4 Classification of Elements

The physical and chemical properties of elements show regular patterns when the elements are arranged systematically. These patterns led to the development of the periodic table, one of the most important tools in chemistry for organising elements, understanding relationships among them, and predicting their properties.

Early attempts at classification, such as Mendeleev’s periodic table, were based on atomic mass. However, some inconsistencies remained unresolved. In 1913, Henry Moseley established that the atomic number (Z), not atomic mass, is the fundamental property governing periodicity.

Main features of the modern periodic table

The modern periodic table is based on the electronic configuration of elements. Its main structural features are as follows:

1. There are seven rows, called periods, indicated by 1, 2, 3, 4, 5, 6, and 7.
2. Across a period, the outermost electronic configuration changes gradually, which causes the gradual change in properties. In the next period, a similar outer electronic configuration reappears. Hence, the properties also repeat.
3. There are 18 columns, called groups, numbered 1 to 18 according to IUPAC.
4. Elements in the same group have the same valence-shell electronic configuration and show similar chemical properties. For example, Li, Na, and K have the same outer electronic configuration of ns¹, and they are placed in group 1.
5. Dissimilar elements do not fall in the same group since subgroups have been separated.
6. The modern periodic table completely separates metals from non-metals.
7. The left-hand side of the table consists of highly reactive metals, and the right-hand side consists of non-metals; the middle portion contains transition elements.
8. Elements have been divided into four distinct blocks: s-block, p-block, d-block, and f-block depending on which orbital receives the last electron
9. Lanthanides and actinides are not placed in the main body of the periodic table.

1. Explains periodicity: Repeating electronic configurations explain why properties recur at regular intervals.
2. Resolves Mendeleev’s anomalies: Atomic number ordering justifies the placement of Ar before K, Co before Ni, and Te before I pairs.
3. No problem with isotopes: Isotopes have the same Z, so they automatically share one position.
4. Clear separation of metals and non-metals: The diagonal boundary (from boron to astatine) makes classification straightforward.
5. Block division: s, p, d, f blocks allow systematic study of chemical properties.

• Position of hydrogen: Hydrogen shares properties with both alkali metals (Group 1) and halogens (Group 17). It does not fit perfectly into either group.
• Position of helium: Helium’s outermost configuration (1s²) resembles Group 2, but it is placed in Group 18 due to its chemical inertness. Also, it lacks a p-orbital, unlike its p-block neighbours.
• Lanthanides and actinides: Placing the f-block elements at the bottom of the table (outside the main body) technically violates the continuous arrangement expected by the periodic law.

Conclusion

The modern periodic table provides a systematic framework for understanding the properties of elements. By studying periodic trends and electronic configurations, chemists can predict the behaviour, reactivity, and bonding patterns of elements with great accuracy.

• There are seven rows, which are called periods and indicated by 1, 2, 3, 4, 5, 6, and 7.
• In a period, the outermost electronic configurations change gradually, which causes the gradual change in properties. In the next period, a similar outer electronic configuration reappears. Hence, the properties also repeat.
• There are 18 columns, which are called groups. Each group consists of a set of elements having the same outer electronic configuration; hence, they show similar properties. For example, Li, Na, and K have the same outer electronic configuration of ns¹, and they are placed under group 1.
• Dissimilar elements do not fall in the same group since subgroups have been separated.
• Elements have been divided into four distinct blocks: s-block, p-block, d-block, and f-block.

(contain s-electrons in valence shell)

Valence configuration: ns¹⁻²

1. All are active metals (except hydrogen).
2. Oxidation states: +1 (Group 1) and +2 (Group 2).
3. They form basic oxides and hydroxides.
4. They impart characteristic colour to the flame.
5. Generally, they form ionic salts with nonmetals.
6. Low ionisation potentials, hence good reducing agents
7. Very small electron affinities.
8. Solids at room temperature (except H₂)

1. Valence electronic configuration: ns² np¹⁻⁶
2. Mostly non-metals; also include metalloids and some metals.
3. Show variable oxidation states.
4. Generally form acidic oxides (except group 18).
5. Generally, they form covalent compounds. Halogens readily form salts with alkali metals
6. They have high ionization potentials, and large electron affinities.
7. They may be solids, liquids, or gases at room temperature (Br₂ is a liquid).

(contain d-electrons in valence shell)

1. General electronic configuration: (n−1)d¹⁻¹⁰ns⁰⁻²
2. Lie between s- and p-block, showing “transitional” properties.
3. Almost all are metals (typically dense, hard, high-melting).
4. Often have variable oxidation states due to partly filled d-orbitals. (Except Sc, Zn, Cd, etc.)
5. Many compounds are brightly coloured — a key identification feature.
6. Many exhibit paramagnetic or ferromagnetic behaviour.
7. Excellent catalysts (e.g., Fe in the Haber process, Ni in hydrogenation).

(contain f-electrons in valence shell)

1. Electronic configuration: ns² (n−) d⁰⁻¹(n−2) f¹⁻¹⁴
2. Also called inner transition elements. include Lanthanides and actinides.

Lanthanides (₅₈Ce – ₇₁Lu)

3. All lanthanides closely resemble lanthanum in properties and are chemically similar to each other.

5. Mostly trivalent. (Except for cerium (III and IV) and europium (III and II))
6. Most lanthanides are widely used in lasers and magnets
7. Commonly used in sunglasses — they absorb UV and IR radiation

Actinides (₉₀Th – ₁₀₃Lr)

8. All actinides are radioactive.
9. These are highly electropositive (show +3, to +6 oxidation states)
10. These metals tarnish in the air.
11. They have several isotopes.
12. They react with boiling water or dilute acids to give H₂ gas.
13. These directly combine with non-metals.

IUPAC (International Union of Pure and Applied Chemistry) numbers the main-body groups from 1 to 18. This replaces the older system that used Roman numerals with A/B suffixes. The table below maps the two systems:

Nuclear Charge (Z)

The total positive charge on the nucleus of an atom, equal to the number of protons. E.g., the nuclear charge of sodium (Na, Z = 11) is +11.

In multi-electron atoms, the inner electrons partially block (or “shield”) the outer electrons from full nuclear attraction. Because of this, the outermost electrons experience a reduced, effective attraction.

Periodic Trend and Periodicity — Periodic Properties of Elements. The properties below show a regular, predictable variation across periods and down groups. Understanding these trends allows us to predict and compare element behaviour without memorising each case separately.

1. Atomic Radius

An atom is considered to be spherical. Hence, the size of an atom is expressed in terms of atomic radius.

Variation of atomic radius across a period — The nuclear charge increases, and electrons are added to the same shell; attraction between the nucleus and the outermost electrons increases. Hence, atomic size decreases with increasing atomic number in a period.

Variation of atomic radius down a group — The nuclear charge increases, which means the size should decrease. But a new electron shell is added away from the nucleus, increasing the size of the atom. The latter effect outweighs the former. Hence, atomic size increases down the group.

On going down in a group, the atomic radius increases because of the increase in an extra shell.

Ionic Radius

Ionic radius may be defined as the distance between the centre of the nucleus and a point at which the nucleus has its influence on its electron in an ion.

1. Cations are smaller than the parent atom. This is generally due to the removal of the valence shell and an increase in the effective nuclear charge.
2. Anions are larger than the parent atom. This can be explained by a decrease in the effective nuclear charge in an anion.
3. Isoelectronic Species — Ions or atoms having the same number of electrons but different nuclear charges are called isoelectronic species. Their size decreases with increasing nuclear charges.

Successive Ionization Energies

It is possible to remove more than one electron from an atom successively.

The minimum amount of energy required to remove the most loosely held electron from an isolated gaseous atom in its ground state is called the first ionization energy of that element.

The minimum amount of energy required to remove the most loosely held electron from an isolated gaseous monopositive ion in its ground state is called the second ionization energy of that element.

The minimum amount of energy required to remove the most loosely held electron from an isolated gaseous dipositive ion in its ground state is called the third ionization energy of that element.

For an element, IE₁ < IE₂ < IE₃

Factors Affecting Ionization Energy

The ionisation energy of an electron depends upon the following factors.

1. Atomic size: As the size of the atom decreases, electrons are more strongly attracted by the nucleus, and more energy is required to remove the electron; hence, the greater the IE.
2. Nuclear charge: As the nuclear charge increases, electrons are more strongly attracted by the nucleus, and more energy is required to remove the electron. Hence, the greater the IP.
3. Electronic configuration: Half-filled and fully filled orbitals are more stable than other configurations; hence, the IE is higher in such cases. For example, Beryllium has a higher IE than Boron for this reason.
4. Screening or Shielding effect: In multi-electron atoms, in addition to attraction between the nucleus and the electrons, there is repulsion between the electrons, so that the valence electron experiences less nuclear attraction. This effect of inner electrons is called the screening or shielding effect. As the screening effect increases, electrons are less strongly attracted by the nucleus, and less energy is required to remove an electron. Hence, ionization energy decreases with an increase in screening effect.

Variation of IE down a group — Going down a group, although nuclear charge increases, the screening effect increases, and the size of the atom increases, so, electrons are less strongly attracted, and less energy is required to remove an electron. Hence, ionization energy decreases.

Variation of IP across a period — Going across a period, nuclear charge increases, and the size of the atom decreases, so electrons are more strongly attracted by the nucleus, and more energy is required to remove the electron. Hence, ionization energy increases.

Main Factors Affecting EA

1. Atomic size: The smaller the size of an atom, the greater the attraction between the outermost electrons and the nucleus, hence the greater the electron affinity.
2. Nuclear charge: The greater the nuclear charge in the atom, the greater the attraction between the outermost electrons and the nucleus, thus increasing the EA.
3. Electronic configurations: Half-filled and fully filled orbitals are more stable than other configurations; hence, the EA is less in such cases.

Variation of electron affinity across a period — Going across a period, nuclear charge increases, and the size of the atom decreases, due to which the tendency to attract incoming electrons increases. Hence, electron affinity increases. However, there are irregularities in the general trend, mainly due to electronic configurations.

Variation of electron affinity down a group — Going down in a group, although the nuclear charge increases, the screening effect and the size of the atom increase, due to which the tendency to attract incoming electrons decreases. Hence, electron affinity decreases.

Variation of electronegativity in a period: Across a period, nuclear charge increases, and the size of the atom decreases, due to which the tendency of the atom to attract a shared pair of electrons increases; hence, electronegativity increases. e.g.

Variation of electronegativity in a group: Down a group, although nuclear charge increases, the screening effect increases, and the size of the atom increases; due to this, electronegativity decreases down the group. e.g.

(to be studied after metallic bonding)

The tendency of an element to lose electron(s) to form a cation is a characteristic of a metal. The elements with low ionization energy and low electronegativity (electropositive) show such characteristics. The elements on the left and lower left side of the periodic table are such elements.

Variation of metallic character in a period: On moving from left to right of the periodic table, the ionization energies of elements increase; hence the metallic character also decreases gradually.

Variation of metallic character in a group: As moving down in a group of the periodic table, ionization energy decreases, hence metallic character increases.

* Electron affinity shows exceptions at Groups 2, 5, and 18 due to stable electronic configurations.