Unit 9: Physical Properties of Metals

While all metals share common properties due to metallic bonding, the strength of this bond varies significantly, leading to differences in properties like melting point and hardness. The strength of the metallic bond is determined by the force of attraction between the positive cations and the delocalized electrons. Key factors include:

• Charge on the Cations: A greater positive charge on the metal ions leads to a stronger electrostatic attraction with the electron sea. For example, a Mg²⁺ ion will attract the electrons more strongly than a Na⁺ ion.
• Size of the Cations (Atomic Radius): Smaller cations can get closer to the electron sea, resulting in a stronger attraction and a more tightly packed lattice. For example, the attraction in sodium (Na) is stronger than in potassium (K) because the Na⁺ ion is smaller.
• Number of Delocalized Electrons: A higher number of delocalized valence electrons per atom creates a more negative "sea," increasing the overall attraction. Magnesium (Mg) contributes two valence electrons per atom, while Sodium (Na) only contributes one, resulting in stronger bonding in magnesium.

Generally, metallic bonding becomes stronger as you move from left to right across a period and weaker as you move down a group. The d-block metals typically exhibit very strong metallic bonding due to the involvement of d-orbital electrons.

The unique nature of the delocalized electron sea gives all metals a set of characteristic physical properties.

• Electrical and Thermal Conductivity: The mobile, delocalized electrons are free to move throughout the metal lattice. When a voltage is applied, they flow towards the positive terminal, creating an electric current. Similarly, they can efficiently transfer kinetic energy (heat) from one part of the metal to another.
• Malleability and Ductility: Malleability is the ability to be hammered into thin sheets, and ductility is the ability to be drawn into wires. In a metal, the layers of cations can slide over one another without breaking the non-directional attraction to the electron sea. This allows the metal to change shape without shattering.
• Lustre: The free electrons at the surface of a metal can absorb and re-emit photons of light, giving metals their characteristic shiny appearance.
• Sonority: This is the quality of producing a deep, ringing sound when struck. The delocalized electrons and uniform lattice structure allow sound waves to travel efficiently through the metal with little loss of energy.

Some physical properties are not uniform across all metals but depend directly on the strength of the metallic bond.

• Melting Point: Melting requires providing enough energy to overcome the forces holding the lattice together, allowing the cations to move freely. A stronger metallic bond requires more energy to break, resulting in a higher melting point. This is why d-block metals like tungsten (melting point 3422 °C) have extremely high melting points, while s-block metals like sodium (melting point 98 °C) have low ones.
• Tensile Strength: This measures a material's resistance to being pulled apart. Metals with stronger metallic bonds have higher tensile strength because more force is required to separate the cations from the electron sea. Steel, an alloy of iron with strong metallic bonding, has a very high tensile strength, making it ideal for construction.

Not all properties of metals are directly tied to the strength of their metallic bonds.

• Magnetic Properties: The strong magnetic property known as ferromagnetism is only found in a few metals, most notably Iron (Fe), Cobalt (Co), and Nickel (Ni). It is caused by the presence of unpaired electrons in their d-orbitals, which create tiny magnetic fields that can align under the influence of an external magnetic field. While other transition metals have unpaired electrons, the specific atomic structure of Fe, Co, and Ni allows these alignments to persist, creating permanent magnets.
• Density: Density is the ratio of mass to volume ($Density = Mass/Volume$). It depends on two main factors: the atomic mass of the element (heavier atoms lead to higher density) and the atomic radius (larger atoms packed less efficiently lead to lower density). Since atomic mass generally increases across a period and down a group more significantly than atomic radius, the densest metals (like Osmium and Iridium) are found in the lower part of the d-block.