Chemistry (9701) | Metallic bonding

Welcome to the study of Metallic Bonding! This chapter explains how metals stick together and why they behave so uniquely—from conducting electricity to being bent and hammered without shattering. Understanding this bonding is key to explaining why metals are essential in our everyday lives, from the wires that power our homes to the coins in our pockets.

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1. Defining Metallic Bonding (The Core Concept)

1.1 The Definition (Syllabus 3.3.1)

The definition of metallic bonding is simple but precise. You must know this verbatim:

Metallic bonding is the electrostatic attraction between positive metal ions (cations) and a surrounding "sea" of delocalised electrons.

Breaking down the definition:

• Positive Metal Ions (Cations): Metal atoms (like Sodium, Na, or Copper, Cu) lose their outermost (valence) electrons easily. These atoms become positive ions, retaining their inner electron shells and nucleus.
• Delocalised Electrons: The valence electrons lost by the atoms are no longer attached to any specific ion or nucleus. Instead, they are free to move throughout the entire structure. They are "delocalised."
• Electrostatic Attraction: This is the powerful force holding everything together. It’s the attraction between the positively charged ions and the collectively shared negatively charged sea of electrons.

💡 Analogy: The Raisin Pudding Model (or Sea of Electrons Model)
Think of the positive metal ions as raisins (or golf balls) fixed in a regular pattern.
The delocalised electrons are the pudding (or glue) holding the raisins together. The electrons are flowing freely around the fixed positive cores, constantly providing the strong attraction that locks the structure into place.

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2. The Giant Metallic Structure (Lattice)

2.1 Structure and Arrangement (Syllabus 4.2.1d)

Metallic bonding results in a specific type of structure called a giant metallic lattice.

• This is a repeating, highly ordered 3D arrangement of the positive ions.
• The delocalised electrons exist in the spaces between these ions, acting as a flexible cement.

Example: Copper (Cu)
In a piece of solid copper (Cu), you have an endless repeating pattern of Cu²⁺ ions surrounded by a sea of mobile electrons, holding the entire structure together.

Did you know? The number of delocalised electrons per atom affects the bond strength. Elements like Sodium (Group 1) only contribute one electron, resulting in a weaker bond and lower melting point compared to transition metals like Iron, which contribute two or more electrons, leading to much stronger bonding.

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3. Explaining Physical Properties of Metals

The properties of metals are easily explained by referring back to the "positive ions in a sea of electrons" model (Syllabus 4.2.2).

3.1 High Melting and Boiling Points

Metals generally have very high melting and boiling points (M.P. and B.P.) because:

1. The electrostatic attraction between the cations and the delocalised electrons is extremely strong.
2. A vast amount of thermal energy is required to overcome these strong attractive forces and break the giant lattice structure, allowing the ions to move past one another (melting) or separate completely (boiling).

Think of it this way: It takes a huge amount of heat energy to boil water (a simple molecular structure) compared to breaking the strong walls of a giant skyscraper (the metallic lattice).

3.2 Excellent Electrical Conductivity

Metals are outstanding electrical conductors in both the solid and liquid (molten) states. This is a direct consequence of the bonding model.

1. Electricity is the flow of charge carriers.
2. In a metal, the delocalised electrons are mobile charge carriers.
3. When a voltage is applied across the metal, these electrons are free to move rapidly through the lattice, carrying the electrical current.

Note: Unlike ionic compounds, which only conduct when molten or dissolved (when the ions become mobile), metals conduct electricity when solid because the electrons are always mobile.

3.3 Good Thermal Conductivity

Metals are also excellent conductors of heat (thermal energy).

• When one end of a metal is heated, the delocalised electrons gain kinetic energy.
• These fast-moving, mobile electrons rapidly collide with the metal ions throughout the lattice, effectively transferring thermal energy quickly from the hot region to the colder regions.

3.4 Malleability and Ductility

Malleability means metals can be hammered into sheets; ductility means they can be drawn into wires. Most other solid structures (like giant ionic or simple molecular) are brittle and shatter when hit.

Why are metals flexible?

1. Metallic bonds are non-directional. The attractive force is uniform in all directions throughout the sea of electrons.
2. When a force is applied (e.g., hammering), the layers of positive ions are able to slide past each other.
3. Crucially, because the delocalised sea of electrons is uniform, new electrostatic attractions are instantly formed, and the positive ions never repel each other (unlike in an ionic lattice, where sliding layers causes positive ions to meet positive ions, leading to shattering).

Analogy: A stack of lubricated pipes.
If you push a stack of dry bricks (ionic), they shatter. If you push a stack of pipes covered in oil (the oil represents the electron sea), the pipes slide smoothly without breaking the overall structure.

3.5 Solubility (Syllabus 4.2.2)

Metals are generally insoluble in standard solvents (like water or organic solvents).
This is because metallic bonding is so strong. The energy released by forming bonds with the solvent (solvation or hydration energy) is insufficient to overcome the massive amount of energy required to break the strong electrostatic attractions in the giant metallic lattice.