Reactions of Metal

Metals react with oxygen to form metal oxides. The rate of reaction depends on the metal’s reactivity. Highly reactive metals, such as sodium and potassium, react at room temperature, while others, like magnesium, require heating.

1. Highly reactive metals (e.g., sodium, potassium) react at room temperature:

4Na(s) + O₂(g) → 2Na₂O(s)

Sodium is stored in kerosene to prevent accidental combustion.

2. Moderately reactive metals (e.g., magnesium) react upon heating:

Na₂O(s) + H₂O(l) → 2NaOH(aq)

Magnesium oxide reacts with water to form magnesium hydroxide:

2Mg(s) + O₂(g) → 2MgO(s)

MgO + H₂O → Mg(OH)₂​

The reaction of metals with water varies depending on their reactivity. Some metals react violently, some react slowly, and others do not react at all.

1. Highly reactive metals (e.g., sodium, potassium) react explosively, producing hydrogen gas and heat:

2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g) + heat

2K(s) + 2H₂O(l) → 2KOH(aq) + H₂(g) + heat

2. Moderately reactive metals (e.g., calcium) react slowly, producing hydrogen gas in the form of bubbles:

2Ca(s) + 2H₂O(l) → 2Ca(OH)₂(aq) + H₂(g)

3. Less reactive metals (e.g., aluminium, zinc, iron) do not react with cold or hot water but react with steam:

2Al(s) + 3H₂O(g) → Al₂O₃(s) + 3H₂(g)

3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g)

Zn(s) + H₂O(g) → ZnO(s) + H₂(g)​​

1. Aim: To observe how different metals react with water.

2. Requirements: Beakers, water, and metal samples (excluding sodium for safety).

3. Procedure

• Drop a small piece of each metal into separate beakers filled with water.
• Observe whether the metal reacts, floats, or sinks.
• Record reaction rate, gas formation, and heat generation.

Reaction of a metal with water

4. Observations

• Sodium and potassium react violently.
• Calcium reacts slowly, forming bubbles.
• Zinc, iron, and aluminium react only with steam.

5. Conclusion: Metals have different reactivity levels with water. Highly reactive metals produce hydrogen gas and form alkalis, while less reactive metals react only with steam.

Most metals react with dilute acids to form salt and hydrogen gas. The reactivity of metals varies, with some reacting vigorously and others reacting slowly.

Reactivity order: Mg > Al > Zn > Fe

1. Highly reactive metals (e.g., magnesium) react vigorously:

Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g)

2. Moderately reactive metals (e.g., aluminium, zinc, iron) react at different speeds:

2Al(s)+6HCl(aq)→2AlCl₃(aq)+3H₂(g)

Fe(s) + 2HCl(aq) → FeCl₂(aq) + H₂(g)

Zn(s) + HCl(aq) → ZnCl₂(aq) + H₂(g)

The hydrogen gas released can be tested by bringing a burning splint near it, which produces a ‘pop’ sound.

1. Aim: To study the reaction of metals with dilute acids.

2. Requirements: Test tubes, dilute hydrochloric acid, and samples of magnesium, aluminium, zinc, and iron.

3. Procedure

• Add a small piece of each metal to separate test tubes containing dilute hydrochloric acid.
• Observe gas bubble formation.
• Test the gas with a burning splint.

Reaction of metals with dilute acid

4. Observations

• Magnesium reacts vigorously, producing hydrogen gas.
• Aluminium, zinc, and iron react at different speeds.
• The gas burns with a ‘pop’ sound.

5. Conclusion: Metals react with acids to form salts and hydrogen gas. The reaction speed varies with different metals.

Metals react with nitric acid to form metal nitrates and release nitrogen oxides (NO, NO₂, or N₂O) instead of hydrogen gas. This is because nitric acid is a strong oxidising agent, which oxidises the hydrogen gas into water. However, magnesium and manganese react with very dilute nitric acid to produce hydrogen gas.

1. Reaction with concentrated nitric acid

Copper reacts with concentrated nitric acid to form copper nitrate and nitrogen dioxide gas:

Cu(s) + 4HNO₃(aq) → Cu(NO₃)₂(aq) + 2NO₂(g) + 2H₂O(l)

NO gas is colourless but turns brown in air due to oxidation into NO₂.

2. Reaction with dilute nitric acid

Copper reacts with dilute nitric acid to form copper nitrate and nitric oxide (NO) gas:

3Cu(s) + 8HNO₃(aq) → 3Cu(NO₃)₂(aq) + 2NO(g) + 4H₂O(l)

Aqua Regia:  It is a highly corrosive mixture of concentrated hydrochloric acid (HCl) and nitric acid (HNO₃) in a 3:1 ratio. It dissolves noble metals like gold and platinum, which do not react with nitric acid alone.

Au(s) + HNO₃ + HCl → H[AuCl₄] + NO

A more reactive metal can displace a less reactive metal from its salt solution. This is known as a displacement reaction.

Iron displaces copper from copper sulphate solution:

Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)

The iron nail gets coated with copper, showing that iron is more reactive than copper.

1. Aim: To observe displacement reactions between metals and metal salts.
2. Requirements: Test tubes, iron nails, copper wire, solutions of copper sulphate and ferrous sulphate.
3. Procedure

• Place an iron nail in copper sulphate solution.
• Place a copper wire in ferrous sulphate solution.
• Observe changes over time.

Reaction of metal with solution of salts of other metals

4. Observations

• The iron nail develops a copper coating, indicating a reaction.
• No reaction is seen in the copper wire with ferrous sulphate solution.

5. Conclusion: A more reactive metal can replace a less reactive metal from its salt solution, confirming the reactivity series.

Metals react with nonmetals to form ionic compounds. This reaction occurs because metals lose electrons to achieve a stable electron configuration, while nonmetals gain electrons to complete their outermost shell. Noble gases (like helium, neon, and argon) do not react because they already have a complete octet, making them chemically inert.

When a metal reacts with a nonmetal, an electron transfer takes place, forming an ionic bond. The metal donates electrons, becoming a cation (+), while the nonmetal accepts electrons, becoming an anion (-).

Example,

1. Formation of Sodium Chloride (NaCl)

2Na(s) + Cl₂(g) → 2NaCl(s)

• Sodium (Na) loses one electron to form Na⁺.
• Chlorine (Cl₂) gains electrons to form Cl⁻.
• The oppositely charged ions attract, forming an ionic bond in NaCl.

Similarly, magnesium and potassium also form ionic compounds with chlorine:

2. Formation of Magnesium Chloride (MgCl₂)

Mg(s) + Cl₂(g) → MgCl₂(s)

• Magnesium loses two electrons to form Mg²⁺.
• Each chlorine atom gains one electron, forming Cl⁻ ions.

3. Formation of Potassium Chloride (KCl)

K(s) + Cl₂(g) → 2KCl(s)

• Potassium loses one electron to form K⁺.
• Chlorine gains an electron to form Cl⁻.