Periodic Trends in Properties of Elements

Full Notes

3.7.1 Trends in Physical Properties

Periodic physical properties of elements change in predictable patterns as you move across periods or down groups. These arise due to variations in atomic structure and nuclear attraction.

Atomic Radius

The atomic radius refers to the distance from the nucleus to the outermost shell. It helps explain trends in bonding and size.

• Atomic radius decreases across a period due to increased nuclear charge.

• Atomic radius increases down a group due to addition of new shells and shielding effect.

Ionic Radius

Ionic radius refers to the size of an ion (distance from nucleus to outermost shell). It varies across a period, depending on whether the ion is a cation or an anion.

• Cations are smaller than parent atoms.
• Anions are larger than parent atoms.
• Among isoelectronic species, greater nuclear charge means smaller radius.

First Ionization Enthalpy

Ionization enthalpy is the energy required to remove the outermost electron from a gaseous atom, always positive since energy is absorbed.

Equation: X(g) → X⁺(g) + e^–

• Increases across a period due to stronger nuclear attraction.

• Decreases down a group as outer electrons are further from the nucleus and more shielded.
• Exceptions exist due to half-filled and fully-filled subshell stability. Across period 2 and 3, there is a decrease in 1st ionisation energy from group 2 to group 3 and from group 5 to group 6.

Exceptions to 1st Ionization Energy trend exceptions (Period 3 (Na to Ar))

Aluminium (Al):
Lower than magnesium (Mg), meaning electron removed must have less attraction to the nucleus (be further away and higher energy). This is now explained as the outer electron being in a 3p orbital (higher in energy than the 3s orbital the outer electron in magnesium is in).

Sulfur (S):
Lower than phosphorus (P), meaning outermost electron removed must be higher in energy and easier to remove than for sulfur. This is now explained as being due to electron pairing in a 3p orbital, causing repulsion and giving evidence that the 3p sub shell must contain 3 orbitals (as the 4th electron in the sub shell has to pair up with another electron in an orbital).

Electron Gain Enthalpy

This is the energy released or absorbed when a gaseous atom of an element gains an electron. It indicates the tendency of an atom to accept electrons.

Equation: X(g) + e^– → X^–(g)

• Becomes more negative across a period (stronger attraction).
• Becomes less negative down a group (weaker attraction).
• Halogens have the most negative values.

Electronegativity

Electronegativity measures an atom’s ability to attract shared electrons in a bond.

• Increases across a period.
• Decreases down a group.
• Highest in fluorine, lowest in alkali metals.

3.7.2 Periodic Trends in Chemical Properties

These trends define how atoms behave in reactions, including what kinds of bonds they form and their oxidation states.

Periodicity of Valence or Oxidation States

Valency refers to the combining capacity of an element. It depends on the number of electrons lost, gained, or shared to achieve stability.

• Valency shows a periodic pattern across s- and p-blocks.
• Transition elements show variable oxidation states due to partially filled d-orbitals.

Anomalous Properties of Second Period Elements

The first element in each group often shows unique properties unlike its group members.

• Caused by small size, high electronegativity, and absence of d-orbitals.
• Leads to diagonal relationships (e.g., Li ↔ Mg).
• Allows multiple bonding

3.7.3 Periodic Trends and Chemical Reactivity

Reactivity of elements correlates with how easily they can lose or gain electrons. This affects whether they behave as metals or non-metals.

• Metals: More reactive down a group, less across a period.
• Non-metals: More reactive across a period, less down a group.

Trends are governed by ionization enthalpy, electron gain enthalpy, and electronegativity.

Summary

• Atomic and ionic radii decrease across periods and increase down groups.
• Ionization enthalpy increases across a period but decreases down a group.
• Electron gain enthalpy becomes more negative across a period.
• Electronegativity rises across a period and falls down a group.
• Valency patterns repeat periodically and explain reactivity trends.
• Second-period elements show unique behaviour due to small size and lack of d-orbitals.