Properties of metals Revision Notes

Metals are generally excellent conductors of heat. This is due to the presence of free electrons (delocalised electrons) within the metallic lattice, which transfer kinetic energy rapidly throughout the metal. When one part of a metal is heated, these free electrons gain energy and move quickly, passing energy to other parts of the metal, resulting in efficient heat conduction.

In contrast, non-metals are usually poor conductors of heat. Non-metals lack free electrons; their electrons are held tightly in covalent bonds or as lone pairs. Heat transfer in non-metals occurs mainly by vibrations of atoms or molecules, which is less efficient than electron movement. For example, diamond, a non-metal, is an exception with high thermal conductivity due to its strong covalent bonding and rigid lattice, but most non-metals like sulfur or phosphorus conduct heat poorly.

• Metals conduct heat well due to free electrons transferring kinetic energy.
• Non-metals conduct heat poorly because they lack free electrons; heat transfer occurs via atomic vibrations.

Metals conduct electricity in solid and molten states because of their free electrons. These delocalised electrons can move freely through the metallic lattice, allowing electric current to flow when a potential difference is applied.

Non-metals generally do not conduct electricity in solid or molten states because they lack free electrons or ions. Their electrons are involved in covalent bonds and are not free to move. Exceptions include graphite, a form of carbon, which conducts electricity due to delocalised electrons within its layered structure.

Electrical conductivity can be tested using a simple circuit with a bulb or ammeter. Metals will complete the circuit and light the bulb or show current, whereas most non-metals will not.

• Metals conduct electricity due to free-moving delocalised electrons.
• Non-metals are usually electrical insulators; exceptions like graphite conduct due to delocalised electrons in layers.

Metals are malleable and ductile. Malleability means metals can be hammered or rolled into thin sheets without breaking. Ductility means metals can be drawn into wires.

These properties arise because metal atoms are arranged in layers within a regular lattice. When force is applied, layers can slide over each other without breaking metallic bonds, as the delocalised electrons maintain the bond between layers.

Non-metals are generally brittle when solid. They tend to shatter or break rather than deform when force is applied. This is because their atoms are held in fixed positions by covalent or ionic bonds that do not allow layers to slide. For example, sulfur crystals are brittle and break easily.

• Metals are malleable and ductile due to layers of atoms sliding over each other in the metallic lattice.
• Non-metals are brittle and break rather than deform under stress.

Metals typically have high melting and boiling points. This is because strong metallic bonds exist between metal atoms, formed by the attraction between positive metal ions and delocalised electrons. A large amount of energy is required to overcome these bonds to change state.

Non-metals usually have lower melting and boiling points compared to metals. Many non-metals exist as simple molecules with weak intermolecular forces (e.g., van der Waals forces) that require less energy to overcome. For example, oxygen (O₂) and nitrogen (N₂) are gases at room temperature due to very weak intermolecular forces.

Some non-metals with giant covalent structures, such as diamond or silicon dioxide, have very high melting points due to strong covalent bonds throughout the lattice, but these are exceptions rather than the rule.

• Metals have high melting and boiling points due to strong metallic bonding.
• Non-metals generally have lower melting and boiling points due to weak intermolecular forces, except for giant covalent structures.

When metals react with dilute acids, they typically produce a salt and hydrogen gas. This reaction occurs because metals lose electrons (oxidation) to hydrogen ions (H⁺) from the acid, which are reduced to hydrogen gas (H₂).

The general equation for a metal M reacting with dilute hydrochloric acid (HCl) is:

M (s) + 2HCl (aq) → MCl₂ (aq) + H₂ (g) (for metals forming +2 ions) or M (s) + 2HCl (aq) → MCl (aq) + H₂ (g) (for metals forming +1 ions). The exact salt depends on the metal and acid used.

• Metals such as magnesium, zinc, and iron react with dilute acids to produce hydrogen gas and a salt.
• Copper, silver, and gold do not react with dilute acids because they are less reactive than hydrogen.
• Observations include effervescence (bubbling) due to hydrogen gas and the metal dissolving.
• The reaction is exothermic and faster with more reactive metals.

Only some metals react with cold water to form metal hydroxides and hydrogen gas. Alkali metals (Group 1) such as sodium and potassium react vigorously with cold water, while some alkaline earth metals like calcium react slowly.

The general reaction is: M (s) + 2H₂O (l) → M(OH)₂ (aq) + H₂ (g) (for alkaline earth metals) or 2M (s) + 2H₂O (l) → 2MOH (aq) + H₂ (g) (for alkali metals).

Less reactive metals like iron, zinc, and magnesium do not react with cold water but may react with steam.

• Alkali metals react vigorously with cold water, producing hydroxides and hydrogen gas.
• Calcium reacts slowly with cold water, forming calcium hydroxide and hydrogen.
• Magnesium and zinc do not react with cold water but react with steam.
• Observations include fizzing and the metal dissolving or disappearing.
• The solution formed is alkaline due to metal hydroxide formation.

Some metals that do not react with cold water react with steam (hot water vapour) to form metal oxides and hydrogen gas.

For example, zinc and iron react with steam as follows:

Zn (s) + H₂O (g) → ZnO (s) + H₂ (g) (zinc oxide formed) 2Fe (s) + 3H₂O (g) → Fe₂O₃ (s) + 3H₂ (g) (iron(III) oxide formed) Magnesium also reacts with steam to form magnesium oxide and hydrogen gas.

• Steam provides the energy needed for less reactive metals to react.
• Metal oxides formed are solids and often appear as a coating on the metal surface.
• Hydrogen gas is released and can be tested by its characteristic 'pop' with a lit splint.
• This reaction is important in metal extraction and corrosion processes.

Metals react with oxygen to form metal oxides. The nature of the oxide depends on the metal's reactivity and position in the reactivity series.

Highly reactive metals like sodium and calcium form basic oxides (e.g., Na₂O, CaO). Aluminium forms an amphoteric oxide (Al₂O₃) that can react with both acids and bases.

Less reactive metals like copper form oxides that are less reactive and often protective layers preventing further oxidation (e.g., CuO).

• Sodium reacts rapidly with oxygen to form sodium oxide: 4Na (s) + O₂ (g) → 2Na₂O (s).
• Magnesium burns with a bright white flame forming magnesium oxide: 2Mg (s) + O₂ (g) → 2MgO (s).
• Iron reacts slowly with oxygen and moisture to form rust (hydrated iron(III) oxide).
• Aluminium forms a thin, hard oxide layer that protects it from further corrosion.
• Observations include colour changes, flame colour, and formation of solid oxides.

The reactivity series ranks metals by their tendency to lose electrons and react. Metals higher in the series react more vigorously with acids, water, and oxygen.

For example, potassium and sodium react explosively with water, magnesium reacts slowly with cold water but faster with steam, and copper does not react with acids or water under normal conditions.

Understanding these reactions helps predict corrosion, metal extraction methods, and safe handling of metals.

• Reactivity decreases down the series: alkali metals > alkaline earth metals > transition metals > less reactive metals.
• More reactive metals displace hydrogen from acids; less reactive metals do not.
• Corrosion involves reaction with oxygen and moisture, forming oxides or hydroxides.
• Protective oxide layers can prevent further reaction (e.g., aluminium).
• Reactivity influences choice of metals for construction, storage, and chemical use.