Periodic table

Learning objectives

Objectives and contextOverview

By the end of this chapter you should be able to:

• state the group and period of an element from the periodic table and deduce its number of outer-shell electrons and shells A
• distinguish metals from non-metals by physical and chemical properties A
• describe the characteristic properties of transition metals and noble gases and explain the unreactivity of noble gases A
• describe and explain the trends in reactivity down group 1 and group 17 and across a period A
• predict the properties of an unfamiliar element from its position A
• use data about elements to choose a material for a purpose and to evaluate its impacts C D

In 1869 Dmitri Mendeleev arranged the roughly sixty elements then known in order of increasing atomic mass and noticed that elements with similar chemical properties appeared at regular intervals. He did something bolder than tidy the data: he left gaps. Where no known element fitted the pattern, he assumed an element had not yet been discovered and predicted its properties from its neighbours. When gallium, scandium and germanium were isolated in the following years, their properties matched his predictions closely. A good classification does not just store information — it predicts. That is the idea to carry through this chapter: once you can read an element's position, you can predict how it will behave.

Mendeleev's table differs from the modern one in two instructive ways. Some groups were missing altogether. The noble gases had not been discovered because they are unreactive and are not found in compounds, so there was no chemical evidence for them; the lanthanides and actinides were largely unknown because many exist in only tiny quantities or do not occur naturally. Secondly, the modern table is ordered by atomic number (proton number), not atomic mass. For almost all elements the two orders agree, but ordering by proton number removes the few pairs that Mendeleev had to swap by hand to keep similar elements together.

02 / Structure of the table

Groups, periods and blocksPeriods and groups

The modern periodic table arranges elements in order of increasing atomic number. The horizontal rows are periods; the vertical columns are groups.

The link between position and electrons is the key to the whole topic. Sodium (atomic number 11) has 11 electrons arranged 2,8,1: three occupied shells, so period 3; one outer electron, so group 1. Chlorine (17) is 2,8,7: period 3, group 17 (the old numbering calls this group 7, because it has seven outer electrons). For the main groups the rule is simple:

• the period number equals the number of occupied shells;
• for groups 1 and 2, the group number equals the number of outer electrons; for groups 13–17 the number of outer electrons is the group number minus 10 (13 → 3, 14 → 4, … 17 → 7);
• group 18 atoms have a full outer shell (2 for helium, 8 for the others).

The table is also divided into blocks. Groups 1 and 2 (plus helium) form the s-block, groups 13–18 the p-block, the ten central columns the d-block (the transition metals), and the two separated rows the f-block (lanthanides and actinides). At MYP level you need to recognise the d-block as the transition metals and group 18 as the noble gases; the block letters are useful labels rather than required theory.

03 / Metals and non-metals

Metals and non-metalsMetals and non-metals

About three-quarters of the elements are metals. The difference between a metal and a non-metal is seen at three levels. At the macroscopic level we observe properties: shiny or dull, conducting or insulating, malleable or brittle. At the particle level metals consist of positive ions in a “sea” of delocalised electrons, while non-metals exist as molecules (O₂, Cl₂, S₈), single atoms (noble gases) or giant covalent networks (carbon, silicon). At the symbolic level metals form positive ions (Na⁺, Mg²⁺) and basic oxides, and non-metals form negative ions (Cl⁻, O²⁻) or share electrons, and form acidic oxides.

The properties explain uses. Cooking pans are made from metals because metals conduct heat well, are malleable enough to shape, and have high melting points; marking guidance did not accept “conductivity” without saying thermal, because the relevant property is heat conduction. Gold survives unchanged for thousands of years in ancient mirrors because it is very unreactive, not because it is malleable or conductive — both true, but neither explains the lack of corrosion.

Elements next to the staircase — boron, silicon, germanium, arsenic — are metalloids. They look metallic but are brittle and are semiconductors: they conduct electricity a little, and better as temperature rises. That intermediate behaviour is why silicon, gallium, arsenic and tellurium are used in solar cells and microchips.

04 / Group 1 and group 2

Trends in the metal groupsTrends

Lithium, sodium and potassium are the first three alkali metals (group 1). They are soft enough to cut with a knife, have low densities (lithium, sodium and potassium float on water) and low melting points for metals. They are stored under oil because they react rapidly with oxygen and water.

2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)

Observations when a small piece is dropped into water: the metal floats, fizzes (hydrogen), moves around the surface, and — for sodium — melts into a ball because the reaction is exothermic. Potassium reacts so vigorously that the hydrogen ignites with a lilac flame. The solution becomes alkaline (hydroxide ions), turning universal indicator purple.

Reactivity increases down group 1. Every alkali metal reacts by losing its single outer electron to form a 1+ ion. Going down the group, each atom has one more occupied shell, so the outer electron is further from the nucleus and more shielded by the inner shells. It is attracted less strongly and is lost more easily, so the reaction is faster and more vigorous.

Group 2 (the alkaline earth metals: beryllium, magnesium, calcium, strontium, barium) follows the same pattern with two outer electrons, forming 2+ ions. They are less reactive than the group 1 metal in the same period because two electrons have to be removed. Several group 2 compounds give characteristic flame colours used in fireworks: calcium compounds orange-red, strontium bright red and barium green.

05 / Group 17

The halogensTrends

The halogens (group 17) are non-metals that exist as diatomic molecules: F₂, Cl₂, Br₂, I₂. Down the group the colour darkens (pale yellow fluorine, pale green chlorine, red-brown bromine, grey-black iodine that gives a purple vapour) and the state changes from gas to liquid to solid, because the larger molecules have stronger forces between them.

Reactivity decreases down group 17. A halogen atom reacts by gaining one electron to complete its outer shell. Down the group the outer shell is further from the nucleus and more shielded, so the attraction for an incoming electron is weaker. The trend runs in the opposite direction to group 1 because the process is the opposite — gaining rather than losing an electron.

The trend is demonstrated by displacement: a more reactive halogen displaces a less reactive one from a solution of its salt.

Cl₂(aq) + 2KBr(aq) → 2KCl(aq) + Br₂(aq)    colourless solution turns orange

Bromine will displace iodine from potassium iodide (the solution turns brown), but iodine cannot displace bromine or chlorine.

06 / The d-block

Transition metalsTransition metals

The transition metals occupy the d-block between groups 2 and 13. They include the everyday structural metals — iron, copper, nickel, chromium, zinc (often grouped with them) — and the precious metals gold, silver and platinum. Compared with group 1 metals they share a clear set of properties:

Because a transition metal can form ions with different charges, the charge is shown in a compound's name with a Roman numeral: iron(II) oxide is FeO, iron(III) oxide is Fe₂O₃, cobalt(II) chloride is CoCl₂. Cobalt(II) chloride paper is the test for water: it turns from blue to pink.

The colours of transition-metal compounds have been used for thousands of years. Cobalt compounds produce the deep blue of pottery glazes and glass; oxides of transition metals colour stained glass; copper chloride gives a blue-green flame in fireworks and iron chloride a red-brown one.

07 / Group 18

Noble gasesNoble gases

Helium, neon, argon, krypton, xenon and radon form group 18. They are colourless gases that exist as single atoms (monatomic) and are extremely unreactive. The explanation is electronic: each has a full outer shell — two electrons for helium, eight for the others — so there is no tendency to gain, lose or share electrons. The atoms do not need to bond to become stable, which is also why they do not form diatomic molecules as hydrogen, nitrogen, oxygen and the halogens do.

Their uses follow from this unreactivity and from their low density:

• Argon fills filament light bulbs. The filament is white-hot; in air it would react with oxygen and burn out. Argon is inert, so the hot metal is protected. If air leaks into the bulb, the oxygen in that air reacts with the filament.
• Helium fills balloons and airships because it is much less dense than air and, unlike hydrogen, is not flammable.
• Neon glows red when electricity passes through it at low pressure, which is used in advertising signs.

08 / Across a period

Trends across a periodTrends

Crossing period 3 from sodium to argon, each element has one more proton and one more outer electron than the last, but the same number of shells. The character of the elements changes gradually:

The pattern — metals on the left losing electrons, non-metals on the right gaining or sharing them, a noble gas at the end — repeats in every period. That repetition is what “periodic” means. It explains why carbon and silicon, one above the other in group 14, both form four covalent bonds and have hydrides with the same type of formula, CH₄ and SiH₄ (Topic 6).

09 / Practice

Examination practicePast examination tasks

Use the interactive table above if a question says “use the periodic table”. Attempt each question before opening the marking guidance.

A Criterion A · Knowing and understanding

B Criterion B · Inquiring and designing

None of the available examination tasks assessed criterion B through this topic.

C Criterion C · Processing and evaluating

None of the available examination tasks assessed criterion C through this topic.

D Criterion D · Reflecting on the impacts of science

10 / Examination feedback

Examiner's overall observationEvidence from examination feedback

11 / Summary

Summary and knowledge organiserRevision

Essential knowledge

• Elements are ordered by atomic number. Groups are columns (same outer electrons, similar chemistry); periods are rows (same number of shells).
• Metals: left of the staircase, conduct, malleable, form positive ions and basic oxides. Non-metals: right, insulators, brittle, form negative ions or share electrons, acidic oxides.
• Group 1 reactivity increases down the group (outer electron further away, lost more easily). Group 17 reactivity decreases down the group (weaker attraction for an electron to be gained).
• Transition metals: high melting point and density, variable oxidation states, coloured compounds, catalysts.
• Noble gases: full outer shell → monatomic and unreactive; argon in bulbs, helium in balloons.