1_6 Periodic Trends in Halides, Hydrides and Oxides

The periodic trends in the elements carry straight over into their compounds. The key idea is the electronegativity difference between the two elements that bond. A large difference gives ionic bonding, a small one gives covalent bonding, and the type of bonding sets the melting point, the conductivity and how the compound behaves in water. Halides, hydrides and oxides show this pattern most clearly.

Halides: From Ionic Lattices to Molecules

Halides are binary compounds of a halogen with another element. Strongly electropositive metals, such as those of group 1, form ionic halides: giant three-dimensional lattices of ions with high melting points, which conduct when molten. Moving right across a period, the electronegativity gap narrows and the bonding becomes increasingly covalent. Elements in the middle, such as beryllium and aluminium, form polymeric halides, in which halogen atoms bridge metal atoms in chains or layers. Further right come covalent halides made of discrete molecules held by weak intermolecular forces.

Chlorides of period 3

Chloride	NaCl	MgCl₂	AlCl₃	SiCl₄	PCl₃	S₂Cl₂
Melting point (°C)	801	714	192	−69	−94	−80
Bonding	Ionic	Mostly ionic	Polymeric / covalent	Covalent	Covalent	Covalent

The change in bonding also shows up in water. Sodium chloride simply dissolves to give hydrated ions. Covalent chlorides are hydrolysed instead: water attacks the central atom and hydrogen chloride is released, which is why a bottle of phosphorus trichloride fumes in moist air.

ClP(Cl)Cl.O.O.O>>[H]P(=O)(O)O.Cl.Cl.Cl

Going down the halogen group, ionic character falls from fluoride to iodide, because the large iodide ion is easily distorted (polarized) by a metal cation, which pulls electron density into the bond. Aluminium fluoride is ionic and stays solid to about 1290 °C, while aluminium iodide is mainly covalent and melts at about 190 °C. Oxidation state matters too: a metal in a high oxidation state polarizes the halide more strongly. Lead(II) chloride (

\mathrm{PbCl_2}

) is a largely ionic solid, while lead(IV) chloride (

\mathrm{PbCl_4}

) is a covalent liquid.

Hydrides: Ionic, Intermediate and Covalent

Hydrides are binary compounds of hydrogen. With group 1 and the heavier group 2 metals, hydrogen gains an electron to form the hydride ion,

\mathrm{H^-}

, giving ionic hydrides such as NaH and

\mathrm{CaH_2}

. These are crystalline solids that conduct when molten, and they react with water to release hydrogen gas:

\mathrm{NaH(s) + H_2O(l) \rightarrow NaOH(aq) + H_2(g)}

Beryllium and magnesium hydrides are intermediate: polymeric solids with largely covalent bonding. Elements further right, with electronegativities close to or above hydrogen's (2.20), form covalent hydrides such as

\mathrm{CH_4}

,

\mathrm{NH_3}

,

\mathrm{H_2O}

and HF. These are gases or volatile liquids that do not conduct. Their stability rises across a period and falls down a group, so HF is very stable while the hydrides of lead and bismuth decompose easily.

Boiling points of group 14 and group 16 hydrides (°C)

Group 14	CH₄ −162	SiH₄ −112	GeH₄ −88	SnH₄ −52
Group 16	H₂O 100	H₂S −60	H₂Se −41	H₂Te −2

In group 14, boiling points rise steadily down the group as the molecules get larger and their London forces stronger. Group 16 follows the same rise from

\mathrm{H_2S}

to

\mathrm{H_2Te}

, but water is wildly out of line: the trend predicts a boiling point near −80 °C, about 180 °C below the real value. The O–H bond is so polar that water molecules form hydrogen bonds with one another, and HF and

\mathrm{NH_3}

are anomalous for the same reason.

Oxides: Basic, Amphoteric and Acidic

Almost every element forms an oxide, and the most useful way to classify them is by acid–base behaviour. Metal oxides are basic. Group 1 and 2 oxides (except BeO) contain the oxide ion,

\mathrm{O^{2-}}

, which grabs a proton from water at once to form hydroxide, so they dissolve to give alkalis. Non-metal oxides are acidic and give acids in water:

O=C=O.O>>OC(=O)O

Between the two are amphoteric oxides, such as BeO,

\mathrm{Al_2O_3}

and ZnO, which react with both acids and bases. Zinc oxide behaves as a base towards sulfuric acid and as an acid towards sodium hydroxide:

\begin{aligned} \mathrm{ZnO + H_2SO_4} &\rightarrow \mathrm{ZnSO_4 + H_2O} \\ \mathrm{ZnO + 2NaOH + H_2O} &\rightarrow \mathrm{Na_2[Zn(OH)_4]} \end{aligned}

• Across a period: strongly basic to amphoteric to strongly acidic, e.g.

  \mathrm{Na_2O}

  , MgO,

  \mathrm{Al_2O_3}

  ,

  \mathrm{SiO_2}

  ,

  \mathrm{P_4O_{10}}

  ,

  \mathrm{SO_3}

  ,

  \mathrm{Cl_2O_7}

  .
• Down a main group: basic character increases, BeO < MgO < CaO < SrO < BaO, as the metals become more electropositive.
• With oxidation state: acidity rises as the metal's oxidation state rises. MnO is basic, but

  \mathrm{Mn_2O_7}

  (manganese +7) is acidic.

Basic and acidic oxides also react directly with each other to form salts, with no water involved, as in

\mathrm{Na_2O + SO_3 \rightarrow Na_2SO_4}

. This is a reminder that "acidic" and "basic" describe how an oxide behaves towards a partner, not just whether it changes the pH of water.

Tips for remembering

• Left ionic, right covalent: as the electronegativity gap with the partner shrinks across a period, halides, hydrides and oxides all change from ionic solids to covalent molecules.
• Ionic chlorides dissolve, covalent ones hydrolyse: NaCl just dissolves, but PCl₃ and SiCl₄ react with water and give off HCl.
• BAA across the period: oxides go Basic → Amphoteric → Acidic, from Na₂O through Al₂O₃ to SO₃ and Cl₂O₇.
• Higher oxidation state, more acidic oxide: MnO is basic but Mn₂O₇ is an acid anhydride, because a highly charged centre pulls on oxygen's electrons.