Question 1

Why does first ionisation enthalpy generally increase across a period?

Accepted Answer

Across a period, electrons are added while the number of principal shells remains the same. Increasing nuclear attraction generally makes the outer electron harder to remove, so the first ionisation enthalpy rises. Be to B and N to O are important local exceptions in period 2.

Question 2

Why is the first ionisation enthalpy of Be greater than that of B?

Accepted Answer

Beryllium has a filled 2s subshell, which is relatively stable. After the 2s electron is removed, boron starts electron removal from the higher-energy 2p subshell. Therefore Be has 899 kJ/mol while B has 801 kJ/mol.

Question 3

Why is the first ionisation enthalpy of N greater than that of O?

Accepted Answer

Nitrogen has a stable half-filled 2p3 arrangement. Oxygen has a paired electron in one 2p orbital, and repulsion between the paired electrons makes that electron easier to remove. Thus N has 1402 kJ/mol while O has 1314 kJ/mol.

Question 4

Why is chlorine's electron gain enthalpy more negative than fluorine's?

Accepted Answer

Chlorine's electron gain enthalpy is -349 kJ/mol, compared with -328 kJ/mol for fluorine. Fluorine is smaller, so its compact 2p shell is more crowded. The repulsion faced by an incoming electron reduces the energy released in fluorine.

Question 5

How do covalent radius and electronegativity change across these periods?

Accepted Answer

Covalent radius generally decreases from left to right as stronger nuclear attraction pulls the electrons closer. Pauling electronegativity generally increases from left to right, with fluorine reaching 4.0. Noble gases do not receive covalent-radius or usual Pauling electronegativity bars in this lab.

Periodic Trends in 3D - Ionisation, Radius, Electronegativity & EGE

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