Acids and Bases in Chemistry
Acids and bases are defined by what they do with protons and electron pairs. An acid donates hydrogen ions (H⁺, protons) or accepts a pair of electrons, while a base accepts hydrogen ions or donates a pair of electrons. Three complementary theories (Arrhenius, Brønsted–Lowry, and Lewis) widen that one idea from water solutions to almost any reaction, and the pH and pKa scales let you measure exactly how strong each one is.
Key takeaways
- Three theories describe acids and bases at increasing generality: Arrhenius (H⁺/OH⁻ in water), Brønsted–Lowry (proton donors and acceptors), and Lewis (electron-pair acceptors and donors).
- pH measures the hydrogen-ion concentration of a solution; pKa measures an acid's intrinsic strength, so the lower the pKa, the stronger the acid.
- Strong acids and bases ionize completely in water; weak ones ionize only partially.
- Acid + base → salt + water (neutralization), and the salt's own acidity depends on the parent acid and base.
- Buffers resist pH change and keep systems such as blood near a stable pH.
The three theories of acids and bases
Chemistry uses three definitions of acids and bases, each broader than the last. Arrhenius describes what happens in water, Brønsted–Lowry generalizes to any proton transfer, and Lewis extends the idea to reactions where no proton moves at all. They do not contradict each other, and each is a wider lens on the same behavior.
Arrhenius theory
The Arrhenius definition is the classic one taught first: an acid is a substance that produces hydrogen ions (H⁺) when dissolved in water, and a base is a substance that produces hydroxide ions (OH⁻). Hydrochloric acid (HCl) is an Arrhenius acid because it releases H⁺, and sodium hydroxide (NaOH) is an Arrhenius base because it releases OH⁻. The limitation is that it only works for aqueous solutions and cannot explain why a substance like ammonia is basic even though it contains no hydroxide.
Brønsted–Lowry theory
The Brønsted–Lowry definition is the most useful for everyday chemistry: an acid is a proton (H⁺) donor and a base is a proton acceptor. This shifts the focus from the solution to the transfer itself. When HCl dissolves in water, HCl donates a proton to water, so HCl is the acid and water is the base:
HCl + H₂O → H₃O⁺ + Cl⁻
Because the reaction is a proton transfer, ammonia now makes sense as a base, because it accepts a proton from water to form the ammonium ion. See Brønsted–Lowry acid and Brønsted–Lowry base for the precise definitions.
Lewis theory
The Lewis definition is the broadest: a Lewis acid is an electron-pair acceptor, and a Lewis base is an electron-pair donor. Every Brønsted acid is also a Lewis acid, but the Lewis view also covers reactions with no proton at all, such as boron trifluoride (BF₃) accepting a lone pair from ammonia, or a metal ion binding water molecules. This is why the Lewis definition matters far beyond simple acid–base chemistry, reaching into catalysis and coordination compounds.
Conjugate acid–base pairs
Every Brønsted–Lowry reaction produces a conjugate acid–base pair. When an acid donates its proton, what remains is its conjugate base; when a base accepts a proton, the result is its conjugate acid. In the reaction of acetic acid with water, acetic acid and the acetate ion form one pair, while water and the hydronium ion (H₃O⁺) form the other. A strong acid always has a weak conjugate base, and vice versa.
Some substances can act as either an acid or a base depending on their partner, a property called amphoterism. Water is the classic example: it donates a proton to ammonia but accepts one from HCl.
Measuring strength: pH and pKa
There are two different things worth measuring, and students often confuse them. pH tells you how acidic a particular solution is right now; pKa tells you how strong an acid is by its nature, regardless of concentration.
The pH scale
pH is defined as the negative logarithm of the hydrogen-ion concentration, pH = −log[H⁺]. In water at 25 °C the scale runs from 0 to 14:
- pH below 7: acidic (lower numbers are more strongly acidic)
- pH of 7: neutral (pure water)
- pH above 7: basic, or alkaline
Because the scale is logarithmic, each unit is a tenfold change: a solution of pH 3 has ten times the hydrogen-ion concentration of one at pH 4. The 0–14 range comes from the ion-product of water, Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25 °C, which fixes neutral water at pH 7.
pKa and Ka
An acid's intrinsic strength is captured by its acid ionization constant, Ka, the equilibrium constant for the acid giving up its proton. Because Ka values span many orders of magnitude, chemists use pKa = −log Ka instead. The lower the pKa, the stronger the acid. Hydrochloric acid has a pKa near −6, so it ionizes essentially completely, while acetic acid has a pKa of about 4.76, so most of it stays intact in solution.
Strong vs weak acids and bases
The words *strong* and *weak* describe how completely a substance ionizes, not how concentrated or dangerous it is. A strong acid is a strong electrolyte that ionizes completely in water, while a weak acid is a weak electrolyte that only partially ionizes and sits at equilibrium with its un-ionized form.
- Strong acids: hydrochloric (HCl), nitric (HNO₃), and sulfuric (H₂SO₄) acids.
- Weak acids: acetic acid (in vinegar), carbonic acid, and citric acid.
- Strong bases: sodium hydroxide (NaOH) and potassium hydroxide (KOH).
- Weak bases: ammonia and most amines.
Note that "strong" is not the same as "concentrated." A dilute solution of a strong acid can have a higher pH than a concentrated solution of a weak acid. In water, all strong acids appear equally strong because the solvent cannot tell them apart, a phenomenon called the leveling effect.
Acids that can donate more than one proton are classified by how many they give up: a monoprotic acid such as HCl donates one, a diprotic acid such as H₂SO₄ donates two, and a polyprotic acid such as phosphoric acid donates three.
Neutralization and salts
When an acid reacts with a base, they cancel each other in a reaction called neutralization, and the products are always a salt and water:
acid + base → salt + water
Hydrochloric acid and sodium hydroxide, for instance, react to give sodium chloride (ordinary table salt) and water. The salt that forms is not always neutral, though. Its pH depends on its parents: a salt of a strong acid and a weak base is an acidic salt (ammonium chloride), while a salt of a weak acid and a strong base is a basic salt (sodium acetate). Only a salt from a strong acid and a strong base, like NaCl, gives a neutral solution.
Chemists find the exact amount of acid or base in a sample by titration, adding a measured base to an acid (or vice versa) until they exactly cancel at the equivalence point. An indicator such as litmus or phenolphthalein changes color to signal the endpoint.
Buffers: resisting pH change
A buffer solution is a mixture that resists changes in pH when small amounts of acid or base are added. It usually contains a weak acid together with its conjugate base (or a weak base with its conjugate acid), so it can soak up added H⁺ or OH⁻ without much shift in pH. Buffers are essential to living things: human blood is buffered by a carbonic-acid/bicarbonate system that holds its pH close to 7.4, and even a small deviation is dangerous. Buffers also keep fermentation, drug formulations, and countless industrial processes running at a controlled pH.
Acids and bases in everyday life and industry
Acid–base chemistry is all around you. Your stomach uses hydrochloric acid to digest food, and antacids (mild bases) neutralize the excess. Citric acid gives citrus fruit its sharpness, acetic acid makes vinegar sour, and carbonic acid puts the fizz in soft drinks. On the basic side, soaps and many cleaning products are alkaline, and baking soda (a mild base) reacts with acids to make cakes rise.
Industry depends on them even more heavily. Sulfuric acid is one of the most-produced chemicals in the world, used to make fertilizers, refine petroleum, and process metals. Ammonia, a weak base, is the feedstock for the fertilizers that support modern agriculture. Understanding whether a substance donates or accepts protons is the first step to predicting how it will behave in any of these settings.
Frequently asked questions
What is the simplest definition of an acid and a base?
The simplest working definition is the Brønsted–Lowry one: an acid is a proton (H⁺) donor and a base is a proton acceptor. This explains why acids and bases neutralize each other, since the proton simply moves from the acid to the base.
What is the difference between a strong acid and a concentrated acid?
"Strong" describes how completely an acid ionizes, while "concentrated" describes how much acid is dissolved. A strong acid ionizes fully but can still be dilute, and a weak acid can be concentrated yet only partly ionized. The two properties are independent.
Is pH the same as acid strength?
No. pH measures the hydrogen-ion concentration of a specific solution, whereas acid strength (pKa) is an intrinsic property of the acid itself. A concentrated weak acid and a dilute strong acid can even share the same pH despite very different strengths.
Can a substance be both an acid and a base?
Yes. Substances that can act as either are called amphoteric. Water is the best example: it donates a proton when reacting with a base and accepts one when reacting with an acid. Bicarbonate ion and several metal oxides behave the same way.
What makes a buffer resist pH change?
A buffer contains a weak acid and its conjugate base in similar amounts. Added acid is consumed by the conjugate base and added base is consumed by the weak acid, so the hydrogen-ion concentration, and therefore the pH, barely moves.
Want precise, exam-ready definitions of the terms above? Explore the full A–Z Chemistry Dictionary for entries like acid, base, pH, and buffer solution.