Ionisation enthalpies of elements of second period are given below Ionisation enthalpy/ $\mathrm{~k} \mathrm{~cal} \mathrm{~mol}^{-1}$ : 520, 899, 801, 1086, 1402, 1314, 1681, 2080. Match the correct enthalpy with the elements and complete the graph given in figure. Also write symbols of elements with their atomic number.
To match the correct enthalpy with the elements and to complete the graph the following points are taken into consideration. As we move from left to right across a period, the ionisation enthalpy keeps on increasing due to increased nuclear charge and simultaneous decrease in atomic radius.
However, there are some exceptions given below
(a) In spite of increased nuclear charge, the first ionisation enthalpy of B is lower than that of Be . This is due to the presence of fully filled $2 s$ orbital of $\mathrm{Be}\left[1 s^2 2 s^2\right]$ which is a stable electronic arrangement. Thus, higher energy is required to knock out the electron from fully filled $2 s$ orbitals. While B $\left[1 s^2 2 s^2 2 P^1\right]$ contains valence electrons in $2 s$ and $2 p$ orbitals. It can easily lose its one $e^{-}$from $2 p$ orbital in order to achieve noble gas configuration. Thus, first ionisation enthalpy of B is lower than that of Be .
Since, the electrons in $2 s$-orbital are more tightly held by the nucleus than these present in $2 p$-orbital, therefore, ionisation enthalpy of B is lower than that of Be .
(b) The first ionisation enthalpy of N is higher than that of O though the nuclear charge of O is higher than that of N . This is due to the reason that in case of N , the electron is to be removed from a more stable exactly half-filled electronic configuration $\left(1 s^2 2 s^2 2 p_x^1 2 p_y^1 2 p_z^1\right)$ which is not present in $\mathrm{O}\left(1 s^2 2 s^2 2 p_x^2 2 p_y^1 2 p_z^1\right)$.
Therefore, the first ionisation enthalpy of N is higher than that of O . The symbols of elements along with their atomic numbers are given in the following graph
Among the elements $\mathrm{B}, \mathrm{Al}, \mathrm{C}$ and Si ,
(a) which element has the highest first ionisation enthalpy?
(b) which element has the most metallic character?
Justify your answer in each case.
The placing of elements are as
| Period | Group-13 | Group-14 |
|---|---|---|
| 2nd Period | Boron | Carbon |
| 3rd Period | Aluminium | Silicon |
(a) Ionisation enthalpy increases along a period (as we move from left to right in a period) with decrease in atomic size and decreases down the group with increase in atomic size. Hence, carbon has the highest first ionisation enthalpy.
(b) Metallic character decreases across a period but increases on moving down the group. Hence, aluminium has the most metallic character.
Write four characteristic properties of $p$-block elements.
The four important characteristic properties of p-block elements are the following
(a) $p$-Block elements include both metals and non-metals but the number of non-metals is much higher than that of metals. Further, the metallic character increases from top to bottom within a group and non-metallic character increases from left to right along a period in this block.
(b) Their ionisation enthalpies are relatively higher as compared to s-block elements.
(c) They mostly form covalent compounds.
(d) Some of them show more than one (variable) oxidation states in their compounds. Their oxidising character increases from left to right in a period and reducing character increases from top to bottom in a group.
Choose the correct order of atomic radii of fluorine and neon (in pm) out of the options given below and justify your answer.
Illustrate by taking examples of transition elements and non-transition elements that oxidation states of elements are largely based on electronic configuration.
Oxidation state of an element depends upon the electrons present in the outermost shell or eight minus the number of valence shell electrons (outermost shell electrons). e.g.,
Alkali metals (Group 1 elements) General valence shell electronic configuration — $n s^1$; Oxidation state $=+1$.
Alkaline earth metals (Group 2 elements) General valence shell electronic configuration $-n s^2 ;$ Oxidation state $=+2$.
Alkali metals and alkaline earth metals belong to s-block elements and elements of group 13 to group 18 are known as $p$-block elements.
Group 13 elements General valence shell electronic configuration - $n s^2 n p^1$; Oxidation states $=+3$ and +1 .
Group 14 elements General valence shell electronic configuration $-n s^2 n p^2$; Oxidation states $=+4$ and +2 .
Group 15 elements General valence shell electronic configuration $-n s^2 n p^3$; Oxidation states $=-3,+3$ and +5 . Nitrogen shows $+1,+2,+4$ oxidation states also.
Group 16 elements General valence shell electronic configuration $-n s^2 n p^4$; Oxidation states $=-2,+2,+4$ and +6.
Group 17 elements General valence shell electronic configuration $-n s^2 n p^5$; Oxidation states $=-1 . \mathrm{Cl}, \mathrm{Br}$ and I also show $+1,+3,+5$ and +7 oxidation states.
Group 18 elements General valence shell configuration $-n s^2 n p^6$. Oxidation state $=$ zero.
Transition elements or d-block elements General electronic configuration $-(n-1) d^{1-10} n s^{1-2}$. These elements show variable oxidation states due to involvement of not only $n s$ electrons but $d$ or $f$-electrons (inner-transition elements) as well. Their most common oxidation states are +2 and +3 .