Maharashtra Board Class 11 Chemistry Chapter 9 Elements of group 13 14 and 15 PDF Download

9. Elements Of Group 13, 14 And 15

9.1 Introduction

You have learnt in Chapter 7 that in the p-block elements the differentiating electron (the last filling electron) enters the outermost p orbital. You also know that maximum six electrons can be accommodated in p-subshell (or three p orbitals). This gives rise to six groups, group 13 to 18, in the p-block. The p-block elements show greater variation in the properties than 's' block, which you learnt in the previous chapter. In this chapter you are going to study the elements of the groups 13, 14 and 15 in some details.

The elements boron (B), aluminium (Al), gallium (Ga), indium (In) and thallium (Tl) constitute the group 13, called the boron family. The elements carbon (C), silicon (Si), germanium (Ge), tin (Sn) and lead (Pb) form the group 14 called the carbon family. The elements nitrogen (N), phosphorous (P), arsenic (As), antimony (Sb) and bismuth (Bi) belong to group 15 of the periodic table called the nitrogen family.

Teacher's Note

Group 13, 14, and 15 elements are very important in chemistry. For example, carbon is used to make diamonds and graphite, which are found in India.

Exam Trick

Remember: Group 13 has 3 valence electrons, Group 14 has 4, and Group 15 has 5. Think 13→3, 14→4, 15→5.

Points to Remember

Boron family is Group 13.

Carbon family is Group 14.

Nitrogen family is Group 15.

9.2 Electronic Configuration Of Elements Of Groups 13, 14 And 15

The general outer electronic configuration of the group 13 elements is \(ns^2 np^1\), those of the group 14 elements is \(ns^2 np^2\) while the group 15 elements are shown as \(ns^2 np^3\). These electronic configurations differ from their nearest inert gas by 3 or 4 electrons. These elements do not occur in free monoatomic state and found as compounds with other elements or as polyatomic molecules (such as N₂, P₄, C₆₀) or polyatomic covalent arrays (such as graphite, diamond).

Teacher's Note

Electronic configuration helps us understand how elements behave. For example, boron in Group 13 always has 3 electrons in its outer shell, just like aluminum.

Exam Trick

Remember: All Group 13 elements end in np¹, Group 14 ends in np², and Group 15 ends in np³. The number tells you how many p-electrons they have!

Points to Remember

Group 13 outer configuration is ns²np¹.

Group 14 outer configuration is ns²np².

Group 15 outer configuration is ns²np³.

These elements do not occur as free atoms.

Problem 9.1: Ionic Radii Of M³⁺ Ions

Atomic numbers of the group 13 elements are in the order B < Al < Ga < In < Tl. Arrange these elements in increasing order of ionic radii of M³⁺.

Solution: The given elements are in an increasing order of atomic number. As we go down the group 13, their general outer electronic configuration is \(ns^2 np^1\). M³⁺ is formed by removal of three electrons from the outermost shell 'n'. In the M³⁺ the 'n-1' shell becomes the outermost. Size of the added 'n-1' shell increases down the group. Therefore the ionic radii of M³⁺ also increase down the group as follows:

\(B^{3+} < Al^{3+} < Ga^{3+} < In^{3+} < Tl^{3+}\)

Teacher's Note

When we remove electrons from an atom, the remaining electrons are pulled closer to the nucleus. This makes the ion smaller. The pattern is same down the group.

Exam Trick

Remember: When you go down a group and form M³⁺ ions, the ionic radius increases because the new outermost shell is larger.

Points to Remember

M³⁺ ions are formed by removing 3 electrons.

The n-1 shell becomes the outermost shell.

Ionic radius increases down the group.

Problem 9.2: Atomic Radius Of Gallium

Why the atomic radius of Gallium is less than that of aluminium?

Solution: Atomic radius increases down the group due to added new shell. 'Al' does not have 'd' electrons. As we go from Al down to 'Ga' the nuclear charge increase by 18 units. Out of the 18 electrons added, 10 electrons are in the inner 3d subshell. 'd' Electrons offer poor shielding effect. Therefore, the effects of attraction due to increased nuclear charge is experienced prominently by the outer electrons of 'Ga' and thus its atomic radius becomes smaller than that of 'Al'.

Teacher's Note

The d-electrons do not protect the outer electrons well from the nucleus. This is why Ga is smaller than Al even though it comes after Al in the periodic table.

Exam Trick

Remember: When inner d-electrons are added, they provide poor shielding. So the outer electrons feel more pull from the nucleus, making the atom smaller.

Points to Remember

D-electrons offer poor shielding effect.

Increased nuclear charge pulls electrons closer.

Ga is smaller than Al.

Problem 9.3: First Ionization Enthalpy Trend

The values of the first ionization enthalpy of Al, Si and P are 577, 786 and 1012 kJ/mol respectively. Explain the observed trend.

Solution: The trend shows increasing first ionization enthalpy from Al to Si to P. Al, Si and P belong to 13 period in the periodic table. They have same valence shell. Due to the increased nuclear charge electrons in the valence shell are more tightly held by the nucleus as we go from Al to Si to P. Therefore more energy is required to remove an electron from its outermost shell.

Teacher's Note

As we move right in a period, the nucleus has more positive charge. This holds the electrons tighter, so more energy is needed to remove them.

Exam Trick

Remember: Across a period from left to right, ionization enthalpy increases because the nuclear charge increases while the shell number stays the same.

Points to Remember

Ionization enthalpy increases across a period.

Nuclear charge pulls electrons tighter.

More energy is needed to remove electrons.

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