Reading the table as a map

The periodic table is more than a list of elements. It works like a map, arranged so that an element’s position tells you something about its chemistry. A period is a horizontal row of elements in the periodic table. A group is a vertical column of elements in the periodic table. Groups are numbered 1 to 18, and in assessments you should use the periodic table in the data booklet as your reference.

A block is a region of the periodic table where the highest-energy electrons are being placed into the same type of sublevel. The four blocks are the s-block, p-block, d-block and f-block, matching the s, p, d and f sublevels. The s-block sits on the left, the p-block on the right. The d-block forms the central block, while the f-block is usually shown separately below the main table to save space.

Metals, non-metals and metalloids

A metal is an element that typically forms giant metallic structures with delocalized electrons and is found on the left and centre of the periodic table. A non-metal is an element that typically forms covalent molecular or network structures, or monatomic gases, and is found on the upper right of the periodic table. Hydrogen is the awkward exception: many tables place it above group 1 because it has one electron, but it is not an alkali metal.

A metalloid is an element with properties intermediate between metals and non-metals, often able to show both metallic and covalent behaviour. Metalloids lie along the stepped boundary between the metals and non-metals, including elements such as boron, silicon and germanium. Don’t treat the boundary as a brick wall; metallic and non-metallic character form a continuum.

Period number and outer energy level

The principal quantum number, n, is a dimensionless number used to label an electron energy level. In main-group elements, the period number shows the outer occupied energy level. Period 3 elements, for instance, have their valence electrons in the third energy level, so those outer electrons occupy 3s and/or 3p sublevels.

Valence electrons are electrons in the outer occupied energy level of an atom that are available for bonding. Elements in the same group share the same pattern of valence electrons. That’s why a group often behaves like a chemical family, not just a vertical column.

You should know these named classifications:

For groups 1 and 2, the group number gives the number of valence electrons. For groups 13–18, use the final digit instead. Group 16 elements therefore have six valence electrons; group 18 elements have eight, except helium, which has two.

Deducing electron configurations up to argon and krypton

The atomic number, Z, is the number of protons in the nucleus of an atom, measured as a dimensionless count. You need to deduce electron configurations up to Z = 36 from an element’s position in the periodic table, and also work backwards from a configuration to its position.

Keep the maximum occupancies ready: s holds 2 electrons, p holds 6, d holds 10 and f holds 14. Up to Z = 36, use this filling order:

1s 2s 2p 3s 3p 4s 3d 4p

Take bromine as an example. It is in period 4 and group 17. Its electrons fill to argon, then 4s², then 3d¹⁰, then 4p⁵, giving:

Br: [Ar] 4s² 3d¹⁰ 4p⁵

Going the other way matters just as much. If an atom has configuration [Ne] 3s² 3p⁴, its outer level is 3, placing it in period 3. It has six valence electrons, so it belongs to group 16: sulfur.

Classification and discovery

The organization of elements has a strong nature-of-science story. Early periodic tables arranged known elements so that similar properties recurred. Gaps were more than blank spaces; they acted as predictions. If an undiscovered element was expected to fit a gap, chemists could estimate its likely properties from the elements around it. That is what a good classification can do: it does not just tidy existing knowledge, it suggests what to look for next.

Classification is also a model, not a law of nature. Hydrogen’s position and the exact membership of group 3 show that a useful classification can still have disputed edges.

What periodicity means

Periodicity is the repeating pattern in the physical and chemical properties of elements when they are arranged by atomic number. The patterns appear because electron configurations repeat in a regular way. When you explain a trend, keep coming back to three ideas: nuclear charge, number of occupied energy levels, and shielding.

Shielding is the reduction in attraction between the nucleus and an outer electron caused by inner electrons repelling that outer electron. Effective nuclear charge is the net positive attraction an electron experiences after shielding has been taken into account.

Atomic radius

Atomic radius is a measure of the size of an atom, usually taken as half the distance between the nuclei of two bonded atoms of the same element. Across a period, atomic radius decreases. Nuclear charge increases because each successive atom has one more proton, but the added electrons enter the same main energy level, so shielding by inner shells changes very little. The nucleus pulls the outer electrons closer.

Down a group, atomic radius increases. Each step down adds a new occupied energy level. The outer electrons sit further from the nucleus and are more shielded, so increased distance and shielding outweigh the increased nuclear charge.

Ionic radius

Ionic radius is a measure of the size of an ion. A cation is a positively charged ion formed when an atom loses electrons. Cations are smaller than their parent atoms because an electron shell may be lost, and the same nucleus attracts the remaining electrons more strongly.

An anion is a negatively charged ion formed when an atom gains electrons. Anions are larger than their parent atoms because the extra electrons increase electron–electron repulsion and reduce the attraction per electron.

For ions with the same electron configuration, nuclear charge has a strong effect on size. An isoelectronic series is a set of atoms or ions with the same number and arrangement of electrons. In an isoelectronic series such as O²⁻, F⁻, Ne, Na⁺ and Mg²⁺, radius decreases as proton number increases, because the same electron cloud is pulled in more strongly.

Ionization energy

First ionization energy is the energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions. It is represented by:

X(g) → X⁺(g) + e⁻

Across a period, first ionization energy generally increases. Nuclear charge increases, shielding changes little, and atomic radius decreases, so removing an outer electron takes more energy. Down a group, first ionization energy decreases because the outer electron is further from the nucleus and more shielded.

Electron affinity

First electron affinity is the energy change when one electron is added to each atom in one mole of gaseous atoms to form one mole of gaseous 1− ions. It is represented by:

X(g) + e⁻ → X⁻(g)

Across a period, electron affinity generally becomes more exothermic because increasing nuclear charge attracts an added electron more strongly. Down a group, the pattern is less tidy than the atomic radius and ionization energy trends. Increased shielding would suggest a less exothermic value, but electron–electron repulsions and sublevel arrangements also matter. In exams, if data are provided, use the data rather than forcing a perfect trend that is not there.

Electronegativity

Electronegativity is a measure of the ability of an atom in a bond to attract the bonding electrons towards itself. Electronegativity increases across a period because nuclear charge increases and atomic radius decreases. It decreases down a group because the bonding electrons are further from the nucleus and more shielded.

Summary of periodic trends across periods and down groups.

These trends link together. Small atoms with high effective nuclear charge tend to have high ionization energies and high electronegativities. Large atoms with well-shielded outer electrons tend to have lower ionization energies and more metallic behaviour.

Group 1: increasing metallic character

Metallic character is the tendency of an element to show metallic behaviour, especially by losing electrons and forming positive ions. Group 1 elements, the alkali metals, each have one valence electron. As you move down the group, first ionization energy decreases because the outer electron is further from the nucleus and is more shielded. The atom loses that electron more easily, so metallic character and reactivity increase.

With water, a group 1 metal produces a metal hydroxide and hydrogen gas. Using M as a general symbol for a group 1 metal, the balanced equation is:

2M(s) + 2H₂O(l) → 2MOH(aq) + H₂(g)

For example:

2K(s) + 2H₂O(l) → 2KOH(aq) + H₂(g)

The solution turns alkaline because hydroxide ions are produced. The reactions get faster and more vigorous down the group: lithium reacts steadily, sodium more vigorously, and potassium more violently. Rubidium and caesium are usually shown by video or simulation rather than in an open classroom.

Group 17: decreasing non-metallic character

Non-metallic character is the tendency of an element to show non-metal behaviour, especially by gaining or attracting electrons. Group 17 elements, the halogens, have seven valence electrons. Down the group, electronegativity decreases as atomic radius and shielding increase. This makes the halogens less effective at gaining an electron, so their reactivity decreases down the group.

A more reactive halogen displaces the halide ion of a less reactive halogen. A halide ion is a 1− ion formed from a halogen atom by gaining one electron. For instance, chlorine displaces bromide ions:

Cl₂(aq) + 2Br⁻(aq) → 2Cl⁻(aq) + Br₂(aq)

Bromine, however, does not displace chloride ions because bromine is less reactive than chlorine:

Br₂(aq) + Cl⁻(aq) → no reaction

The order to remember is:

F₂ > Cl₂ > Br₂ > I₂

So fluorine displaces chloride, bromide and iodide ions; chlorine displaces bromide and iodide ions; bromine displaces iodide ions only; iodine displaces none of the common halides above it.

Why simulations are used

Teachers often use simulations and online reaction videos for group 1 and group 17 reactivity because the real reactions can be hazardous, fast, corrosive or difficult to control fairly. They let students compare trends side by side, slow down observations, repeat conditions, and access reactions that would be unsafe in a school laboratory. The trade-off is that a simulation is a model: it helps with patterns, but it cannot replace careful experimental evidence when practical work is safe and appropriate.

Oxides across the periodic table

An oxide is a compound containing oxygen bonded to one or more other elements. Moving across a period, the oxides change broadly from basic metal oxides, through amphoteric oxides, to acidic non-metal oxides. It’s a tidy bit of evidence that metallic and non-metallic character fades gradually across the table rather than fitting into two fixed boxes.

A basic oxide is an oxide that reacts with acids or water to produce a base. Group 1 metal oxides react with water and form soluble metal hydroxides:

M₂O(s) + H₂O(l) → 2MOH(aq)

For example:

Na₂O(s) + H₂O(l) → 2NaOH(aq)

Group 2 metal oxides react with water to form metal hydroxides:

MO(s) + H₂O(l) → M(OH)₂(aq or s)

For example:

CaO(s) + H₂O(l) → Ca(OH)₂(aq)

Some group 2 hydroxides dissolve only sparingly, so the state symbols can change depending on the metal and the conditions.

An acidic oxide is an oxide that reacts with water to produce an acid, or reacts with bases to form salts. Carbon and sulfur form non-metal oxides that dissolve in water and form acids:

CO₂(g) + H₂O(l) ⇌ H₂CO₃(aq)

SO₂(g) + H₂O(l) → H₂SO₃(aq)

SO₃(g) + H₂O(l) → H₂SO₄(aq)

An amphoteric oxide is an oxide that can react both as an acid and as a base. Aluminium oxide is the standard example: it reacts with acids, and it also reacts with strong bases. For that reason, the period 3 pattern is usually written as basic Na₂O and MgO, amphoteric Al₂O₃, then acidic SiO₂, P₄O₁₀, SO₂ and SO₃.

Period 3 oxides change from basic to amphoteric to acidic across the period.

Bonding link: why metal and non-metal oxides behave differently

Metal oxides usually contain ionic lattices with oxide ions, O²⁻. Oxide ions are strongly basic because they accept protons from water or acids, forming hydroxide ions or water. Non-metal oxides are more often covalent molecules. When these oxides react with water, they form molecules that release H⁺ in aqueous solution, which makes their solutions acidic. The bonding isn’t just extra description here; it explains the difference in their chemical properties.

Acid rain and ocean acidification

Acid rain is precipitation with a pH below the normal value for rainwater, caused by acidic gases dissolving and reacting in atmospheric water. Pure water at 298 K is neutral at pH 7, but normal rain is already slightly acidic because carbon dioxide dissolves to form carbonic acid:

CO₂(g) + H₂O(l) ⇌ H₂CO₃(aq)

Sulfur dioxide and nitrogen oxides make rain more acidic. Sulfur dioxide may come from volcanic activity and from burning sulfur-containing fossil fuels. It dissolves and is oxidized in air, forming stronger acids, including sulfuric acid:

SO₂(g) + H₂O(l) → H₂SO₃(aq)

2H₂SO₃(aq) + O₂(g) → 2H₂SO₄(aq)

Nitrogen dioxide also reacts with water to form acids:

2NO₂(g) + H₂O(l) → HNO₃(aq) + HNO₂(aq)

Acid rain damages limestone buildings, corrodes metals, acidifies lakes and harms plants. Ocean acidification is the decrease in ocean pH caused mainly by increasing atmospheric CO₂ dissolving in seawater and forming carbonic acid. This is especially serious for organisms such as corals and shellfish because carbonate chemistry is disturbed, making calcification more difficult.

What oxidation state means

An oxidation state is a signed number assigned to an atom in a species. It shows the charge that atom would have if all bonds were treated as fully ionic. The term oxidation number means the same thing in this course. Write oxidation states with the sign first, then an Arabic numeral: +2, −1, +7, and so on.

Think of this as bookkeeping. It doesn’t mean every compound is ionic. In water, oxygen is more electronegative than hydrogen, so for oxidation-state purposes the bonding electrons are assigned to oxygen. Oxygen is assigned −2, and each hydrogen is assigned +1.

Rules for deducing oxidation states

Use the rules in a sensible order: fixed rules first, then the sum rule.

1. An atom in a free element has oxidation state 0. This includes single atoms such as Ne, metals such as Mg, and elemental molecules such as H₂, O₂, Cl₂ and S₈.
2. The sum of oxidation states in a neutral compound is 0.
3. The sum of oxidation states in a polyatomic ion equals the charge on the ion.
4. Fluorine has oxidation state −1 in compounds.
5. Group 1 metals have oxidation state +1; group 2 metals have oxidation state +2.
6. Oxygen usually has oxidation state −2. In peroxides such as H₂O₂ or Na₂O₂, oxygen is −1. In OF₂, oxygen is +2 because fluorine is more electronegative.
7. Hydrogen is usually +1 when bonded to non-metals, but −1 in metal hydrides such as NaH and CaH₂.

An element has oxidation state zero because atoms of the same element have the same electronegativity. Take Cl₂. The bonding electrons are shared equally between identical atoms, so neither atom is assigned a gain or loss of electrons.

A clean method

For KMnO₄, put in the known values first: K is +1 and each O is −2. The compound is neutral, so:

(+1) + Mn + 4(−2) = 0

Mn = +7

For SO₄²⁻, oxygen is −2 and the total must equal −2:

S + 4(−2) = −2

S = +6

For NaH, sodium is +1, so hydrogen must be −1. This is the metal hydride exception students often miss.

For Na₂O₂, sodium is +1. The two sodium ions give +2 overall, so the two oxygen atoms must give −2 overall. Each oxygen is −1 because this is a peroxide.

Oxyanion names

An oxyanion is a polyatomic anion containing oxygen and at least one other element. Systematic names often use Roman numerals in brackets to show the oxidation state of the non-oxygen atom. For example, SO₄²⁻ contains sulfur in oxidation state +6, so it may be named sulfate(VI). SO₃²⁻ contains sulfur in oxidation state +4, so it may be named sulfate(IV).

Common names are also accepted for familiar oxyanions:

Oxidation states are especially useful in redox chemistry because rises and falls in oxidation state show where oxidation and reduction have occurred. That leads into analysing redox reactions later.

Why the graph is not a smooth climb

Across a period, first ionization energy generally rises because nuclear charge increases, while shielding by inner shells changes little. The graph still has two clear dips: from group 2 to group 13, and from group 15 to group 16. Those breaks show that electrons occupy sublevels and orbitals with different energies.

Group 2 to group 13: s to p

Compare beryllium and boron:

Be: 1s² 2s²

B: 1s² 2s² 2p¹

Boron loses a 2p electron. A 2p electron is higher in energy than a 2s electron, and the 2s electrons shield it slightly more from the nucleus. So although boron has a greater nuclear charge than beryllium, it takes less energy to remove boron’s outer electron. The drop shows that the p sublevel is higher in energy than the s sublevel in the same main energy level.

Group 15 to group 16: paired-electron repulsion

Now look at nitrogen and oxygen:

N: 1s² 2s² 2p³

O: 1s² 2s² 2p⁴

In nitrogen, the three 2p electrons sit in separate p orbitals. In oxygen, one 2p orbital has a pair of electrons. The electron removed from oxygen is one of that pair, and paired electrons repel each other because they occupy the same region of space. That repulsion increases the energy of the electron being removed, so less energy is needed than expected.

Watch the wording here. The stronger explanation is not the slogan “half-filled sublevels are specially stable”. Focus on the energy of the electron removed: higher-energy p electrons, and higher-energy paired electrons, are easier to remove.

What counts as a transition element

A transition element is an element whose atoms have an incomplete d-sublevel, or one that forms at least one stable ion with an incomplete d-sublevel. So, not every d-block element qualifies as a transition element.

Zinc is the standard comparison. A zinc atom is [Ar] 4s² 3d¹⁰, and the common Zn²⁺ ion is [Ar] 3d¹⁰. Its d-sublevel is full in the atom and still full in the ion, so zinc doesn’t meet the definition. Scandium is less clear-cut: Sc³⁺ is [Ar], with no d electron, but compounds containing Sc²⁺ with a 3d¹ configuration are known. You can argue for including scandium because it can form an ion with an incomplete d-sublevel; you can argue against it because its overwhelmingly common chemistry is Sc³⁺, which has no d electrons.

Characteristic properties

Transition elements commonly show these properties:

• Variable oxidation state: they can form ions with different charges because 4s and 3d electrons are relatively close in energy.
• High melting points: strong metallic bonding comes from many delocalized electrons and small, highly charged metal ions.
• Magnetic properties: many have unpaired d electrons. You don’t need to learn types of magnetism, but you should know that unpaired electrons are the key idea.
• Catalytic properties: they often provide surfaces or variable oxidation states, giving reactions alternative pathways.
• Formation of coloured compounds: incomplete d-sublevels allow electronic transitions involving visible light.
• Formation of complex ions with ligands: transition metal ions can accept electron pairs from surrounding ions or molecules.

A catalyst is a species that increases the rate of a reaction by providing an alternative pathway with lower activation energy, without being consumed overall. A heterogeneous catalyst is a catalyst in a different physical phase from the reactants. Transition metals often work well as heterogeneous catalysts because reactant molecules can adsorb onto their surfaces, react, and then desorb.

Catalytic converters in vehicles use transition metals such as platinum, palladium and rhodium to help convert harmful exhaust gases into less harmful products. You don’t need to memorize the industrial details here; for classification, the key point is that incomplete d-sublevels help produce a recognizable family of properties.

Electron configurations of first-row transition metal ions

The first-row transition elements are the period 4 d-block elements from scandium to zinc. In neutral atoms, the 4s sublevel fills before 3d, so the usual pattern is [Ar] 4s² 3dˣ. Watch for the two standard exceptions:

Cr: [Ar] 4s¹ 3d⁵

Cu: [Ar] 4s¹ 3d¹⁰

When ions form, the 4s electrons come off before the 3d electrons. It can seem the wrong way round, since 4s fills first, but after the 3d sublevel has electrons in it, 4s is the sublevel that loses electrons first.

Examples:

Why variable oxidation states are possible

Successive ionization energies are the energies needed to remove electrons one at a time from the same species. Transition elements can often lose different numbers of electrons because the ionization energies for removing 4s electrons and several 3d electrons are fairly close. Several cations can therefore be energetically accessible, especially when forming a compound releases energy through bonding or lattice formation.

Chromium shows this well. It can use its 4s electron and several 3d electrons in bonding, so oxidation states from +1 to +6 are known. Copper has fewer common oxidation states because its successive ionization energies rise more sharply at an earlier stage; Cu⁺ and Cu²⁺ are much more common than higher states.

The general pattern you should be able to apply is:

1. Write the neutral atom configuration, remembering Cr and Cu.
2. Remove 4s electrons first.
3. Remove 3d electrons next.
4. Relate the number of electrons removed to the oxidation state of the ion.

For example, vanadium is [Ar] 4s² 3d³. If it forms V⁵⁺, all five 4s and 3d electrons have been removed, giving [Ar]. That is why +5 is a plausible high oxidation state for vanadium compounds.