1_5 Periodic Trends in Melting Points, Oxidation States, Conductivity and Hydration Energy

Ionization energy and atomic size explain how single atoms behave. The properties in this lesson describe elements in bulk and ions in solution, and here the deciding factor is often structure: how the particles are held together. Once you ask "what bonds have to break?", the trends make sense.

Melting and Boiling Points

Across period 3, melting points rise and then fall sharply, and the reason is the change in structure from left to right:

Melting points across period 3

Element	Na	Mg	Al	Si	P (white)	S	Cl	Ar
Structure	Metallic	Metallic	Metallic	Giant covalent	P₄ molecules	S₈ molecules	Cl₂ molecules	Atoms
Melting point (°C)	98	650	660	1414	44	115	−101	−189

• Sodium to aluminium: each atom gives one, two, then three electrons to the metallic bond, and the ions get smaller, so metallic bonding strengthens and melting points rise.
• Silicon: a giant covalent lattice in which every atom is bonded to four others, so melting means breaking a huge number of strong covalent bonds. Carbon in the same group behaves the same way: diamond does not melt until above 3500 °C.
• Phosphorus to argon: small separate molecules or single atoms. Melting overcomes only weak intermolecular forces, not covalent bonds, so the values are low. Sulfur's S₈ is larger than P₄ and so melts higher.

Down a group, the trend depends on the bonding. In group 1, metallic bonding weakens as the atoms grow, so melting points fall steadily from lithium (181 °C) to caesium (28 °C). Group 2 metals melt much higher than group 1, because each atom gives two electrons to the metallic bond. Down group 2, though, the trend is not smooth: magnesium (650 °C) melts lower than calcium (842 °C) because its crystal structure differs. In the halogens the trend reverses: melting and boiling points increase from fluorine to iodine. Larger molecules have more electrons, which are more easily distorted (more polarizable), so the London dispersion forces between them are stronger. That is why chlorine is a gas, bromine a liquid and iodine a solid at room temperature.

Oxidation States and Group Number

The oxidation state of an atom is the charge it would carry if all its bonds were ionic, with shared electrons assigned to the more electronegative atom. It is zero in the free element, +1 for sodium and −1 for chlorine in NaCl, and +4 for tin in the covalent

\mathrm{SnCl_4}

. For main-group elements it is tied closely to the group number:

• Groups 1 to 14 (IA–IVA): the usual positive state equals the number of valence electrons: Na +1, Mg +2, Al +3, Si +4. Heavier members of groups 13 and 14 also show a state two lower, such as Tl +1 and Pb +2.
• Group 15 (VA): +5 using all five valence electrons, +3 (N, P, As and Sb show both), or −3 by filling the three vacancies, as in

  \mathrm{NH_3}

  .
• Group 16 (VIA): sulfur is +6 in

  \mathrm{H_2SO_4}

  (all six valence electrons) and −2 in

  \mathrm{H_2S}

  (two vacancies filled).
• Group 17 (VIIA): usually −1, one vacancy filled. Chlorine, bromine and iodine also show +1, +3, +5 and +7 with oxygen.
• Group 18: normally 0 because the shell is full, though xenon forms compounds such as

  \mathrm{XeF_2}

  (+2).

Transition elements can reach their group number too, as in manganese +7 in

\mathrm{KMnO_4}

and chromium +6 in

\mathrm{K_2Cr_2O_7}

. Because both their outer s and inner d electrons can take part in bonding, most of them show several oxidation states.

Electrical Conductivity

Metals conduct because their outer electrons are delocalised: they are free to move through the lattice when a voltage is applied. The group 11 coinage metals, copper, silver and gold, are outstanding conductors, and silver is the best of all. Non-metals in groups 16 and 17 conduct so poorly that they count as insulators. Transition metals show no simple trend.

Carbon shows that structure matters more than the element itself. In diamond, all four valence electrons are locked in localised bonds, so it is an insulator. In graphite, each carbon bonds to only three others, and the fourth electron is delocalised across the layers, so graphite conducts. Further down group 14, tin and lead are true metals and conduct well.

Hydration Energy

The hydration energy of an ion is the enthalpy change when one mole of gaseous ions dissolves in water to give a very dilute solution. It is always negative (exothermic), because polar water molecules cluster around the ion: the δ− oxygen faces a cation, the δ+ hydrogens face an anion. How much energy is released depends on the ion's charge density, its charge relative to its size.

Approximate hydration energies (kJ/mol); published tables differ slightly

Ion	Li⁺	Na⁺	K⁺	Mg²⁺	Ca²⁺	Al³⁺	F⁻	Cl⁻	I⁻
ΔH(hyd)	−499	−390	−305	−1891	−1562	−4613	−457	−384	−307

Down a group, ions grow and their charge density falls, so hydration energy becomes less negative:

\mathrm{Li^+}

releases far more than

\mathrm{K^+}

, and

\mathrm{F^-}

more than

\mathrm{I^-}

. Across a period, cations become smaller and more highly charged, so hydration energy rises steeply from

\mathrm{Na^+}

to

\mathrm{Mg^{2+}}

to

\mathrm{Al^{3+}}

. The tiny hydrogen ion is the extreme case: it releases over 1000 kJ/mol and exists in water as

\mathrm{H_3O^+}

.

Tips for remembering

• Period 3 melting points peak at silicon: structure decides it: metals climb Na to Al, giant covalent Si tops the row, then small molecules (P₄, S₈, Cl₂, Ar) drop low.
• Halogens climb, alkali metals fall: bigger molecules mean stronger London forces; bigger metal atoms mean weaker metallic bonding.
• Oxidation state = electrons in or holes to fill: main-group elements use their valence electrons (+group number) or their vacancies (negative), e.g. S is +6 in H₂SO₄ and −2 in H₂S.
• Small and highly charged hydrates hard: hydration energy follows charge density, so Al³⁺ beats Mg²⁺ beats Na⁺, and Li⁺ beats Cs⁺.