Hydrides and halides of groups 13 and 14

The same family of compounds, two different behaviours

Open a bottle of silicon tetrachloride in a damp room, and white fumes of hydrogen chloride come out. Carbon tetrachloride stands next to it for years and nothing happens. Boron trifluoride is a gas that takes ammonia, but it is less eager than boron tribromide. These differences are the subject of this lesson.

Hydrides

Hydrides are put in three classes by the number of electrons they have for bonding.

• Electron deficient: not enough electrons for ordinary bonds. $\text{B}_{2}\text{H}_{6}$ and the polymeric $(\text{AlH}_{3})\text{n}$ of group 13.
• Electron precise: just enough. $\text{CH}_{4}$, $\text{SiH}_{4}$, $\text{GeH}_{4}$ and $\text{SnH}_{4}$ of group 14.
• Electron rich: lone pairs left over. $\text{NH}_{3}$, $\text{H}_{2}\text{O}$ and $\text{HF}$.

Group 13. Boron forms many boranes, such as $\text{B}_{2}\text{H}_{6}$, $\text{B}_{4}\text{H}_{10}$ and $\text{B}_{5}\text{H}_{9}$. Aluminium hydride is a solid polymer. The hydrides of gallium and indium are unstable, and thallium forms none that lasts. The stability falls down the group. Two useful complex hydrides are lithium aluminium hydride, $\text{LiAlH}_{4}$, and sodium borohydride, $\text{NaBH}_{4}$. They contain the tetrahedral ions $[\text{AlH}_{4}]^{-}$ and $[\text{BH}_{4}]^{-}$, and they are reducing agents. They are made by $4\text{LiH} + \text{AlCl}_{3} \to \text{LiAlH}_{4} + 3\text{LiCl}$ in ether. Lithium aluminium hydride reacts with water at once: $\text{LiAlH}_{4} + 4\text{H}_{2}\text{O} \to \text{LiOH} + \text{Al}(\text{OH})_{3} + 4\text{H}_{2}$.

Group 14. The hydrides $\text{CH}_{4}$, $\text{SiH}_{4}$, $\text{GeH}_{4}$, $\text{SnH}_{4}$ and $\text{PbH}_{4}$ become less stable down the group. The silanes are more reactive than the alkanes. Silane takes fire in air, $\text{SiH}_{4} + 2\text{O}_{2} \to \text{SiO}_{2} + 2\text{H}_{2}\text{O}$, and it is hydrolysed by water that has a trace of alkali, $\text{SiH}_{4} + 2\text{H}_{2}\text{O} \to \text{SiO}_{2} + 4\text{H}_{2}$. Methane does neither. It is made from magnesium silicide and acid, $\text{Mg}_{2}\text{Si} + 4\text{HCl} \to 2\text{MgCl}_{2} + \text{SiH}_{4}$, and the product is a mixture of silanes. Another route is $\text{SiCl}_{4} + \text{LiAlH}_{4} \to \text{SiH}_{4} + \text{LiCl} + \text{AlCl}_{3}$.

Halides of group 13

Boron halides. The trihalides $\text{BX}_{3}$ are flat triangles. The boron atom has only six electrons, so they are Lewis acids. Their strength is not in the order of the electronegativity of the halogen. It is $\text{BF}_3 < \text{BCl}_3 < \text{BBr}_3 < \text{BI}_3$. The reason is back bonding. In $\text{BF}_{3}$ the fluorine atom has a filled $2p$ orbital of the same size as the empty $2p$ orbital of boron. A lone pair flows back to boron, making the B-F bond partly double, and the boron is less eager for a new pair. The larger halogens overlap badly with boron, so there is less back bonding, and more acid strength.

The halides are hydrolysed: $\text{BCl}_{3} + 3\text{H}_{2}\text{O} \to \text{B}(\text{OH})_{3} + 3\text{HCl}$, and for the fluoride $4\text{BF}_{3} + 3\text{H}_{2}\text{O} \to \text{H}_{3}\text{BO}_{3} + 3\text{HBF}_{4}$.

Aluminium chloride is a dimer, $\text{Al}_{2}\text{Cl}_{6}$, with a high affinity for water. Aluminium fluoride is an ionic solid. Thallium(III) is oxidising, so thallium forms TlCl rather than TlCl$_3$.

Halides of group 14

• CCl$_4$ is made by $\text{CS}_{2} + 3\text{Cl}_{2} \to \text{CCl}_{4} + \text{S}_{2}\text{Cl}_{2}$. It is not hydrolysed by water. Carbon has no d orbitals, and the small carbon atom is crowded by four large chlorine atoms, so water cannot attack it.
• SiCl$_4$ is made by $\text{SiO}_{2} + 2\text{C} + 2\text{Cl}_{2} \to \text{SiCl}_{4} + 2\text{CO}$. It is hydrolysed at once, $\text{SiCl}_{4} + 4\text{H}_{2}\text{O} \to \text{Si}(\text{OH})_{4} + 4\text{HCl}$. Silicon can accept the lone pair of water into its empty 3d orbitals, and the atom is larger. The product loses water and gives silica. Silicon tetrafluoride gives $\text{SiF}_{4} + 2\text{F}^{-} \to [\text{SiF}_{6}]^{2-}$.
• GeCl$_4$ and SnCl$_4$ are volatile liquids that fume in air. SnCl$_2$ is a solid that reduces.
• PbCl$_2$ is a white solid that dissolves a little in cold water and more in hot water. PbCl$_4$ is unstable. PbI$_2$ is a yellow solid that dissolves in hot water and comes out as bright golden crystals on cooling.

Choose the four chlorides in turn, with the same mass. The pH is lowest for the chloride with the lowest molar mass for each chlorine. That is boron trichloride, with the smallest molar mass for 3 chlorine atoms.

Practise with new numbers

Try it in Python

The program hydrolyses 1.00 g of each chloride and finds the pH of the acid that forms in 1.00 L of water.

Common mistakes

• Saying that BF$_3$ is the strongest Lewis acid because F is the most electronegative. Back bonding makes it the weakest of the four.
• Saying that CCl$_4$ is hydrolysed. It is not.
• Forgetting that SiF$_4$ gives $[\text{SiF}_{6}]^{2-}$ with fluoride. The covalency of silicon goes from 4 to 6.
• Writing PbBr$_4$ or PbI$_4$ as stable compounds. They do not exist.
• Counting one HCl for each chloride. Count one for each Cl atom.

Remember these

• Hydrides: electron deficient (B, Al), electron precise (C, Si, Ge, Sn). Stability falls down both groups.
• $\text{BF}_3 < \text{BCl}_3 < \text{BBr}_3 < \text{BI}_3$ in Lewis acid strength, because of back bonding.
• $\text{CCl}_{4}$ resists water. $\text{SiCl}_{4}$ is hydrolysed to silica and HCl.
• $\text{PbCl}_{4}$ loses chlorine. $\text{PbBr}_{4}$ and $\text{PbI}_{4}$ do not exist.

Check yourself

Try each one on your own first. If you are stuck, open a hint. Then check with Show the answer.

1. Why is BI$_3$ a stronger Lewis acid than BF$_3$?

2. Why is silicon tetrachloride hydrolysed but carbon tetrachloride is not?

3. Name the three classes of hydrides and place $\text{B}_{2}\text{H}_{6}$, $\text{CH}_{4}$ and $\text{NH}_{3}$.

4. What happens when lithium aluminium hydride is added to water?

References

• OpenStax, Chemistry 2e, Chapter 18 (Representative Metals, Metalloids and Nonmetals).
• NCERT Chemistry for Class 11, the chapter The p-Block Elements (older edition), the sections on hydrides and halides.
• Housecroft and Sharpe, Inorganic Chemistry, the chapters on the group 13 and group 14 elements.