CBSE Class 10 Chemistry Periodic Classification of Elements Notes Set 24

Early Attempts at Classification

1. Dobereiner's triads: groups of three elements with similar properties in which the atomic mass of the middle element is roughly the average of the other two. Examples: Li, Na, K (7, 23, 39); Ca, Sr, Ba (40, 87.6, 137); Cl, Br, I (35.5, 80, 127). Only a few triads could be found, so the idea failed as a general scheme.
2. Newlands' Law of Octaves: when elements are arranged in increasing atomic mass, every eighth element repeats the properties of the first, like the notes of a musical scale. It worked up to calcium only, assumed that only 56 elements existed, and forced unlike elements into the same slot.
3. Lothar Meyer plotted atomic volume against atomic mass and found that similar elements occupy similar positions on the curve.

Mendeleev's Periodic Table

Mendeleev's periodic law: the properties of elements are a periodic function of their atomic masses. He arranged about 63 known elements in horizontal rows (periods) and vertical columns (groups), so that elements with similar properties fell in the same group.
Strong points: he left gaps for undiscovered elements and predicted their properties (eka-aluminium became gallium, eka-silicon became germanium, eka-boron became scandium); he reversed the order of some elements (for example Te and I) to keep similar elements together; the noble gases fitted in later as a new group without disturbing the table.
Defects: the position of hydrogen was uncertain; isotopes had no place of their own; the atomic mass does not always increase steadily (Co and Ni, Ar and K); similar elements sometimes fell in different groups, and unlike elements in the same group.

Modern Periodic Law

Moseley showed that atomic number, not atomic mass, is the fundamental property. The modern periodic law states that the properties of elements are a periodic function of their atomic numbers. This removes the anomalies of reversed pairs, and isotopes (same Z) occupy the same position.

Structure of the Long Form Table

1. There are 7 horizontal rows called periods and 18 vertical columns called groups.
2. Period number = number of occupied shells. Periods 1 to 7 contain 2, 8, 8, 18, 18, 32 and 32 (incomplete) elements.
3. For groups 1 and 2, group number = number of valence electrons. For groups 13 to 18, valence electrons = group number minus 10.
4. Blocks: s-block (groups 1, 2), p-block (groups 13 to 18), d-block (groups 3 to 12, the transition elements) and f-block (lanthanides and actinides, the inner transition elements). The s- and p-block elements are the representative (normal) elements.
5. Special family names: group 1 alkali metals, group 2 alkaline earth metals, group 16 chalcogens, group 17 halogens, group 18 noble gases.

Periodic Trends

Atomic and Ionic Size

1. A cation is smaller than its parent atom (fewer electrons, same nuclear charge); an anion is larger than its parent atom.
2. For isoelectronic species, a higher nuclear charge gives a smaller radius. Example: Mg²⁺ < Na⁺ < F^– < O^2– < N^3–.
3. The van der Waals radius is half the distance between nuclei of two non-bonded atoms and is always larger than the covalent radius. The metallic radius is larger than the covalent radius of the same element.

Worked Examples

Example 1. An element has atomic number 16. Find its group and period.
Configuration: 2, 8, 6. Three shells, so period 3; six valence electrons, so group 16. The element is sulphur.

Example 2. What valency does an element of atomic number 17 show?
Configuration 2, 8, 7. It needs one electron to complete the octet, so its valency is 1 (chlorine).

Example 3. Atomic masses of X and Z are 39 and 132. Find the mass of the intermediate element Y if X, Y, Z form a triad.
Mass of Y = (39 + 132)/2 = 85.5.

Example 4. Arrange Cl^–, P^3–, S^2– and F^– in increasing order of size.
Cl^–, P^3– and S^2– are isoelectronic (18 electrons), and size increases as the nuclear charge falls: Cl^– < S^2– < P^3–. F^– has only two shells, so it is the smallest. Order: F^– < Cl^– < S^2– < P^3–.

Example 5. An element is in the second period and group 14. Find its atomic number.
Two shells, 4 valence electrons: configuration 2, 4, so Z = 6 (carbon).

Example 6. Which of N and P is larger, and why?
Phosphorus is larger because it has three shells, while nitrogen has only two.

Example 7. An element X in the third group (modern group 13) forms an oxide. Write the formula.
The element has valency 3 and oxygen has valency 2, so the oxide is X₂O₃.

Diagonal Relationship

Some elements of period 2 resemble the element one group to the right in period 3. The pairs are Li and Mg, Be and Al, B and Si. The similarity arises because the increase in charge on moving one group to the right is balanced by the increase in size on moving one period down.