Questions: Periodic Trends

Question 1

Moving from sodium (Na) to chlorine (Cl) across period 3 of the periodic table, what happens to atomic radius?

• A. Increases steadily because more electrons are added
• B. Decreases steadily because nuclear charge increases faster than shielding
• C. Stays roughly constant because each new electron cancels the new proton
• D. First decreases then increases due to d-orbital effects

Answer: (B)

Across a period, protons are added to the nucleus one at a time, but the new electrons go into the same principal energy level (same shell) and provide only weak shielding of each other. The effective nuclear charge (Zeff) therefore increases across the period, pulling all electrons inward and shrinking the atomic radius. This is why Cl (atomic radius ~99 pm) is much smaller than Na (~186 pm) despite having more electrons.

Question 2

Electronegativity and electron affinity measure the same property of an atom.

• T. True
• F. False

Answer: False

They are related but distinct. Electron affinity is a measured thermodynamic quantity: the energy change when a gaseous atom gains one electron (in kJ/mol). Electronegativity is a dimensionless relative scale (like the Pauling scale) describing an atom's tendency to attract electron density within a covalent bond. An atom can have a high electron affinity but electronegativity is defined only in a bonding context, and the two scales are not numerically equivalent.

Question 3

Explain why atomic radius increases going down a group in the periodic table, using effective nuclear charge and electron shielding.

Model answer: Each successive element in a group adds a new electron shell. The inner shells (core electrons) shield the outer electrons from the nucleus very effectively, so the effective nuclear charge felt by the outermost electrons stays roughly constant down the group. With each new shell farther from the nucleus, the valence electrons are simply further away, making the atom larger.

The key is that inner shells shield very well — each core electron roughly cancels one unit of nuclear charge for the valence electrons. So even though the nucleus gains more protons going down a group, the valence electrons see approximately the same Zeff (roughly +1 for alkali metals, for example). The dominant effect is that valence electrons occupy higher and higher principal quantum number shells (n=2, 3, 4...), which are intrinsically larger orbitals farther from the nucleus.