Exercises: Periodic table and trends

At the end of the 19th century, Dmitri Mendeleev arranged the chemical elements in order of increasing atomic mass, which caused certain discrepancies in pairs such as tellurium ($Z = 52$, atomic mass $127.6\text{ u}$) and iodine ($Z = 53$, atomic mass $126.9\text{ u}$). Henry Moseley later resolved this anomaly using X-ray spectroscopy.

1. a) Explain the basis of Moseley's law and why the atomic number $Z$, and not the atomic mass, is the correct criterion for periodic ordering.
2. b) Justify the relative positions of tellurium and iodine in the current periodic table in terms of their valence electron configurations.

The ground-state electron configurations determine the position of chemical elements in the periodic table. Consider three neutral elements called A, B and C, whose valence shells are: A: $3s^2\,3p^1$; B: $[\mathrm{Ar}]\,4s^2$ (complete configuration except for the abbreviated core); C: $3s^2\,3p^5$.

1. a) Deduce the period and group (using the IUPAC numbering from 1 to 18) to which each of the three elements belongs.
2. b) Classify each element, giving reasons, as a metal, non-metal or metalloid.

Electronegativity quantifies the relative tendency of a bonded atom to attract the electron density of the chemical bond towards itself.

1. a) Explain how electronegativity varies across a period of the periodic table from left to right, justifying the behaviour in terms of effective nuclear charge and atomic radius.
2. b) Justify why noble gases generally lack defined electronegativity values on the usual Pauling scale.

Electron affinity measures the energy change when a neutral gaseous atom gains an electron: $\mathrm{X}_{\text{(g)}} + \mathrm{e}^- \to \mathrm{X}^-_{\text{(g)}}$. According to the usual thermodynamic convention, an exothermic process has $\Delta H_{\text{ae}} < 0$. The thermodynamic values for the halogens chlorine and fluorine are $\Delta H_{\text{ae}}(\mathrm{Cl}) = -349\text{ kJ}\times\text{mol}^{-1}$ and $\Delta H_{\text{ae}}(\mathrm{F}) = -328\text{ kJ}\times\text{mol}^{-1}$, while for gaseous nitrogen it is $\Delta H_{\text{ae}}(\mathrm{N}) \ge 0$.

1. a) Explain why adding an electron to fluorine releases less energy than for chlorine, even though fluorine has higher electronegativity and a smaller atomic radius.
2. b) Justify in detail why nitrogen has an unfavourable electron affinity (endothermic or zero) compared with carbon ($\Delta H_{\text{ae}} = -122\text{ kJ}\times\text{mol}^{-1}$).