Carbon family is placed towards the right side of the periodic table. They are called the group 14 elements consisting of carbon (C), silicon (Si), germanium (Ge), tin (Sn), lead (Pb), and flerovium (Fl). They belong to the p-block of elements in the periodic table with the electronic configuration of ns2np2. In this article, we are going to discuss the electronic configurations, ionization enthalpy, electronegativity and covalent radius of the group 14 elements.
Elements of Carbon Family
- Carbon is the first element in the group which is one of the most abundant elements found on earth. It is found in the free state as well as in combined states. It is found in air, organic compounds, polymers, carbonates, etc. It has three isotopes, namely, 12C, 13C, and 14C where 14C is radioactive.
- Silicon is found in dust, sand, clay, stone, silica and silicate minerals. It is rarely found as a pure element. It is classified as neither a nonmetal or a metal but as a metalloid.
- Germanium is a rare element which is used in the manufacturing semiconductor devices. Pure germanium is an excellent semiconductor. However, it only occurs in traces as it is too reactive to be found in the elemental state.
- Tin is a soft, malleable metal with a low melting point. Tin is mainly obtained from the mineral cassiterite. It has two main allotropes at regular pressure and temperature.
- Lead is also known as plumbate. It is obtained from Galena and is used in making lead-acid batteries, oxidizing agents, and alloys. Lead is toxic to humans.
Electronic configuration of an atom can be defined as an illustration of the layout of electrons distributed among the sub-shells and orbitals.This configuration of electrons helps in understanding the physical and chemical properties of elements.The chemistry behind the elements can be determined by studying the number of valence electrons in the outermost shells.
Before understanding the electronic configuration of elements one has to understand the rules for assigning the electrons into the orbitals such as Pauli’s exclusion principle, Hund’srule of maximum multiplicity and Aufbau principle. Electrons fill the orbitals in such a way that the energy of the atom is minimized.Hence, the electrons of an element fill the energy levels in an increasing order as per the Aufbau principle. Pauli defined a set of unique quantum numbers for each electron. Pauli exclusion principle states that all the four quantum numbers for any two electrons in an atom can never be same. As perHund’s rule, the pairing of electrons in an orbital takes place only when all the sub-shells have one electron each.
The group 14 elements have a general electronic configuration of ns2np2. These elements have 2 electrons in the outermost p orbitals. The electronic configuration of group 14 elements is shown below:
|4||Germaniun||Ge||32||[Ar]3d10 4s2 4p2|
|5||Tin||Sn||50||[Kr]4d10 5s2 5p2|
|6||Lead||Pb||82||[Xe]4f14 5d10 6s2 6p2|
As all the elements in group 14 have4 electrons in the outermost shell, the valency of Group-14 elements is 4. They use these electrons in the bond formation in order to obtain octet configuration.
- The covalent radius increases down the group.
- There is a substantial increase in radius from carbon to silicon.
- Post that, the difference is less considerable. The reason can be credited to the d and f orbitals which are completely filled with the heavier members.
- The ionization enthalpies decrease down the group since the distance from the nucleus increases.
- There is a substantial decrease of ionization enthalpy from carbon to silicon. Post that, the difference is less considerable.
- There is a slight increase in ionization enthalpy from tin to lead due to the poor shielding effect of the d and f orbitals.
In general, electronegativity decreases down the group. The reason behind this irregularity is because of the filling of intervening d and f atomic orbitals. However, the electronegativity is almost the same from silicon to lead.