The periodic table arranges all known elements by increasing atomic number, grouping elements with similar electron configurations and chemical properties in the same column. Understanding the trends in atomic radius, ionisation energy and electronegativity allows you to predict how elements will react.
12.1 The Periodic Table and Atomic Trends
Definition
Atomic radius
The atomic radius of an element is defined as half the distance between the nuclei of two identical bonded atoms. Atomic radius increases down a group (more electron shells) and decreases across a period from left to right (increasing nuclear charge pulls electrons closer).
Definition
Ionisation energy (IE)
The first ionisation energy of an element is the minimum energy required to remove one electron from a neutral atom in the gaseous state. Ionisation energy increases across a period (stronger nuclear attraction) and decreases down a group (outer electrons are further from the nucleus and more shielded).
Definition
Electronegativity
Electronegativity is the ability of an atom in a molecule to attract the shared electrons of a covalent bond towards itself. Electronegativity increases across a period and decreases down a group. Fluorine (F) has the highest electronegativity value of 4,0.
Definition
Electron affinity
Electron affinity is the energy released when an electron is added to a neutral atom in the gaseous state to form a negative ion. Elements with high electronegativity generally have high electron affinity.
Periodic Trends — Summary
| Property | Down a Group | Across a Period (left → right) |
|---|---|---|
| Atomic radius | Increases (more shells added) | Decreases (more protons, same shell — stronger pull) |
| Ionisation energy | Decreases (outer e⁻ further, more shielded) | Increases (harder to remove from larger nuclear charge) |
| Electronegativity | Decreases (less attraction for bonding e⁻) | Increases (more nuclear pull on shared e⁻) |
| Metallic character | Increases (elements become more reactive metals) | Decreases (moves from metals → metalloids → non-metals) |
WHY do these trends occur? Down a group: each successive element has one more electron shell. The outermost electrons are further from the nucleus and shielded by the inner shells, so the nuclear attraction is weaker. This makes the atom larger, easier to ionise (lower IE) and less electronegative. Across a period: the number of protons increases while electrons are added to the SAME shell. More protons mean stronger nuclear charge, which pulls the outer electrons closer (smaller radius), holds them more tightly (higher IE) and attracts bonding electrons more strongly (higher electronegativity).
Exam Tip
Fluorine (top-right of the main group, Period 2, Group 17) has the highest electronegativity (4,0) of all elements. Caesium (bottom-left, Period 6, Group 1) has the lowest first ionisation energy. These are useful anchor points for all trend questions.
Worked Example
Arrange the following Group 1 elements in order of INCREASING first ionisation energy: Na, K, Li, Cs.
Given
- All elements are in Group 1 (alkali metals)
Find
Order of increasing first ionisation energy
Solution
- 1Group 1 elements: Li (Period 2), Na (Period 3), K (Period 4), Cs (Period 6).
- 2First ionisation energy DECREASES DOWN a group (outer electron further from nucleus, more shielded).
- 3So IE increases UP the group: Cs (lowest) < K < Na < Li (highest).
Practice Question
Explain why the atomic radius of lithium (Li) is greater than the atomic radius of fluorine (F), even though both elements are in Period 2.
(4 marks)