Electronegativity (HSC SSCE Chemistry): Revision Notes

Electronegativity

What is electronegativity?

Electronegativity indicates how strongly an atom pulls bonding electrons towards itself when it forms compounds. The higher an element's electronegativity value, the more powerfully it attracts electrons in a chemical bond.

Electronegativity values typically range from just under $1$ (for elements like sodium, potassium and barium) to nearly $4.0$ (for fluorine and chlorine). These values follow predictable patterns across the periodic table.

Electronegativity values across the periodic table

The table below shows electronegativity values for common elements:

Group 1	Group 2	Group 13	Group 14	Group 15	Group 16	Group 17	Group 18
H: 2.20							He: 0
Li: 0.98	Be: 1.57	B: 2.04	C: 2.55	N: 3.04	O: 3.44	F: 3.98	Ne: 0
Na: 0.93	Mg: 1.31	Al: 1.61	Si: 1.90	P: 2.19	S: 2.58	Cl: 3.16	Ar: 0
K: 0.82	Ca: 1.00			As: 2.18	Se: 2.55	Br: 2.96	Kr: 0
Rb: 0.82	Sr: 0.95					I: 2.66	Xe: 0
Cs: 0.79	Ba: 0.89						Rn: 0

Periodic trends in electronegativity

Electronegativity follows two clear patterns across the periodic table:

Across a period (left to right): Electronegativity increases as you move from left to right. For example, in Period 2, lithium ($0.98$) has much lower electronegativity than fluorine ($3.98$).

Down a group (top to bottom): Electronegativity decreases as you move down a group. For instance, in Group 17 (halogens), fluorine ($3.98$) has higher electronegativity than iodine ($2.66$).

Elements that form negative ions

Electronegativity serves as an approximate indicator of an atom's ability to gain electrons and form negative ions (anions). Elements with the highest electronegativity values are most likely to form these ions.

The elements most likely to form negative ions are:

• Fluorine forms fluoride ions ($\text{F}^-$)
• Chlorine forms chloride ions ($\text{Cl}^-$)
• Oxygen forms oxide ions ($\text{O}^{2-}$)
• Bromine forms bromide ions ($\text{Br}^-$)
• Sulphur forms sulphide ions ($\text{S}^{2-}$)

Notice that these are all non-metals located on the right side of the periodic table.

Why do these trends occur?

The explanations for electronegativity trends are similar to those for first ionisation energy, but work in the opposite direction.

Trend across a period: As you move from left to right across a period, atoms become progressively smaller. This happens because the nuclear charge increases (more protons) whilst electrons are being added to the same outer shell. The smaller atomic radius means the nucleus can attract additional electrons more strongly, resulting in higher electronegativity.

Trend down a group: As you move down a group, atoms become larger because electrons occupy shells that are further from the nucleus. Additionally, there is increased shielding (or screening) from the completely filled inner electron shells. These inner shells block some of the nuclear attraction, reducing the nucleus's ability to attract extra electrons. Therefore, electronegativity decreases and there is less tendency for the atom to form a negative ion.

Remember!