Group 14 — Carbon Family (General)

The Family at a Glance

The members of Group 14 are carbon (C), silicon (Si), germanium (Ge), tin (Sn) and lead (Pb), together with the synthetic radioactive element flerovium (Fl), whose chemistry has not been established. As one descends this column, the character of the elements shifts steadily from non-metal to metal: carbon and silicon are non-metals, germanium is a metalloid, and tin and lead are soft metals with low melting points. This non-metal-to-metal gradation is the unifying theme of the family, and almost every general trend below traces back to it.

Their natural occurrence underlines their importance. Carbon is the seventeenth most abundant element by mass in the crust and is found free as coal, graphite and diamond and combined as carbonates, hydrocarbons and atmospheric $\ce{CO2}$ (about $0.03\%$). Silicon is the second most abundant crustal element ($27.7\%$ by mass), present as silica and silicates. Germanium occurs only in traces; tin occurs mainly as cassiterite ($\ce{SnO2}$) and lead as galena ($\ce{PbS}$). Ultrapure germanium and silicon are the basis of transistors and semiconductor devices.

Electronic Configuration

The defining feature of the family is its valence-shell electronic configuration, $ns^2np^2$ — four electrons in the outermost shell. What changes down the group is the inner core: silicon onward, completely filled $d$ orbitals appear, and lead additionally carries a filled $4f^{14}$ set. Because these inner $d$ and $f$ electrons screen the nuclear charge poorly, the effective nuclear charge experienced by the valence electrons rises more than a simple shell count would suggest, and this perturbs every property that depends on it — size, ionisation enthalpy and the relative stability of the two oxidation states.

Atomic and Physical Trends

Three atomic properties shape the chemistry of Group 14, and each carries a small irregularity worth memorising.

Covalent radius

There is a considerable increase in covalent radius from C to Si, but thereafter, from Si to Pb, only a small increase is observed. The reason is the completely filled $d$ and $f$ orbitals in the heavier members: their poor screening lets the rising nuclear charge pull the outer electrons inward, partly offsetting the size increase expected from adding new shells.

Ionisation enthalpy

The first ionisation enthalpy of every Group 14 element is higher than that of the corresponding Group 13 element, reflecting the smaller size and the influence of the inner core. In general $\Delta_i H$ decreases down the group, but not smoothly: there is a small decrease from Si to Ge to Sn and then a slight increase from Sn to Pb, again the consequence of poor shielding by the intervening $d$ and $f$ orbitals together with the growing atomic size.

Electronegativity

Because of their small size, Group 14 elements are slightly more electronegative than the Group 13 elements. Carbon is distinctly the most electronegative member (2.5 on the Pauling scale), while the values for Si through Pb are almost the same (close to 1.8–1.9). All members of the group are solids; melting and boiling points are much higher than those of the corresponding Group 13 elements.

Oxidation States and the Inert Pair

With four valence electrons, the family exhibits the common oxidation states $\mathbf{+4}$ and $\mathbf{+2}$ (carbon additionally shows negative states, as in $\ce{CH4}$). Because the sum of the first four ionisation enthalpies is very high, $+4$ compounds are essentially covalent, not ionic; for example, $\ce{CCl4}$ has eight electrons around carbon and is electron-precise.

The decisive trend is the growing stability of the lower $+2$ state on descending the group, expressed as $\ce{Ge < Sn < Pb}$. This is the inert pair effect: the $ns^2$ pair of the valence shell becomes increasingly reluctant to participate in bonding. NIOS Chapter 18 ascribes this to two factors — the rising promotion energy needed to take an atom from the ground state $ns^2np^2$ to a bonding state, and the poorer overlap of orbitals for the larger atoms, which yields weaker bonds that no longer repay the cost of unpairing. The net result is the lesser stability of the higher oxidation state with increasing atomic number.

The practical consequences are sharp. Carbon and silicon show essentially only the $+4$ state. Germanium forms stable $+4$ compounds and only a few in the $+2$ state. Tin uses both states, and $\ce{Sn^2+}$ is a reducing agent (readily oxidised to $\ce{Sn^4+}$). Lead is the opposite extreme: $\ce{Pb^2+}$ compounds are stable, whereas $\ce{Pb^4+}$ compounds are strong oxidising agents.

Catenation Down the Group

Catenation is the tendency of identical atoms to link through covalent bonds into chains and rings. Carbon does this prolifically — it is the foundation of the whole of organic chemistry — because the C–C bond is very strong. Down the group the atomic size increases and electronegativity decreases, so the M–M bond enthalpy falls and the catenating tendency declines steeply. The order established by bond enthalpies is:

$$\ce{C} \;\gg\; \ce{Si} \;>\; \ce{Ge} \;\approx\; \ce{Sn}$$

and lead does not show catenation at all. (The chapter summary writes the same order as $\ce{C} > \ce{Si} > \ce{Ge} \approx \ce{Sn} > \ce{Pb}$.)

Catenation, combined with carbon's ability to form $p\pi$–$p\pi$ multiple bonds, is precisely what gives carbon its allotropy — the diamond, graphite and fullerene structures explored in the sibling note on the allotropes of carbon.

Anomalous Behaviour of Carbon

Like the first member of every p-block group, carbon stands apart from the rest of its family. The supplementary NCERT text and NIOS Chapter 18 attribute this to four interlinked causes: carbon's small size, higher electronegativity, higher ionisation enthalpy and the unavailability of $d$ orbitals in its valence shell. Two consequences are tested repeatedly.

First, with only $2s$ and $2p$ orbitals available, carbon can accommodate at most four electron pairs and so is limited to a maximum covalence of four. Heavier members expand their covalence using vacant $d$ orbitals — hence species such as $\ce{[SiF6]^2-}$, $\ce{[GeCl6]^2-}$ and $\ce{[Sn(OH)6]^2-}$ exist (with $sp^3d^2$ hybridisation of the central atom), while no analogous $\ce{[CF6]^2-}$ can form. The same limitation explains why $\ce{CCl4}$ resists hydrolysis whereas $\ce{SiCl4}$ hydrolyses readily.

Second, only carbon (and the other small second-period elements) forms strong $p\pi$–$p\pi$ multiple bonds, both to itself and to small, electronegative partners: $\ce{C=C}$, $\ce{C#C}$, $\ce{C=O}$, $\ce{C=S}$ and $\ce{C#N}$. The heavier elements do not form such $p\pi$–$p\pi$ bonds because their atomic orbitals are too large and diffuse for effective sideways overlap.

General Chemical Reactivity

The general reactions of the family fall into three blocks, and the trends within each track the oxidation-state and acidity ideas above.

With oxygen

All members form oxides on heating, chiefly monoxides ($\ce{MO}$) and dioxides ($\ce{MO2}$). Oxides in the higher oxidation state are generally more acidic than those in the lower state. Among dioxides, $\ce{CO2}$, $\ce{SiO2}$ and $\ce{GeO2}$ are acidic, while $\ce{SnO2}$ and $\ce{PbO2}$ are amphoteric. Among monoxides, $\ce{CO}$ is neutral, $\ce{GeO}$ is distinctly acidic, and $\ce{SnO}$ and $\ce{PbO}$ are amphoteric.

With water

Carbon, silicon and germanium are not affected by water. Tin decomposes steam to give the dioxide and dihydrogen, while lead is unaffected, probably because of a protective oxide film.

$$\ce{Sn + 2H2O ->[\Delta] SnO2 + 2H2}$$

With halogens

These elements form halides of types $\ce{MX2}$ and $\ce{MX4}$. Most tetrahalides are covalent and tetrahedral with $sp^3$ hybridisation of the central atom; exceptions $\ce{SnF4}$ and $\ce{PbF4}$ are ionic. Among the dihalides, stability increases down the group — consistent with the inert pair effect — so $\ce{GeX4}$ is more stable than $\ce{GeX2}$, whereas $\ce{PbX2}$ is more stable than $\ce{PbX4}$. Notably $\ce{PbI4}$ does not exist, because the $\ce{Pb-I}$ bond energy is too small to unpair and excite the $6s^2$ electrons needed for four bonds. Except $\ce{CCl4}$, the tetrachlorides hydrolyse, the central atom accepting the water oxygen's lone pair into a $d$ orbital, as illustrated by silicon:

$$\ce{SiCl4 + 4H2O -> Si(OH)4 + 4HCl}$$