Chemical Properties of Metal

1. Reactivity of Metals: Metals are very reactive and easily combine with other substances like oxygen, water, or acids.

2. Losing Electrons: When metals react, they lose electrons (tiny negatively charged particles in an atom). By losing electrons, the metal becomes positively charged.

3. Positively charged ions: Metals lose electrons and turn into positively charged ions. For example, when sodium reacts with water, it loses electrons and becomes a positively charged sodium ion.

4. Electropositive Elements: Metals are called electropositive elements because they easily lose electrons and become positively charged. This means metals have a tendency to give away their electrons in a chemical reaction.

5. Electronic Configuration: Electronic configuration is the basis of the chemical behaviour of elements. Most metals have up to three electrons in their outermost shell.

6. Formation of Ions: Metals have a tendency to lose their valence electrons to form positively charged ions, called cations.
      Na → Na⁺ + e⁻
      (2, 8, 1) (2, 8)
      Sodium → Sodium ion

      Mg → Mg²⁺ + 2e⁻
      (2, 8, 2) (2, 8)
      Magnesium → Magnesium ion

      Al → Al³⁺ + 3e⁻
      (2, 8, 3) (2, 8)
      Aluminium → Aluminium ion

7. Reaction with Oxygen: Metals combine with oxygen to form their oxides.

Metal + Oxygen → Metal oxide

The metal oxides are basic in nature. Metal oxides react with acids to form salt and water.

Metal oxide + Acid → Salt + Water

8. Reaction with Acid: Take dilute hydrochloric acid in a test tube. Add zinc dust to it. Take a glowing splinter near the mouth of the tube and observe. Most metals react with dilute acids to form metal salts, releasing hydrogen gas.

Metal + Dilute acid → Salt + Hydrogen gas.

9. Reaction with Water: Most metals do not react visibly and rapidly with cold water. However, some metals like sodium and potassium react with cold water to produce their hydroxides and hydrogen gas. Magnesium requires steam for a similar reaction.

1. Aim: To observe the reaction of different metals when heated in a flame.

2. Requirements

• Apparatus: Pair of tongs or spatula, knife, burner
• Chemicals: Aluminium, copper, iron, lead, magnesium, zinc, and sodium (handle sodium carefully under supervision)

3. Procedure

• Hold each metal sample at the top of the flame using tongs or a spatula.
• Observe which metal catches fire readily.
• Note any changes in the metal’s surface and the flame’s colour while burning.

Combustion of metal

4. Conclusion: Some metals, like magnesium and sodium, catch fire quickly, while others do not burn easily. The surface of the burning metal may change, forming an oxide layer. Different metals produce different flame colours when burnt.

Metals are reactive. They lose electrons easily and become positively charged ions. That is why metals are called electropositive elements.

• Most metals form basic oxides with oxygen; aluminium and zinc oxides are amphoteric.
• Reactivity with oxygen varies; sodium and potassium react quickly and are kept in kerosene.
• Metals may react with water to form hydroxides and release hydrogen gas, depending on temperature.
• Metals react with dilute acids to give salt and hydrogen gas; copper and nitric acid are exceptions.
• More reactive metals can displace less reactive ones; this forms the basis of the reactivity series.