Overview
Almost every fluid in the body, from stomach acid to blood plasma, has a precise degree of acidity that the body works constantly to control. This lesson explains what acids and bases actually are, how chemists measure acidity with the pH scale, and how buffers protect delicate biological systems from dangerous swings in pH. These ideas are foundational for understanding blood-gas results, drug behavior, and the body’s response to illness.
What Are Acids and Bases?
Two complementary definitions are worth knowing.
The Arrhenius model is the simplest: an acid releases hydrogen ions (H⁺) in water, and a base releases hydroxide ions (OH⁻). Hydrochloric acid (HCl) is a classic Arrhenius acid; sodium hydroxide (NaOH) is a classic base.
The Bronsted-Lowry model is broader and more useful in biology. Here an acid is a proton (H⁺) donor and a base is a proton acceptor. When an acid donates its proton, what remains is its conjugate base. For example, carbonic acid (H₂CO₃) donates a proton to become bicarbonate (HCO₃⁻), its conjugate base. This donor-acceptor view explains reactions that do not involve hydroxide at all, which is exactly what happens throughout the bloodstream.
The pH Scale
pH is a measure of hydrogen-ion concentration, defined as the negative base-10 logarithm of [H⁺]:
pH = −log₁₀[H⁺]
Because the scale is logarithmic, each whole-number step represents a tenfold change in H⁺ concentration. A solution at pH 4 has ten times more H⁺ than one at pH 5 and one hundred times more than one at pH 6.
| pH range | Nature | Everyday example |
|---|---|---|
| 0–6 | Acidic | Stomach acid (~1.5–3.5), lemon juice (~2) |
| 7 | Neutral | Pure water |
| 8–14 | Basic (alkaline) | Baking soda (~9), bleach (~13) |
Pure water is neutral at pH 7 because it contains equal, tiny amounts of H⁺ and OH⁻. Values below 7 are acidic and above 7 are basic.
Strong vs. Weak, and Neutralization
A strong acid or strong base dissociates almost completely in water. HCl separates fully into H⁺ and Cl⁻, so nearly every molecule contributes to acidity. A weak acid, such as acetic acid or carbonic acid, only partially dissociates and sits in equilibrium with its conjugate base. Strength is about the degree of dissociation, not how concentrated or “dangerous” a solution feels.
Neutralization occurs when an acid reacts with a base to yield water and a salt:
HCl + NaOH → H₂O + NaCl
The H⁺ from the acid and the OH⁻ from the base combine to form water, and the leftover ions form a salt. This is the same principle behind antacids, which neutralize excess stomach acid.
Buffers: Resisting pH Change
A buffer is a solution that resists changes in pH when small amounts of acid or base are added. Buffers are made of a weak acid paired with its conjugate base. When acid enters, the conjugate base mops up the extra H⁺; when base enters, the weak acid releases H⁺ to compensate. The net effect is a remarkably stable pH.
The Bicarbonate Buffer System
The body’s most important extracellular buffer is the bicarbonate buffer system, based on this equilibrium:
CO₂ + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻
If acid builds up, bicarbonate (HCO₃⁻) binds the excess H⁺ to form carbonic acid, which converts to CO₂ and water. The lungs then exhale the extra CO₂, and the kidneys adjust how much bicarbonate is retained or excreted. Because breathing rate can change within seconds and kidney handling over hours to days, this system provides both fast and slow control.
Clinical relevance
Human arterial blood is held within a narrow window of about 7.35 to 7.45, and even small deviations impair enzyme function, oxygen delivery, and heart rhythm. When blood pH falls below 7.35 the state is called acidosis; above 7.45 it is called alkalosis. These can arise from lung problems (respiratory, involving CO₂) or metabolic problems (involving bicarbonate and other acids). For example, a patient hyperventilating blows off too much CO₂ and drifts toward respiratory alkalosis, while uncontrolled diabetes can flood the blood with acids and cause metabolic acidosis. Clinicians read arterial blood-gas panels, tracking pH, CO₂, and bicarbonate, precisely because the bicarbonate buffer and its lung and kidney partners tell the story of how the body is coping. Understanding acids, bases, and buffers is therefore not abstract chemistry; it is the language of managing a critically ill patient.