10.5 Periodicity of Properties

The electronic configuration of elements exhibits a periodic variation with increasing atomic number. Consequently, the physical and chemical properties of elements also vary in a periodic manner. Elements with similar valence shell electronic configurations are placed in the same group.

• Chemical properties are primarily determined by the valence shell electronic configuration. Since elements within the same group share similar valence configurations, they exhibit similar chemical behaviors.
• Physical properties are often dependent on the size of the atoms. As atomic size changes gradually down a group and across a period, physical properties show a corresponding gradation.

In a period, the number of valence electrons increases from left to right, leading to a gradual change in both chemical and physical properties.

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10.5.1 Electron Affinity

Electron affinity is a measure of the energy change that occurs when an electron is added to a neutral atom in the gaseous state to form a negative ion (anion).

Definition: Electron affinity is the amount of energy released when an electron is added to the valence shell of an isolated gaseous atom to form a uninegative gaseous ion. This is also referred to as the first electron affinity.

The general reaction is:

${\mathrm{X}}_{(g)}+e^{-}\to {\mathrm{X}}_{(g)}^{-}+\text{Energy (Electron Affinity)}$

A more negative electron affinity value (more energy released) indicates a greater attraction for an electron.

Factors Affecting Electron Affinity

Three primary factors influence an atom's electron affinity:

1. Nuclear Charge: A higher nuclear charge (more protons) attracts the incoming electron more strongly, increasing electron affinity.
2. Atomic Size: A smaller atomic radius means the valence shell is closer to the nucleus, strengthening the attraction for an incoming electron.
3. Shielding Effect: Greater shielding from inner-shell electrons reduces the effective nuclear charge ($Z_{eff}$) felt by the incoming electron, decreasing electron affinity.

Ionization Energy→

Trends in Electron Affinity

1. Across a Period (Left to Right)

• Trend: Electron affinity generally increases (becomes more negative) from left to right across a period.
• Reasoning:
  • Nuclear Charge increases across a period → stronger attraction for incoming electron.
  • Atomic Size decreases → valence shell is closer to nucleus.
  • Shielding Effect remains relatively constant (electrons added to same principal shell).
• Conclusion: The increased nuclear charge and smaller size make it more energetically favorable for the atom to accept an electron, releasing more energy.

2. Down a Group (Top to Bottom)

• Trend: Electron affinity generally decreases (becomes less negative) from top to bottom in a group.
• Reasoning:
  • Atomic Size increases significantly as new energy shells are added → incoming electron placed farther from nucleus.
  • Shielding Effect increases due to additional inner electron shells → nucleus's attraction is weakened.
• Conclusion: The increased distance and shielding weaken the attraction of the nucleus for an incoming electron. Less energy is released when the anion is formed.

Important Exception: Cl vs F

Although the general trend predicts fluorine (Period 2) should have a higher electron affinity than chlorine (Period 3), chlorine has the highest electron affinity of all elements:

Reason: Fluorine's very small atomic size causes significant electron–electron repulsion in its compact 2p subshell when an extra electron is added. This repulsion reduces the energy released, making fluorine's electron affinity slightly lower than chlorine's.

Electron Affinities of Main Group Elements (kJ/mol)

Note: More negative values = higher electron affinity. Positive or near-zero values indicate the atom does not readily accept an electron.

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10.5.2 Electronegativity

Definition: Electronegativity is the power of an atom in a molecule to attract the shared pair of electrons (bonding electrons) towards itself.

Electronegativity is measured on the Pauling scale (dimensionless), where fluorine is assigned the highest value of 4.0.

Factors Affecting Electronegativity

1. Nuclear Charge: Higher nuclear charge → stronger pull on bonding electrons → higher electronegativity.
2. Atomic Radius: Smaller atomic radius → bonding electrons closer to nucleus → higher electronegativity.
3. Shielding Effect: Greater shielding → reduced effective nuclear charge → lower electronegativity.

Trends in Electronegativity

1. Across a Period (Left to Right)

• Trend: Electronegativity increases from left to right across a period.
• Reason: Nuclear charge increases, atomic radius decreases, and shielding remains roughly constant → $Z_{eff}$ increases → stronger attraction for bonding electrons.
• Example: Li (1.0) → Be (1.5) → B (2.0) → C (2.5) → N (3.0) → O (3.5) → F (4.0)

2. Down a Group (Top to Bottom)

• Trend: Electronegativity decreases down a group.
• Reason: Atomic radius increases and shielding increases as new shells are added → $Z_{eff}$ felt by bonding electrons decreases → weaker attraction.
• Example: F (4.0) → Cl (3.0) → Br (2.8) → I (2.5)

Electronegativity and Bond Type

The difference in electronegativity ($\Delta \chi$) between two bonded atoms predicts bond character:

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