S3.1 The periodic table: classification of elements

The layout of the periodic table

In the modern periodic table the elements are arranged in order of increasing atomic number, $Z$, the number of protons in the nucleus. The table is printed in the data booklet, with the atomic number and the relative atomic mass of every element.

Each period begins with a reactive metal and ends with an unreactive noble gas. Period 1 has only two elements (hydrogen and helium), periods 2 and 3 have eight each, periods 4 and 5 have eighteen each, and periods 6 and 7 have thirty-two each. The lengths 2, 8, 18 and 32 come directly from the number of electrons that the sublevels being filled in each period can hold.

Elements in the same group have similar chemical properties. Several groups have names that are used throughout chemistry:

The two rows printed separately under the main table hold the f-block elements cerium to lutetium and thorium to lawrencium. Together with lanthanum and actinium, which the data booklet places in group 3, they make up the lanthanoids and the actinoids. They belong in periods 6 and 7 and are printed separately only so that the table fits on a page. Hydrogen is placed at the top of group 1 because its atom has one outer electron, but it is a non-metal and is not an alkali metal.

The four blocks

The periodic table is also divided into four blocks, named after the sublevel that receives the last electron when the atom’s electrons are added in Aufbau order. The block of an element therefore depends on its electron configuration, not on its group number or its reactivity.

The widths of the blocks match the capacities of the sublevels. An s sublevel holds 2 electrons, so the s block is two columns wide; a p sublevel holds 6 (six columns), a d sublevel 10 (ten columns) and an f sublevel 14 (fourteen columns). Helium ($1\text{s}^2$) belongs to the s block, but it is placed in group 18 with the other noble gases because it has a full outer shell and very similar properties.

Metals, metalloids and non-metals

More than three-quarters of the elements are metals. They fill the left-hand side and centre of the table, including the whole of the d and f blocks. The non-metals occupy the upper right, with hydrogen as the one non-metal on the left. Between them, along a diagonal "staircase" running from boron down to tellurium, lie the metalloids.

Silicon is a typical metalloid. It is shiny and grey like a metal, but it is brittle and has a giant covalent structure. It conducts electricity only slightly at room temperature, and its conductivity increases as the temperature rises; this makes silicon and germanium semiconductors, the materials of computer chips and solar cells. Chemically, silicon forms an acidic oxide, ${{\mathrm{SiO}{\vphantom{X}}_{\smash[t]{2}}}}$, like a non-metal.

The physical properties in the table come from bonding. Metals contain a lattice of cations in a sea of delocalized electrons, which carry charge and heat and let layers of cations slide over each other. Non-metal elements are covalently bonded, either into small molecules held by weak intermolecular forces or into giant covalent structures, so their electrons are not free to move.

Electron configuration and position in the table

The periodic table is a map of electron configurations. Two simple rules link an element’s position to the arrangement of its electrons.

Valence electrons are the electrons in the outer main energy level, the ones involved in bonding. Sodium is in period 3, group 1, so its outer electrons are in the third main level and it has one valence electron: $1\text{s}^2\,2\text{s}^2\,2\text{p}^6\,3\text{s}^1$. Chlorine is in period 3, group 17, so it has seven valence electrons in the third level: $[\text{Ne}]\,3\text{s}^2\,3\text{p}^5$. Elements of the same group have the same number of valence electrons in the same type of sublevel, which is why their chemistry is similar.

Reading across a period from left to right follows the order in which sublevels fill. The map of sublevels for periods 1–4 shows where each sublevel lies.

The 3d sublevel fills in period 4, after 4s, even though its main energy level is 3. This is why the d block is always one main level behind the period number: the first row of the d block (scandium to zinc) is in period 4 but fills the 3d sublevel. The 4f sublevel likewise fills in period 6, in the lanthanoids.

The reverse problem uses the same rules. The configuration $[\text{Ne}]\,3\text{s}^2\,3\text{p}^3$ has its highest occupied level $n = 3$, so the element is in period 3; it has five valence electrons, with the last one in a p sublevel, so it is in group 15 of the p block. The element is phosphorus.

Chromium and copper are the two exceptions up to $Z = 36$: their atoms are $[\text{Ar}]\,3\text{d}^{5}\,4\text{s}^{1}$ and $[\text{Ar}]\,3\text{d}^{10}\,4\text{s}^{1}$ rather than the configurations the filling order predicts. Electron configurations, including these two, are built up in full with the Aufbau principle, Hund’s rule and orbital diagrams in the study of electron configurations.

Periodicity and the prediction of new elements

The periodic table was built from evidence before anything was known about electrons. In 1869 Dmitri Mendeleev arranged the 63 elements then known in order of atomic mass, placing elements with similar properties in the same column. Where the pattern demanded it, he left gaps for elements not yet discovered and changed the order of some pairs. Using the trends in the columns and rows around each gap, in 1871 he predicted the properties of the missing elements.

The element below silicon, which he called eka-silicon, was discovered by Clemens Winkler in 1886 and named germanium. Its measured properties were remarkably close to the predictions.

Predicted values from Mendeleev’s published predictions; measured values for germanium from standard data tables.

The success of these predictions, and of his prediction of gallium (eka-aluminium, discovered 1875), persuaded chemists that periodicity is a real property of the elements and not a coincidence. A theory that makes testable predictions which are then confirmed is far stronger than one that only describes what is already known. The ordering problems in Mendeleev’s table, such as tellurium having a greater atomic mass than iodine but needing to come before it, were later resolved when Henry Moseley showed in 1913 that the true ordering property is atomic number. The same predictive use of the table continues: the superheavy elements 113–118, made in nuclear reactions and the last of them named in 2016, completed period 7 in the places the table had reserved for them.

Explaining periodic trends: nuclear charge and shielding

Almost every trend on the periodic table comes from the attraction between the nucleus and the outer electrons. Two factors control how strongly an outer electron is held.

• Nuclear charge: the total positive charge of the nucleus, equal to the number of protons, $Z$. A greater nuclear charge attracts the outer electrons more strongly.

• Shielding: the inner electrons lie between the nucleus and the outer electrons and repel the outer electrons, so the outer electrons feel less than the full nuclear charge. Electrons in inner main levels shield effectively; electrons in the same main level shield each other only slightly.

The distance of the outer electrons from the nucleus also matters, because electrostatic attraction weakens rapidly with distance. Electrons in a higher main level are on average further from the nucleus.

Across a period, protons are added to the nucleus one at a time, while the extra electrons go into the same main energy level. The number of inner (shielding) electrons stays the same, so the attraction felt by the outer electrons increases steadily from left to right. Down a group, each element has one more occupied main energy level than the one above. The outer electrons are further from the nucleus and are shielded by an extra full level of inner electrons. Although the nuclear charge increases down a group, the extra distance and shielding outweigh it, and the outer electrons are held less strongly.

Atomic radius

An atom has no sharp edge, so its size is measured from the distance between bonded nuclei. The data booklet lists atomic radii in units of $10^{-12}\ \text{m}$ (picometres). The graph of atomic radius against atomic number shows a clear repeating pattern.

Atomic radius decreases across a period. The nuclear charge increases by one at each step, while the electrons are added to the same main energy level and shielding hardly changes. The outer electrons are pulled closer to the nucleus, so the atom gets smaller. Across period 3 the radius falls from $160 \times 10^{-12}\ \text{m}$ for sodium to $100 \times 10^{-12}\ \text{m}$ for chlorine.

Atomic radius increases down a group. Each element down a group has its outer electrons in a higher main energy level, further from the nucleus and shielded by more inner electrons. Down group 1 the radius increases from lithium (130) through sodium (160) and potassium (200) to caesium ($238 \times 10^{-12}\ \text{m}$).

Across the first row of the d block the radius changes less from one element to the next than it does across the s and p blocks. The electrons being added go into the inner 3d sublevel, where they shield the outer 4s electrons fairly well from the extra nuclear charge.

Ionic radius

The ionic radius of an element differs from its atomic radius, because forming an ion changes the balance between protons and electrons.

• Cations are smaller than their atoms. A metal atom such as sodium loses its outer electron, usually emptying its outer main level completely. The ion ${{\mathrm{Na}{\vphantom{X}}^{+}}}$ ($1\text{s}^2\,2\text{s}^2\,2\text{p}^6$) has one main level fewer, and its 11 protons now attract only 10 electrons, so the remaining electrons are pulled in more tightly. The radius falls from $160$ to $102 \times 10^{-12}\ \text{m}$.

• Anions are larger than their atoms. A non-metal atom such as chlorine gains an electron into its outer level. The nuclear charge is unchanged but there are more electrons, so electron–electron repulsion increases and the electron cloud expands. The radius rises from $100$ for Cl to $181 \times 10^{-12}\ \text{m}$ for ${{\mathrm{Cl}{\vphantom{X}}^{-}}}$

Species that have the same number of electrons are isoelectronic. In an isoelectronic series the radius depends only on the nuclear charge: the more protons, the more strongly the same number of electrons is attracted and the smaller the ion. The ions ${{\mathrm{N}{\vphantom{X}}^{3-}}}$, ${{\mathrm{O}{\vphantom{X}}^{2-}}}$, ${{\mathrm{F}{\vphantom{X}}^{-}}}$, ${{\mathrm{Na}{\vphantom{X}}^{+}}}$, ${{\mathrm{Mg}{\vphantom{X}}^{2+}}}$ and ${{\mathrm{Al}{\vphantom{X}}^{3+}}}$ all have the configuration of neon.

Across period 3 the ionic radius therefore falls from ${{\mathrm{Na}{\vphantom{X}}^{+}}}$ (102) to ${{\mathrm{Mg}{\vphantom{X}}^{2+}}}$ (72) to ${{\mathrm{Al}{\vphantom{X}}^{3+}}}$ ($54 \times 10^{-12}\ \text{m}$): these cations are isoelectronic, and the nuclear charge increases. At group 15 the radius jumps sharply, because the anions ${{\mathrm{P}{\vphantom{X}}^{3-}}}$, ${{\mathrm{S}{\vphantom{X}}^{2-}}}$ and ${{\mathrm{Cl}{\vphantom{X}}^{-}}}$ have one more occupied main energy level than the cations (the configuration of argon, not neon). Within the anions the radius then falls again, ${{\mathrm{P}{\vphantom{X}}^{3-}}} > {{\mathrm{S}{\vphantom{X}}^{2-}}} > {{\mathrm{Cl}{\vphantom{X}}^{-}}}$, for the same isoelectronic reason. The data booklet gives $184 \times 10^{-12}\ \text{m}$ for ${{\mathrm{S}{\vphantom{X}}^{2-}}}$ and $181 \times 10^{-12}\ \text{m}$ for ${{\mathrm{Cl}{\vphantom{X}}^{-}}}$.

Down a group the ionic radius increases, just as the atomic radius does, because each ion has one more occupied main energy level: ${{\mathrm{Li}{\vphantom{X}}^{+}}}$ 76, ${{\mathrm{Na}{\vphantom{X}}^{+}}}$ 102, ${{\mathrm{K}{\vphantom{X}}^{+}}}$ 138; ${{\mathrm{F}{\vphantom{X}}^{-}}}$ 133, ${{\mathrm{Cl}{\vphantom{X}}^{-}}}$ 181, ${{\mathrm{Br}{\vphantom{X}}^{-}}}$ 196, ${{\mathrm{I}{\vphantom{X}}^{-}}}$ $220 \times 10^{-12}\ \text{m}$.

First ionization energy

The first ionization energy measures how strongly an atom holds its outermost electron. It is always endothermic (positive), because energy is needed to pull an electron away from the attraction of the nucleus. The data booklet lists first ionization energies in $\text{kJ mol}^{-1}$.

First ionization energy generally increases across a period. The nuclear charge increases while the electron being removed stays in the same main level, with similar shielding, and the atom gets smaller. The outer electron is held more strongly and more energy is needed to remove it. The lowest value in each period belongs to the group 1 metal and the highest to the noble gas.

First ionization energy decreases down a group. The outer electron is in a higher main energy level, further from the nucleus and more shielded by inner electrons, so it is held less strongly. Down group 1 the value falls from $520\ \text{kJ mol}^{-1}$ for lithium to $376\ \text{kJ mol}^{-1}$ for caesium; down group 18 it falls from $2372\ \text{kJ mol}^{-1}$ for helium to $1170\ \text{kJ mol}^{-1}$ for xenon.

The general rise across a period is not perfectly smooth. In period 3, aluminium has a lower first ionization energy than magnesium and sulfur a lower value than phosphorus, so the correct order for sodium, magnesium and aluminium is $\text{Na} < \text{Al} < \text{Mg}$. These small dips are evidence for sublevels, an HL idea explained by the energy of the sublevel from which the electron is removed.

Electron affinity

For most elements the first electron affinity is exothermic (negative), because the incoming electron is attracted by the nucleus and energy is released when it is captured. The more strongly an atom attracts an extra electron, the more negative its electron affinity. The data booklet gives the value for each element that has one.

Data: IB Chemistry data booklet (2025).

Across a period, electron affinity generally becomes more negative towards group 17. The incoming electron enters the same main level in each atom, but the nuclear charge is greater and the atom smaller, so the electron is attracted more strongly. The halogens release the most energy, because in each period they have the greatest nuclear charge and the smallest atoms that can still accept an electron into their outer p sublevel. Down a group, electron affinity generally becomes less negative: the incoming electron enters a main level further from the nucleus and is more shielded, so less energy is released. Down group 17 it changes from $-349$ (Cl) to $-325$ (Br) to $-295\ \text{kJ mol}^{-1}$ (I).

The data show exceptions to these general trends. Fluorine’s electron affinity ($-328\ \text{kJ mol}^{-1}$) is less negative than chlorine’s ($-349\ \text{kJ mol}^{-1}$): the fluorine atom is so small that the incoming electron is strongly repelled by the electrons already crowded into its 2p sublevel. The group 2 elements, nitrogen and the noble gases have small or positive electron affinities (the booklet gives no value for most of them), because the incoming electron must enter a new sublevel, a half-filled sublevel or a new main level.

Adding a second electron is always endothermic. The second electron affinity of oxygen, ${{\mathrm{O}{\vphantom{X}}^{-}\mskip2mu (\mathrm{g}) {}+{} \mathrm{e}{\vphantom{X}}^{-}\xrightarrow{}\mathrm{O}{\vphantom{X}}^{2-}\mskip2mu (\mathrm{g})}}$, is $+753\ \text{kJ mol}^{-1}$, because the electron is being added to an ion that is already negative and repels it. Energy must be supplied to force it on. Oxide ions are stable in ionic compounds only because the large lattice enthalpy released when they form a crystal more than repays this energy.

Electronegativity

Electronegativity is measured on the Pauling scale, which has no units. Fluorine is the most electronegative element (4.0), and the values decrease towards the bottom left of the table, where caesium has 0.8. Most noble gases have no value, because they form few or no bonds.

Across a period electronegativity increases. The nuclear charge rises while the bonding electrons stay in the same main level with similar shielding, and the atom gets smaller, so the shared pair is closer to a more positive nucleus and is attracted more strongly. Down a group electronegativity decreases: the bonding pair is further from the nucleus and more shielded by inner electrons. Down group 17 it falls from F 4.0 to Cl 3.2, Br 3.0 and I 2.7 (data booklet).

Differences in electronegativity decide the polarity of a bond. Because the halogen’s electronegativity falls down group 17 while carbon’s stays at 2.6, the carbon–halogen bond becomes less polar from C–F to C–I. In C–I the two electronegativities (2.6 and 2.7) are almost equal.

Summary of periodic trends

Metallic and non-metallic character

Metallic character describes how readily an element shows metallic behaviour: losing electrons to form cations, forming basic oxides, and having metallic bonding. Because ionization energy decreases down a group and increases across a period, metallic character increases down a group and decreases across a period. The most metallic elements are at the bottom left of the table and the most non-metallic at the top right (excluding the noble gases).

Group 14 shows the whole range in a single column: carbon is a non-metal, silicon and germanium are metalloids, and tin and lead are metals. Lead, at the bottom, has the most metallic character.

Group 1: the alkali metals

The alkali metals lithium, sodium, potassium, rubidium and caesium (and the radioactive francium) all have one valence electron in an s orbital. They lose this electron easily to form ${{\mathrm{M}{\vphantom{X}}^{+}}}$ ions, so they are among the most reactive metals and are stored under oil to keep them away from air and water.

Physical properties

The alkali metals are soft (they can be cut with a knife) and have low densities: lithium, sodium and potassium float on water. Their melting points are low for metals and decrease down the group. Each atom contributes only one delocalized electron to the metallic bonding, and the cations get larger down the group, so the attraction between the cations and the sea of delocalized electrons gets weaker.

Data: IB Chemistry data booklet (2025).

Reaction with water

Every alkali metal reacts with water to form a metal hydroxide and hydrogen gas. The general equation, and the equation for sodium, are:

The hydroxide dissolves to give an alkaline solution (pH above 7), which is how the group gets its name. The reaction is a redox reaction: the metal is oxidized from 0 to +1, losing its valence electron, and hydrogen in water is reduced from +1 to 0.

The reaction becomes more vigorous down the group:

In every case the solution turns universal indicator blue or purple and phenolphthalein pink, and the temperature rises: the reactions are exothermic. Rubidium and caesium react so violently that their reactions are studied with video or computer simulations rather than in the laboratory, and the trend is extended to them by prediction.

Reactivity increases down group 1. Every alkali metal reacts by losing its single valence electron. Down the group the atoms get larger and the valence electron is further from the nucleus and more shielded by inner electrons, so the first ionization energy falls and the electron is lost more easily. Caesium is therefore the most reactive of the stable alkali metals. This increase in reactivity is the increase in metallic character down a group.

Group 17: the halogens

The halogens fluorine, chlorine, bromine and iodine (and the radioactive astatine) have seven valence electrons, $n\text{s}^2\,n\text{p}^5$. They exist as diatomic molecules, ${{\mathrm{X}{\vphantom{X}}_{\smash[t]{2}}}}$, with a single covalent bond, and they react by gaining one electron to form halide ions, ${{\mathrm{X}{\vphantom{X}}^{-}}}$, or by sharing one electron in a covalent bond.

Physical properties

Melting and boiling points: IB Chemistry data booklet (2025).

Melting and boiling points increase down the group. The halogen molecules are non-polar and are held together only by London (dispersion) forces. Down the group the molecules have more electrons and larger, more easily polarized electron clouds, so the London forces get stronger and more energy is needed to separate the molecules.

Reactivity and oxidizing strength

Reactivity decreases down group 17. A halogen reacts by gaining an electron, so the strongest attraction for an extra electron gives the most reactive halogen. Down the group the atoms get larger, and the incoming electron enters a main level further from the nucleus and more shielded by inner electrons. It is attracted less strongly, so the halogen is less able to gain electrons and is less reactive. Fluorine is the most reactive of all the non-metals. This decrease in reactivity is the decrease in non-metallic character down a group.

Because halogens react by gaining electrons, they are oxidizing agents: they remove electrons from other species. Their oxidizing strength decreases down the group, ${{\mathrm{F}{\vphantom{X}}_{\smash[t]{2}}}} > {{\mathrm{Cl}{\vphantom{X}}_{\smash[t]{2}}}} > {{\mathrm{Br}{\vphantom{X}}_{\smash[t]{2}}}} > {{\mathrm{I}{\vphantom{X}}_{\smash[t]{2}}}}$.

Displacement reactions with halide ions

The order of reactivity can be shown in the laboratory by adding aqueous solutions of the halogens (chlorine water, bromine water and iodine solution) to solutions of the potassium halides. A more reactive halogen oxidizes the halide ions of a less reactive halogen, and the less reactive halogen is released. This is a displacement reaction.

In each reaction the halogen gains electrons (it is reduced, from 0 to −1) and acts as the oxidizing agent, while the halide ion loses electrons (it is oxidized, from −1 to 0) and acts as the reducing agent. The potassium ions are spectator ions and are left out of the ionic equations. No reaction occurs when a less reactive halogen is added to the halide ion of a more reactive one: bromine does not react with chloride ions, and iodine reacts with neither chloride nor bromide ions.

The colours of bromine and iodine in water are similar at low concentration. Shaking the mixture with a non-polar solvent such as cyclohexane makes the product easier to identify: the halogen dissolves in the upper organic layer, where bromine is orange and iodine is a distinctive purple.

Oxides: from basic to acidic

The elements do not divide sharply into metals and non-metals; their properties form a continuum across the table. This continuum is clearest in the oxides. Metal oxides are basic, non-metal oxides are acidic, and between them lie oxides that are amphoteric.

The change from basic to acidic oxides follows the change in bonding. The metal oxides of groups 1 and 2 are ionic and contain oxide ions, ${{\mathrm{O}{\vphantom{X}}^{2-}}}$. An oxide ion is a strong base: it accepts a proton from a water molecule to form two hydroxide ions, ${{\mathrm{O}{\vphantom{X}}^{2-} {}+{} \mathrm{H}{\vphantom{X}}_{\smash[t]{2}}\mathrm{O}\xrightarrow{}2\,\mathrm{OH}{\vphantom{X}}^{-}}}$, so these oxides give alkaline solutions. Non-metal oxides such as ${{\mathrm{CO}{\vphantom{X}}_{\smash[t]{2}}}}$, ${{\mathrm{SO}{\vphantom{X}}_{\smash[t]{2}}}}$ and ${{\mathrm{SO}{\vphantom{X}}_{\smash[t]{3}}}}$ are covalent molecules. They react with water to form acids, which release hydrogen ions. Aluminium oxide, ${{\mathrm{Al}{\vphantom{X}}_{\smash[t]{2}}\mathrm{O}{\vphantom{X}}_{\smash[t]{3}}}}$, is ionic but its small, highly charged ${{\mathrm{Al}{\vphantom{X}}^{3+}}}$ ion gives it considerable covalent character, and it reacts with both acids and alkalis.

Typical pH values for concentrated solutions; the exact pH depends on concentration.

Silicon dioxide does not dissolve in water, but it is classed as acidic because it reacts with hot concentrated alkalis to form silicates. Aluminium oxide is also insoluble, and its amphoteric nature shows only in its reactions with acids and with alkalis. Aluminium is therefore the one metal in period 3 that does not form a basic oxide.

Equations for oxides reacting with water

The oxides of group 1 metals dissolve in water to form the metal hydroxide. Each ${{\mathrm{M}{\vphantom{X}}_{\smash[t]{2}}\mathrm{O}}}$ gives two ${{\mathrm{MOH}}}$:

The oxides of group 2 metals react to form the hydroxide ${{\mathrm{M}(\mathrm{OH}){\vphantom{X}}_{\smash[t]{2}}}}$. Magnesium hydroxide is only slightly soluble, so the solution is weakly alkaline; calcium hydroxide is more soluble (its solution is limewater).

Carbon dioxide dissolves in water and reacts to a small extent to form carbonic acid, a weak acid. The reaction is reversible:

Sulfur forms two acidic oxides. Sulfur dioxide gives sulfurous acid and sulfur trioxide gives sulfuric acid:

The acid–base nature of oxides is tested by shaking a small amount of each oxide with distilled water and measuring the pH with universal indicator or a pH probe. A pH probe is calibrated with buffer solutions first and rinsed with distilled water between readings. Insoluble oxides give a pH close to 7 and are tested instead by whether they react with dilute hydrochloric acid or dilute sodium hydroxide solution.

Acid deposition

Pure water has a pH of 7, but natural, unpolluted rain is slightly acidic, with a pH of about 5.6. Carbon dioxide in the air dissolves in rainwater to form carbonic acid, which partly dissociates to release hydrogen ions:

Sulfur oxides come mainly from burning coal and other fossil fuels that contain sulfur compounds, and from smelting sulfide ores. The sulfur burns to sulfur dioxide, which forms sulfurous acid in water or is oxidized further in the air to sulfur trioxide and then sulfuric acid:

Nitrogen oxides (${{\mathrm{NO}}}$ and ${{\mathrm{NO}{\vphantom{X}}_{\smash[t]{2}}}}$, together written ${{\mathrm{NO}{\vphantom{X}}_{\smash[t]{x }}}}$) form when nitrogen and oxygen from the air combine at the very high temperatures inside car engines, power stations and furnaces. Nitrogen dioxide then reacts with water and oxygen to form nitric acid:

Acid deposition has several harmful effects:

• it dissolves limestone and marble (${{\mathrm{CaCO}{\vphantom{X}}_{\smash[t]{3}}}}$) in buildings and statues: ${{\mathrm{CaCO}{\vphantom{X}}_{\smash[t]{3}}\mskip2mu (\mathrm{s}) {}+{} \mathrm{H}{\vphantom{X}}_{\smash[t]{2}}\mathrm{SO}{\vphantom{X}}_{\smash[t]{4}}\mskip2mu (\mathrm{aq})\xrightarrow{}\mathrm{CaSO}{\vphantom{X}}_{\smash[t]{4}}\mskip2mu (\mathrm{aq}) {}+{} \mathrm{H}{\vphantom{X}}_{\smash[t]{2}}\mathrm{O}\mskip2mu (\mathrm{l}) {}+{} \mathrm{CO}{\vphantom{X}}_{\smash[t]{2}}\mskip2mu (\mathrm{g})}}$;

• it corrodes metal structures such as bridges and railways;

• it lowers the pH of lakes and rivers, killing fish and other aquatic life, partly by releasing toxic ${{\mathrm{Al}{\vphantom{X}}^{3+}}}$ ions from soils into the water;

• it damages trees directly and leaches nutrients such as calcium and magnesium out of soils.

The causes are reduced at source. Sulfur can be removed from fuels before they are burned, and the waste gases of power stations can be "scrubbed" by flue-gas desulfurization, in which a slurry of calcium oxide or calcium carbonate reacts with the sulfur dioxide. Catalytic converters on cars reduce nitrogen oxides back to nitrogen. Lakes and soils that have already been acidified can be treated with lime (calcium oxide or calcium hydroxide) to neutralize them. These controls have been effective: annual sulfur dioxide emissions from power stations in the United States fell by 96% between 1990 and 2023, largely through desulfurization, low-sulfur fuels and emissions trading.

Ocean acidification

The oceans absorb roughly a quarter of the carbon dioxide released by human activities, mainly the burning of fossil fuels. As the concentration of carbon dioxide in the air rises, more dissolves in seawater, forms carbonic acid and releases hydrogen ions, lowering the pH of the ocean:

Seawater is slightly alkaline, so the oceans are not becoming acidic in the everyday sense; their pH is falling, which is what ocean acidification describes. The extra hydrogen ions react with carbonate ions dissolved in seawater, converting them into hydrogencarbonate ions:

Corals, shellfish and many plankton build their shells and skeletons from calcium carbonate, ${{\mathrm{CaCO}{\vphantom{X}}_{\smash[t]{3}}}}$, using carbonate ions from seawater. As the carbonate concentration falls, shell building becomes harder, and in more acidic water existing calcium carbonate can begin to dissolve. Damage at the base of the food chain and to coral reefs affects whole marine ecosystems. The pH changes behind ocean acidification are treated quantitatively with the pH scale in the study of proton transfer reactions.

Oxidation states

The oxidation state is a bookkeeping device. In an ionic compound such as ${{\mathrm{NaCl}}}$ it is simply the charge on each ion: +1 for sodium and −1 for chlorine. In a covalent compound the electrons are shared, but the oxidation state treats the shared pair as if it belonged entirely to the more electronegative atom. In water, oxygen is more electronegative than hydrogen, so each O–H pair is counted as oxygen’s: oxygen is −2 and each hydrogen +1, although no full charges exist.

Oxidation states are written with the sign first and then the number: +2, −1, +6. The terms "oxidation state" and "oxidation number" are used interchangeably. In names, the oxidation state of an element that can have more than one is shown with Roman numerals in brackets: iron(II) chloride, ${{\mathrm{FeCl}{\vphantom{X}}_{\smash[t]{2}}}}$, and iron(III) chloride, ${{\mathrm{FeCl}{\vphantom{X}}_{\smash[t]{3}}}}$.

The oxidation state of an element is zero

In an uncombined element every atom has an oxidation state of zero. In a molecule such as ${{\mathrm{Cl}{\vphantom{X}}_{\smash[t]{2}}}}$ or ${{\mathrm{O}{\vphantom{X}}_{\smash[t]{2}}}}$ the bonded atoms are identical, so they have the same electronegativity and share the bonding electrons equally; no electrons are transferred towards either atom. In a metal such as solid iron or sodium the atoms are all the same and none has gained or lost electrons to another element, so they are counted as neutral. The same holds for elements made of many atoms, such as ${{\mathrm{S}{\vphantom{X}}_{\smash[t]{8}}}}$, ${{\mathrm{P}{\vphantom{X}}_{\smash[t]{4}}}}$ and diamond.

Rules for assigning oxidation states

The exceptions for hydrogen and oxygen follow from electronegativity. In a metal hydride such as sodium hydride, ${{\mathrm{NaH}}}$, hydrogen is bonded to a metal that is less electronegative than itself, so it holds the bonding electrons and is −1 (the compound contains the hydride ion, ${{\mathrm{H}{\vphantom{X}}^{-}}}$). In a peroxide such as hydrogen peroxide, ${{\mathrm{H}{\vphantom{X}}_{\smash[t]{2}}\mathrm{O}{\vphantom{X}}_{\smash[t]{2}}}}$, the two oxygen atoms are bonded to each other (H–O–O–H); the O–O pair is shared equally, so each oxygen gains only one electron’s worth of charge from its hydrogen and is −1.

Nitrogen shows an even wider range, every value from −3 to +5. The number line of nitrogen compounds and ions places each species at its oxidation state.

The same rules apply to carbon in organic molecules, where hydrogen is +1 and oxygen −2. Carbon is −4 in methane, ${{\mathrm{CH}{\vphantom{X}}_{\smash[t]{4}}}}$, −2 in methanol, ${{\mathrm{CH}{\vphantom{X}}_{\smash[t]{3}}\mathrm{OH}}}$, 0 in methanal, ${{\mathrm{HCHO}}}$, +2 in methanoic acid, ${{\mathrm{HCOOH}}}$, and +4 in carbon dioxide. Each step along this series is an oxidation of the carbon atom, which is how combustion and the oxidation of alcohols are analysed as redox reactions.

Naming compounds and oxyanions

Roman numerals in a name give the oxidation state of the element just before them. Copper(I) oxide is ${{\mathrm{Cu}{\vphantom{X}}_{\smash[t]{2}}\mathrm{O}}}$ (Cu +1) and copper(II) oxide is ${{\mathrm{CuO}}}$ (Cu +2); manganese(IV) oxide is ${{\mathrm{MnO}{\vphantom{X}}_{\smash[t]{2}}}}$. Titanium(II) phosphate is ${{\mathrm{Ti}{\vphantom{X}}_{\smash[t]{3}}(\mathrm{PO}{\vphantom{X}}_{\smash[t]{4}}){\vphantom{X}}_{\smash[t]{2}}}}$: three ${{\mathrm{Ti}{\vphantom{X}}^{2+}}}$ ions balance two ${{\mathrm{PO}{\vphantom{X}}_{\smash[t]{4}}{\vphantom{X}}^{3-}}}$ ions.

Negative ions that contain oxygen, oxyanions, can also be named with oxidation states, but their traditional names are still widely used and accepted:

Oxidation states are the basis of redox chemistry: an increase in oxidation state is an oxidation and a decrease is a reduction. In the reaction of sodium with water, sodium changes from 0 to +1 (oxidized) and hydrogen from +1 to 0 (reduced). Oxidation states are used throughout the study of electron transfer reactions to identify oxidizing and reducing agents and to balance redox equations.

First ionization energy and evidence for sublevels (HL)

The first ionization energy rises across a period because the nuclear charge increases while the electrons enter the same main level. If electrons simply occupied main levels, the rise would be smooth. In fact it falls twice in every period, between groups 2 and 13 and between groups 15 and 16.

From group 2 to group 13

Boron ($801\ \text{kJ mol}^{-1}$) has a lower first ionization energy than beryllium ($900\ \text{kJ mol}^{-1}$), and aluminium ($578$) a lower value than magnesium ($738\ \text{kJ mol}^{-1}$). The electron removed from beryllium is a 2s electron, $1\text{s}^{2}\,2\text{s}^{2}$; the electron removed from boron is a 2p electron, $1\text{s}^{2}\,2\text{s}^{2}\,2\text{p}^{1}$. The 2p sublevel is higher in energy than 2s (and its electron is partly shielded by the 2s electrons), so less energy is needed to remove the 2p electron despite boron’s greater nuclear charge.

This dip is evidence that the second main level is divided into sublevels of different energy: s, which fills first, and p, which is higher in energy.

From group 15 to group 16

Oxygen ($1314\ \text{kJ mol}^{-1}$) has a lower first ionization energy than nitrogen ($1402\ \text{kJ mol}^{-1}$), and sulfur ($1000$) a lower value than phosphorus ($1012\ \text{kJ mol}^{-1}$). In nitrogen, $1\text{s}^{2}\,2\text{s}^{2}\,2\text{p}^{3}$, the three 2p electrons occupy the three 2p orbitals singly, with parallel spins (Hund’s rule). In oxygen, $1\text{s}^{2}\,2\text{s}^{2}\,2\text{p}^{4}$, the fourth 2p electron must pair up with another electron in an orbital that is already occupied. The two electrons in the same orbital repel each other strongly, which raises the energy of the paired electron, so it is easier to remove.

This dip is evidence that the p sublevel is made of three orbitals, each holding a maximum of two electrons, and that electrons occupy the orbitals singly before pairing. The pattern within each period, two values then a dip, three values then a dip, then three more, matches $\text{s}^2$, then $\text{p}^3$ (single electrons in three orbitals), then $\text{p}^3$ (pairing up).

Transition elements (HL)

The first-row d-block elements run from scandium ($Z = 21$) to zinc ($Z = 30$), filling the 3d sublevel. By the definition, titanium to copper are transition elements. Zinc is not: its atom is $[\text{Ar}]\,3\text{d}^{10}\,4\text{s}^{2}$ and its only ion, ${{\mathrm{Zn}{\vphantom{X}}^{2+}}}$, is $[\text{Ar}]\,3\text{d}^{10}$, so it never has an incomplete d sublevel. Zinc compounds are white and zinc shows only one oxidation state, +2. Copper counts because, although its atom is $[\text{Ar}]\,3\text{d}^{10}\,4\text{s}^{1}$, its common ion ${{\mathrm{Cu}{\vphantom{X}}^{2+}}}$ is $[\text{Ar}]\,3\text{d}^{9}$.

The partly filled d sublevel gives the transition elements a set of characteristic properties.