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S1.1 - Introduction to the particulate nature of matter S1.2 - The nuclear atom S1.3 - Electron configurations S1.4 - Counting particles by mass - The mole S1.5 - Ideal gases S2.1 - The ionic model S2.2 - The covalent model S2.3 - The metallic model S2.4 - From models to materials S3.1 - The periodic table - Classification of elements S3.2 - Functional groups - Classification of organic compounds R1.1 - Measuring enthalpy changes R1.2 - Energy cycles in reactions R1.3 - Energy from fuels R1.4 - Entropy and spontaneity AHL R2.1 - How much? The amount of chemical change R2.2 - How fast? The rate of chemical change R2.3 - How far? The extent of chemical change R3.1 - Proton transfer reactions R3.2 - Electron transfer reactions R3.3 - Electron sharing reactions R3.4 - Electron-pair sharing reactions

S2.3 - The metallic model

2.3.1 Metallic Bonding and Properties of Metal 2.3.2 Strength of Metallic Bonding 2.3.3 Transition Metal (AHL)

Strength of Metallic Bonding and Melting Points

Specification Reference S2.3.2

Quick Notes:

  • Metallic bond strength depends on:
    • The charge of the metal cation
    • The radius of the metal ion
  • Stronger metallic bonds:
    • Occur with higher charge (more delocalized electrons)
    • Occur with smaller ionic radius (stronger electrostatic attraction)
  • Melting point trends in the s- and p-block:
    • Increase across a period (greater charge, smaller radius)
    • Decrease down a group (larger ions, weaker attraction)

Full Notes:

What Affects the Strength of a Metallic Bond?

The metallic bond is the electrostatic attraction between positive metal ions and delocalized electrons.

Its strength depends on:

Trends in Melting Points (s- and p-block Metals)

Across a Period (Left to Right)

Down a Group (Top to Bottom)

Summary Table: Factors Affecting Metallic Bonding

Factor / Trend Effect on Bond Strength Reason
Increased Cation Charge Stronger More delocalized electrons and higher electrostatic attraction to the electron sea
Decreased Ionic Radius Stronger Electrons are closer to nuclei, increasing attraction
Across a period (Group 1 → Group 3) Melting point increases Higher charge density and smaller radii
Down a group Melting point decreases Larger ions reduce attraction to delocalized electrons

Note

s- and p-block metals show clear trends based on simple electrostatic principles.

Transition metals follow more complex patterns due to d-electrons and variable oxidation states, and are not the focus here.

Summary