Group 14 trends and the allotropes of carbon

Pencil lead and diamond

A pencil lead is soft and black, and it writes on paper. A diamond is clear and is the hardest natural material. Both are pure carbon. The atoms are the same, but they are joined in different ways. Such different forms of the same element are called allotropes.

Carbon heads group 14 of the periodic table. The group has carbon (C), silicon (Si), germanium (Ge), tin (Sn) and lead (Pb). Each atom has four electrons in its outer shell, in the pattern $ns^2\,np^2$.

The trends of group 14

• C: covalent radius $77\ \text{pm}$, first ionisation enthalpy $1086\ \text{kJ/mol}$, electronegativity $2.5$.
• Si: $118\ \text{pm}$, $786\ \text{kJ/mol}$, $1.8$.
• Ge: $122\ \text{pm}$, $761\ \text{kJ/mol}$, $1.8$.
• Sn: $140\ \text{pm}$, $708\ \text{kJ/mol}$, $1.8$.
• Pb: $146\ \text{pm}$, $715\ \text{kJ/mol}$, $1.9$.

Size and ionisation enthalpy. The radius grows down the group and the ionisation enthalpy falls, with one small break: lead is a little above tin, because of poor shielding by the 4f and 5d electrons.

Metallic character. Carbon is a non-metal. Silicon and germanium are metalloids, and they are semiconductors. Tin and lead are metals.

Oxidation states. The usual state is $+4$, and $+2$ is also found. The $+2$ state becomes more stable down the group, and $+4$ less stable. The third course of this chapter explains this inert pair effect.

Catenation. This is the ability of an element to join its own atoms in chains and rings. The bond enthalpy of the E-E bond falls down the group: C-C $348$, Si-Si $297$, Ge-Ge $260$ and Sn-Sn $240\ \text{kJ/mol}$. Carbon forms very long chains, silicon only short ones, and the rest hardly any.

Why carbon is different

1. Small size, high electronegativity. Carbon forms short, strong bonds.
2. No d orbitals. The largest covalency of carbon is 4. Silicon can use its 3d orbitals and form $[\text{SiF}_{6}]^{2-}$.
3. Multiple bonds. Carbon forms strong double and triple bonds, $\text{C}=\text{C}$, $\text{C}\equiv \text{C}$ and $\text{C}=\text{O}$. The $p\pi\text{-}p\pi$ overlap of the larger atoms is poor, so Si=Si bonds are rare.
4. Catenation. The C-C bond is strong, so carbon makes chains and rings. This is why there are millions of organic compounds.

The allotropes of carbon

Diamond. Every carbon is bonded to four others at the corners of a tetrahedron. It is sp$^3$ hybridised, and the C-C bond length is $154\ \text{pm}$. The bonds are strong and the same everywhere, in a three-dimensional network. So diamond is very hard, has a very high melting point, and does not conduct electricity, because all four outer electrons of each atom are in bonds. The unit cell is cubic and holds 8 atoms. Its density is $3.51\ \text{g/cm}^3$.

Graphite. Every carbon is bonded to three others in a flat layer of hexagons, with a bond length of $141.5\ \text{pm}$. It is sp$^2$ hybridised. The fourth electron of each atom is in a $p$ orbital above and below the layer, and these electrons spread over the whole layer. So graphite conducts electricity along the layers. The layers lie $335\ \text{pm}$ apart and are held by weak van der Waals forces, so they slide over each other. This makes graphite soft and slippery, and it is used in pencils, as a lubricant and for electrodes. Its density is $2.26\ \text{g/cm}^3$. NCERT prints $340\ \text{pm}$ for the layer spacing, but $335\ \text{pm}$ is the accepted value. Graphite is slightly more stable than diamond, at ordinary pressure.

Fullerenes. A fullerene is a closed cage of carbon atoms. The best known is C$_{60}$, buckminsterfullerene, found by Kroto, Curl and Smalley in 1985. It is made by heating graphite with an electric arc in an atmosphere of helium or argon. The soot is dissolved in benzene or toluene, and the cage dissolves, which graphite and diamond do not. C$_{60}$ looks like a football. It has 60 atoms, 12 pentagons and 20 hexagons. Each carbon is bonded to three others. There are 30 double bonds, which sit between two hexagons, and 60 single bonds. The double bonds are a little shorter, near $138\ \text{pm}$, than the single bonds, near $145\ \text{pm}$.

Other forms. Carbon nanotubes are rolled-up graphite layers. Graphene is a single graphite layer. Coke, charcoal and carbon black are made of tiny crystals and are called amorphous carbon.

Set the slider to 60. You get 20 hexagons, 90 edges and 30 double bonds, as in the football. Now move to 70. The cage has 5 more hexagons and is longer, like a rugby ball. The check at the bottom always shows 2.

Practise with new numbers

Try it in Python

The program takes some fullerenes and counts the pentagons, hexagons, edges and double bonds, and then checks Euler's rule.

Common mistakes

• Saying that graphite is hard. It is soft, because the layers slide.
• Saying that diamond conducts. All four electrons of each atom are in bonds.
• Giving every fullerene 12 hexagons. Every fullerene has 12 pentagons. The hexagons are $n/2 - 10$.
• Counting 4 atoms in the cell of diamond. It has 8.
• Writing sp$^2$ for the carbon of diamond. It is sp$^3$.

Remember these

• Group 14: $ns^2\,np^2$, oxidation states $+4$ and $+2$. The $+2$ state grows more stable down the group.
• Covalent radius (pm): C 77, Si 118, Ge 122, Sn 140, Pb 146. Bond enthalpy E-E falls from 348 (C) to 240 kJ/mol (Sn).
• Carbon: no d orbitals, so a covalency of 4, strong multiple bonds and catenation.
• Diamond: sp$^3$, 3D network, 8 atoms in the cell. Graphite: sp$^2$, layers, conducts. Fullerene C$_{60}$: 12 pentagons, 20 hexagons, 90 bonds.

Check yourself

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

1. Why does graphite conduct electricity and diamond does not?

2. How many pentagons, hexagons and bonds are there in a cage of 70 carbon atoms?

3. Why can silicon form $[\text{SiF}_{6}]^{2-}$ but carbon cannot form $[\text{CF}_{6}]^{2-}$?

4. Name the solvent in which C$_{60}$ dissolves, and how it is separated from soot.

References

• OpenStax, Chemistry 2e, Chapter 18 (Representative Metals, Metalloids and Nonmetals), the section on the structure and properties of the nonmetals, carbon.
• NCERT Chemistry for Class 11, the chapter The p-Block Elements (older edition), the section on group 14.
• Housecroft and Sharpe, Inorganic Chemistry, the chapter on the group 14 elements.