Acid-Base Balance and Buffering

• 2025 is the year the learning objectives are current

General Terms

• It is important to define the acid-base state of an organism and regulate pH.
• Distinguish between volatile and nonvolatile acids and bases.
• Diet composition determines which acids and bases are generated.
• Define and apply partial pressure, equilibrium, buffering, and buffer value.
• Distinguish buffering from compensation.

Types of Buffers

• Identify the body fluid compartments in acid-base equilibrium and how they relate.
• Hemoglobin plays a role in buffering systemic pH.
• Understand the HCO3¯/CO2 system's operation and its pivotal role in acid-base balance.
• Learn the modification of the Henderson-Hasselbalch equation relating to body pH.
• Be aware of the role of other acid-base buffers in regulating systemic pH.

Davenport Diagram

• Become familiar with applying the acid-base map to conditions.
• Show how the body titrates acids and bases at a fixed pCO2 along the isobar.
• Distinguish how body fluids are titrated with carbonic acid vs. non-carbonic acid buffers.
• Identify and quantitate base deficit and base excess.

Acid Base and its Regulation

• The Davenport diagram is a four-dimensional acid-base map.
• The X-axis is pH.
• The Y-axis is HCO3.
• pCO2 is on the isobars (curved lines).
• The hemoglobin titration curve is the slanted line.
• The body has separate but interacting fluid compartments: bone, plasma, interstitial fluid, intracellular fluid, and RBCs.
• The final acid-base state is a composite of these compartments, which may change over time.
• The acid-base state of the body is usually taken to mean the state of arterial blood or extracellular fluid (ECF).
• The pH of ECF reflects the pH of other compartments, with the simplification that ECF or arterial blood pH reflects body pH.
• pH is important because the functional activities of proteins (enzymes, transport molecules, etc.) depend on it.

Acid Production

• The body produces acid through metabolism.
• CO2 metabolic production is approximately 13,000 mmol/day.
• CO2 is removed by exhaling 300 L CO2/day and is a Class 1 volatile acid.
• Oxidation of P and S in proteins and nucleotides forms H3PO4 and H2SO4 (~70-100 mmol/day).
• These are excreted by the kidney and are Class 2 nonvolatile acids.
• The body is constantly producing acid, with more volatile acid produced than nonvolatile acid.
• Class 1 acids are volatile (carbonic acid), and Class 2 acids are non-volatile (all other acids).
• Excess acid is removed by excretion.
• Carbonic acid becomes volatile CO2 and is excreted by the lungs.
• Non-carbonic acids are excreted by the kidneys.

The CO2/HCO3 System

• pCO2 is the partial pressure of CO2, and H2CO3 is an acid.
• The body is an open system.
• Volatile acid, CO2, is controlled by intake/output by the lungs.
• Nonvolatile acid excretion is controlled by the kidneys.
• pCO2 is regulated in body fluids because it measures dissolved CO2 concentration.
• CO2 in its hydrated form, H2CO3, is an acid, with its concentration regulated as part of the pH regulatory process.

H2CO3 Regulation

• CO2 + H2O ↔ H2CO3 ↔ H+ + HCO3-

• The rate of CO2 excretion is proportional to alveolar ventilation.

• Increased pCO2 leads to increased respiratory center activity, while decreased pH has the same effect.

• Increased alveolar ventilation leads to decreased alveolar pCO2.

• Alveolar pCO2 is in rapid equilibrium with arterial pCO2, also decreasing pH will increase toward normal.

• The body regulates carbonic and non-carbonic acid by controlling CO2 and nonvolatile acid intake/output.

• Increased pCO2 or decreased pH increases alveolar ventilation.

• Increased alveolar ventilation decreases alveolar pCO2.

• An increase in alveolar ventilation decreases alveolar pCO2, and decreases arterial pCO2 which inceases pH toward normal

• A decrease in alveolar ventilation increases pCO2

• Increased acid ingestion or metabolism increases H+ production.

• Increased acid secretion leads to increased acid excretion into urine.

• Non-volatile acid regulation is that increased H⁺ or decreased pH increases acid excretion by kidneys

• Non-volatile acid content is determined by the rate of acid intake, production by metabolism, and excretion of H+ by kidneys.

• An increase in volatile acid increases acid excretion by kidneys.

• The rate of acid secretion by the kidney is controlled by pCO2 and pH.

• Increased pCO2 or decreased pH increases non-volatile acid secretion/excretion by the kidneys.

Acid-Base Unbalance

• COPD (chronic obstructive pulmonary disease) is an impaired gas exchange in lungs; this leads to carbonic acid accumulation.
• Diabetic ketoacidosis occurs with the addition of β-hydroxybutyrate and acetoacetate leads to surplus of non-carbonic acid in the body.
• COPD causes a build-up of CO2 and carbonic acid.
• Uncontrolled diabetes mellitus with ketoacidosis causes addition of non-carbonic acids to the body.

Buffering

• Buffering is the property of a substance to reduce pH change from adding a strong acid or base.
• A weak acid or base buffers against changes in pH where acid and conjugate base concentrations are comparable (H+ + B- ↔ HB).

Definitions

• Strong electrolytes completely ionize in solution (e.g., NaCl, HCI).
• Weak electrolytes partially ionize in solution (e.g. lactic acid, ketoacids).
• Acid is a substance that can supply H+.
• Base is a substance that can take up H+ ions.
• Acetate is the conjugate base of acetic acid.

Dissociation Constant of a Weak Acid

• Weak acids or bases buffer against pH changes where the acid and conjugate base are comparable.
• A solution containing weak acid and its conjugate base acts as a buffer against added acid or base.
• Substances with a pK in the normal body pH range (7.40) possess a imidazole ring of the amino acid histidine or are phosphates.
• Imidazole in histidine has acidic and basic forms.
• The pKa of ≈7.0 is very close to the physiologic pH.

Titration Curve

• A titration curve plots the concentration of added strong acid/base against pH.
• Buffer value is the slope of the titration curve. It is used to calculate the titration curve for Imidazole.
• The curve is steepest at pH = pK.
• Little change in pH occurs for each mEq of added acid/base when pH is closest to pK.
• Titration curve is approximately linear between pH 6.5 and 7.5.
• Body fluid pH is accurately approximated by straight lines.

Titration With Imidazole

• If 4 mmol of HCl is added to 1 liter of imidazole solution, all added base combines with HIm+ and converts to Im+ water.
• Since the initial solution contains 10 mmol Im and 10 mmol HImt, basic form Im will decrease by 4 mmol to 6 mmol and the conjugate acid HImt will increase by 4 mmol to 14 mmol.
• [H+] concentration can be estimated from the table to be 2.33 nEq, giving a pH = 6.63.
• If 4 mmol of NaOH is added to 1 liter of the imidazole solution, it converts to Im+ water then HImt decreases by 4 mmol to 6 mmol and the conjugate base Im will increase by 4 mmol to 14 mmol.
• [H+] concentration can be estimated from the table to be 42.9 nEq, giving a pH = 7.37.

Body Acid-Base Review

• Large amounts of acid/base can be added to the body without lethal pH changes due to buffer substances.
• Most buffers are located intracellularly; this process is relatively slow.
• Intracellular buffers mainly consist of proteins and phosphates.
• Hemoglobin in red cells buffers acids/bases added to ECF and has a high histidine content: 36 histidines per hemoglobin tetramer. Histidine has the imidazole group and accounts for 90% of the blood's buffer value.

Hemoglobin as Buffer

• High histidine content = 36 histidines in a tetramer

• Most CO2 carried in blood is as HCO3 with

• Most of the H+ is in the form of HImt in RBCS

• All measurements of pH in blood are corrected to 100% saturation of Hgb with O2

• ECF also includes buffers such as plasma proteins and phosphates.

• Bone is a large sink for acids/bases, is a very slow buffer, and participates in chronic acidosis conditions.

• Most CO2 carried by the blood is in the form of HCO3¯.

• The H+ ions are carried in the form of HImt by histidines in hemoglobin.

• Extracellular blood pH reflects intracellular pH.

• Blood pH reflects pH in all parts of the body.

• The HCO3-/CO2 buffer system is very important.

• The pK=6.1 (not that close to physiologic).

• Dual control comes from the lungs and the kidneys.

CO2 System

• Concentration of dissolved CO2 is [CO2]dis which is 0.03 mM/mmHg * pCO2
• pCO2 = 40 mmHg
• Dissolved CO2: all gases dissolve to some extent in water.
• Plasma [CO2]dis is proportional to the plasma pCO2
• Arterial blood is normally equilibrated with alveolar air.
• pCO2 is equal to the partial pressure of CO2 in alveolar air, normally 40 mmHg.
• [CO2]dis = 0.03 * 40 = 1.2 mM
• CO2 + H2O ↔ H+ + HCO3-

Henderson-Hasselbach Equation

• Doubling pCO2 form 40 to 80 mmHg increase the [HCO3¯] from 24 to 27 mEq, which is 3 mEq, and lowers the pH to 7.15.
• Henderson Hasselbach DOES NOT tell us the acid base disorder and is ALWAYS obeyed!

Buffering in Body – Using the Davenport Diagram

• The Davenport diagram is the section in the box of the titration curve for imidazole and can be used to assess buffering.
• The curve is steepest at neighborhood pH = pK.
• There is a little change in pH for each mEq of added acid/base when the pH is the closest to the pK.
• The titration curve is approximately linear between pH 6.5 and 7.5.
• The range of body fluid pH values compatible with life is less than one pH unit.
• Titration curves for body fluids are accurately approximated by straight lines.
• The HCO3 concentration is given on the Y-axis.
• The pH is given on the X-axis.
• Isobars are the curved lines, and they are constructed by drawing smooth curves through the points where pCO2 is constant.
• For measured pH, pCO2 and [HCO3¯], the third is calculated.
• Titration curves at fixed pCO2 is the pCO2 isobars and it is when Carbonic Acid/Bicarbonate Solution is added to fixed pCO2 values.
• Strong non-carbonic acid/base incrementally titrates to the solution.

Titration at Fixed pCO2

• Point O is normal for sea level humans = 40 mmHg
• The pCO2 is held constant at 40 mmHg
• Added HCl is buffered by NaHCO3.
• HCO3- decreases from 24 to 19 mM
• When very nearly all added HCl combines with NaHCO3 it forms CO2, H2O, and NaCl
• The pH drops to 7.3
• [HCO3] decreases by 5 mEq to a final value of 19 mM (mEq). The pH calculated from Henderson-Hasselbach is now 7.30. Point A is reached by adding 5 mEq of strong acid (HCI) to each liter of solution with the CO2 held constant at 40 mmHg.

Titration of Body Fluid With H2CO3

• Where Hgb-Im represents titratable protein buffers and Hgb is on histidine with mostly imidazole.
• H2CO3 + Hgb-Im ↔ Hgb-Hlm+ + HCO3¯
• For each H+ buffered by Hab-Im a HCO3 ion appears
• = concentration of added H2CO3 gives The increase in [HCO3]
• The reaction of carbonic acid with buffers, the is When increasing dissolved CO2 gives increasing pCO2.
• The main buffering reaction, H2CO3 + Hgb-Im↔ Hgb-HImt + HCO3¯

Body Fluids of Non-Carbonic Acids

• The addition of strong non-carbonic acids/HCI is the combination of all added H+ ions with buffers.
• Acid in body fluids is buffered by the Im form of imidazole and the HCO3¯ to form carbonic acid
• Added carbonic acid is buffered by bicarbonate.
• Therefore, the change in [HCO3] does NOT represent the total amount of added acid.

Titration with Non-Carbonic Acid: Base Deficit

• Determination of the amount of added non-carbonic acid or base
• Point O the the initial condition where point A the condition after is added an unknown non-carbonic
• Most of the H+ ions combine with HCO3¯ to form CO2 and H2O.
• The remainder combine with Hgb and plasma proteins.
• 1. L1 is the amount of HCl buffered by Hgb (HIm) and plasma proteins
• the A HCO3 of 2. L2 is the amount buffered by carbonic acid, from normal [HCO3] .
• 2 + L1 = base deficit
• The addition of a small amount of strong non-carbonic acid such as HCI to a buffered solution results in the combination of virtually all added H+ ions with the buffers.
• Acid is buffered by the im form of imidazole and the HCO3¯ to form carbonic acid
• Non carbonic acid is buffered by BOTH bicarbonate and hemoglobin
• Therefore, the change in HCO3 does NOT represent the total. amount added acid

Acid Base Compensation

• Acid base disturbances are classified first by pH:

• Acidemia – condition where pH < 7.40 (normal at sea level)

• Alkalemia - condietion where pH > 7.40 (normal at sea level)

• Then, sub-classified based on changes in carbonic acid or non-carbonic acid content:

• Respiratory imbalance – changes in pCO2/carbonic acid due to changes in ventilation

• ▲ in volatile acid: pCO2 = measure of carbonic acid unbalance • pCO2 excess = respiratory acidosis • pCO2 deficit = respiratory alkalosis

• Metabolic imbalance - changes in non-carbonic acid with base excess/deficit • ▲ in nonvolatile acid • Base deficit = metabolic acidosis • Base excess = metabolic alkalosis

• If an arterial blood sample has abnormal pH, pCO2 and/or [HCO3-] values, then there is an acid base imbalance

• The fact that the values are abnormal indicates a disturbance, but what is the SIZE of the disturbance?

• the value of pCO2 is a good measure of the size of a carbonic acid imbalance but

• [HCO3] is not a direct measure of the size of a non-carbonic acid imbalance

• need to evaluate the base excess/base deficit

• Normal Range: Acid-Base Values BOLD hexagon delineates the ranges of pH, pCO2 and [HCO3-] values in normal humans living near sea level with a body temperature of 37°C pH ≈ 7.35 to 7.45 (really more closely regulated to within 7.38 - 7.42 than can be shown on the graph easily) p CO2 ≈ 35 - 45 mmHg (as for pH, really more closely regulated to 37 - 42 mmHg) [HCO3-] ≈ 22 - 26 mEq

• Normal values at sea level and 37°C are pH 7.40, HCO3- 24 mM, pCO2 40 mmHg

• Obtain sample of arterial blood

• Measure pH, pCO2 and [HCO3-]

• Plot these values on the Davenport diagram

• Make a diagnosis based on the position of the piont

• If at all possible getting relevant history about the patient Focus on single acid-base disorder

Factors Affecting HCO3 Reabsorption

• It will be good to review the factors affecting reabsorption at that time
• CO2 + H2O ↔ H2CO3 ↔ H+ + HCO3-
• Carbonic anhydrase (CA) activity: ↓carbonic anhydrase activity→
• ↓H+ secretion →↓ reabsorption of HCO3: Partial pressure (PCO₂): ↓PCO2→↓ HCO3- reabsorption
• Amount filtered HCO3↑ ↑HCO3filtered→HCO3 reabsorption
• Amounts of filtered/secreted buffers (distal): and that be protonated to and H2PO

Compensation

• Respiratory Acidosis (Hypercapnia): Carbonic Acid Excess
• [H+] * [HCO3- ] = 0.03 Ka pCO2
• Normal
• Uncompensated
• Compensated (physiologic response)
• A disorder is when high and low pH is move to higher isobar.

Respiratory Alkalosis

• Compensation is when kidney increases excretion will manifest as further increase in along the same isobar/Renal compensated

Malfunctions Causing Respiratory Acidosis (Hypercapnia)

• Pulmonary disease: Asthma, COPD
• Neuromuscular disorders: Myasthenia gravis, Poliomyelitis,Spinal cord lesions
• Restricted thoracic movement: Obesity, Arthritis
• Chemoreceptor depressants: Anesthetics, Sedatives, Narcotics

Acute vs Chronic Respiratory Acidosis (or Hypercapnia)

• [H+] = 0.03 Ka
• Acutely the pCO2 increased and over time the compensatory physiologic response

Respiratory Alkalosis: Carbonic Acid Deficit

• is a deficit carbonic acid (content- the elevated
• The decrease CO2 decreased
• Cells increase production
• The kidney + (acid), -is

Malfunctions Causing Respiratory Alkalosis (Hypocapnia)

• The primary problem is a deficit carbonic acid
• Initial/ Low
• Decreases the + + Cell will ncrease production thekidney then

Metabolic Acidosis

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Malfunctions Causing Metabolic Alkalosis vomiting

Overview/Summary

• primary = metabolic acidosis the pCO2 high then with primary/