exacerbation of COPD. 

Metabolic Acid-Base Balance

Thankfully, respiratory acid-base changes are explained simply enough using our well known carbonic acid dissociation equation as above.

Metabolic acid-base disturbance is more complex and contested.

During the Traditional era, two useful concepts were introduced to help address this black-box of metabolic acid-base interpretation:

Base excess and the anion gap.

Base Excess

Base excess is available on a blood gas as a number. The normal is zero, with a physiologic range of -2 to +2. 

What is it?

The number given represents the amount of excess or deficit of base in the blood. A positive base excess implies an excess of base (metabolic alkalosis), whereas if it is negative, this suggests a deficit of base (and a metabolic acidosis, given that base is being used up to buffer protons). 

Base excess (BE) is an aggregate measure of metabolic changes, that considers all buffering capacity; this includes HCO3- and non-bicarbonate buffers (haemoglobin, plasma proteins such as albumin, and phosphates).

Importantly, BE is unaffected by respiratory acidosis/alkalosis, as it is calculated using a ‘normal’ pCO2 (40 mmHg / 5.3 kPa), and so base excess specifically reflects only metabolic (non-respiratory) changes in bicarbonate and related buffers, but not the immediate respiratory-driven buffering processes. However, a chronic respiratory acidosis or alkalosis would affect the HCO3-, which would alter the base excess.

Limitations:

There are some limitations to BE. It only measures blood/serum buffers, and so won’t pick up on the frantic intracellular buffering of H+ from excess CO2 in a respiratory acidosis, for example.

The fact it is an aggregate measure of all buffering systems also means it doesn’t say what specifically is going on, but just points in a direction for metabolic acidosis/alkalosis. 

Because BE is influenced by anions and proteins such as albumin, as well as phosphate, if these are low (like hypoalbuminaemia in liver disease, or hypophosphataemia), the base excess may be altered.

Anion Gap

The anion gap (AG) complements BE by providing a bit more detail about what could be causing a metabolic disturbance. 

AG = ([Na+] + [K+]) − ([Cl−] + [HCO3−])

Anion gap serves to differentiate a non-anion gap metabolic acidosis (NAGMA) from an anion gap metabolic acidosis (AGMA).

• NAGMA is essentially a hyperchloraemic acidosis
• AGMA is either:
  • an increase in endogenous acid eg lactate, ketones etc
  • ingestion of exogenous acid eg toxic alcohols, salicylate etc

Anion gap recognises the need for electroneutrality in the body.

As previously explained, this is a fundamental law that states every solution must have a net charge of zero, with an equal balance between all the positive ions (cations) and negative ions (anions). The anion gap is calculated by measuring the difference in charge between the commonly measured extracellular ions; Na+, K+, Cl- and HCO3-. When calculated, this reveals a normal ‘gap’ of around 8±4 mmol/l (around 4-12 if K is not included).

This gap is made up of the unmeasured anions such as lactate, sulfate, ketones, phosphate, and weakly negatively charged proteins such as albumin. This 8±4 mmol/l represents a ‘normal’ amount of this unmeasured stuff in a healthy individual. The range of a normal anion gap is reasonably wide, which is an inherent weakness in this calculation. But clinically significant acid load from unmeasured ions will generally be large enough to fall above this normal range and be picked up on AG calculation.

If any of these unmeasured anions increases, this will subsequently reduce HCO3- as constrained by the law of electroneutrality. This reduction in HCO3- will give a raised anion gap value.

Therefore, if there is an acidosis with a widened gap, we can assume that there is an increase in the hidden, unmeasured anions contributing to the metabolic acidosis. This could be ketoacidosis, lactic acidosis, uraemia, salicylate poisoning, etc. It’s useful to have a mnemonic to sieve through the potential culprits in a raised anion gap acidosis

CAT MUDPILES

1. Carbon monoxide/cyanide
2. Alcoholic ketoacidosis
3. Toluene
4. Methanol
5. Uremia
6. Diabetic ketoacidosis
7. Paracetamol/acetaminophen
8. Iron/isoniazid
9. Lactic acidosis
10. Ethanol/ethylene glycol
11. Salicylate/ASA

Limitations:

Similar to BE, the anion gap calculation can become less reliable in states of low albumin and phosphate, where a ‘normal’ gap may be hard to define for such individuals. A low albumin may hide a raised anion gap acidosis, by giving a falsely normal/low anion gap. As a rough estimate, every 1 g/L decrease in albumin will decrease the anion gap by 0.25 mmol/L

A correction for albumin in calculating the anion gap can be as follows:

Corrected AG = AG + 0.25 x (40 – albumin (g/L))

This corrected anion gap formula can be used when there are significant deviations from the norm in albumin levels.

A further limitation is that a high anion gap acidosis and normal anion gap acidosis can co-exist, reducing the utility of the anion gap measurement.

Respiratory compensation

Applying Winter’s Formula can help to establish whether or not there is adequate respiratory compensation in a metabolic acidosis. An apparently ‘normal’ pCO2 in the context of a metabolic acidosis may actually represent a mixed respiratory-metabolic acidosis.

Expected pCO2 = (1.5 x [HCO3^–]) + 8 ±2

Stewart Approach – further refining metabolic acidosis analysis

As stated, the Stewart approach holds that a rise or fall in H+ concentration is dependent on variations in electrochemical charge, rather than the direct addition/subtraction of H+ per se. 

This physiochemical approach assigns H+ and HCO3- as dependent variables that are only altered by other external and independent variables; namely strong ions, carbon dioxide, and weak acids.

The primary dependent variables that can be altered are:

• H+ (pH)
• HCO3-

The 3 independent variables that can influence H+ and HCO3- concentration are:

• Carbon dioxide / pCO2
• Strong ion difference (SID)
• Total nonvolatile weak acids (Atot)

We have already established that carbon dioxide / pCO2 is a widely agreed model for respiratory acid-base balance, which leaves us with SID and Atot:

Strong ions and Strong Ion Difference (SID)

Strong ions are those ions that are fully dissociated in solution, both cations (positive ions) and anions (negative ions). 

In the body, these cations are primarily sodium, potassium, magnesium and calcium, whilst the anions include chloride, lactate and ketones. 

So the strong ion (SID) difference is calculated by:

([Na+]+[K+]+[Ca2+]+[Mg2+])−([Cl−]+[lactate−])

Due to the relatively low numbers of the other ions, this equation can be simplified for clinical application to [Na+] – [Cl-]. Potassium, magnesium and calcium are kept in such a tight range that they can clinically be excluded from the calculation. Lactate similarly is normally minimal.

Reference ranges usually for Na (135-145) and Cl (95-108).

A typically normal SID (for Na-Cl) is approximately 38 to 42 mEq/L. This difference comprises HCO3- and weak acids (plasma proteins (mostly albumin), phosphate and other unmeasured anions).

When calculating a simplified SID, any increase in chloride, or reduction in sodium, will reduce this difference, with a low SID, such as <38 mEq/L, representing a hyperchloaremic acidosis. 

This is analogous to a NAGMA (non-anion gap metabolic acidosis).

Wait, how does this rise in anions cause an acidosis?

A decrease in strong ion difference (SID) can result from an increase in strong anions, such as chloride, or a decrease in strong cations relative to strong anions.

In Stewart’s framework, a decrease in SID shifts the acid–base equilibrium towards a higher H⁺ concentration and a lower bicarbonate concentration, resulting in a metabolic acidosis.

These changes occur while electroneutrality is maintained throughout; H⁺ is not generated as a compensatory response to an electrical imbalance. Rather, H⁺ and HCO₃⁻ are dependent variables whose concentrations are determined by the altered SID, total weak acid concentration, PaCO₂ and the relevant chemical equilibria.

Physiological compensation may subsequently occur. For example, increased ventilation lowers PaCO₂, partially correcting the fall in pH. Renal responses may also contribute, depending on the underlying disturbance.

The full SID calculation ([Na+]+[K+]+[Ca2+]+[Mg2+])−([Cl−]+[lactate−]) would reveal an acidosis/alkalosis due to fluctuations in those ions within the equation. 

Total Weak Acids (Atot)

The final independent variable is total weak acids (Atot), comprising mainly albumin and phosphate. In the absence of other anions (such as lactate, ketones etc), Atot is hidden in the anion gap, given this gap is normally comprised of albumin and phosphate, as well as small amounts of other anions.

The amount of serum phosphate is quite small, and so fluctuations rarely produce clinically significant changes in acid-base status. The contribution of albumin to acid-base is significantly larger than phosphate.

Why is albumin a weak acid?

Albumin has many amino acid side chains. Some amino acid side chains (such as those of lysine and arginine) are positively charged. However, the larger amount of carboxyl groups on side chains of other amino acids are negatively charged, as they lose their H+ (-COOH → COO- + H+). This imbalance leaves albumin with a net negative charge, and it acts as weak acid to influence the total hydrogen ion. 

Due to this, fluctuations in albumin can influence acid base balance.

Hypoalbuminaemia (often seen in critical care, cirrhosis, etc) will reduce the total amount of weak acid, reducing net negative charge, increasing HCO3- and reducing H+, to give a raised pH with subsequent increased alkalosis.

The reverse would be true for hyperalbuminaemia, such as in administration of albumin solution. ↑ albumin = ↑ weak acid = ↑ negative charge = ↑ H+ = ↓ pH / ↑ acidosis. 

With this in mind, if plasma protein levels are normal, then acid-base disturbances can be solely analysed in terms of changes in the other 2 independent variables; pCO2 and SID. 

When there are significant deviations in normal plasma proteins (mainly albumin), these should be taken into account when analysing acid-base. 

What is the difference between the strong ion difference and anion gap?

The strong ion difference is the gap between cations and anions. The full equation includes the following:

([Na+]+[K+]+[Ca2+]+[Mg2+])−([Cl−]+[lactate−]). 

Therefore, any change in the measured SID variables will be captured in this equation. These would primarily be hyperchloraemic and lactic acidoses, given acid-base disturbances from fluctuations in the other ions is rare.

On the other hand, the anion gap is measured thus:

(Na+ + K+) – (Cl- + HCO3-).

The key difference is that the AG includes a dependent variable; HCO3-.

This dependent variable can be influenced by any of the independent variables mentioned earlier (pCO2, SID and Atot). In the context of an acute metabolic acidosis, any increased anion will reduce the measured serum HCO3- in order to maintain electroneutrality. This decrease in HCO3- will result in a raised anion gap, indicating unmeasured anions’ indirect influence on ‘squeezing out’ HCO3- due to their anionic negative charge.

So the inclusion of HCO3- in the anion gap equation allows us to see if there is any increase in the unmeasured anions, by the indirect influence on HCO3- altering the anion gap. Whereas the strong ion difference is only influenced by changes in the ions directly in the equation, given they are independent variables.

My approach:

All this begs the very reasonable question: why overly complicate things?

We have a traditional, workable approach to acid-base interpretation that is easily applied to the vast majority of clinical case. Such ramblings on the topic fly in the face of Occam’s razor, where the simplest explanation is best.

“Interpretation of acid-base disorders will always remain partly an art, one that combines an intelligent synthesis of the clinical history, physical examination, and other ancillary laboratory data taken together in the context of the individual patient and the nature and temporal course of his or her disease.” – unknown

The physiochemical approach provides a more narrative framework for why acid-bases disturbances occur, but I’m not going to apply its numerous equations to the bedside, because it’s absurdly impractical for me.

Trying to understand the theory, whilst clinically applying it in the simplest way possible is my aim.

As deranged as it may be, this is my website, and therefore the following is my own approach to acid-base interpretation:

1. pH: is there a clear acid-base disturbance?
2. pCO2: Is there a respiratory contribution?
   • pCO2 and pH go in opposite directions:
     • Resp acidosis: low pH, high pCO2
     • Resp alkalosis: high pH, low pCO2
3. Metabolic snapshot; Base excess, HCO₃⁻, lactate
   • BE corrects for any respiratory contribution and provides a pure metabolic impression. Abnormal HCO3- also helps point toward acid-base disturbance.
   • Eg. a bicarb of 10 or BE of -14.0 is useful in highlighting a metabolic acidosis, but neither bicarb nor BE give the cause. 
   • Lactate: check for lactic acidosis
4. Anion Gap: Check for unmeasured contributors
   • If raised, consider CATMUDPILES
   • If normal anion gap, consider causes (GIT losses, renal tubular acidosis, iatrogenic chloride)
5. Consider albumin: Adjust anion gap if necessary
   1. If there is hypoalbuminaemia present, correct for this with the following equation: Corrected AG = AG + 0.25 x (40 – albumin (g/L))

Simplified:

1. pH
2. pCO2
3. Metabolic snapshot; BE & HCO3-, Lactate
4. Anion gap
5. Consider albumin

This entire article barely scrapes the surface for more in-depth acid-base analysis, and it doesn’t even cover differentials and causes for abnormalities. My main aim was to establish a grounding in the fundamentals. I’d suggest reading further from the references below to get a better understanding, if that’s the kind of thing you are after.

As always, all criticisms/corrections welcome.

References

1. National Center for Biotechnology Information. (2019). “Stewart Acid-Base Balance.” In StatPearls. Available at: https://www.ncbi.nlm.nih.gov/books/NBK507807/
2. Anaesthesia MCQ. “Acid-Base Physiology.” Available at: http://www.anaesthesiamcq.com/AcidBaseBook/ABindex.php
3. Acid-Base. “The Acid-Base Physiology Community.” Available at: https://acidbase.org
4. Weingart, S. “AG Metabolic Acidosis.” Internet Book of Critical Care (IBCC). EmCrit Project. Available at: https://emcrit.org/ibcc/agma/
5. Weingart, S. (2011). “Understanding Acid-Base Disorders Using the Stewart Approach.” EmCrit Project. Available at: https://emcrit.org/wp-content/uploads/2011/04/acid-base-stewart.pdf
6. Maseeh, A., & Saleh, A. (2014). “Facing acid-base disorders in the third millennium – the Stewart approach revisited.” International Journal of Nephrology and Renovascular Disease, 7, 349-354. doi:10.2147/IJNRD.S61538
7. Jones, D. “Stewart Acid-Base: A Simplified Bedside Approach.” The Open Mind – Anaesthesia and Analgesia. Available at: https://www.openmindanaesthesia.org/stewart-acid-base-bedside-approach
8. The Physiological Society. (2020). “An Introduction to Stewart Acid-Base.” The Physiological Society Magazine. Available at: https://www.physoc.org/magazine-articles/an-introduction-to-stewart-acid-base/
9. National Center for Biotechnology Information. (2021). “Metabolic Acidosis in Critical Care and Anesthesia.” In StatPearls. Available at: https://www.ncbi.nlm.nih.gov/books/NBK559139/
10. Park, M.A.J., Cave, G., & Freebairn, R.C. (2020). “Metabolic Acidosis in Anaesthesia and Critical Care.”Anaesthesia Intensive Care.
11. Moviat, M.A.M., van Haren, F.M.P., & van der Hoeven, J.G. (2003). “Stewart’s approach: Just a heresy or another lens into acid-base physiology?” Critical Care, 7(3), 245-246.
12. Morgan, T.J. (2009). “Acid-base analysis: a critique of the Stewart and bicarbonate-centered approaches.”Anaesthesia Intensive Care, 37(4), 482-487.
13. Fencl, V., & Leith, D.E. (2004). “Bench-to-bedside review: Fundamental principles of acid-base physiology.” Critical Care, 7(3), 160-164.
14. Deranged Physiology. “Traditional and Physico-chemical Approaches to Acid-Base.” Deranged Physiology – CICM Primary Exam Required Reading. Available at: https://derangedphysiology.com/main/cicm-primary-exam/required-reading/acid-base-physiology/Chapter%20412/traditional-and-physico-chemical-approaches-acid-base