Electronegativity, valence and reactivity

Electronegativity

Electronegativity is the ability of an atom in a chemical compound to attract a shared pair of electrons towards itself. Unlike the ionisation enthalpy, it is not a measured energy. It is a number on a scale. The best known is the Pauling scale, on which fluorine, the most electronegative element, is $4.0$.

It rises across a period, because the atoms shrink and $Z_{\text{eff}}$ grows. It falls down a group, because the shared pair is farther from the nucleus.

Metals and non-metals

The trends go together. On the left the atoms are large and hold their electrons loosely: they lose electrons easily and behave as metals. On the right the atoms are small and electronegative: they gain or share electrons and behave as non-metals. Metallic character falls across a period and rises down a group.

The periodicity of valence

The valence of a main-group element is usually the number of its outer electrons, or eight minus that number.

So the formulas of compounds repeat down each group. Lithium and sodium both form oxides of the type $\text{M}_2\text{O}$. Aluminium (valence 3) and oxygen (valence 2) give $\text{Al}_2\text{O}_3$.

The second period is different

The first element of groups 1, 2 and 13 to 17 (lithium, beryllium, boron and so on) often behaves differently from the rest of its group. Its atom is small, its charge-to-radius ratio is high, and it is very electronegative. Its second shell also has only four orbitals, $2s$ and three $2p$. So its maximum covalency is 4. Boron forms $[\text{BF}_4]^-$, while aluminium can form $[\text{AlF}_6]^{3-}$.

The first element of a group often resembles the second element of the next group, one step down and to the right. This is the diagonal relationship:

• lithium resembles magnesium;
• beryllium resembles aluminium.

Reactivity and the nature of oxides

The most reactive elements sit at the two ends of a period. On the left, alkali metals lose an electron easily, because their ionisation enthalpy is low. On the right, halogens take an electron readily, because their electron gain enthalpy is very negative.

The oxides follow the same pattern across a period:

• Basic on the left. Sodium oxide gives an alkali with water: $\text{Na}_2\text{O} + \text{H}_2\text{O} \rightarrow 2\text{NaOH}$.
• Amphoteric in the middle. Aluminium oxide reacts with both acids and bases.
• Acidic on the right. $\text{Cl}_2\text{O}_7 + \text{H}_2\text{O} \rightarrow 2\text{HClO}_4$, a strong acid.

Practise with new numbers

Try it in Python

The program lists the highest oxide of each period 3 element, its nature, and how far the element is from oxygen in electronegativity.

Common mistakes

• Treating electronegativity as an energy. It is a number on a scale, for an atom in a compound.
• Putting the highest electronegativity on the left. It rises towards fluorine, top right.
• Expecting every group to have one valence. Groups 15 to 17 can show two, such as 3 or 5 for phosphorus.
• Giving boron a covalency of 6. The second shell has only four orbitals.
• Calling every oxide basic. Oxides become acidic across a period.

Remember these

• Electronegativity: the pull of an atom in a compound on a shared pair. Pauling scale, fluorine $4.0$.
• It rises across a period and falls down a group. Metallic character goes the other way.
• Valence $=$ outer electrons, or $8$ minus them.
• Second-period elements: small, covalency at most 4. Diagonal pairs: Li and Mg, Be and Al.
• Oxides: basic on the left, amphoteric in the middle, acidic on the right.

Check yourself

Try each one on your own first. If you are stuck, open a hint. Then check with Show the answer.

1. Which is more electronegative: nitrogen or phosphorus? Why?

2. Element E is in group 13. What is the formula of its chloride?

3. Why can aluminium form $[\text{AlF}_6]^{3-}$ when boron forms only $[\text{BF}_4]^-$?

4. Is the oxide of sulphur, $\text{SO}_3$, acidic or basic? Write what happens in water.

5. Which metal of group 2 does lithium resemble, and what is this called?