Types of Reactions: Decomposition Reactions
🔄 Quick Recap
In the previous section, we learned about combination reactions where two or more substances combine to form a single product. Now, we'll explore decomposition reactions, which are essentially the opposite!
📚 What are Decomposition Reactions?
A decomposition reaction occurs when a single compound breaks down into two or more simpler substances (elements or compounds).
The general form is:
C → A + B
It's like breaking down a complex dish into its individual ingredients!
🔍 Types of Decomposition Reactions
Decomposition reactions require energy to break the bonds. This energy can come in different forms:
1. Thermal Decomposition
Decomposition caused by heat.
Example: Calcium carbonate (limestone) decomposes when heated:
CaCO₃(s) Heat→ CaO(s) + CO₂(g)
This reaction is used to produce quicklime (CaO) for cement manufacturing.
2. Photochemical Decomposition
Decomposition caused by light.
Example: Silver chloride decomposes in sunlight:
2AgCl(s) Sunlight→ 2Ag(s) + Cl₂(g)
This reaction is the basis for black and white photography!
3. Electrolytic Decomposition
Decomposition caused by electricity.
Example: Water decomposes into hydrogen and oxygen:
2H₂O(l) Electricity→ 2H₂(g) + O₂(g)
This process, called electrolysis, is used to produce hydrogen fuel.
🧪 Activity Time! Observing Decomposition Reactions
Activity 1: Thermal Decomposition of Ferrous Sulphate
Materials Needed:
- Ferrous sulphate crystals (FeSO₄·7H₂O)
- Test tube
- Burner or spirit lamp
- Test tube holder
Procedure:
- Take about 2g of ferrous sulphate crystals in a test tube
- Note the color (green)
- Heat the test tube over a flame
- Observe the changes
Observations:
- The green color of the crystals changes
- A smell of burning sulphur is noticed
- Water droplets form on the cooler parts of the test tube
What Happened?
Ferrous sulphate crystals decompose when heated:
2FeSO₄·7H₂O(s) Heat→ Fe₂O₃(s) + SO₂(g) + SO₃(g) + 14H₂O(g)
The brown solid formed is ferric oxide (Fe₂O₃), and the gases produced include sulphur dioxide (SO₂) and sulphur trioxide (SO₃).
Activity 2: Electrolysis of Water
Materials Needed:
- Water with a few drops of sulphuric acid (as an electrolyte)
- Two test tubes
- Two carbon electrodes
- Battery (6V)
- Wires
Procedure:
- Set up the apparatus as shown in your textbook
- Connect the electrodes to the battery
- Observe the bubbles forming at both electrodes
- Collect the gases in the test tubes
- Test the gases with a burning splint
Observations:
- Bubbles form at both electrodes
- The gas collected at the negative electrode (cathode) pops with a flame - it's hydrogen
- The gas collected at the positive electrode (anode) rekindles a glowing splint - it's oxygen
- The volume of hydrogen collected is twice the volume of oxygen
What Happened?
Water decomposes into hydrogen and oxygen:
2H₂O(l) Electricity→ 2H₂(g) + O₂(g)
The water molecules split, with hydrogen collecting at the negative electrode and oxygen at the positive electrode. The ratio is 2:1 because water's formula is H₂O.
🧮 Mathematical Corner: Balancing Decomposition Reactions
Let's balance a decomposition reaction step-by-step:
Example: Lead nitrate decomposes when heated
Pb(NO₃)₂ → PbO + NO₂ + O₂
Step 1: Count the atoms on each side
- Left side: 1 Pb, 2 N, 6 O
- Right side: 1 Pb, 1 N, 3 O + 2 O = 1 Pb, 1 N, 5 O
Step 2: Balance nitrogen by adding a coefficient to NO₂
Pb(NO₃)₂ → PbO + 2NO₂ + O₂
Step 3: Recount
- Left side: 1 Pb, 2 N, 6 O
- Right side: 1 Pb, 2 N, 1 O + 4 O + 2 O = 1 Pb, 2 N, 7 O
Step 4: Balance oxygen by adjusting O₂ coefficient
Pb(NO₃)₂ → PbO + 2NO₂ + ½O₂
Step 5: Multiply all coefficients by 2 to eliminate the fraction
2Pb(NO₃)₂ → 2PbO + 4NO₂ + O₂
Final balanced equation:
2Pb(NO₃)₂ → 2PbO + 4NO₂ + O₂
🌍 Real-Life Applications of Decomposition Reactions
1. Digestion
When we eat food, complex carbohydrates, proteins, and fats decompose into simpler substances through enzymatic reactions.
2. Photography
Traditional black and white photography relies on the decomposition of silver bromide by light:
2AgBr Sunlight→ 2Ag + Br₂
3. Bleaching Agents
Hydrogen peroxide decomposes to produce oxygen, which acts as a bleaching agent:
2H₂O₂ → 2H₂O + O₂
4. Antacids
Sodium bicarbonate (baking soda) decomposes in acidic conditions to neutralize stomach acid:
NaHCO₃ + HCl → NaCl + H₂O + CO₂
5. Fireworks
Many colorful effects in fireworks come from metal salts that decompose when heated.
⚖️ Quick Comparison: Combination vs. Decomposition Reactions
| Aspect | Combination Reaction | Decomposition Reaction |
|---|---|---|
| Direction | Multiple substances → One product | One compound → Multiple products |
| Energy change | Usually exothermic (releases energy) | Usually endothermic (absorbs energy) |
| Examples | Burning, rusting, formation of water | Electrolysis, photodecomposition |
| Natural occurrence | More common in nature | Less common without energy input |
⚠️ Common Misconceptions
-
Misconception: Decomposition reactions always form elements. Truth: They can produce elements OR simpler compounds.
-
Misconception: Decomposition happens spontaneously. Truth: Most decomposition reactions require energy input (heat, light, electricity).
🧠 Memory Trick
Remember "DEC-omposition = DEC-onstruction" - breaking down a compound into simpler parts.
💡 Key Points to Remember
- Decomposition reactions break a single compound into multiple simpler substances
- They require energy input (heat, light, or electricity)
- The general form is C → A + B
- Three main types: thermal, photochemical, and electrolytic decomposition
- These reactions are important in manufacturing, photography, and many biochemical processes
🤔 Think About It!
- Why do most decomposition reactions require energy input, while combination reactions usually release energy?
- Can you think of any decomposition reactions that might occur in your kitchen?
- How might the decomposition of water through electrolysis contribute to future energy solutions?
🔜 What Next?
In the next section, we'll explore displacement reactions, where one element "displaces" or replaces another element in a compound!