Well, you are aware of carbon. Isn’t it? Be it in your chapter on respiration or environmental protection, you have heard enough about carbon. However, that is actually not enough! There is so much more to the story. In this chapter, we will look at the carbon family or element 14. We will look at their physical and chemical properties, with examples. Let’s begin.
The Carbon Family
We can find the Carbon family towards the right side of the periodic table. We refer to them as the group 14 elements. The members of this family include carbon (C), silicon (Si), germanium (Ge), tin (Sn), lead (Pb), and flerovium (Fl).
These elements belong to the p-block of elements in the periodic table. We can, thus know, their electronic configuration is ns2np2. Let us first look at all the members of this group in greater detail.
Elements of Carbon Family
- Carbon is the first element in this 14th group of elements. It is one of the most plentily available elements present on our earth. We can find it in combined as well as free states. We usually find it in air, polymers, organic compounds, carbonates etc. It has three isotopes, namely, 12C, 13C, and 14C where 14C is radioactive.
- Silicon is a common element in dust, sand, clay, stone, silica and silicate minerals. We can hardly find it as a pure element. It is neither a nonmetal or a metal. In fact, it is a metalloid.
- Germanium is a rare element which we use in the manufacturing of 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. It is mainly obtained from the mineral cassiterite. It has two main allotropes at regular pressure and temperature.
- Lead, also plumbate, is obtained from Galena. We find its common use in the making of lead-acid batteries, oxidizing agents, and alloys. Lead is toxic for us, the humans.
Learn more about Group 16 Elements here.
Electronic Configuration of the Carbon Family
Electronic configuration of an atom is nothing but an illustration of the layout of electrons distributed among the sub-shells and orbitals. By this configuration of electrons, we can understand the various physical and chemical properties of the 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, we must understand the rules for assigning the electrons into the orbitals. There are many principles that help us in doing so. These include 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 at the minimum. 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 per Hund’s rule, the pairing of electrons in an orbital takes place only when all the sub-shells have one electron each. The general electronic configuration of these group 14 elements is ns2np2. These elements have 2 electrons in the outermost p orbitals. The electronic configuration of group 14 elements is as follows:
|4||Germanium||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 have 4 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.
Learn more about Group 17 Elements here.
Properties and Trends in Element 14
1) Covalent Radius
As we move down the group, the covalent radius increases. Therefore, 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.
2) Ionisation Enthalpy
Moving down the group, we notice that the ionization enthalpies decrease. This is because of the increase in the distance from the nucleus. 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.
Learn more about s-Block Elements here.
Solved Example for You
Q: How does the electronegativity vary along the Group 14 elements?
Ans: As we move down the group, the electronegativity decreases in general. 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.