Skip to main content
ScotlandChemistrySyllabus dot point

How does the position of an element in the periodic table explain its properties?

The arrangement of elements in the periodic table by atomic number into groups and periods, the link between electron arrangement and chemical behaviour, and the meaning of covalent radius, ionisation energy and electronegativity.

An SQA Higher Chemistry answer on periodicity, covering how elements are arranged by atomic number into groups and periods, how electron arrangement explains chemical behaviour, and the three trends of covalent radius, ionisation energy and electronegativity.

Generated by Claude Opus 4.810 min answer

Reviewed by: AI editorial process; not yet individually human-reviewed

Have a quick question? Jump to the Q&A page

Jump to a section
  1. What this key area is asking
  2. Arranging the elements
  3. Covalent radius
  4. Ionisation energy
  5. Worked example: ionising a sample
  6. Electronegativity
  7. Examples in context
  8. Try this

What this key area is asking

The SQA wants you to describe how the periodic table arranges elements by atomic number into groups and periods, explain how an element's electron arrangement governs its chemistry, and define the three periodic trends you must know: covalent radius, ionisation energy and electronegativity. This key area pairs closely with periodic trends, and the ionisation-energy definition is examined as a written equation almost every year.

Arranging the elements

An element's chemistry is set by its electron arrangement, especially the outer electrons. Elements in Group 1 each have one outer electron and react similarly; the noble gases in Group 0 have full outer shells and are unreactive.

Covalent radius

The covalent radius is a measure of the size of an atom.

  • Across a period it decreases, because the nuclear charge increases while electrons are added to the same shell, pulling the outer electrons closer.
  • Down a group it increases, because each element has an extra occupied shell, so the outer electrons are further from the nucleus.

Ionisation energy

  • Across a period the first ionisation energy generally increases, because the rising nuclear charge holds the outer electrons more tightly.
  • Down a group it decreases, because the outer electron is in a higher shell, further from the nucleus and screened by more inner electrons.

The energy to remove successive electrons rises, and the second ionisation energy is always larger than the first because an electron is being pulled from a positive ion.

Worked example: ionising a sample

Electronegativity

Electronegativity differences explain bond polarity, which you meet in the structure and bonding key area: a large difference gives a polar covalent bond, and a very large difference gives ionic bonding.

Examples in context

Periodicity underpins how chemists predict the behaviour of unfamiliar elements. When a new Group 1 metal such as francium is considered, its place at the foot of the group lets chemists predict it should be the most reactive metal of all, with the lowest ionisation energy, because the single outer electron is the furthest from the nucleus and the most screened. In the semiconductor industry, the position of silicon in the same group as carbon explains why it forms four covalent bonds and a giant network, making it ideal for the lattice of a microchip. The regular march of electronegativity across period 3, from metallic sodium to non-metallic chlorine, is mirrored by the change from ionic to covalent character in their oxides and chlorides.

Try this

Q1. State and explain the trend in covalent radius across a period. [2 marks]

  • Cue. It decreases because the increasing nuclear charge pulls the electrons in the same shell closer.

Q2. Why does the first ionisation energy decrease down Group 1? [2 marks]

  • Cue. The outer electron is in a higher shell, further from the nucleus and more screened, so it is easier to remove.

Q3. The first ionisation energy of lithium is 526 kJ mol−1526 \text{ kJ mol}^{-1}. Calculate the energy to ionise 0.400 mol0.400 \text{ mol} of gaseous lithium atoms. [2 marks]

  • Cue. E=0.400×526=210 kJE = 0.400 \times 526 = 210 \text{ kJ}.

Exam-style practice questions

Practice questions written in the style of SQA exam questions on this dot point, with worked answer explainers. The year tag is the paper they imitate, not the source.

SQA Higher 20173 marksThe first ionisation energy of sodium is 502 kJ mol−1502 \text{ kJ mol}^{-1}. (a) Write the equation, including state symbols, that represents the first ionisation energy of sodium. (b) Calculate the energy required to ionise 0.300 mol0.300 \text{ mol} of gaseous sodium atoms.
Show worked answer →

Part (a) tests the formal definition; part (b) is a straightforward scaling calculation.

(a) The first ionisation energy is for one mole of gaseous atoms losing one mole of electrons:

Na(g)→Na+(g)+e−Na(g) \rightarrow Na^+(g) + e^-

(b) The quoted value is per mole, so for 0.300 mol0.300 \text{ mol}:

E=0.300×502=151 kJE = 0.300 \times 502 = 151 \text{ kJ}

Markers reward the state symbols (a common loss) and correct rounding. Quoting the answer without a unit also loses a mark.

SQA Higher 20203 marksState and explain the trend in electronegativity (a) across period 3 from sodium to chlorine and (b) down Group 7 from fluorine to iodine.
Show worked answer →

Each trend needs the direction and a reason linked to nuclear charge, size and screening.

(a) Across period 3 the electronegativity increases. The nuclear charge rises while electrons fill the same shell, so the atom gets smaller and attracts the bonding electrons more strongly.

(b) Down Group 7 the electronegativity decreases. Each element has an extra occupied shell and more screening, so the bonding electrons are held further from the nucleus and feel less attraction.

Markers want the direction stated and then explained; just naming the direction earns at most one mark. Fluorine being the most electronegative element is a useful anchor point.

Related dot points

Sources & how we know this