# Periodic table trends for OCR A-level Chemistry 1. Electronegativity
2. Ionisation energy
3. Electron configuration

The electronegativity of an element describes how readily it attracts a shared pair of electrons towards itself. For example, if an atom of the least electronegative element, francium, were to form a bond with an atom of the most electronegative element, fluorine, electrons that participate in bonding between them would be primarily surrounding the fluorine atom.

The electronegativity of an element increases as you move from the bottom left to the top right corner of the periodic table. There are many mathematical definitions of electronegativity, the most commonly used is the Pauling scale. The 1st ionisation energy of an element is the amount of energy required to remove the first valence electron from a gaseous atom. 1st ionisation can be presented using a half-equation that shows a neutral atom X forming X⁺ and an electron: X(g) → X⁺(g) + e⁻.

As the atomic radius decreases, the nucleus has a tighter hold on its electrons, so more energy is required to remove them. Therefore, the 1st ionisation energy increases going towards the right across a period and decreases going down a group.

There are exceptions to the general trend. Electrons present in a sub-shell further away from the nucleus or at a higher energy would require less energy to be removed from an atom. Paired electrons are easier to remove from an atom than an unpaired ones from the same sub-level. Reduced energy required results in a lower 1st ionisation energy in both cases.

You may use the following link to visualise the trends of the 1st ionisation energy in the periodic table: https://ptable.com/#Properties/Ionization/1st Electron shells in an atom are made up of atomic orbitals, which are regions around the nucleus that can hold up to two electrons, with opposite spins. There are s, p and d orbitals, each with a different shape. Here are all available orbitals for the first 4 energy levels:
1st → 1s
2nd → 2s 2p
3rd → 3s 3p 3d
4th → 4s 4p 4d 4f

The other 3 energy levels have a similar pattern to the 4th energy level. The fact that particular shells exist in a level doesn’t mean that they will be filled for an element in a corresponding period. For example, d-orbitals that are found in 3rd level will only start to be filled for 4th period elements as there are no d-block elements in period 3. s orbital exits alone, p orbitals in groups of 3, d in groups of 5 and f in groups of 7. Each orbital can fit 2 electrons.

Shells fill up in order s → p → d → f by going through each element from hydrogen to the element of interest and assigning them to corresponding orbitals. An electron configuration is written by using the aforementioned templates for the energy levels and by including superscript letters representing the number of electrons in each orbital. For example, calcium would have an electronic configuration of 1s² 2s² 2p⁶ 3s² 3p⁶ 4s². # 1

How does electronegativity change going down a group of the periodic table?

Electronegativity decreases going down a group of the periodic table.

decreases # 2

Explain why this sketch is an incorrect depiction of a bond between fluorine and boron.

Fluorine is more electronegative than boron. The shared electrons should be held closer to fluorine.

Due to higher electronegativity of fluorine, electrons would be closer to it. # 3

What is the electronic configuration of the atom shown in the diagram?

This element has 2 + 3 = 5 electrons.
For atoms the number of electrons and the atomic number is the same.
The atomic number of the element is 5 and it is boron.

To get to boron in the periodic table we need to ggoet through 2 s elements in period 1, 2 s elements in period 2 and 1 p element in period 2 (boron).

1s² 2s² 2p¹ # 4

What is the electronic configuration for this atom?

This element has 2 + 8 = 10 electrons.
For atoms the number of electrons and the atomic number is the same.
The atomic number of the element is 10 and it is neon.

To get to neon in the periodic table we need to go through 2 s elements in period 1, 2 s elements in period 2 and 6 p elements in period 2 (including neon).

1s² 2s² 2p⁶ # 5

Write a half-equation for the 1st ionisation of silver.

In the first ionisation half-equation silver would go from a neutral gaseous atom to a gaseous Ag⁺ and would produce an electron.

Ag(g) → Ag⁺(g) + e⁻ End of page