CBSE Class 11 Chemistry Classification Of Elements And Periodicity In Properties Of Elements Notes Set 05

These revision notes for Class 11 Chemistry Chapter 3, Classification of Elements and Periodicity in Properties, cover the modern periodic law, the IUPAC naming of elements beyond atomic number 100, the four blocks of the periodic table and the periodic trends in atomic radius, ionic radius, ionisation enthalpy, electron gain enthalpy, electronegativity and valency. Use them for quick revision before tests and the annual examination.

1. Modern Periodic Law and the Present Form of the Periodic Table

1. In 1913, the English physicist Henry Moseley modified Mendeleev’s periodic table; the law on which it is based is known as the modern periodic law.
2. This law states: “the physical and chemical properties of the elements are periodic functions of their atomic numbers.”
3. The modern periodic law is essentially the consequence of the periodic variation in electronic configurations, which determine the physical and chemical properties of elements and their compounds.
4. The long form of the periodic table consists of horizontal rows called periods and vertical columns called groups or families. It contains 7 periods and 18 groups.
5. Elements having similar outer electronic configuration in their atoms are arranged in the same group. The period number corresponds to the highest principal quantum number (n) of the elements in that period.
6. The first period contains 2 elements. The subsequent periods contain 8, 8, 18, 18 and 32 elements respectively. The seventh period, like the sixth, can hold a maximum of 32 elements.
7. In this form of the periodic table, 14 elements of both the sixth and the seventh periods (lanthanoids and actinoids respectively) are placed in separate panels at the bottom.

2. Nomenclature of Elements with Atomic Numbers > 100

The names are derived by using roots for the three digits in the atomic number of the element, followed by adding ‘–ium’ at the end. The roots for the digits are:

For example, element 109 = un + nil + enn + ium = unnilennium (Une); element 112 = un + un + bi + ium = ununbium (Uub).

3. Important Properties of s, p, d and f-block Elements

(i) s-block elements

General electronic configuration of the elements of this block is ns^1–2 (where n = 1, 2, …).
1. Group 1 (alkali metals) and group 2 (alkaline earth metals) elements belong to this block.
2. These are reactive metals with low ionisation enthalpies.
3. Their metallic and reactive character increases on going down a group.
4. The compounds of s-block elements, with the exception of those of lithium and beryllium, are predominantly ionic.

(ii) p-block elements

The general configuration of these elements can be written as ns²np^1–6 (where n = 2, 3, …).
1. Groups 13 to 18 (excluding He) belong to this block. Their last electron enters a p-orbital.
2. The s- and p-block elements together are known as representative elements or main group elements.
3. Group 15 members are called pnicogens, group 16 members are called chalcogens and group 17 members are called halogens.

(iii) d-block elements

1. Elements of groups 3 to 12 of the periodic table belong to the d-block.
2. General electronic configuration of elements of this block is (n − 1)d^1–10 ns^0–2 (where n = 4 to 7).
3. They are all metals. They mostly form coloured ions, exhibit variable valence and paramagnetism, and are used as catalysts.
4. Zn, Cd and Hg have the electronic configuration (n − 1)d¹⁰ ns² and do not show most of the properties of transition elements.
5. They are also called transition elements as they form a bridge between the chemically active metals of the s-block and the less active elements of groups 13 and 14.

(iv) f-block or inner transition elements

1. The last electron enters an f-orbital.
2. General configuration is (n − 2)f^1–14 (n − 1)d^0–1 ns², where n = 6 to 7.
3. Two series, 4f (lanthanoids) and 5f (actinoids), belong to this block.

4. Valence Electrons

The electrons present in the outermost shell are called valence electrons.
1. For s-block elements, group number = number of valence electrons.
2. For p-block elements, group number = 10 + number of valence electrons.
3. For d-block elements, group number = number of electrons in the (n − 1)d subshell + number of electrons in the valence (ns) shell.

5. Periodic Trends in Properties of Elements

(i) Atomic radius

The atomic radius is defined as “the distance from the centre of the nucleus to the outermost shell of electrons.” Depending upon the nature of the combining atoms, atomic radius can be of the following types:
(a) Covalent radius: \( r_{\text{covalent}} = \frac{1}{2} \) × [internuclear distance between two covalently bonded atoms]
(b) Metallic radius: one-half of the internuclear distance separating the metal cores in a metallic crystal.
(c) van der Waals’ radius: one-half of the distance between the nuclei of two identical non-bonded isolated atoms (of neighbouring molecules).

Variation of atomic radii in the periodic table:
1. Variation along a period: The atomic radii generally decrease from left to right along a period. Within a period, electrons are added to the same valence shell, so the effective nuclear charge increases and the outer electrons are held more tightly by the nucleus, which results in smaller radii.
2. Variation in a group: The atomic radii generally increase on moving down a group. The electrons are added in successive shells (the principal quantum number n increases), and the inner-shell electrons shield the valence electrons from the pull of the nucleus. The distance between the valence electrons and the nucleus therefore increases, and so does the atomic radius.

(ii) Ionic radius

The ionic radius is defined as the effective distance from the centre of the nucleus of an ion up to which it has an influence in the ionic bond.
Internuclear distance = radius of cation + radius of anion
(a) Variation of ionic radii: It shows the same trend in the periodic table as the atomic radii, i.e. it decreases across a period and increases down a group. A cation is smaller and an anion is larger than its parent atom.
(b) Isoelectronic species and their radii: These are atoms or ions which have the same number of electrons but different nuclear charges. The radii of isoelectronic species increase as the magnitude of the nuclear charge decreases, e.g. Al³⁺ < Mg²⁺ < Na⁺ < F⁻ < O²⁻.

(iii) Ionisation enthalpy (IE)

The minimum amount of energy required to remove the most loosely bound electron from an isolated gaseous atom (X) in its ground state is called the ionisation enthalpy.
X(g) → X⁺(g) + e⁻; Δ_iH = +ve
or X(g) + Δ_iH → X⁺(g) + e⁻

1. Variation of IE in the periodic table: Generally, IE increases on moving from left to right in a period and decreases on moving down a group. Atoms with half-filled and fully-filled orbitals are more stable and thus have high IE.
2. Factors on which IE depends: (a) atomic size – varies inversely; (b) screening effect – varies inversely; (c) nuclear charge – varies directly.
3. Exceptions: IE of group 2 elements is higher than that of the corresponding group 13 elements because of the fully-filled configuration (ns²).
4. Similarly, IE of group 15 elements is higher than that of the corresponding group 16 elements because of the half-filled configuration (np³).

(iv) Electron gain enthalpy (Δ_egH)

When an electron is added to a neutral gaseous atom in its ground state to convert it into a negative ion, the enthalpy change accompanying the process is called the electron gain enthalpy (Δ_egH).
X(g) + e⁻ → X⁻(g)
Variation in the periodic table: Generally, electron gain enthalpy becomes more negative on moving from left to right in a period and less negative on moving down a group.

Factors affecting electron gain enthalpy:
1. Atomic size: varies inversely.
2. Nuclear charge: varies directly.
3. Configuration: Half-filled and fully-filled orbitals are stable, therefore the electron gain enthalpy of such atoms is low, and sometimes energy has to be supplied rather than released (positive Δ_egH).
4. Exception: The order of electron gain enthalpy (negative values) is Cl > F > S > O. F and O atoms have a small size and high charge density, so the incoming electron suffers more repulsion and they have lower electron gain enthalpy than Cl and S.
5. Chlorine has the highest (most negative) electron gain enthalpy, but the oxidising power of fluorine is greater than that of chlorine.

(v) Electronegativity

It is the ability of an atom in a compound to attract the shared pair of electrons towards itself.

6. Variation of Electronegativity in the Periodic Table

1. Variation along a period: On moving from left to right across a period, the effective nuclear charge increases and size decreases, so electronegativity increases due to the greater attraction between the shared pair of electrons and the nucleus.
2. Variation along a group: On moving down a group, the atomic size increases, the force of attraction between the shared pair of electrons and the nucleus decreases, and hence electronegativity decreases.

Factors which affect the magnitude of electronegativity:
1. Atomic radius: As the atomic radius increases, the electronegativity decreases, i.e. \( \text{Electronegativity} \propto \frac{1}{\text{Atomic radius}} \)
2. Effective nuclear charge: Electronegativity increases as the effective nuclear charge on the nucleus increases, i.e. Electronegativity ∝ Z_eff
3. Oxidation state: Electronegativity increases as the oxidation state (i.e. the positive charge) of the atom increases.
4. s-character: If the s-character in the hybridisation state of the central atom increases, electronegativity also increases.

Different scales for calculating electronegativity: (a) Pauling scale (b) Mulliken–Jaffe scale.

7. Valency or Oxidation States

Valency is the most characteristic property of the elements and can be understood in terms of their electronic configuration. It is the combining power of an element. Transition elements and actinoids exhibit variable valency.

Variation of Valency in the Periodic Table

1. Along a period from left to right, valency with respect to hydrogen increases gradually from 1 to 4 and then decreases to 1.
2. Valency of noble gases is taken as zero.
3. In a group, all the elements have the same valency because they have the same number of valence electrons.