Regulation of hydrogen ion (H⁺) balance is similar to the regulation of other body ions.
There must be a balance between H⁺ intake or production and H⁺ removal.
This balance is necessary to maintain homeostasis.
The kidneys play a major role in removing H⁺ from the body.
Regulation of extracellular fluid H⁺ concentration requires more than kidney excretion alone.
Acid–base buffering mechanisms are also essential.
The blood contains important buffering mechanisms.
Body cells contain important buffering mechanisms.
The lungs also help regulate H⁺ concentration.
These mechanisms maintain normal H⁺ concentration in extracellular fluid.
These mechanisms also maintain normal H⁺ concentration in intracellular fluid.
This chapter explains the mechanisms that regulate H⁺ concentration.
Special emphasis is given to renal H⁺ secretion.
Special emphasis is given to renal H⁺ reabsorption.
Special emphasis is given to renal production of HCO₃⁻.
Special emphasis is given to renal excretion of HCO₃⁻.
HCO₃⁻ is a key component of the body’s acid–base control system.

HYDROGEN ION CONCENTRATION IS PRECISELY REGULATED

Precise regulation of H⁺ concentration is essential.
Almost all enzyme systems are affected by H⁺ concentration.
Changes in H⁺ concentration alter enzyme activity.
Changes in H⁺ concentration affect almost all cell functions.
Changes in H⁺ concentration affect almost all body functions.
H⁺ concentration in body fluids is normally very low.
Extracellular fluid sodium concentration is 142 mEq/L.
Normal H⁺ concentration is only 0.00004 mEq/L.
Sodium concentration is about 3.5 million times greater than H⁺ concentration.
Normal variation in H⁺ concentration is extremely small.
The normal variation in H⁺ concentration is only about one-millionth of the normal variation in sodium (Na⁺) concentration.
This precise regulation shows the importance of H⁺ concentration for normal cellular function.

KEY CONCEPT

H⁺ balance depends on a balance between H⁺ production/intake and H⁺ removal.
The kidneys, blood, cells, and lungs work together to regulate H⁺ concentration.
The kidneys regulate H⁺ secretion and HCO₃⁻ reabsorption, production, and excretion.
H⁺ concentration strongly affects enzyme activity and cell function.
Normal Na⁺ concentration = 142 mEq/L.
Normal H⁺ concentration = 0.00004 mEq/L.
Na⁺ concentration is about 3.5 million times greater than H⁺ concentration.

ACIDS AND BASES—DEFINITIONS AND MEANINGS

A hydrogen ion (H⁺) is a single free proton released from a hydrogen atom.
Molecules that release H⁺ in solution are called acids.
Hydrochloric acid (HCl) is an example of an acid.
HCl ionizes in water.
HCl forms H⁺ and Cl⁻ ions.
Carbonic acid (H₂CO₃) is also an acid.
H₂CO₃ ionizes in water.
H₂CO₃ forms H⁺ and HCO₃⁻.
A base is an ion or molecule that accepts H⁺.
HCO₃⁻ is a base.
HCO₃⁻ combines with H⁺ to form H₂CO₃.
HPO₄²⁻ is also a base.
HPO₄²⁻ accepts H⁺ to form H₂PO₄⁻.
Body proteins also act as bases.
Some amino acids in proteins carry negative charges.
These negative charges readily accept H⁺.
Hemoglobin in red blood cells is an important body base.
Proteins in other body cells are also important bases.
The terms base and alkali are often used interchangeably.
An alkali is formed by combining an alkaline metal with a highly basic ion.
Examples of alkaline metals are sodium, potassium, and lithium.
A common highly basic ion is OH⁻.
The base portion quickly reacts with H⁺.
This reaction removes H⁺ from the solution.
Therefore, alkalis are typical bases.
Alkalosis means excessive removal of H⁺ from body fluids.
Acidosis means excessive addition of H⁺ to body fluids.

Strong and Weak Acids and Bases

Strong acids dissociate rapidly.
Strong acids release large amounts of H⁺.
HCl is a strong acid.
Weak acids dissociate less readily.
Weak acids release H⁺ more slowly.
H₂CO₃ is a weak acid.
Strong bases react rapidly with H⁺.
Strong bases quickly remove H⁺ from solution.
OH⁻ is a strong base.
OH⁻ combines with H⁺ to form H₂O.
HCO₃⁻ is a weak base.
HCO₃⁻ binds H⁺ less strongly than OH⁻.
Most acids and bases in extracellular fluid are weak.
These weak acids and bases regulate normal acid–base balance.
The most important buffer pair is H₂CO₃ / HCO₃⁻.

Normal H⁺ Concentration and pH of Body Fluids and Changes That Occur in Acidosis and Alkalosis

Normal blood H⁺ concentration is about 0.00004 mEq/L.
This is equal to 40 nEq/L.
Normal variation is only 3–5 nEq/L.
Under extreme conditions, H⁺ concentration may fall to 10 nEq/L.
Under extreme conditions, H⁺ concentration may rise to 160 nEq/L.
These extreme values may still be compatible with life.
H⁺ balance depends on a balance between H⁺ production or intake and H⁺ removal.
The kidneys play a major role in removing H⁺.
Kidney excretion alone is not enough to regulate H⁺.
Blood buffering systems also regulate H⁺.
Cells also regulate H⁺.
The lungs also regulate H⁺.
These systems maintain normal extracellular H⁺ concentration.
These systems also maintain normal intracellular H⁺ concentration.
This chapter focuses on H⁺ regulation.
It emphasizes renal H⁺ secretion.
It emphasizes renal H⁺ reabsorption.
It emphasizes renal production of HCO₃⁻.
It emphasizes renal excretion of HCO₃⁻.
HCO₃⁻ is a major component of acid–base regulation.

HYDROGEN ION CONCENTRATION IS PRECISELY REGULATED

Precise regulation of H⁺ is essential.
Almost all enzyme systems are affected by H⁺ concentration.
Changes in H⁺ concentration alter enzyme activity.
Changes in H⁺ concentration affect almost all cell functions.
Changes in H⁺ concentration affect almost all body functions.
H⁺ concentration in body fluids is normally very low.
Normal extracellular sodium concentration is 142 mEq/L.
Normal H⁺ concentration is 0.00004 mEq/L.
Sodium concentration is about 3.5 million times greater than H⁺ concentration.
Normal variation in H⁺ concentration is only about one-millionth of the normal variation in Na⁺ concentration.
This precise regulation highlights the importance of H⁺ in cell function.

ACIDS AND BASES—DEFINITIONS AND MEANINGS

Because H⁺ concentration is very low, pH is used to express it.
pH uses a logarithmic scale.

Formula

$\boxed{\text{pH}=\log\left(\frac{1}{[H^+]}\right)=-\log[H^+]}$ pH=log([H+]1)=−log[H+]

H⁺ concentration is expressed in equivalents per liter (Eq/L).

Mathematical Calculation

Given:

H⁺ = 40 nEq/L
40 nEq/L = 0.00000004 Eq/L = 4 × 10⁻⁸ Eq/L

Calculation: $\text{pH}=-\log(4\times10^{-8})$ pH=−log(4×10−8) $=-[\log4+\log10^{-8}]$ =−[log4+log10−8] $=-(0.6-8)$ =−(0.6−8) $=7.4$ =7.4

Final Answer

Normal pH = 7.4
pH is inversely related to H⁺ concentration.
Low pH means high H⁺ concentration.
High pH means low H⁺ concentration.
Normal arterial blood pH is 7.4.
Venous blood pH is about 7.35.
Interstitial fluid pH is about 7.35.
Venous blood contains more CO₂ than arterial blood.
Extra CO₂ forms more H₂CO₃.
More H₂CO₃ slightly lowers pH.
These values are summarized in Table 31.1.
Acidemia occurs when arterial blood pH falls significantly below 7.4.
Alkalemia occurs when arterial blood pH rises above 7.4.
The lower limit of survival is about pH 6.8.
The upper limit of survival is about pH 8.0.
Intracellular pH is usually lower than plasma pH.
Cell metabolism produces acids.
H₂CO₃ is an important intracellular acid.
Intracellular pH ranges from 6.0 to 7.4.
Hypoxia causes acid accumulation.
Poor tissue blood flow causes acid accumulation.
These conditions decrease intracellular pH.
Acidosis is the process that leads to acidemia.
Alkalosis is the process that leads to alkalemia.
Urine pH ranges from 4.5 to 8.0.
Urine pH depends on extracellular acid–base status.
The kidneys correct abnormal H⁺ concentration.
The kidneys excrete acids or bases as needed.
Gastric HCl is an example of an extremely acidic body fluid.
HCl is secreted by oxyntic (parietal) cells of the stomach.
This process is discussed in Chapter 65.
H⁺ concentration in these cells is about 4 million times greater than in blood.
Gastric HCl has a pH of 0.8.
The remainder of the chapter discusses regulation of extracellular H⁺ concentration.

KEY CONCEPT

Acids donate H⁺; bases accept H⁺.
Strong acids/bases dissociate rapidly; weak acids/bases dissociate slowly.
The most important extracellular buffer system is H₂CO₃ / HCO₃⁻.
Normal H⁺ = 40 nEq/L (0.00004 mEq/L).
Normal arterial pH = 7.4.
Venous/interstitial pH = 7.35.
Intracellular pH = 6.0–7.4.
Urine pH = 4.5–8.0.
Survival pH range = 6.8–8.0.
pH = −log[H⁺].
Calculated pH = 7.4.
Table Mentioned: Table 31.1.

DEFENDING AGAINST CHANGES IN H⁺ CONCENTRATION: BUFFERS, LUNGS, AND KIDNEYS

Three primary systems regulate H⁺ concentration in body fluids.
The first system is the chemical acid–base buffer system.
Buffer systems immediately combine with acids or bases.
Buffer systems prevent excessive changes in H⁺ concentration.
The second system is the respiratory center.
The respiratory center regulates the removal of CO₂.
Removal of CO₂ also removes H₂CO₃ from the extracellular fluid.
The third system is the kidneys.
The kidneys can excrete acidic urine.
The kidneys can excrete alkaline urine.
The kidneys return extracellular H⁺ concentration toward normal during acidosis.
The kidneys return extracellular H⁺ concentration toward normal during alkalosis.
When H⁺ concentration changes, buffer systems respond within seconds.
Buffer systems minimize changes in H⁺ concentration.
Buffer systems do not remove H⁺ from the body.
Buffer systems do not add H⁺ to the body.
Buffer systems temporarily bind H⁺ until balance is restored.
The respiratory system is the second line of defense.
The respiratory system responds within a few minutes.
It removes CO₂ from the body.
Removal of CO₂ also removes H₂CO₃.
The first two defense systems limit changes in H⁺ concentration.
They protect the body until the kidneys respond.
The kidneys are the third line of defense.
The kidneys remove excess acid from the body.
The kidneys remove excess base from the body.
The kidneys respond more slowly than buffers and lungs.
Kidney responses take several hours to several days.
The kidneys are the most powerful acid–base regulatory system.

BUFFERING OF H⁺ IN THE BODY FLUIDS

A buffer is any substance that can reversibly bind H⁺.

Buffer Reaction

$\boxed{\text{Buffer} + \text{H}^+ \;\rightleftharpoons\; \text{H-Buffer}}$ Buffer+H+⇌H-Buffer

A free H⁺ combines with the buffer.
This forms a weak acid called H-Buffer.
H-Buffer may remain as an undissociated molecule.
H-Buffer may dissociate back into Buffer and H⁺.
When H⁺ concentration increases, the reaction moves to the right.
More H⁺ binds to the buffer.
This continues as long as buffer is available.
When H⁺ concentration decreases, the reaction moves to the left.
H⁺ is released from the buffer.
These reactions minimize changes in H⁺ concentration.
Body fluid buffers are very important.
H⁺ concentration in body fluids is normally very low.
The body produces relatively large amounts of acid every day.
About 80 mEq of H⁺ is ingested or produced daily by metabolism.
Normal H⁺ concentration in body fluids is only about 0.00004 mEq/L.
Without buffers, daily acid production would rapidly cause life-threatening changes in H⁺ concentration.
The bicarbonate buffer system is the most important extracellular buffer.
The bicarbonate buffer system is explained next.

KEY CONCEPT

The body has three lines of defense against changes in H⁺ concentration:
- 1st: Chemical buffer systems (seconds).
- 2nd: Respiratory system (minutes).
- 3rd: Kidneys (hours to days, most powerful).
Buffers temporarily bind or release H⁺ to minimize changes in H⁺ concentration.
Buffer Reaction: Buffer + H⁺ ⇌ H-Buffer
When H⁺ increases → Reaction shifts right → More H⁺ binds to buffer.
When H⁺ decreases → Reaction shifts left → Buffer releases H⁺.
The body produces about 80 mEq of H⁺ per day, while normal H⁺ concentration is only 0.00004 mEq/L.
The bicarbonate buffer system is the most important extracellular buffer.
Figure Mentioned: None in the provided text.
Table Mentioned: Table 31.1
Mathematical Equations:
- Buffer + H⁺ ⇌ H-Buffer (Buffering reaction)

BICARBONATE BUFFER SYSTEM

The bicarbonate buffer system contains two components.
Component 1: A weak acid (H₂CO₃).
Component 2: A bicarbonate salt, usually NaHCO₃.

Formation of Carbonic Acid (H₂CO₃)

Biochemical Equation $\boxed{\mathrm{CO_2 + H_2O \xrightleftharpoons[\;]{Carbonic\ Anhydrase} H_2CO_3}}$ CO2+H2OCarbonic AnhydraseH2CO3

Easy Concept

CO₂ combines with H₂O.
Carbonic anhydrase speeds up this reaction.
H₂CO₃ (carbonic acid) is formed.
Without carbonic anhydrase, this reaction is very slow.
Only a very small amount of H₂CO₃ is formed without the enzyme.
Carbonic anhydrase is abundant in the walls of the lung alveoli.
In the lungs, CO₂ is released.
Carbonic anhydrase is also present in renal tubular epithelial cells.
In the kidneys, CO₂ reacts with H₂O to form H₂CO₃.

Ionization of Carbonic Acid

Biochemical Equation $\boxed{\mathrm{H_2CO_3 \xrightleftharpoons[]{} H^+ + HCO_3^-}}$ H2CO3H++HCO3−

Easy Concept

Carbonic acid breaks into:
- H⁺
- HCO₃⁻
This ionization is weak.
Therefore, only a small amount of H⁺ is produced.
H₂CO₃ ionizes weakly.
Small amounts of H⁺ are formed.
Small amounts of HCO₃⁻ are formed.

Ionization of Sodium Bicarbonate

Biochemical Equation $\boxed{\mathrm{NaHCO_3 \xrightleftharpoons[]{} Na^+ + HCO_3^-}}$ NaHCO3Na++HCO3−

Easy Concept

NaHCO₃ separates into:
- Na⁺
- HCO₃⁻
This ionization is almost complete.
NaHCO₃ is the main bicarbonate salt in extracellular fluid.
It dissociates almost completely.
It produces Na⁺.
It produces HCO₃⁻.

Complete Bicarbonate Buffer System

Biochemical Equation $\boxed{\mathrm{CO_2 + H_2O \xrightleftharpoons[]{} H_2CO_3 \xrightleftharpoons[]{} H^+ + HCO_3^-}}$ CO2+H2OH2CO3H++HCO3−

Easy Concept

Step 1: CO₂ + H₂O → H₂CO₃
Step 2: H₂CO₃ → H⁺ + HCO₃⁻
H₂CO₃ dissociates only slightly.
Therefore, H⁺ concentration remains very low.
Because H₂CO₃ dissociates weakly, H⁺ concentration stays extremely low.

Addition of a Strong Acid (HCl)

Biochemical Equation $\boxed{\mathrm{HCl \rightarrow H^+ + Cl^-}}$ HCl→H++Cl−

Easy Concept

HCl releases a large amount of H⁺.

Buffer Reaction $\boxed{\mathrm{H^+ + HCO_3^- \rightarrow H_2CO_3 \rightarrow CO_2 + H_2O}}$ H++HCO3−→H2CO3→CO2+H2O

Easy Concept

Step 1: HCl releases H⁺.
Step 2: HCO₃⁻ immediately binds H⁺.
Step 3: H₂CO₃ is formed.
Step 4: H₂CO₃ breaks into CO₂ and H₂O.
Step 5: CO₂ is removed by the lungs.
H⁺ released from HCl is buffered by HCO₃⁻.
More H₂CO₃ is formed.
More CO₂ is produced.
More H₂O is produced.
The extra CO₂ strongly stimulates respiration.
Increased respiration removes CO₂ from the extracellular fluid.

Addition of a Strong Base (NaOH)

Biochemical Equation $\boxed{\mathrm{NaOH + H_2CO_3 \rightarrow NaHCO_3 + H_2O}}$ NaOH+H2CO3→NaHCO3+H2O

Easy Concept

Step 1: NaOH provides OH⁻.
Step 2: OH⁻ combines with H₂CO₃.
Step 3: NaHCO₃ and H₂O are formed.
Step 4: The strong base becomes a weak base.
OH⁻ combines with H₂CO₃.
More HCO₃⁻ is formed.
NaHCO₃ replaces the strong base NaOH.
H₂CO₃ concentration decreases.

Replacement of Carbonic Acid

Biochemical Equation $\boxed{\mathrm{CO_2 + H_2O \rightarrow H_2CO_3}}$ CO2+H2O→H2CO3

Easy Concept

As H₂CO₃ decreases,
More CO₂ combines with H₂O.
New H₂CO₃ is formed.
More CO₂ reacts with H₂O to replace H₂CO₃.
Blood CO₂ tends to decrease.
Low blood CO₂ inhibits respiration.
CO₂ expiration decreases.
Blood HCO₃⁻ concentration increases.
The kidneys excrete more HCO₃⁻.

Quantitative Dynamics of the Bicarbonate Buffer System

All acids ionize to some extent.
H₂CO₃ also ionizes.

Equation 1

$\boxed{K’=\frac{[H^+]\times[HCO_3^-]}{[H_2CO_3]}}$ K′=[H2CO3][H+]×[HCO3−]

Easy Concept

K′ is the dissociation constant.
It relates H⁺, HCO₃⁻, and H₂CO₃ concentrations.

Equation 2

$\boxed{[H^+]=K’\times\frac{[H_2CO_3]}{[HCO_3^-]}}$ [H+]=K′×[HCO3−][H2CO3]

Easy Concept

H⁺ concentration depends on:
- H₂CO₃ concentration.
- HCO₃⁻ concentration.
More H₂CO₃ → More H⁺.
More HCO₃⁻ → Less H⁺.
H₂CO₃ concentration cannot be measured directly.
H₂CO₃ rapidly changes into CO₂ and H₂O.
H₂CO₃ also rapidly dissociates into H⁺ and HCO₃⁻.
Dissolved CO₂ is directly proportional to H₂CO₃.

Equation 3

$\boxed{[H^+]=K\times\frac{CO_2}{[HCO_3^-]}}$ [H+]=K×[HCO3−]CO2

Easy Concept

H₂CO₃ is replaced by dissolved CO₂.
H⁺ depends on:
- Dissolved CO₂.
- HCO₃⁻ concentration.
The dissociation constant K is about 1/400 of K′.
This is because the H₂CO₃ : CO₂ ratio is 1 : 400.
Clinical laboratories usually measure PCO₂ instead of dissolved CO₂.
Dissolved CO₂ is proportional to PCO₂.
The solubility coefficient of CO₂ is 0.03 mmol/L/mm Hg.
This value applies at body temperature.
Each 1 mm Hg PCO₂ corresponds to 0.03 mmol/L CO₂.

Equation 4

$\boxed{[H^+]=K\times\frac{(0.03\times PCO_2)}{[HCO_3^-]}}$ [H+]=K×[HCO3−](0.03×PCO2)

Easy Concept

H⁺ concentration depends on:
- PCO₂ (controlled by lungs).
- HCO₃⁻ (controlled by kidneys).

KEY CONCEPT

The bicarbonate buffer system contains:
- Weak acid: H₂CO₃
- Weak base: NaHCO₃ (HCO₃⁻)
Carbonic anhydrase rapidly converts CO₂ + H₂O into H₂CO₃.
H₂CO₃ ⇌ H⁺ + HCO₃⁻.
Strong acid (HCl) is buffered by HCO₃⁻, producing CO₂ + H₂O.
Strong base (NaOH) reacts with H₂CO₃, forming NaHCO₃ + H₂O.
The lungs regulate PCO₂.
The kidneys regulate HCO₃⁻.
H⁺ concentration depends on the ratio of CO₂ to HCO₃⁻.
Mathematical/Biochemical Equations Solved:
- CO₂ + H₂O ⇌ H₂CO₃
- H₂CO₃ ⇌ H⁺ + HCO₃⁻
- NaHCO₃ ⇌ Na⁺ + HCO₃⁻
- H⁺ + HCO₃⁻ → H₂CO₃ → CO₂ + H₂O
- NaOH + H₂CO₃ → NaHCO₃ + H₂O
- $K’=\frac{[H^+][HCO_3^-]}{[H_2CO_3]}$ K′=[H2CO3][H+][HCO3−]
- $[H^+]=K’\times\frac{[H_2CO_3]}{[HCO_3^-]}$ [H+]=K′×[HCO3−][H2CO3]
- $[H^+]=K\times\frac{CO_2}{[HCO_3^-]}$ [H+]=K×[HCO3−]CO2
- $[H^+]=K\times\frac{0.03\times PCO_2}{[HCO_3^-]}$ [H+]=K×[HCO3−]0.03×PCO2

Mathematical/Biochemical Equations Solved (SUPERFAST SIMPLIFIED)

1. Formation of Carbonic Acid

Equation

$\boxed{\mathrm{CO_2 + H_2O \xrightleftharpoons[\;]{Carbonic\ Anhydrase} H_2CO_3}}$ CO2+H2OCarbonic AnhydraseH2CO3

Easiest Understanding

Think of it as:

➡️ CO₂ + Water = Carbonic Acid

Step-by-Step

CO₂ enters the blood.
CO₂ meets water (H₂O).
Carbonic anhydrase makes the reaction very fast.
Carbonic acid (H₂CO₃) is produced.

Memory Trick

CO₂ + Water = Carbonic Acid2. Breakdown of Carbonic Acid

Equation

$\boxed{\mathrm{H_2CO_3 \xrightleftharpoons{} H^+ + HCO_3^-}}$ H2CO3H++HCO3−

Easiest Understanding

Carbonic acid breaks into two pieces:

✅ Hydrogen ion (H⁺)

✅ Bicarbonate ion (HCO₃⁻)

Memory Trick

Carbonic Acid → Acid (H⁺) + Buffer (HCO₃⁻)

3. Sodium Bicarbonate Dissociation

Equation

$\boxed{\mathrm{NaHCO_3 \xrightleftharpoons{} Na^+ + HCO_3^-}}$ NaHCO3Na++HCO3−

Easiest Understanding

Sodium bicarbonate separates into:

Sodium (Na⁺)
Bicarbonate (HCO₃⁻)

Memory Trick

NaHCO₃ = Sodium + Bicarbonate

4. What Happens When a Strong Acid (HCl) Enters Blood?

Step 1

Strong acid releases H⁺ $\boxed{\mathrm{HCl \rightarrow H^+ + Cl^-}}$ HCl→H++Cl−

↓

Step 2

Buffer immediately catches H⁺ $\boxed{\mathrm{H^+ + HCO_3^- \rightarrow H_2CO_3}}$ H++HCO3−→H2CO3

↓

Step 3

Carbonic acid breaks $\boxed{\mathrm{H_2CO_3 \rightarrow CO_2 + H_2O}}$ H2CO3→CO2+H2O

↓

Step 4

Lungs remove CO₂ $\boxed{\mathrm{CO_2 \uparrow}}$ CO2↑

Whole Story

Strong Acid

↓

Releases H⁺

↓

HCO₃⁻ catches H⁺

↓

Makes H₂CO₃

↓

Breaks into CO₂ + H₂O

↓

Lungs remove CO₂

↓

Blood becomes normal again

Memory Trick

Acid → HCO₃⁻ → H₂CO₃ → CO₂ → Lungs

5. What Happens When a Strong Base (NaOH) Enters Blood?

Step 1

NaOH releases OH⁻

↓

Step 2

OH⁻ attacks carbonic acid $\boxed{\mathrm{NaOH + H_2CO_3 \rightarrow NaHCO_3 + H_2O}}$ NaOH+H2CO3→NaHCO3+H2O

↓

Step 3

Carbonic acid decreases

↓

Step 4

CO₂ + Water make more H₂CO₃ $\boxed{\mathrm{CO_2 + H_2O \rightarrow H_2CO_3}}$ CO2+H2O→H2CO3

↓

Step 5

Respiration slows

↓

CO₂ is retained

↓

Carbonic acid returns to normal

Whole Story

Strong Base

↓

Uses H₂CO₃

↓

More CO₂ is saved

↓

More H₂CO₃ forms

↓

Blood becomes normal again

Memory Trick

Base → Uses H₂CO₃ → CO₂ Saved → H₂CO₃ Restored

Quantitative Equations (Super Easy)

Equation 1

$\boxed{K’=\frac{[H^+]\times[HCO_3^-]}{[H_2CO_3]}}$ K′=[H2CO3][H+]×[HCO3−]

Meaning

This equation tells us:

How much H⁺ is present compared with carbonic acid.

Think of it as

Acid Strength Formula

Equation 2

$\boxed{[H^+]=K’\times\frac{[H_2CO_3]}{[HCO_3^-]}}$ [H+]=K′×[HCO3−][H2CO3]

Easiest Meaning

Hydrogen ions depend on two things

Numerator

H₂CO₃

⬆ More Carbonic Acid

⬇

More H⁺

Denominator

HCO₃⁻

⬆ More Bicarbonate

⬇

Less H⁺

Easy Rule

$\boxed{\uparrow H_2CO_3=\uparrow H^+}$ ↑H2CO3=↑H+ $\boxed{\uparrow HCO_3^-=\downarrow H^+}$ ↑HCO3−=↓H+

Equation 3

Since H₂CO₃ is difficult to measure,

Guyton replaces it with CO₂ $\boxed{[H^+]=K\times\frac{CO_2}{HCO_3^-}}$ [H+]=K×HCO3−CO2

Easiest Meaning

Hydrogen ions depend on

CO₂

divided by

Bicarbonate

Easy Rule

More CO₂

↓

More H⁺

↓

More Acidic

More HCO₃⁻

↓

Less H⁺

↓

More Alkaline

Equation 4

Clinically we measure PCO₂, not dissolved CO₂.

CO₂ dissolved in blood

0.03 × PCO₂

Therefore $\boxed{[H^+]=K\times\frac{0.03\times PCO_2}{HCO_3^-}}$ [H+]=K×HCO3−0.03×PCO2

Easiest Understanding

Blood acidity depends on only TWO things

① Lungs

Measure

PCO₂

↑ PCO₂

↓

↑ H⁺

↓

Acidosis

② Kidneys

Control

HCO₃⁻

↑ HCO₃⁻

↓

↓ H⁺

↓

Alkalosis

One-Line Memory Formula

\boxed{\textbf{Acidity=\frac{CO_2}{HCO_3^-}}}

or\boxed{\textbf{H^+\propto\frac{CO_2}{HCO_3^-}}}

Super Memory Flow Chart

CO₂ + H₂O
      │
      ▼
   H₂CO₃
      │
      ▼
 H⁺ + HCO₃⁻

If Acid Comes

H⁺
 │
 ▼
HCO₃⁻ catches it
 │
 ▼
H₂CO₃
 │
 ▼
CO₂ + H₂O
 │
 ▼
Lungs remove CO₂

If Base Comes

OH⁻
 │
 ▼
Uses H₂CO₃
 │
 ▼
CO₂ combines with H₂O
 │
 ▼
New H₂CO₃ formed
 │
 ▼
Balance restored

Final Golden Concept (Guyton)

The bicarbonate buffer system works because:

Lungs control CO₂ (acid part).
Kidneys control HCO₃⁻ (base part).
Blood pH depends on the ratio:

$\boxed{\mathbf{\frac{CO_2}{HCO_3^-}}}$ HCO3−CO2

Easy memory sentence:

“CO₂ is controlled by the lungs, HCO₃⁻ is controlled by the kidneys, and together they determine blood pH.”

Henderson-Hasselbalch Equation

H⁺ concentration is usually expressed in pH units instead of actual H⁺ concentration.
Recall: pH = −log(H⁺).

Equation

\boxed{\textbf{pH = -log[H^+]}}

Easiest Understanding

High H⁺ = Low pH = More acidic
Low H⁺ = High pH = More alkaline

The dissociation constant (pK) is also expressed using a logarithm.

Equation

$\boxed{\textbf{pK = -log K}}$ pK = -log K

Easiest Understanding

K = Dissociation constant.
pK = Logarithmic form of K.

Equation 4 can be converted into pH units.
This is done by taking the negative logarithm of Equation 4.

Equation (5)

Mathematical Equation

$\boxed{ -\log[H^+] = -\log K – \log\left(\frac{0.03\times PCO_2}{HCO_3^-}\right) }$ −log[H+]=−logK−log(HCO3−0.03×PCO2)

Step-by-Step Understanding

Start with Equation 4: $[H^+] = K \times \frac{0.03\times PCO_2}{HCO_3^-}$ [H+]=K×HCO3−0.03×PCO2

↓

Take −log on both sides.

↓

Replace −log(H⁺) with pH.

↓

Replace −log(K) with pK.

Equation (6)

Mathematical Equation

$\boxed{ pH = pK – \log \left( \frac{0.03\times PCO_2}{HCO_3^-} \right) }$ pH=pK−log(HCO3−0.03×PCO2)

Easiest Understanding

pH depends on:
- pK
- PCO₂
- HCO₃⁻

Instead of using a negative logarithm, the numerator and denominator are inverted.
This follows the law of logarithms.

Equation (7)

Mathematical Equation

$\boxed{ pH = pK + \log \left( \frac{HCO_3^-} {0.03\times PCO_2} \right) }$ pH=pK+log(0.03×PCO2HCO3−)

Step-by-Step Simplification

Equation (6)

↓ $pH = pK – \log \left( \frac{0.03\times PCO_2} {HCO_3^-} \right)$ pH=pK−log(HCO3−0.03×PCO2)

↓

Using the logarithm rule $-\log\left(\frac{A}{B}\right) = +\log\left(\frac{B}{A}\right)$ −log(BA)=+log(AB)

↓

Final Equation $pH = pK + \log \left( \frac{HCO_3^-} {0.03\times PCO_2} \right)$ pH=pK+log(0.03×PCO2HCO3−)

Equation (8)

For the bicarbonate buffer system,
pK = 6.1.

Henderson–Hasselbalch Equation

$\boxed{ pH = 6.1 + \log \left( \frac{HCO_3^-} {0.03\times PCO_2} \right) }$ pH=6.1+log(0.03×PCO2HCO3−)

Easiest Understanding

Blood pH depends on only TWO things

Numerator

HCO₃⁻

↓

Controlled by Kidneys

↓

More HCO₃⁻

↓

Higher pH

↓

Alkalosis

Denominator

PCO₂

↓

Controlled by Lungs

↓

More PCO₂

↓

Lower pH

↓

Acidosis

Golden Memory Formula

$\boxed{ \textbf{Blood pH} = \frac{\textbf{Kidney (HCO₃⁻)}} {\textbf{Lung (PCO₂)}} }$ Blood pH=Lung (PCO₂)Kidney (HCO₃⁻)

Equation 8 is called the Henderson–Hasselbalch equation.
It is used to calculate pH.
HCO₃⁻ concentration must be known.
PCO₂ must also be known.

An increase in HCO₃⁻ raises pH.
Increased HCO₃⁻ shifts acid–base balance toward alkalosis.

Easy Rule

⬆ HCO₃⁻

↓

⬆ pH

↓

Alkalosis

An increase in PCO₂ lowers pH.
Increased PCO₂ shifts acid–base balance toward acidosis.

Easy Rule

⬆ PCO₂

↓

⬇ pH

↓

Acidosis

The Henderson–Hasselbalch equation explains normal pH regulation.
It also explains acid–base balance in extracellular fluid.
It explains physiological control of acids and bases.
HCO₃⁻ concentration is mainly regulated by the kidneys.
PCO₂ is mainly regulated by respiration.

Easy Concept

Kidneys

↓

Control HCO₃⁻

Lungs

↓

Control PCO₂

Increased respiration removes more CO₂.
Plasma CO₂ decreases.

Easy Rule

⬆ Respiration

↓

⬇ CO₂

↓

⬆ pH

Decreased respiration increases PCO₂.

Easy Rule

⬇ Respiration

↓

⬆ CO₂

↓

⬇ pH

A primary decrease in HCO₃⁻ causes metabolic acidosis.

Memory

⬇ HCO₃⁻

↓

Metabolic Acidosis

A primary increase in HCO₃⁻ causes metabolic alkalosis.

Memory

⬆ HCO₃⁻

↓

Metabolic Alkalosis

An increase in PCO₂ causes respiratory acidosis.

Memory

⬆ PCO₂

↓

Respiratory Acidosis

A decrease in PCO₂ causes respiratory alkalosis.

Memory

⬇ PCO₂

↓

Respiratory Alkalosis

KEY CONCEPT

pH = −log(H⁺)
pK = −log(K)
Henderson–Hasselbalch Equation:

$\boxed{ pH = 6.1 + \log \left( \frac{HCO_3^-} {0.03\times PCO_2} \right) }$ pH=6.1+log(0.03×PCO2HCO3−)

Kidneys regulate HCO₃⁻ (base).
Lungs regulate PCO₂ (acid).
↑ HCO₃⁻ → ↑ pH → Metabolic Alkalosis
↓ HCO₃⁻ → ↓ pH → Metabolic Acidosis
↑ PCO₂ → ↓ pH → Respiratory Acidosis
↓ PCO₂ → ↑ pH → Respiratory Alkalosis

Mathematical/Biochemical Equations Solved

pH = −log(H⁺)
pK = −log(K)
Equation (5): $-\log(H^+)=-\log(K)-\log\left(\frac{0.03\times PCO_2}{HCO_3^-}\right)$ −log(H+)=−log(K)−log(HCO3−0.03×PCO2)
Equation (6): $pH=pK-\log\left(\frac{0.03\times PCO_2}{HCO_3^-}\right)$ pH=pK−log(HCO3−0.03×PCO2)
Equation (7): $pH=pK+\log\left(\frac{HCO_3^-}{0.03\times PCO_2}\right)$ pH=pK+log(0.03×PCO2HCO3−)
Equation (8) (Henderson–Hasselbalch Equation): $pH=6.1+\log\left(\frac{HCO_3^-}{0.03\times PCO_2}\right)$ pH=6.1+log(0.03×PCO2HCO3−)

Bicarbonate Buffer System Titration Curve

Figure Mentioned: Fig. 31.1

Fig. 31.1 shows how the pH of extracellular fluid changes when the HCO₃⁻/CO₂ ratio changes.
Changing the HCO₃⁻/CO₂ ratio changes the pH of the extracellular fluid.
When HCO₃⁻ and CO₂ concentrations are equal, the right side of Equation 8 becomes log(1).
log(1) = 0.

Mathematical Solution

When, $\frac{HCO_3^-}{CO_2}=1$ CO2HCO3−=1

Then, $\log(1)=0$ log(1)=0

Using Henderson-Hasselbalch Equation, $pH=6.1+\log(1)$ pH=6.1+log(1) $pH=6.1+0$ pH=6.1+0 $\boxed{pH=6.1}$ pH=6.1

Therefore, when HCO₃⁻ = CO₂, the pH is equal to the pK (6.1) of the bicarbonate buffer system.
When a base is added, some dissolved CO₂ is converted into HCO₃⁻.

Biochemical Equation

$CO_2 \longrightarrow HCO_3^-$ CO2⟶HCO3−

Easy Concept

Base Added

↓

CO₂ decreases

↓

HCO₃⁻ increases

↓

HCO₃⁻/CO₂ ratio increases

↓

pH increases

↓

Solution becomes more alkaline

Increasing the HCO₃⁻/CO₂ ratio increases the pH.
This is explained by the Henderson-Hasselbalch equation.
When an acid is added, it is buffered by HCO₃⁻.

Biochemical Equation

$H^+ + HCO_3^- \rightarrow H_2CO_3 \rightarrow CO_2 + H_2O$ H++HCO3−→H2CO3→CO2+H2O

Easy Concept

Acid Added

↓

HCO₃⁻ binds H⁺

↓

H₂CO₃ forms

↓

CO₂ forms

↓

HCO₃⁻ decreases

↓

CO₂ increases

↓

HCO₃⁻/CO₂ ratio decreases

↓

pH decreases

↓

Solution becomes more acidicffer Power Determined By Amount and Relative Concentrations of Buffer Components

Figure Mentioned: Fig. 31.1

Fig. 31.1 demonstrates several important features of the bicarbonate buffer system.
First, the pH equals the pK when HCO₃⁻ and CO₂ each make up 50% of the total buffer concentration.

Easy Concept

50% HCO₃⁻

50% CO₂

↓

Ratio = 1

↓

pH = pK = 6.1

Second, the buffer system is most effective in the middle of the titration curve.
The buffer system works best when the pH is close to the pK.

Easy Concept

pH ≈ pK

↓

Maximum buffering

↓

Minimum change in pH

When the pH is near the pK, adding acid or base causes the smallest change in pH.
The bicarbonate buffer system remains reasonably effective for 1 pH unit above and below the pK.

Mathematical Solution

Given: $pK=6.1$ pK=6.1

Lower limit: $6.1-1=5.1$ 6.1−1=5.1

Upper limit: $6.1+1=7.1$ 6.1+1=7.1

Effective Buffer Range

$\boxed{pH=5.1\ \text{to}\ 7.1}$ pH=5.1 to 7.1

The bicarbonate buffer system works effectively between pH 5.1 and 7.1.
Beyond pH 5.1–7.1, the buffering power rapidly decreases.
When all CO₂ has been converted into HCO₃⁻, the buffer system cannot buffer any more base.

Easy Concept

All CO₂ used

↓

No acid component left

↓

No more buffering

When all HCO₃⁻ has been converted into CO₂, the buffer system cannot buffer any more acid.

Easy Concept

All HCO₃⁻ used

↓

No base component left

↓

No more buffering

The total concentration of buffer also determines buffering power.
Higher buffer concentration provides greater buffering power.
Lower buffer concentration provides weaker buffering power.
When buffer concentration is low, even a small amount of acid or base causes a large change in pH.

Easy Concept

High Buffer Concentration

↓

Strong buffering

↓

Small pH change

Low Buffer Concentration

↓

Weak buffering

↓

Large pH change

KEY CONCEPT

Figure Mentioned: Fig. 31.1
Fig. 31.1 shows how pH changes when the HCO₃⁻/CO₂ ratio changes.
When HCO₃⁻ = CO₂, Ratio = 1, log(1) = 0, therefore pH = pK = 6.1.
Adding Base → ↑ HCO₃⁻ → ↑ HCO₃⁻/CO₂ ratio → ↑ pH.
Adding Acid → HCO₃⁻ buffers H⁺ → ↑ CO₂ → ↓ HCO₃⁻/CO₂ ratio → ↓ pH.
Maximum buffering occurs when pH ≈ pK.
Effective buffering range is pH 5.1–7.1.
Buffering power decreases rapidly outside this range.
Higher buffer concentration provides stronger buffering.
Lower buffer concentration provides weaker buffering.

Mathematical/Biochemical Equations Solved

$\displaystyle \frac{HCO_3^-}{CO_2}=1$ CO2HCO3−=1
$\displaystyle \log(1)=0$ log(1)=0
$\displaystyle pH=6.1+\log(1)=6.1$ pH=6.1+log(1)=6.1
$\displaystyle CO_2 \rightarrow HCO_3^-$ CO2→HCO3−
$\displaystyle H^+ + HCO_3^- \rightarrow H_2CO_3 \rightarrow CO_2 + H_2O$ H++HCO3−→H2CO3→CO2+H2O
Effective buffer range:

$\boxed{pH=5.1\ \text{to}\ 7.1}$ pH=5.1 to 7.1

Figure 31.1: Bicarbonate Buffer Titration Curve (SUPERFAST Explanation)This graph explains how the bicarbonate buffer system (HCO₃⁻/H₂CO₃) keeps the blood pH stable when acid or base is added.

Step 1: Understand the Axes

X-Axis (Horizontal)

This shows how acidic or alkaline the blood is.

Left side (pH 4) = Very acidic
Middle (pH 6.1) = Equal amounts of HCO₃⁻ and H₂CO₃
Normal blood = pH 7.4
Right side (pH 8) = Very alkaline (basic)

As you move right → pH increases → Blood becomes more alkaline.

Left Y-Axis

Percent of Buffer in the Form of H₂CO₃ (or CO₂)

This tells us:

How much of the bicarbonate buffer exists as acid (H₂CO₃).

Top = 0%

Bottom = 100%

Notice the arrow on the left:

⬇ Acid Added

Meaning:

When acid is added,
more buffer converts into H₂CO₃ (acid form).

Right Y-Axis

Percent of Buffer in the Form of HCO₃⁻

This tells us:

How much buffer exists as bicarbonate (base form).

Bottom = 0%

Top =100%

Arrow on the right:

⬆ Base Added

Meaning:

When base is added,

more buffer converts into HCO₃⁻ (base form).

The Red S-Shaped Curve

This curve is called the titration curve.

It tells us

At each pH, what percentage of buffer is acid (H₂CO₃) and what percentage is base (HCO₃⁻).

Start at pH = 4 (Far Left)

Look at the left end of the curve.

Here:

Left axis = nearly 100%

Right axis = nearly 0%

Meaning

Almost all buffer is H₂CO₃ (acid form).

Very little bicarbonate remains.

Why?

Because blood contains lots of acid.

The buffer has accepted hydrogen ions.

Reaction:

HCO₃⁻ + H⁺ → H₂CO₃

Almost everything becomes H₂CO₃.

Moving from pH 4 → 5

The curve begins rising slowly.

This means

Some H₂CO₃ changes back into HCO₃⁻.

Still,

Most buffer is acid.

pH = 6.1 (Point Marked “pK”)

This is the most important point.

The graph labels it pK.

Here,

Left axis = 50%

Right axis = 50%

Meaning

Exactly half the buffer is

H₂CO₃

and

Half is HCO₃⁻.

So,

HCO₃⁻ = H₂CO₃

This is why,

According to the Henderson–Hasselbalch equation:

When pH = pK, the ratio HCO₃⁻ : H₂CO₃ = 1 : 1.

Why is pK Important?

At pH = pK,

the buffer can

accept acid
accept base

equally well.

This is the point of maximum buffering efficiency.

From pH 6.1 → 7

The curve rises steeply.

Now,

More H₂CO₃ changes into HCO₃⁻.

The buffer becomes mostly bicarbonate.

Normal Operating Point (pH ≈ 7.4)

This dot is extremely important.

It represents the normal blood pH.

At this point,

Right axis ≈ 95%

Left axis ≈ 5%

Meaning

About

95% of buffer is HCO₃⁻ (base)

Only

5% is H₂CO₃ (acid).Why Does the Body Stay Here?

Because blood normally contains much more bicarbonate than carbonic acid.

Approximately,

HCO₃⁻ : H₂CO₃ = 20 : 1

This ratio gives

Normal blood pH ≈ 7.4.

From pH 7.4 → 8

The curve becomes flat again.

Almost all buffer is now bicarbonate.

Nearly

100%

is HCO₃⁻.

Very little acid remains.Understanding the Left Arrow (Acid Added)

Suppose acid enters the blood.

Example:

Lactic acid

Hydrochloric acid

Sulfuric acid

Immediately,

Hydrogen ions react with bicarbonate.

HCO₃⁻ + H⁺ → H₂CO₃ → CO₂ + H₂O

Result:

✅ HCO₃⁻ decreases

✅ H₂CO₃ increases

The graph moves

← toward the left.

Blood becomes more acidic.

Understanding the Right Arrow (Base Added)

Suppose a base enters blood.

Example:

NaOH

Extra bicarbonate

The base removes hydrogen ions.

Then,

H₂CO₃ breaks apart.

H₂CO₃ → H⁺ + HCO₃⁻

More bicarbonate is formed.

Result

✅ HCO₃⁻ increases

✅ H₂CO₃ decreases

The graph moves

→ toward the right.

Blood becomes more alkaline.

Why is the Curve S-Shaped?

The curve has three regions, each with a different meaning.

1. Left Flat Region (Low pH)

Almost all buffer is H₂CO₃.
There is very little HCO₃⁻ left to neutralize additional acid.
Buffering capacity against further acid is low.

2. Middle Steep Region (Around pK)

3. Right Flat Region (High pH)

Almost all buffer is HCO₃⁻.
Very little H₂CO₃ remains to neutralize additional base.
Buffering capacity against further base is low.

Easy Story to Remember

Imagine the buffer exists in two forms:

H₂CO₃ = Acid Team
HCO₃⁻ = Base Team

At different pH values:

Low pH (acidic blood): Acid Team is much larger.
pH 6.1 (pK): Both teams are equal (50% each).
Normal pH 7.4: Base Team dominates (about 95%), while Acid Team is small (about 5%).
High pH (alkaline blood): Nearly everyone is on the Base Team.

KEY CONCEPT

X-axis: Blood pH (acidic → alkaline).
Left Y-axis: Percentage of buffer present as H₂CO₃/CO₂ (acid form).
Right Y-axis: Percentage of buffer present as HCO₃⁻ (base form).
Red S-shaped curve: Shows how the proportions of H₂CO₃ and HCO₃⁻ change with pH.
pK = 6.1: 50% H₂CO₃ and 50% HCO₃⁻ (1:1 ratio), giving maximum buffering efficiency.
Normal blood pH = 7.4: About 95% HCO₃⁻ and 5% H₂CO₃, corresponding to the physiological 20:1 bicarbonate-to-carbonic acid ratio that keeps blood pH normal.

Bicarbonate Buffer System Is the Most Important Extracellular Buffer

(Figure 31.1) shows the titration curve of the bicarbonate buffer system.
From Figure 31.1, the bicarbonate buffer system does not appear to be a powerful buffer at first glance.
There are two reasons why it seems weak.

Reason 1

The normal extracellular fluid (ECF) pH is about 7.4.
The pK of the bicarbonate buffer system is 6.1.
This means there is about 20 times more bicarbonate (HCO₃⁻) than dissolved CO₂ (or H₂CO₃).
Equation:
- HCO₃⁻ : CO₂ (or H₂CO₃) = 20 : 1
- Solved Ratio = 20 ÷ 1 = 20
- Therefore, HCO₃⁻ is 20 times greater than CO₂ (or H₂CO₃).
Because of this 20:1 ratio, the bicarbonate buffer system works on the part of the titration curve where the slope is low.
A low slope means the buffering power is poor.

Reason 2

The concentrations of both components of the bicarbonate buffer system (CO₂ and HCO₃⁻) are not high.
Therefore, this is another reason why the bicarbonate buffer system appears to be weak.
Despite these characteristics, the bicarbonate buffer system is the most powerful extracellular buffer in the body.
This seems like a paradox (an apparent contradiction).
The main reason is that both components of the buffer system are continuously regulated.
HCO₃⁻ is regulated by the kidneys.
CO₂ is regulated by the lungs.
Because the kidneys regulate HCO₃⁻ and the lungs regulate CO₂, the extracellular fluid pH can be controlled very precisely.
The kidneys control pH by removing or adding HCO₃⁻.
The lungs control pH by removing CO₂.
Together, the kidneys and lungs maintain precise control of extracellular fluid pH.

KEY CONCEPT

Figure 31.1 shows the titration curve of the bicarbonate buffer system.
Normal ECF pH = 7.4 and pK = 6.1.
HCO₃⁻ : CO₂ (or H₂CO₃) = 20 : 1, meaning bicarbonate is 20 times greater than dissolved CO₂.
The bicarbonate buffer works on the low-slope part of the curve, so it appears to have poor buffering power.
The concentrations of CO₂ and HCO₃⁻ are not high.
Despite this, it is the most important extracellular buffer because HCO₃⁻ is regulated by the kidneys and CO₂ is regulated by the lungs.
The kidneys and lungs together precisely regulate extracellular fluid pH.

PHOSPHATE BUFFER SYSTEM

The phosphate buffer system is not a major extracellular fluid buffer.
However, it plays an important role in buffering renal tubular fluid and intracellular fluid.
The two main components of the phosphate buffer system are:
- H₂PO₄⁻ (dihydrogen phosphate)
- HPO₄²⁻ (hydrogen phosphate)
When a strong acid (HCl) is added, the H⁺ ions are accepted by HPO₄²⁻ (the base form).
The reaction is: HCl + Na₂HPO₄ → NaH₂PO₄ + NaCl
Reaction Explanation:
- Strong acid (HCl) is converted into a weak acid (NaH₂PO₄).
- Therefore, the fall in pH is minimized.
As a result, HCl is replaced by an additional amount of the weak acid NaH₂PO₄.
Therefore, the decrease in pH becomes much smaller.
When a strong base (NaOH) is added to the phosphate buffer system, the OH⁻ ions are buffered by H₂PO₄⁻.
This forms more HPO₄²⁻ and water (H₂O).
The reaction is: NaOH + NaH₂PO₄ → Na₂HPO₄ + H₂O
Reaction Explanation:
- Strong base (NaOH) is converted into a weak base (Na₂HPO₄).
- Therefore, the increase in pH is only slight.
The phosphate buffer system has a pK of 6.8.
This pK is close to the normal body fluid pH of 7.4.
Therefore, the phosphate buffer system works near its maximum buffering power.
However, the concentration of phosphate in the extracellular fluid is low.
It is only about 8% of the concentration of the bicarbonate buffer.
Mathematical Calculation:
- Phosphate buffer concentration = 8% of bicarbonate buffer
- 8 ÷ 100 = 0.08
- Therefore, the phosphate buffer concentration is 0.08 times (8%) that of the bicarbonate buffer.
Therefore, the total buffering power of the phosphate buffer in extracellular fluid is much less than that of the bicarbonate buffer system.
In contrast, the phosphate buffer is especially important in the renal tubular fluid.
There are two reasons for this.

Reason 1

Phosphate becomes greatly concentrated in the renal tubules.
Therefore, the buffering power of the phosphate buffer system increases.

Reason 2

The tubular fluid usually has a much lower pH than the extracellular fluid.
Therefore, the working pH of the buffer becomes closer to its pK (6.8).
As a result, the phosphate buffer works more effectively in the renal tubules.
The phosphate buffer system is also important inside cells (intracellular fluid).
The concentration of phosphate inside cells is many times higher than in the extracellular fluid.
The pH of intracellular fluid is lower than the pH of extracellular fluid.
Therefore, the intracellular fluid pH is usually closer to the pK (6.8) of the phosphate buffer system.
As a result, the phosphate buffer system is more effective inside cells than in the extracellular fluid.

KEY CONCEPT

The phosphate buffer system mainly buffers renal tubular fluid and intracellular fluid.
Its two components are H₂PO₄⁻ and HPO₄²⁻.
Strong acid: HCl + Na₂HPO₄ → NaH₂PO₄ + NaCl, converting a strong acid into a weak acid.
Strong base: NaOH + NaH₂PO₄ → Na₂HPO₄ + H₂O, converting a strong base into a weak base.
The pK of the phosphate buffer system is 6.8, which is close to the normal body fluid pH.
The phosphate buffer concentration in extracellular fluid is 8% (0.08 times) that of the bicarbonate buffer.
Therefore, it has less buffering power in extracellular fluid.
It is more effective in renal tubules and intracellular fluid because phosphate concentration is higher and the pH is closer to its pK (6.8).