# Acids and Bases: Definitions, Examples, and pH Basics
An acid is a substance that increases the hydrogen ion concentration of a solution, and a base is a substance that decreases it, either by accepting protons or by releasing hydroxide. These two behaviors govern nearly every reaction in a living cell, from the way an enzyme folds to the way blood carries carbon dioxide.
Every buffer you pipette, every growth medium you adjust, and every protein you purify sits inside a narrow pH window that decides whether the experiment works. A lysosome runs near pH 5 while the cytosol sits near pH 7.2, and that two-unit gap is the difference between a functioning digestive compartment and a dead cell. Understanding acids and bases is therefore not a chemistry detour. It is the operating manual for [molecular biology](/blog/careers/molecular-biology).
## Why Acid-Base Chemistry Matters in the Lab
Three practical realities make this topic unavoidable at the bench.
First, enzymes have a pH optimum. Pepsin works in the stomach at pH 2, while [alkaline phosphatase](/knowledge/molecular-biology/alkaline-phosphatase) in the intestine prefers pH 9 to 10. Move a protein two pH units away from its optimum and activity collapses, often irreversibly, because the folded structure depends on the protonation state of surface and active-site residues.
Second, buffers protect those optima. A buffer is a weak acid paired with its conjugate base, and it resists pH change when you add acid, base, or a concentrated salt solution. Tris, HEPES, phosphate-buffered saline, and MOPS all exist because a reaction needs stable pH for hours.
Third, pH is a diagnostic and analytical variable. Clinical medicine reads blood pH to detect acidosis and alkalosis [1][2]. Analytical chemistry uses acid-base behavior to select matrices for mass spectrometry, where Brønsted-Lowry theory guides which matrix will ionize a metabolite without shattering it [3]. Acid-base theory is also used to rationalize degradation pathways in materials science, for example in methylammonium halide perovskites where the Brønsted acidity of HBr versus HI helps explain relative stability [4].
## The Three Definitions of Acids and Bases
Chemistry did not arrive at one definition. It arrived at three, each broader than the last. The 1923 year is a useful anchor because Brønsted, Lowry, and Lewis published their ideas independently within months of each other, and that historical moment is now used in science education to show how competing models get resolved [5].
### Arrhenius: The Water-Bound Definition
Svante Arrhenius defined an acid as a substance that releases hydrogen ions (H⁺) in water and a base as a substance that releases hydroxide ions (OH⁻) in water.
- Acid example: hydrochloric acid, HCl, which dissociates fully into H⁺ and Cl⁻ in water.
- Base example: sodium hydroxide, NaOH, which dissociates into Na⁺ and OH⁻.
The Arrhenius definition is intuitive and still useful for introductory work with strong mineral acids and bases. Its limitation is that it requires water. Ammonia, NH₃, is unmistakably basic, yet it contains no hydroxide to release. It produces OH⁻ only by reacting with water, which the Arrhenius model cannot describe cleanly.
### Brønsted-Lowry: The Proton-Transfer Definition
In 1923, Johannes Brønsted and Thomas Lowry independently proposed that an acid is a proton (H⁺) donor and a base is a proton acceptor. This is the definition working biologists use most often, because it describes reactions in water and in nonaqueous solvents alike [6][7].
- Acid example: acetic acid, CH₃COOH, donates a proton to water.
CH₃COOH + H₂O ⇌ CH₃COO⁻ + H₃O⁺
- Base example: ammonia, NH₃, accepts a proton from water.
NH₃ + H₂O ⇌ NH₄⁺ + OH⁻
Two features make this model powerful. First, every acid reaction produces a conjugate base, and every base reaction produces a conjugate acid. Acetic acid and acetate are a conjugate acid-base pair. Ammonium and ammonia are another. Second, water is amphiprotic, meaning it can act as either an acid or a base depending on its partner. That flexibility is why pure water supports both reactions above.
The Brønsted-Lowry framework is the backbone of the Henderson-Hasselbalch equation, of buffer design, and of clinical acid-base assessment [1][8]. It is also the theory behind rational matrix selection in MALDI mass spectrometry, where the matrix must protonate or deprotonate the analyte for ionization to occur [3].
One caveat comes from a 2023 study of sterically isolated Brønsted acid and base sites held in a pentanuclear metal scaffold. In that system, the acid site refused to deprotonate and the base site refused to protonate, so the pair coexisted across a wide pH range without interconverting [9]. The result does not overturn Brønsted-Lowry theory, but it shows that when a proton cannot physically reach the site, the usual rules stall.
### Lewis: The Electron-Pair Definition
Gilbert Lewis defined an acid as an electron-pair acceptor and a base as an electron-pair donor. This is the broadest of the three definitions and the one that reaches into coordination chemistry and enzymology.
- Acid example: the proton itself, H⁺, which accepts a lone pair from a base. Boron trifluoride, BF₃, is a classic non-proton Lewis acid because its empty orbital accepts an electron pair.
- Base example: ammonia, NH₃, which donates its lone pair on nitrogen to form a coordinate bond, as in BF₃ + NH₃ → F₃B-NH₃.
Lewis theory explains metal-ligand binding, the catalytic role of zinc and magnesium ions in enzymes, and nucleophilic attack in metabolism. Its limitation is breadth. Because it covers so many reactions, it is less useful for calculating pH, where the proton-transfer view is more direct.
### Comparison Table
| Theory | Proton donor/acceptor? | Example | Limitation |
|--|--|--|--|
| Arrhenius | Acid releases H⁺, base releases OH⁻ in water | HCl (acid), NaOH (base) | Requires water and hydroxide, so NH₃ is not explained |
| Brønsted-Lowry | Acid donates H⁺, base accepts H⁺ | CH₃COOH (acid), NH₃ (base) | Fails when a proton cannot physically access the site [9] |
| Lewis | Acid accepts an electron pair, base donates one | BF₃ (acid), NH₃ (base) | Too broad to calculate pH directly |
## pH, pKa, and What the Numbers Mean
The pH scale runs from 0 to 14, with values below 7 acidic, 7 neutral, and above 7 basic. Image: Christinelmiller, CC BY-SA 4.0, via Wikimedia Commons.
### The pH Scale
pH is defined as the negative base-10 logarithm of hydrogen ion activity.
pH = -log₁₀[H⁺]
The result is dimensionless. pH is a logarithm of a ratio, so it carries no units. Writing "pH 7.4 units" is incorrect. Writing "pH 7.4" is correct.
At 25 °C, pure water has [H⁺] = 1 × 10⁻⁷ M, so pH = 7.00. Solutions below 7 are acidic, solutions above 7 are basic. The scale is logarithmic, so a change of one pH unit means a tenfold change in hydrogen ion concentration. A drop from pH 7.4 to pH 7.0 in blood is a substantial physiological event, not a rounding error.
### pKa and the Half-Neutralization Rule
The acid dissociation constant, Ka, measures how readily an acid gives up its proton.
Ka = [H⁺][A⁻] / [HA]
pKa is the negative logarithm of Ka.
pKa = -log₁₀Ka
A small pKa means a strong acid. A large pKa means a weak acid that holds its proton tightly.
The single most useful fact about pKa is this: pKa equals the pH at which the acid is exactly half-neutralized, meaning half the molecules are protonated (HA) and half are deprotonated (A⁻). At that point [HA] = [A⁻], the ratio inside the logarithm becomes 1, and the logarithm of 1 is zero. This is why pKa is read directly off the midpoint of a titration curve and why buffers work best within one pH unit of their pKa.
### Strong Versus Weak Acids
A strong acid dissociates completely in water. HCl, HBr, HI, HNO₃, H₂SO₄, and HClO₄ are the standard examples. Because dissociation is complete, the hydrogen ion concentration equals the nominal acid concentration, and no equilibrium calculation is needed.
A weak acid dissociates only partially. Acetic acid, carbonic acid, phosphoric acid, and the ammonium ion are common examples. Here the equilibrium matters, and you need Ka or pKa to find pH.
## Worked pH Calculations
### Example 1: A Strong Acid
Calculate the pH of 0.010 M HCl at 25 °C.
HCl dissociates completely.
HCl → H⁺ + Cl⁻
So [H⁺] = 0.010 M = 1.0 × 10⁻² M.
pH = -log₁₀(1.0 × 10⁻²) = 2.00
The answer is pH 2.00. No Ka is required because there is no equilibrium to solve. This is the calculation students should be able to do in their heads, since the exponent gives the pH directly for powers of ten.
### Example 2: A Weak Acid Using Ka
Calculate the pH of 0.100 M acetic acid. The Ka of acetic acid is 1.8 × 10⁻⁵.
Set up the equilibrium.
CH₃COOH ⇌ H⁺ + CH₃COO⁻
Initial concentrations: [CH₃COOH] = 0.100 M, [H⁺] = 0, [CH₃COO⁻] = 0.
Change: [CH₃COOH] loses x, [H⁺] gains x, [CH₃COO⁻] gains x.
Equilibrium: [CH₃COOH] = 0.100 - x, [H⁺] = x, [CH₃COO⁻] = x.
Substitute into Ka.
Ka = x² / (0.100 - x) = 1.8 × 10⁻⁵
Because Ka is small, x is much smaller than 0.100, so the approximation 0.100 - x ≈ 0.100 is valid.
x² = (1.8 × 10⁻⁵)(0.100) = 1.8 × 10⁻⁶
x = 1.34 × 10⁻³ M
So [H⁺] = 1.34 × 10⁻³ M.
pH = -log₁₀(1.34 × 10⁻³) = 2.87
The answer is pH 2.87. Note how much less acidic 0.100 M acetic acid is than 0.010 M HCl, even though the acetic acid is ten times more concentrated. Partial dissociation is the reason.
### Example 3: A Buffer Using Henderson-Hasselbalch
The Henderson-Hasselbalch equation rearranges the Ka expression into a form that works directly with concentrations and pKa.
pH = pKa + log₁₀([A⁻] / [HA])
Calculate the pH of a buffer made from 0.050 M acetic acid and 0.050 M sodium acetate. The pKa of acetic acid is 4.76.
pH = 4.76 + log₁₀(0.050 / 0.050)
pH = 4.76 + log₁₀(1)
pH = 4.76 + 0
pH = 4.76
This is the half-neutralization case. Equal concentrations of acid and conjugate base give pH exactly equal to pKa.
Now change the ratio. Suppose the buffer contains 0.100 M acetate and 0.050 M acetic acid.
pH = 4.76 + log₁₀(0.100 / 0.050)
pH = 4.76 + log₁₀(2)
pH = 4.76 + 0.30
pH = 5.06
Doubling the base relative to the acid raised pH by 0.30 units. This is the practical lever for buffer design. Adjust the ratio to set pH, and adjust the total concentration to set buffering capacity.
### A Note on the Henderson-Hasselbalch Equation
The Henderson-Hasselbalch equation is the traditional backbone of clinical acid-base assessment, used to categorize respiratory and metabolic disturbances through pH, PCO₂, and bicarbonate [1][2]. It is descriptive rather than mechanistic, and it can give misleading conclusions when albumin, total protein, or phosphate concentrations are markedly abnormal [1][10]. The Stewart or strong ion approach was developed partly in response, treating pH and bicarbonate as dependent variables controlled by the strong ion difference, the total weak acid concentration, and PCO₂ [2][11][12]. Both frameworks remain in use, and the Henderson-Hasselbalch equation is still the standard for calculating blood and urinary bicarbonate in routine practice [13][14].
## Buffers in Biological Systems
A buffer resists pH change because the weak acid component neutralizes added base and the conjugate base component neutralizes added acid. Buffering capacity is greatest at pH = pKa and falls off as you move away from it. This is why a buffer is normally chosen to sit within one pH unit of its pKa.
### The Bicarbonate Buffer
The bicarbonate system is the major extracellular buffer in blood and the one clinicians measure most often. It relies on the equilibrium between dissolved carbon dioxide and bicarbonate.
CO₂ + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻
The effective pKa of this system in blood plasma is about 6.1, which is well below the normal blood pH. That seems like a poor design until you account for the fact that the lungs continuously regulate PCO₂ and the kidneys regulate bicarbonate. The open system keeps the ratio of bicarbonate to dissolved CO₂ near 20 to 1, which holds pH in the normal range despite the unfavorable pKa.
### The Phosphate Buffer
The phosphate system is a major intracellular buffer and an important urinary buffer. It works through the equilibrium between dihydrogen phosphate and hydrogen phosphate.
H₂PO₄⁻ ⇌ HPO₄²⁻ + H⁺
The relevant pKa is about 6.8, close to intracellular pH, which makes phosphate effective inside cells. In plasma, phosphate concentration is low, so its contribution to extracellular buffering is modest compared with bicarbonate and plasma proteins [8].
### Physiological pH of Blood
Normal arterial blood pH in humans is 7.35 to 7.45. Values below 7.35 indicate acidemia, and values above 7.45 indicate alkalemia. The body defends this narrow window through respiratory control of CO₂ and renal control of bicarbonate, phosphate, and hydrogen ion excretion. The total concentration of nonvolatile weak acids, primarily inorganic phosphate and albumin, is one of the independent variables in the Stewart-Fencl analysis of acid-base status [2][12].
### Choosing a Buffer at the Bench
For molecular biology work, the common choices follow their pKa values. Acetate (pKa 4.76) suits pH 4 to 5. MES (pKa 6.1) suits pH 5.5 to 6.5. Phosphate (pKa 6.8) suits pH 6 to 7.5. HEPES (pKa 7.5) and Tris (pKa 8.1) suit pH 7 to 9. Glycine (pKa 9.8) suits pH 9 to 10.
Two practical points matter. First, temperature changes pKa, so a Tris buffer prepared at 25 °C will read differently at 4 °C. Second, phosphate buffers can precipitate divalent cations, so phosphate-buffered saline is a poor choice when your reaction needs free calcium or magnesium.
## How Acid-Base Behavior Is Measured
pH is measured with a glass electrode calibrated against standard buffers, typically pH 4.00, 7.00, and 10.00. The electrode produces a voltage proportional to hydrogen ion activity, and the meter converts that voltage to pH. Calibration at two or three points bracketing your expected value is standard practice.
For research applications, acid-base behavior is measured in other ways too.
- Titration with a pH electrode determines pKa and buffering capacity by tracking pH as strong acid or base is added. The inflection midpoint gives pKa.
- Spectrophotometric titration determines pKa when a molecule changes color or absorbance upon deprotonation.
- Photoluminescence methods extend pKa measurement to excited states, where the excited-state pKa can differ sharply from the ground-state value [15].
- Mass spectrometry exploits Brønsted-Lowry behavior for ionization, where matrix selection based on acid-base theory improves detection of small metabolites in blood, tissue, and biofluids [3].
## Common Mistakes and Limitations
**Confusing strong with concentrated.** A strong acid dissociates completely. A concentrated acid has many moles per liter. 0.001 M HCl is a strong acid and a dilute solution. 17.4 M acetic acid is a weak acid and a concentrated solution. The two properties are independent.
**Dropping the approximation check.** The shortcut [H⁺] = √(Ka × C) only works when the acid is weak and the concentration is not extremely low. If the calculated x exceeds roughly 5 percent of the nominal concentration, solve the quadratic instead.
**Treating pH as having units.** pH is dimensionless. It is a logarithm of a ratio of activities. Adding "units" or "molarity" to a pH value is a category error.
**Assuming pKa is fixed.** pKa shifts with temperature, ionic strength, and solvent. Acetic acid behaves as an acid in water but as a basic solution in DMSO when referenced to the aqueous pH scale, which shows how strongly solvent shapes acid-base behavior [6].
**Forgetting that buffers have a working range.** A buffer one and a half pH units from its pKa provides almost no resistance to pH change. Choose the buffer to match the target pH.
**Overextending the Henderson-Hasselbalch equation clinically.** The equation is descriptive, and it produces erroneous conclusions about the cause of an acid-base disturbance when serum protein, albumin, or phosphate concentrations are markedly abnormal [1]. In those settings, the strong ion approach gives a more mechanistic picture [10].
**Ignoring steric and structural limits.** Brønsted-Lowry theory assumes the proton can reach the site. When acid and base sites are sterically isolated, proton transfer can fail entirely and the pair can persist over a wide pH range without interconverting [9].
## Quick Review
1. Arrhenius defines acids and bases by H⁺ and OH⁻ release in water. Brønsted-Lowry defines them by proton donation and acceptance. Lewis defines them by electron-pair acceptance and donation.
2. pH = -log₁₀[H⁺] and is dimensionless.
3. pKa = -log₁₀Ka, and pKa equals the pH at half-neutralization, when [HA] = [A⁻].
4. Strong acids dissociate completely, so [H⁺] equals the nominal concentration. Weak acids need Ka and an equilibrium calculation.
5. Henderson-Hasselbalch, pH = pKa + log([A⁻]/[HA]), is the working equation for buffers.
6. Normal human blood pH is 7.35 to 7.45, defended mainly by the bicarbonate system with phosphate and proteins contributing.
7. Buffer capacity peaks at pH = pKa and falls off within about one pH unit.
## Frequently Asked Questions
### What is the difference between a strong acid and a weak acid?
A strong acid dissociates completely in water, so its hydrogen ion concentration equals its nominal concentration. A weak acid dissociates only partially, so you need its Ka or pKa to calculate pH. HCl is a strong acid and acetic acid is a weak acid.
### Why does pKa equal pH at half-neutralization?
At half-neutralization, half the acid molecules have donated their proton and half have not, so [HA] equals [A⁻]. The ratio inside the Henderson-Hasselbalch logarithm becomes 1, the logarithm of 1 is zero, and pH equals pKa.
### Is pH a unit?
No. pH is dimensionless because it is the negative logarithm of a ratio of hydrogen ion activities. Writing pH 7.4 is correct, and writing pH 7.4 units is not.
### What is the normal pH of human blood?
Normal arterial blood pH is 7.35 to 7.45. The bicarbonate buffer system, supported by respiratory control of carbon dioxide and renal control of bicarbonate, holds pH in that window.
### What is an example of a Lewis acid that is not a proton?
Boron trifluoride, BF₃, is a Lewis acid because it accepts an electron pair into an empty orbital. It has no proton to donate, which is why the Brønsted-Lowry definition does not cover it.
### Why do buffers stop working outside their pKa range?
Buffering depends on having both the acid and conjugate base forms present in useful amounts. More than about one pH unit from the pKa, one form dominates so heavily that added acid or base has little partner to react with, and pH changes almost unchecked.
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