R3.1 Proton transfer reactions

Acids and bases are some of the most familiar substances in chemistry: hydrochloric acid in the stomach, ethanoic acid in vinegar, sodium hydrogencarbonate in baking powder and ammonia in household cleaners. Their reactions all involve the same simple event, the transfer of a hydrogen ion (a proton) from one species to another. This idea explains why solutions are acidic or alkaline, how acidity is measured on the pH scale, and how acids and bases neutralize each other.

The Brønsted–Lowry theory of acids and bases

Acids were first recognized by their properties: they taste sour, turn blue litmus red, react with reactive metals to give hydrogen and react with carbonates to give carbon dioxide. Bases were the substances that cancel out these properties. Explaining why substances behave like this needed a model at the level of particles, and that model has changed as chemists found reactions the older definitions could not explain.

The first particle-level definition, proposed by Svante Arrhenius in the 1880s, described an acid as a substance that produces hydrogen ions, ${{\mathrm{H}{\vphantom{X}}^{+}}}$, in water and a base as one that produces hydroxide ions, ${{\mathrm{OH}{\vphantom{X}}^{-}}}$. It works well for many aqueous reactions, but it cannot describe ammonia as a base (ammonia contains no hydroxide), and it says nothing about reactions that happen without water, such as hydrogen chloride gas reacting with ammonia gas to form a white smoke of ammonium chloride:

In 1923 Johannes Brønsted and Thomas Lowry independently proposed a more general definition based on what actually moves in these reactions: a hydrogen ion.

A hydrogen atom that has lost its only electron is just a proton, which is why ${{\mathrm{H}{\vphantom{X}}^{+}}}$ and "proton" mean the same thing here. In the reaction of hydrogen chloride with ammonia, ${{\mathrm{HCl}}}$ donates a proton and ${{\mathrm{NH}{\vphantom{X}}_{\smash[t]{3}}}}$ accepts it, forming ${{\mathrm{NH}{\vphantom{X}}_{\smash[t]{4}}{\vphantom{X}}^{+}}}$ and ${{\mathrm{Cl}{\vphantom{X}}^{-}}}$ ions. The Brønsted–Lowry definition covers this reaction in the gas phase as well as the familiar reactions in water. It was extended again later: the Lewis definition, in terms of electron pairs, includes reactions in which no proton moves at all (R3.4).

The proton in water: the oxonium ion

A bare proton is far too small and too highly charged to exist on its own in solution. When an acid dissolves in water, the proton is accepted by a lone pair on the oxygen atom of a water molecule, forming the oxonium ion, ${{\mathrm{H}{\vphantom{X}}_{\smash[t]{3}}\mathrm{O}{\vphantom{X}}^{+}}}$ (also called the hydronium ion). In this reaction water acts as a base. The Lewis formulas of the reaction between hydrogen chloride and water show the proton moving from chlorine to oxygen.

Both forms are used to represent the aqueous proton, and both are correct:

The first equation shows the role of water as a base explicitly. The second is shorter and is used when the role of water does not matter, for example in pH calculations, where $[{{\mathrm{H}{\vphantom{X}}^{+}}}]$ and $[{{\mathrm{H}{\vphantom{X}}_{\smash[t]{3}}\mathrm{O}{\vphantom{X}}^{+}}}]$ mean the same concentration.

What makes a species an acid or a base

To act as a Brønsted–Lowry acid, a species needs a hydrogen atom that it can lose as ${{\mathrm{H}{\vphantom{X}}^{+}}}$. This is usually a hydrogen bonded to a more electronegative atom, such as the H in ${{\mathrm{H}{-}\mathrm{Cl}}}$ or in the ${{\mathrm{O}{-}\mathrm{H}}}$ group of a carboxylic acid; the hydrogen atoms bonded to carbon in methane or in the ${{\mathrm{CH}{\vphantom{X}}_{\smash[t]{3}}}}$ group of ethanoic acid are not released. To act as a Brønsted–Lowry base, a species needs a lone pair of electrons that can form a bond to the incoming proton. The nitrogen atom in ammonia and amines, the oxygen atom in water, and anions such as ${{\mathrm{OH}{\vphantom{X}}^{-}}}$, ${{\mathrm{O}{\vphantom{X}}^{2-}}}$ and ${{\mathrm{CO}{\vphantom{X}}_{\smash[t]{3}}{\vphantom{X}}^{2-}}}$ all have lone pairs available.

This explains why some reactions are acid–base reactions even though no substance in them is called an acid. Magnesium nitride reacts with water to give magnesium hydroxide and ammonia. Each nitride ion, ${{\mathrm{N}{\vphantom{X}}^{3-}}}$, accepts three protons from water molecules, so the nitride ion is the base and water is the acid:

In the same way, the oxide ion of a soluble metal oxide accepts a proton from water: ${{\mathrm{O}{\vphantom{X}}^{2-} {}+{} \mathrm{H}{\vphantom{X}}_{\smash[t]{2}}\mathrm{O}\mskip2mu (\mathrm{l})\xrightarrow{}2\,\mathrm{OH}{\vphantom{X}}^{-}\mskip2mu (\mathrm{aq})}}$. The oxide ion cannot exist in water, so this reaction goes to completion. This is why sodium oxide and calcium oxide dissolve in water to form alkaline solutions.

Identifying the acid and the base in an equation

To identify the Brønsted–Lowry acid and base in an equation, compare each reactant with the product it forms. The species that loses an ${{\mathrm{H}{\vphantom{X}}^{+}}}$ (its formula has one H fewer and a charge one lower) is the acid; the species that gains an ${{\mathrm{H}{\vphantom{X}}^{+}}}$ is the base. Many acid–base reactions are reversible, so the reverse reaction also has an acid and a base. In the reaction of ammonia with water, water donates a proton in the forward reaction and the ammonium ion donates one in the reverse reaction:

So ${{\mathrm{H}{\vphantom{X}}_{\smash[t]{2}}\mathrm{O}}}$ and ${{\mathrm{NH}{\vphantom{X}}_{\smash[t]{4}}{\vphantom{X}}^{+}}}$ are the two acids in this equilibrium, and ${{\mathrm{NH}{\vphantom{X}}_{\smash[t]{3}}}}$ and ${{\mathrm{OH}{\vphantom{X}}^{-}}}$ are the two bases.

Bases and alkalis

Many bases are insoluble in water. Copper(II) oxide, magnesium oxide and iron(III) hydroxide all neutralize acids, so they are bases, but they do not dissolve to give hydroxide ions. A base that does dissolve in water and releases ${{\mathrm{OH}{\vphantom{X}}^{-}}}$ ions is called an alkali.

Aqueous ammonia is an alkali even though ammonia contains no hydroxide: it produces ${{\mathrm{OH}{\vphantom{X}}^{-}}}$ ions by accepting protons from water, as in the equilibrium for ammonia and water. Soluble carbonates such as sodium carbonate also give alkaline solutions for the same reason, because the carbonate ion accepts a proton from water.

Conjugate acid–base pairs

When an acid donates a proton, what is left can accept a proton again, so it is a base. When a base accepts a proton, the product can give it back, so it is an acid. Species linked in this way are called a conjugate pair.

Every Brønsted–Lowry acid–base equilibrium contains two conjugate pairs. In the equilibria for ethanoic acid and for ammonia in water, each pair is made up of one species on the left and one on the right.

Writing the formula of a conjugate acid or base

Two simple rules give the formula of any conjugate species:

• Conjugate base: remove one H from the formula and lower the charge by one. ${{\mathrm{HNO}{\vphantom{X}}_{\smash[t]{3}}\xrightarrow{}\mathrm{NO}{\vphantom{X}}_{\smash[t]{3}}{\vphantom{X}}^{-}}}$; ${{\mathrm{HCO}{\vphantom{X}}_{\smash[t]{3}}{\vphantom{X}}^{-}\xrightarrow{}\mathrm{CO}{\vphantom{X}}_{\smash[t]{3}}{\vphantom{X}}^{2-}}}$; ${{\mathrm{NH}{\vphantom{X}}_{\smash[t]{4}}{\vphantom{X}}^{+}\xrightarrow{}\mathrm{NH}{\vphantom{X}}_{\smash[t]{3}}}}$; ${{\mathrm{H}{\vphantom{X}}_{\smash[t]{2}}\mathrm{O}\xrightarrow{}\mathrm{OH}{\vphantom{X}}^{-}}}$.

• Conjugate acid: add one H to the formula and raise the charge by one. ${{\mathrm{NH}{\vphantom{X}}_{\smash[t]{3}}\xrightarrow{}\mathrm{NH}{\vphantom{X}}_{\smash[t]{4}}{\vphantom{X}}^{+}}}$

Only one proton moves at a time. The conjugate base of ${{\mathrm{H}{\vphantom{X}}_{\smash[t]{2}}\mathrm{SO}{\vphantom{X}}_{\smash[t]{4}}}}$ is ${{\mathrm{HSO}{\vphantom{X}}_{\smash[t]{4}}{\vphantom{X}}^{-}}}$, not ${{\mathrm{SO}{\vphantom{X}}_{\smash[t]{4}}{\vphantom{X}}^{2-}}}$, and the conjugate base of ${{\mathrm{H}{\vphantom{X}}_{\smash[t]{3}}\mathrm{PO}{\vphantom{X}}_{\smash[t]{4}}}}$ is ${{\mathrm{H}{\vphantom{X}}_{\smash[t]{2}}\mathrm{PO}{\vphantom{X}}_{\smash[t]{4}}{\vphantom{X}}^{-}}}$. The polyatomic ions met in S2.1 are the conjugate bases of familiar acids, as the table of conjugate pairs shows.

Read from right to left, the same table gives conjugate acids: the conjugate acid of the carbonate ion is ${{\mathrm{HCO}{\vphantom{X}}_{\smash[t]{3}}{\vphantom{X}}^{-}}}$, and the conjugate acid of the hydroxide ion is water.

When the structure of a conjugate species is drawn, its charge belongs on the atom that lost or gained the proton. In the ethanoate ion the negative charge sits on the oxygen that held the acidic hydrogen; in the methylammonium ion the positive charge sits on nitrogen, which used its lone pair to bond to the proton.

Amphiprotic species

Some species can act as an acid in one reaction and as a base in another. To do this they need both a hydrogen that can be released as ${{\mathrm{H}{\vphantom{X}}^{+}}}$ and a lone pair that can accept one.

Water is the most important example. It acts as a base when hydrogen chloride dissolves in it, accepting a proton to become ${{\mathrm{H}{\vphantom{X}}_{\smash[t]{3}}\mathrm{O}{\vphantom{X}}^{+}}}$, and as an acid when ammonia dissolves in it, donating a proton to become ${{\mathrm{OH}{\vphantom{X}}^{-}}}$. Which role it takes depends on the other reactant: water behaves as a base towards stronger acids and as an acid towards stronger bases.

Many anions that still contain an acidic hydrogen are amphiprotic too, because they also carry the negative charge and lone pairs that attract a proton. Each has a conjugate base and a conjugate acid.

The hydrogencarbonate ion shows this behaviour in everyday chemistry. It neutralizes acids, which is why sodium hydrogencarbonate is used in antacids and baking powder (${{\mathrm{HCO}{\vphantom{X}}_{\smash[t]{3}}{\vphantom{X}}^{-} {}+{} \mathrm{H}{\vphantom{X}}^{+}\xrightarrow{}\mathrm{H}{\vphantom{X}}_{\smash[t]{2}}\mathrm{O} {}+{} \mathrm{CO}{\vphantom{X}}_{\smash[t]{2}}}}$), and it also neutralizes alkalis (${{\mathrm{HCO}{\vphantom{X}}_{\smash[t]{3}}{\vphantom{X}}^{-} {}+{} \mathrm{OH}{\vphantom{X}}^{-}\xrightarrow{}\mathrm{CO}{\vphantom{X}}_{\smash[t]{3}}{\vphantom{X}}^{2-} {}+{} \mathrm{H}{\vphantom{X}}_{\smash[t]{2}}\mathrm{O}}}$). Amino acids are amphiprotic for a similar reason: the ${{{-}\mathrm{COOH}}}$ group can donate a proton and the ${{{-}\mathrm{NH}{\vphantom{X}}_{\smash[t]{2}}}}$ group can accept one.

Amphiprotic and amphoteric

A wider term, amphoteric, describes any substance that can react as both an acid and a base. Every amphiprotic species is amphoteric, but not the other way round. Aluminium oxide reacts with acids and with alkalis, so it is amphoteric, but it contains no hydrogen and cannot donate a proton, so it is not amphiprotic. The oxides of the period 3 elements show the trend across a period, from basic sodium and magnesium oxides, through amphoteric aluminium oxide, to the acidic oxides of phosphorus and sulfur (S3.1). The acidic oxides of sulfur and nitrogen released by burning fuels dissolve in rainwater as acids, which is the origin of acid rain (S3.1).

The pH scale

The concentration of hydrogen ions in aqueous solutions covers an enormous range: from about $1\ \text{mol dm}^{-3}$ in a strong acid to about $10^{-14}\ \text{mol dm}^{-3}$ in a strong alkali. Numbers that differ by fourteen powers of ten are awkward to compare, so chemists use a logarithmic scale. Taking the logarithm turns each power of ten into one step on the scale, and the minus sign makes the values positive for most solutions.

Both equations are in the data booklet, which writes the first as $\text{pH} = -\log_{10}[{{\mathrm{H}{\vphantom{X}}_{\smash[t]{3}}\mathrm{O}{\vphantom{X}}^{+}}}] = -\log_{10}[{{\mathrm{H}{\vphantom{X}}^{+}}}]$. Because the scale is logarithmic, pH has three important features:

• A change of one pH unit is a tenfold change in $[{{\mathrm{H}{\vphantom{X}}^{+}}}]$. A solution of pH 3 has ten times the hydrogen ion concentration of one of pH 4, and a thousand times that of one of pH 6.

• A lower pH means a higher $[{{\mathrm{H}{\vphantom{X}}^{+}}}]$. The minus sign reverses the order, so the most acidic solutions have the lowest pH.

• pH has no units. It is the logarithm of the numerical value of a concentration measured in $\text{mol dm}^{-3}$.

A graph of pH against $[{{\mathrm{H}{\vphantom{X}}^{+}}}]$ is therefore not a straight line. pH falls steeply at very low concentrations and only slowly at high ones: going from $10^{-5}$ to $10^{-4}\ \text{mol dm}^{-3}$ lowers the pH by one unit, and going from $0.01$ to $0.1\ \text{mol dm}^{-3}$ also lowers it by just one unit, although the second change is a thousand times larger in absolute terms.

The everyday range of pH runs from about 0 to 14 at 298 K. Values outside this range are possible: a solution of a strong acid more concentrated than $1\ \text{mol dm}^{-3}$ has a negative pH, but such solutions are rarely described with pH.

Approximate values for typical samples; real samples vary.

Calculations with pH

For a solution of a strong monoprotic acid, which releases one proton per molecule, the acid is completely ionized, so $[{{\mathrm{H}{\vphantom{X}}^{+}}}]$ equals the concentration of the acid. The pH follows directly from the definition. Working in the other direction, $[{{\mathrm{H}{\vphantom{X}}^{+}}}] = 10^{-\text{pH}}$ gives the hydrogen ion concentration from a measured pH.

Diluting a strong acid lowers $[{{\mathrm{H}{\vphantom{X}}^{+}}}]$ in proportion to the dilution, so the pH rises by one unit for every tenfold dilution. Diluting $0.100\ \text{mol dm}^{-3}$ hydrochloric acid (pH 1.00) ten times gives pH 2.00, and diluting it a thousand times gives pH 4.00. Dilution alone can never take an acid past pH 7, however: as $[{{\mathrm{H}{\vphantom{X}}^{+}}}]$ from the acid approaches $10^{-7}\ \text{mol dm}^{-3}$, the ions from water itself become important, and the pH approaches 7 from below.

Ocean acidification

Ocean acidification is the fall in the pH of seawater caused by the uptake of carbon dioxide from the atmosphere. About a quarter of the carbon dioxide released by burning fossil fuels dissolves in the oceans. Dissolved carbon dioxide reacts with water to form carbonic acid, a weak acid that releases hydrogen ions:

The average pH of surface seawater has fallen from about 8.2 before the industrial era to about 8.1 today. The change looks small, but on a logarithmic scale a fall of 0.1 units means $[{{\mathrm{H}{\vphantom{X}}^{+}}}]$ has been multiplied by $10^{0.1} = 1.26$, an increase of about 26 %. Measurements that put today's value slightly below 8.1 give the figure of about 30 % that is often quoted. Seawater is still basic, so "acidification" describes the direction of the change, not a change to an acidic solution.

The extra hydrogen ions react with carbonate ions in seawater, which act as a Brønsted–Lowry base: ${{\mathrm{H}{\vphantom{X}}^{+}\mskip2mu (\mathrm{aq}) {}+{} \mathrm{CO}{\vphantom{X}}_{\smash[t]{3}}{\vphantom{X}}^{2-}\mskip2mu (\mathrm{aq})\xrightarrow{}\mathrm{HCO}{\vphantom{X}}_{\smash[t]{3}}{\vphantom{X}}^{-}\mskip2mu (\mathrm{aq})}}$. Corals, oysters, mussels and many plankton build their shells and skeletons from calcium carbonate, so a lower carbonate ion concentration makes it harder for them to grow and can make existing shells start to dissolve.

Measuring pH: indicators and pH meters

The pH of a solution can be estimated with universal indicator, a mixture of several indicators that together change colour gradually across the whole pH range: red in strongly acidic solution, through orange and yellow to green near neutral, then blue and violet in strongly alkaline solution. A drop of the solution on indicator paper, or a few drops of indicator solution added to it, gives a colour that is matched against a chart.

Universal indicator is quick, cheap and needs no equipment, but its readings are only estimates, to about $\pm 1$ pH unit, and they depend on judging a colour by eye. Coloured or cloudy solutions are difficult to read. Natural dyes behave in the same way: the juice of red cabbage contains anthocyanin pigments that are red in acid, purple near neutral and green to yellow in alkali.

A pH meter measures pH precisely. Its probe contains a thin glass membrane that develops a voltage depending on $[{{\mathrm{H}{\vphantom{X}}^{+}}}]$ in the solution, and the meter converts the voltage into a pH reading, usually to two decimal places. The meter must first be calibrated with buffer solutions of known pH (for example pH 4.00, 7.00 and 10.00), because the probe's response drifts with time and temperature. Connected to a data logger, a pH probe can record readings continuously, which is how the pH curves of titrations are measured.

A digital probe is the better choice when precise values, continuous readings over time or many readings are needed, as in following a titration. A quick test of whether a solution is strongly acidic, roughly neutral or strongly alkaline only needs indicator paper.

The ionization of water and Kw

Even the purest water conducts electricity very slightly, because it contains a small number of ions. Water is amphiprotic, so one water molecule can donate a proton to another:

This autoionization of water happens to a tiny extent. At 298 K pure water contains $1.0\times10^{-7}\ \text{mol dm}^{-3}$ of each ion, while the concentration of water molecules is about $55.5\ \text{mol dm}^{-3}$, so only about two molecules in every thousand million are ionized at any moment.

Because the concentration of water is effectively constant, the equilibrium is described by the product of the two ion concentrations alone. This product is the ion product constant of water.

$K_\text{w}$ applies to every aqueous solution, not only to pure water. Adding an acid increases $[{{\mathrm{H}{\vphantom{X}}^{+}}}]$, and the equilibrium shifts to the left until $[{{\mathrm{OH}{\vphantom{X}}^{-}}}]$ has fallen enough for the product to equal $K_\text{w}$ again. The two concentrations are therefore inversely proportional: if one increases tenfold, the other decreases tenfold. This gives a precise way to classify solutions at 298 K.

Every aqueous solution, even a strongly alkaline one, contains some hydrogen ions, and every acidic solution contains some hydroxide ions. The graph of both concentrations across the pH scale shows how they change together.

The effect of temperature on Kw

The autoionization of water is the reverse of neutralization, which releases energy as ${{\mathrm{O}{-}\mathrm{H}}}$ bonds form. Breaking an ${{\mathrm{O}{-}\mathrm{H}}}$ bond to separate the ions therefore absorbs energy: the forward reaction is endothermic. By Le Châtelier's principle (R2.3), raising the temperature shifts the equilibrium to the right, so more water ionizes and $K_\text{w}$ increases.

Data: J. Clark, chemguide.co.uk, from standard tables of the ionization of water (accessed 2026).

Pure water at 50 °C has a pH of 6.63, but it is not acidic. In pure water each ionization produces one ${{\mathrm{H}{\vphantom{X}}^{+}}}$ and one ${{\mathrm{OH}{\vphantom{X}}^{-}}}$, so $[{{\mathrm{H}{\vphantom{X}}^{+}}}] = [{{\mathrm{OH}{\vphantom{X}}^{-}}}]$ at every temperature, and that equality is what defines a neutral solution. A pH of 7 marks neutrality only at 298 K.

Strong and weak acids and bases

Hydrochloric acid and ethanoic acid of the same concentration behave very differently. The difference lies in how far each acid ionizes when it dissolves in water.

The particle pictures of two acids at the same concentration show the difference. In a strong acid every molecule has given up its proton to water; in a weak acid most molecules remain intact.

The equations show the difference with the arrow. A strong acid is written with a single forward arrow, because the reverse reaction does not occur to a measurable extent. A weak acid is written with an equilibrium arrow:

In $0.100\ \text{mol dm}^{-3}$ ethanoic acid only about 1.3 % of the molecules are ionized at 298 K, so $[{{\mathrm{H}{\vphantom{X}}^{+}}}]$ is about $1.3\times10^{-3}\ \text{mol dm}^{-3}$ and the pH is about 2.9, compared with pH 1.0 for hydrochloric acid of the same concentration.

Sulfuric acid is strong only in its first ionization, ${{\mathrm{H}{\vphantom{X}}_{\smash[t]{2}}\mathrm{SO}{\vphantom{X}}_{\smash[t]{4}}\xrightarrow{}\mathrm{H}{\vphantom{X}}^{+} {}+{} \mathrm{HSO}{\vphantom{X}}_{\smash[t]{4}}{\vphantom{X}}^{-}}}$. The hydrogensulfate ion is a weak acid, so its second proton is released only partially.

Strength and concentration

"Strong" and "weak" describe the extent of ionization, a property of the acid or base itself. "Concentrated" and "dilute" describe how much acid or base is dissolved in a given volume, which the person making the solution chooses. The two properties are independent, so all four combinations are possible.

A dilute solution of a strong acid can have a higher pH than a concentrated solution of a weak acid. Concentrated ethanoic acid ($10\ \text{mol dm}^{-3}$, pH about 1.9) is more acidic than $0.001\ \text{mol dm}^{-3}$ hydrochloric acid (pH 3.0), yet ethanoic acid is still the weak acid.

Distinguishing strong and weak acids by experiment

Solutions of a strong and a weak acid with the same concentration contain different concentrations of ions, so they can be told apart by any property that depends on $[{{\mathrm{H}{\vphantom{X}}^{+}}}]$ or on the total concentration of ions.

One thing that is the same for both: equal volumes of the two acids need exactly the same volume of sodium hydroxide for neutralization, and produce the same final volume of gas with excess magnesium. As the base removes ${{\mathrm{H}{\vphantom{X}}^{+}}}$ ions, the weak acid's equilibrium shifts to the right and releases more, until all of the acid has reacted. The amount of acid, not its strength, decides how much base it neutralizes.

The same tests distinguish strong and weak bases. A $0.10\ \text{mol dm}^{-3}$ sodium hydroxide solution has pH 13 and a high conductivity; $0.10\ \text{mol dm}^{-3}$ aqueous ammonia has pH about 11 and a much lower conductivity.

Conjugate strength and the position of acid–base equilibria

The strength of an acid and the strength of its conjugate base are linked. A strong acid gives up its proton readily, so its conjugate base has very little tendency to take a proton back: the chloride ion, conjugate base of hydrochloric acid, is so weak a base that it does not react with water at all. A weak acid holds on to its proton, so its conjugate base accepts protons readily: the ethanoate ion is a noticeably stronger base than the chloride ion.

The ladder of acids and conjugate bases arranges common pairs in order. The oxonium ion, ${{\mathrm{H}{\vphantom{X}}_{\smash[t]{3}}\mathrm{O}{\vphantom{X}}^{+}}}$, is the strongest acid that can exist in water: any stronger acid gives its proton to water completely, which is why the acids above it in the ladder are all strong.

In any acid–base equilibrium, two acids compete for the proton. The stronger acid gives its proton away more readily than the weaker acid does, so at equilibrium most of the proton-donating species is the weaker acid. An acid–base equilibrium therefore lies on the side of the weaker acid and weaker base. For ethanoic acid in water, the two acids are ${{\mathrm{CH}{\vphantom{X}}_{\smash[t]{3}}\mathrm{COOH}}}$ and ${{\mathrm{H}{\vphantom{X}}_{\smash[t]{3}}\mathrm{O}{\vphantom{X}}^{+}}}$; ${{\mathrm{H}{\vphantom{X}}_{\smash[t]{3}}\mathrm{O}{\vphantom{X}}^{+}}}$ is the stronger, so the equilibrium lies to the left, which is what "partially ionized" means.

For strong acids the equilibrium lies so far to the right that the reaction is effectively complete, and the equilibrium constant is very large; for weak acids it is much smaller than 1 (R2.3). These equilibrium constants, $K_\text{a}$ for acids and $K_\text{b}$ for bases, compare the strengths of weak acids and bases quantitatively.

Why the hydrogen halides get stronger down group 17

Hydrochloric, hydrobromic and hydroiodic acids are all strong, but hydrofluoric acid is weak, even though fluorine is the most electronegative halogen. The extent of ionization depends largely on how easily the ${{\mathrm{H}{-}\mathrm{X}}}$ bond breaks. Bond enthalpies fall steeply down the group, from $567\ \text{kJ mol}^{-1}$ for ${{\mathrm{H}{-}\mathrm{F}}}$ to $431$ for ${{\mathrm{H}{-}\mathrm{Cl}}}$, $366$ for ${{\mathrm{H}{-}\mathrm{Br}}}$ and $298\ \text{kJ mol}^{-1}$ for ${{\mathrm{H}{-}\mathrm{I}}}$ (data booklet), as the halogen atom becomes larger and the bonding electrons are further from its nucleus. The weaker the bond, the more readily the proton is released, so acid strength increases from HF to HI.

Neutralization reactions

When an acid reacts with a base, the hydrogen ions of the acid are removed and the characteristic properties of both are lost. The ionic product of such a reaction is a salt.

In aqueous solution, the reaction between any strong acid and any strong base is the same reaction. The metal ions and the anions of the acid are spectator ions: they are present before and after and take no part. Removing them gives the ionic equation for neutralization:

This is why the enthalpy change of neutralization is almost the same, about $-57\ \text{kJ}$ per mole of water formed, for every strong acid with every strong base. Neutralization is exothermic because a new ${{\mathrm{O}{-}\mathrm{H}}}$ bond forms between the hydroxide ion and the proton, releasing energy, while no bonds need to be broken. Calorimetry of neutralization reactions is described in R1.1.

Neutralization can also be followed by measuring the electrical conductivity of the mixture, because the reaction changes the ions present. When dilute sulfuric acid is added to barium hydroxide solution, the conductivity falls steadily and reaches a minimum close to zero at the equivalence point: the ${{\mathrm{H}{\vphantom{X}}^{+}}}$ and ${{\mathrm{OH}{\vphantom{X}}^{-}}}$ ions combine to form water, and the ${{\mathrm{Ba}{\vphantom{X}}^{2+}}}$ and ${{\mathrm{SO}{\vphantom{X}}_{\smash[t]{4}}{\vphantom{X}}^{2-}}}$ ions are removed as a white precipitate of barium sulfate. Adding more acid then raises the conductivity again.

Types of neutralization reaction

Acids react in characteristic ways with each type of base. Each reaction is a proton transfer to the base: to the oxide ion, the hydroxide ion, the carbonate or hydrogencarbonate ion, or the lone pair on the nitrogen of ammonia or an amine.

The observations follow from the products. Black copper(II) oxide dissolves in warm dilute acid to give a blue solution of a copper(II) salt. Carbonates and hydrogencarbonates fizz as carbon dioxide is released; the gas turns limewater milky. The reaction of ammonia or an amine produces no water, because the base only accepts a proton: ${{\mathrm{NH}{\vphantom{X}}_{\smash[t]{3}}\mskip2mu (\mathrm{aq}) {}+{} \mathrm{H}{\vphantom{X}}^{+}\mskip2mu (\mathrm{aq})\xrightarrow{}\mathrm{NH}{\vphantom{X}}_{\smash[t]{4}}{\vphantom{X}}^{+}\mskip2mu (\mathrm{aq})}}$.

Each type of reaction has its own ionic equation. For an insoluble metal oxide or carbonate the solid appears in full, because it is not dissolved as separate ions:

Organic acids are neutralized in the same ways. Ethanoic acid reacts with sodium hydroxide to give sodium ethanoate, and with sodium carbonate to give sodium ethanoate, water and carbon dioxide. Because ethanoic acid is weak, its ionic equation shows the un-ionized molecule: