Chemical Reactions

• (a) $2\text{H}_2\text{O}_2 \to 2\text{H}_2\text{O} + \text{O}_2 \uparrow$: Decomposition Reaction (A single compound breaks down into two simpler substances in the presence of catalyst $\text{MnO}_2$).
• (b) $\text{Zn} + 2\text{HCl} \to \text{ZnCl}_2 + \text{H}_2 \uparrow$: Single Displacement Reaction (Zinc, being more electropositive than hydrogen, displaces hydrogen gas from dilute acid).
• (c) $\text{CaO} + \text{SiO}_2 \to \text{CaSiO}_3$: Direct Combination (Synthesis) Reaction (Two compounds unite to form a single product, calcium silicate).
• (d) $\text{AgNO}_3 + \text{KI} \to \text{AgI} \downarrow + \text{KNO}_3$: Double Decomposition (Precipitation) Reaction (Mutual exchange of ions produces an insoluble yellow precipitate of silver iodide).

• Observations:
1. The intense blue color of the copper sulphate solution gradually fades and turns light emerald green.
2. A reddish-brown coating of metallic copper deposits on the surface of the iron nail.
• Chemical Equation:

$$\mathbf{\text{Fe (s)} + \text{CuSO}_4\text{ (aq) [blue]} \to \text{FeSO}_4\text{ (aq) [green]} + \text{Cu (s) [reddish-brown]} \downarrow}$$

• Type of Reaction: Single Displacement Reaction.
• Explanation: According to the reactivity series, Iron (Fe) is more reactive than Copper (Cu). Iron displaces copper ions from the solution, taking their place to form green ferrous sulphate solution while free copper atoms precipitate out.

• No reaction will occur.
• Reason: In the metal reactivity series, Zinc is more reactive than Copper ($\text{Zn} > \text{Cu}$).
• A less reactive metal (Copper) cannot displace a more reactive metal (Zinc) from its aqueous salt solution.
• Therefore, the solution remains completely unchanged: $\text{Cu} + \text{ZnSO}_4 \to \mathbf{\text{No Reaction}}$.

• (a) Substance Oxidized: Carbon (C) (It gains oxygen to become carbon monoxide: $\text{C} \to \text{CO}$).
• (b) Substance Reduced: Zinc Oxide (ZnO) (It loses oxygen to become metallic zinc: $\text{ZnO} \to \text{Zn}$).
• (c) Oxidizing Agent: Zinc Oxide (ZnO) (It supplies the oxygen for oxidation and is itself reduced).
• (d) Reducing Agent: Carbon (C) (It removes oxygen from zinc oxide and is itself oxidized).

• Photochemical Reaction: A chemical reaction that is initiated or accelerated by the absorption of light energy (photons, $h\nu$).
• Examples:
1. Photosynthesis in Green Plants:

$$6\text{CO}_2 + 6\text{H}_2\text{O} \xrightarrow{\text{Chlorophyll, Sunlight}} \text{C}_6\text{H}_{12}\text{O}_6 + 6\text{O}_2 \uparrow$$

2. Photolytic Decomposition of Silver Chloride:

$$2\text{AgCl (s) [white]} \xrightarrow{\text{Sunlight}} 2\text{Ag (s) [grey]} + \text{Cl}_2\text{ (g)} \uparrow$$

(This is why silver salt solutions are always stored in dark brown amber bottles).

• Precipitation Reaction: A double decomposition reaction in which two clear aqueous salt solutions react to form an insoluble solid residue (precipitate $\downarrow$) that settles out of the solution.
Equation: $\mathbf{\text{BaCl}_2\text{ (aq)} + \text{Na}_2\text{SO}_4\text{ (aq)} \to \text{BaSO}_4 \downarrow \text{ (white ppt)} + 2\text{NaCl}\text{ (aq)}}$
• Neutralisation Reaction: A double decomposition reaction in which an acid reacts with a base to mutually destroy each other's chemical properties, producing a soluble salt and water.
Equation: $\mathbf{\text{HNO}_3\text{ (aq)} + \text{KOH}\text{ (aq)} \to \text{KNO}_3\text{ (aq)} + \text{H}_2\text{O}\text{ (l)}}$.

• Catalyst: A chemical substance that alters the rate of a chemical reaction without itself undergoing any permanent chemical change at the end of the reaction.
• Positive Catalyst: A catalyst that increases / accelerates the speed of a chemical reaction.
Example: Manganese dioxide ($\text{MnO}_2$) in the thermal decomposition of potassium chlorate: $2\text{KClO}_3 \xrightarrow{\text{MnO}_2} 2\text{KCl} + 3\text{O}_2 \uparrow$.
• Negative Catalyst (Inhibitor): A catalyst that decreases / retards the speed of a chemical reaction.
Example: Phosphoric acid (or acetanilide) added to hydrogen peroxide ($\text{H}_2\text{O}_2$) to retard its spontaneous decomposition.

• The electronic theory defines redox reactions in terms of electron transfer:
• OIL: Oxidation Is Loss of electrons. An atom or ion loses one or more valence electrons, resulting in an increase in positive oxidation state:

$$\text{Na} \to \text{Na}^+ + e^- \quad \land \quad \text{Fe}^{2+} \to \text{Fe}^{3+} + e^-$$

• RIG: Reduction Is Gain of electrons. An atom or ion gains one or more electrons, resulting in a decrease in positive oxidation state:

$$\text{Cl} + e^- \to \text{Cl}^- \quad \land \quad \text{Cu}^{2+} + 2e^- \to \text{Cu}$$

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