10.3 Group 14 Elements: The Carbon Family

Carbon, silicon, germanium, tin, lead and flerovium are the members of group 14. Carbon is the seventeenth most abundant element by mass in the earth’s crust. It is widely distributed in nature in free as well as in the combined state. In elemental state it is available as coal, graphite and diamond; however, in combined state it is present as metal carbonates, hydrocarbons and carbon dioxide gas ( \(0.03 \%\) ) in air. One can emphatically say that carbon is the most versatile element in the world. Its combination with other elements such as dihydrogen, dioxygen, chlorine and sulphur provides an astonishing array of materials ranging from living tissues to drugs and plastics. Organic chemistry is devoted to carbon containing compounds. It is an essential constituent of all living organisms. Naturally occurring carbon contains two stable isotopes: \({ }^{12} \mathrm{C}\) and \({ }^{13} \mathrm{C}\). In addition to these, third isotope, \({ }^{14} \mathrm{C}\) is also present. It is a radioactive isotope with half-life 5770 years and used for radiocarbon dating. Silicon is the second ( \(27.7 \%\) by mass) most abundant element on the earth’s crust and is present in nature in the form of silica and silicates. Silicon is a very important component of ceramics, glass and cement. Germanium exists only in traces. Tin occurs mainly as cassiterite, \(\mathrm{SnO}_2\) and lead as galena, PbS. Flerovium is synthetically prepared radioactive element

Ultrapure form of germanium and silicon are used to make transistors and semiconductor devices.

Symbol of Flerovium is Fl. It has atomic number 114 , atomic mass \(289 \mathrm{gmol}^{-1}\) and electronic configuration \([\mathrm{Rn}] 5 f^{14} 6 d^{10} 7 s^2 7 p^2\). It has been prepared only in small amount. Its half life is short and its chemistry has not been established yet. The important atomic and physical properties along with their electronic configuration of the elements of group 14 leaving flerovium are given in Table 10.3. Some of the atomic, physical and chemical properties are discussed below:

Electronic Configuration

The valence shell electronic configuration of these elements is \(n s^2 n p^2\). The inner core of the electronic configuration of elements in this group also differs.

Covalent Radius

There is a considerable increase in covalent radius from \(\mathrm{C}\) to \(\mathrm{Si}\) , thereafter from \(\mathrm{Si}\) to \(\mathrm{Pb}\) a small increase in radius is observed. This is due to the presence of completely filled \(d\) and \(f\) orbitals in heavier members.

Ionization Enthalpy

The first ionization enthalpy of group 14 members is higher than the corresponding members of group 13. The influence of inner core electrons is visible here also. In general the ionisation enthalpy decreases down the group. Small decrease in \(\Delta_i H\) from \(\mathrm{Si}\) to \(\mathrm{Ge}\) to \(\mathrm{Sn}\) and slight increase in \(\Delta_i H\) from \(\mathrm{Sn}\) to \(\mathrm{Pb}\) is the consequence of poor shielding effect of intervening \(d\) and \(f\) orbitals and increase in size of the atom.

Electronegativity

Due to small size, the elements of this group are slightly more electronegative than group 13 elements. The electronegativity values for elements from \(\mathrm{Si}\) to \(\mathrm{Pb}\) are almost the same.

Physical Properties

All members of group 14 are solids. Carbon and silicon are non-metals, germaniumis a metalloid, whereas tin and lead are soft metals with low melting points. Melting points and boiling points of group 14 elements are much higher thanthose of corresponding elements of group 13.

Chemical Properties

Oxidation states and trends in chemical reactivity

The group 14 elements have four electrons in outermost shell. The common oxidation states exhibited by these elements are +4 and +2 . Carbon also exhibits negative oxidation states. Since the sum of the The group 14 elements have four electrons in outermost shell. The common oxidation states exhibited by these elements are +4 and +2 . Carbon also exhibits negative oxidation states. Since the sum of the first four ionization enthalpies is very high, compounds in +4 oxidation state are generally covalent in nature. In heavier members the tendency to show +2 oxidation state increases in the sequence \(\mathrm{Ge}<\mathrm{Sn}<\mathrm{Pb}\). It is due to the inability of \(\mathrm{ns}^2\) electrons of valence shell to participate in bonding. The relative stabilities of these two oxidation states vary down the group. Carbon and silicon mostly show +4 oxidation state. Germanium forms stable compounds in +4 state and only few compounds in +2 state. Tin forms compounds in both oxidation states ( Sn in +2 state is a reducing agent). Lead compounds in +2 state are stable and in +4 state are strong oxidising agents. In tetravalent state the number of electrons around the central atom in a molecule (e.g., carbon in \(\mathrm{CCl}_4\) ) is eight. Being electron precise molecules, they are normally not expected to act as electron acceptor or electron donor species. Although carbon cannot exceed its covalence more than 4, other elements of the group can do so. It is because of the presence of \(d\) orbital in them. Due to this, their halides undergo hydrolysis and have tendency to form complexes by accepting electron pairs from donor species. For example, the species like, \(\mathrm{SiF}_6^{2-},\left[\mathrm{GeCl}_6\right]^{2-}\), \(\left[\mathrm{Sn}(\mathrm{OH})_6\right]^{2-}\) exist where the hybridisation of the central atom is \(s p^3 d^2\).

(i) Reactivity towards oxygen

All members when heated in oxygen form oxides. There are mainly two types of oxides, i.e., monoxide and dioxide of formula \(\mathrm{MO}\) and \(\mathrm{MO}_2\) respectively. \(\mathrm{SiO}\) only exists at high temperature. Oxides in higher oxidation states of elements are generally more acidic than those in lower oxidation states. The dioxides \(-\mathrm{CO}_2, \mathrm{SiO}_2\) and \(\mathrm{GeO}_2\) are acidic, whereas \(\mathrm{SnO}_2\) and \(\mathrm{PbO}_2\) are amphoteric in nature. Among monoxides, \(\mathrm{CO}\) is neutral, \(\mathrm{GeO}\) is distinctly acidic whereas \(\mathrm{SnO}\) and \(\mathrm{PbO}\) are amphoteric.

Example 10.4: Select the member(s) of group 14 that (i) forms the most acidic dioxide, (ii) is commonly found in +2 oxidation state, (iii) used as semiconductor.

Solution: (i) carbon (ii) lead (iii) silicon and germanium

(ii) Reactivity towards water

Carbon, silicon and germanium are not affected by water. Tin decomposes steam to form dioxide and dihydrogen gas.
\(
\mathrm{Sn}+2 \mathrm{H}_2 \mathrm{O} \xrightarrow{\Delta} \mathrm{SnO}_2+2 \mathrm{H}_2
\)
Lead is unaffected by water, probably because of a protective oxide film formation.

(iii) Reactivity towards halogen

These elements can form halides of formula \(\mathrm{MX}_2\) and \(\mathrm{MX}_4\) (where \(\mathrm{X}=\mathrm{F}, \mathrm{Cl}, \mathrm{Br}, \mathrm{I}\)). Except carbon, all other members react directly with halogen under suitable condition to make halides. Most of the \(\mathrm{MX}_4\) are covalent in nature. The central metal atom in these halides undergoes \(s p^3\) hybridisation and the molecule is tetrahedral in shape. Exceptions are \(\mathrm{SnF}_4\) and \(\mathrm{PbF}_4\), which are ionic in nature. \(\mathrm{PbI}_4\) does not exist because \(\mathrm{Pb}-\mathrm{I}\) bond initially formed during the reaction does not release enough energy to unpair \(6 s^2\) electrons and excite one of them to higher orbital to have four unpaired electrons around lead atom. Heavier members Ge to Pb are able to make halides of formula \(\mathrm{MX}_2\). Stability of dihalides increases down the group. Considering the thermal and chemical stability, \(\mathrm{GeX}_4\) is more stable than \(\mathrm{GeX}_2\), whereas \(\mathrm{PbX}_2\) is more than \(\mathrm{PbX}_4\). Except \(\mathrm{CCl}_4\), other tetrachlorides are easily hydrolysed by water because the central atom can accommodate the lone pair of electrons from oxygen atom of water molecule in \(d\) orbital.

Hydrolysis can be understood by taking the example of \(\mathrm{SiCl}_4\). It undergoes hydrolysis by initially accepting lone pair of electrons from water molecule in \(d\) orbitals of \(\mathrm{Si}\) , finally leading to the formation of \(\mathrm{Si}(\mathrm{OH})_4\) as shown below :

Example 10.5: \(\left[\mathrm{SiF}_6\right]^{2-}\) is known whereas \(\left[\mathrm{SiCl}_6\right]^{2-}\) not. Give possible reasons.

Solution: The main reasons are :
(i) six large chloride ions cannot be accommodated around \(\mathrm{Si}^{4+}\) due to limitation of its size.
(ii) interaction between lone pair of chloride ion and \(\mathrm{Si}^{4+}\) is not very strong.