Gold and platinum are placed at the bottom of the reactivity series (noble metals) and are extremely unreactive; hence they do not react with atmospheric oxygen, moisture, or acids.

Aluminium ($\text{Al}$) is the reducing agent because it removes oxygen from $\text{MnO}_2$ and gets oxidised to $\text{Al}_2\text{O}_3$.

During respiration, glucose undergoes oxidation inside cells and releases energy in the form of ATP: $\text{C}_6\text{H}_{12}\text{O}_6 + 6\text{O}_2 \rightarrow 6\text{CO}_2 + 6\text{H}_2\text{O} + \text{Energy}$.

A vigorous exothermic reaction occurs producing slaked lime (calcium hydroxide) with the evolution of a large amount of heat and a hissing sound. Product formula: $\text{Ca(OH)}_2$.

Law of Conservation of Mass: Matter can neither be created nor destroyed in a chemical reaction; hence the total mass of reactants equals the total mass of products.

Decomposition of silver bromide in the presence of sunlight: $2\text{AgBr(s)} \xrightarrow{\text{Sunlight}} 2\text{Ag(s)} + \text{Br}_2\text{(g)}$.

1. Element 'X' is Copper ($\text{Cu}$).

2. The black coloured compound is Copper(II) oxide ($\text{CuO}$).

3. Balanced Chemical Equation: $2\text{Cu(s)} + \text{O}_2\text{(g)} \xrightarrow{\Delta} 2\text{CuO(s)}$.

1. Nitrogen is an unreactive inert gas that acts as an antioxidant and prevents oxygen from coming in contact with fats and oils.

2. This prevents **Rancidity** (विकृतगंधिता), which is the aerial oxidation of unsaturated fats causing unpleasant taste, sour odor, and harmful free radicals.

(a) **Electrolytic Decomposition (विद्युत अपघटन):**

$$2\text{H}_2\text{O(l)} \xrightarrow{\text{Electric Current}} 2\text{H}_2\text{(g)} + \text{O}_2\text{(g)}$$

(b) **Thermal Decomposition (ऊष्मीय वियोजन):**

$$\text{CaCO}_3\text{(s)} \xrightarrow{\Delta} \text{CaO(s)} + \text{CO}_2\text{(g)}$$

1. **Definition:** A reaction in which one reactant undergoes oxidation (loss of electrons / gain of oxygen) while the other undergoes reduction (gain of electrons / loss of oxygen) simultaneously.

2. In $\text{ZnO} + \text{C} \rightarrow \text{Zn} + \text{CO}$:

- **Substance Oxidised:** Carbon ($\text{C}$) $\rightarrow$ Gains oxygen to form $\text{CO}$.

- **Substance Reduced:** Zinc oxide ($\text{ZnO}$) $\rightarrow$ Loses oxygen to form $\text{Zn}$.

1. **Displacement Reaction:** A more reactive element displaces a less reactive element from its aqueous salt solution. Example: $\text{Fe(s)} + \text{CuSO}_4\text{(aq)} \rightarrow \text{FeSO}_4\text{(aq)} + \text{Cu(s)}$.

2. **Double Displacement Reaction:** Mutual exchange of ions occurs between two aqueous salt solutions to form two new compounds (often one insoluble precipitate). Example: $\text{Na}_2\text{SO}_4\text{(aq)} + \text{BaCl}_2\text{(aq)} \rightarrow \text{BaSO}_4\downarrow\text{(s)} + 2\text{NaCl(aq)}$.

1. Chemically, rust is hydrated iron(III) oxide: $\text{Fe}_2\text{O}_3\cdot x\text{H}_2\text{O}$.

2. Two essential conditions:

- Presence of atmospheric oxygen ($\text{O}_2$).

- Presence of moisture or water vapor ($\text{H}_2\text{O}$).

(a) **Observations:**

1. The pale green crystals turn into a reddish-brown solid (ferric oxide).

2. Water droplets appear on the cooler upper walls of the test tube.

(b) **Gases released:** Sulphur dioxide ($\text{SO}_2$) and sulphur trioxide ($\text{SO}_3$). They have the suffocating odor of burning sulphur.

(c) **Chemical Equation:**

$$2\text{FeSO}_4\text{(s)} \xrightarrow{\Delta} \text{Fe}_2\text{O}_3\text{(s)} + \text{SO}_2\text{(g)} + \text{SO}_3\text{(g)}$$

(a) **Electrodes:**

- At Cathode ($-ve$): Hydrogen gas ($\text{H}_2$)

- At Anode ($+ve$): Oxygen gas ($\text{O}_2$)

(b) **2:1 Volume Ratio:** Water contains two parts of hydrogen and one part of oxygen by volume ($2\text{H}_2\text{O} \rightarrow 2\text{H}_2 + \text{O}_2$). Hence, the volume of hydrogen evolved at the cathode is exactly double that of oxygen at the anode.

(c) **Confirmatory Tests:**

- $\text{H}_2$: Brings a burning splinter near mouth $\rightarrow$ burns with a characteristic 'pop' sound.

- $\text{O}_2$: Brings a glowing wooden splinter $\rightarrow$ rekindles/reignites brilliantly.

(a) **Residue:** Yellow residue of Lead(II) oxide ($\text{PbO}$).

(b) **Brown Fumes:** Nitrogen dioxide gas ($\text{NO}_2$).

(c) **Equation & Type:**

$$2\text{Pb(NO}_3)_2\text{(s)} \xrightarrow{\Delta} 2\text{PbO(s)} + 4\text{NO}_2\text{(g)} + \text{O}_2\text{(g)}$$

Type: **Thermal Decomposition Reaction** (ऊष्मीय वियोजन अभिक्रिया).

(a) **Photochemical decomposition:** Silver chloride decomposes in sunlight to form grey silver metal:

$$2\text{AgCl(s)} \xrightarrow{\text{Sunlight}} 2\text{Ag(s)} + \text{Cl}_2\text{(g)}$$

(b) **Lime water reactions:**

- Turns milky due to insoluble calcium carbonate precipitate:

$$\text{Ca(OH)}_2\text{(aq)} + \text{CO}_2\text{(g)} \rightarrow \text{CaCO}_3\downarrow\text{(s)} + \text{H}_2\text{O(l)}$$

- Turns clear on excess $\text{CO}_2$ due to soluble calcium hydrogen carbonate:

$$\text{CaCO}_3\text{(s)} + \text{H}_2\text{O(l)} + \text{CO}_2\text{(g)} \rightarrow \text{Ca(HCO}_3)_2\text{(aq)}$$

(a) $$\text{ZnCO}_3\text{(s)} \xrightarrow{\Delta} \text{ZnO(s)} + \text{CO}_2\text{(g)}$$

(b) $$3\text{BaCl}_2\text{(aq)} + \text{Al}_2(\text{SO}_4)_3\text{(aq)} \rightarrow 2\text{AlCl}_3\text{(aq)} + 3\text{BaSO}_4\downarrow\text{(s)}$$

1. **Corrosion:** The slow destruction and degradation of metals caused by chemical or electrochemical reaction with atmospheric gases (oxygen, moisture, carbon dioxide, etc.).

2. **Prevention Methods:** (i) Galvanisation (coating with zinc), (ii) Painting or Greasing / Electroplating.

3. **Green coating on copper:** Copper reacts slowly with moist atmospheric $\text{CO}_2$ and oxygen to form basic copper carbonate ($\text{CuCO}_3\cdot\text{Cu(OH)}_2$):

$$2\text{Cu(s)} + \text{H}_2\text{O(g)} + \text{CO}_2\text{(g)} + \text{O}_2\text{(g)} \rightarrow \text{CuCO}_3\cdot\text{Cu(OH)}_2\text{(s)}$$

(a) **Types of Chemical Reactions:**

1. **Combination Reaction:** Two or more reactants combine to form a single product.

$$\text{CaO(s)} + \text{H}_2\text{O(l)} \rightarrow \text{Ca(OH)}_2\text{(aq)} + \text{Heat}$$

2. **Decomposition Reaction:** A single reactant splits into two or more simpler products.

$$2\text{KClO}_3\text{(s)} \xrightarrow{\Delta} 2\text{KCl(s)} + 3\text{O}_2\text{(g)}$$

3. **Displacement Reaction:** A more reactive element replaces a less reactive element from its compound.

$$\text{Zn(s)} + \text{CuSO}_4\text{(aq)} \rightarrow \text{ZnSO}_4\text{(aq)} + \text{Cu(s)}$$

4. **Double Displacement Reaction:** Mutual ion exchange between two compounds in solution.

$$\text{AgNO}_3\text{(aq)} + \text{NaCl(aq)} \rightarrow \text{AgCl}\downarrow\text{(s)} + \text{NaNO}_3\text{(aq)}$$

(b) **Opposite Nature:** Decomposition is the exact opposite of combination because combination joins multiple substances into one product, whereas decomposition breaks a single substance apart into multiple products with the absorption of energy.

(a) **Balancing Equations:**

(i) Left: $\text{H}=3, \text{N}=1, \text{O}=5, \text{Ca}=1$. Right: $\text{H}=2, \text{N}=2, \text{O}=7, \text{Ca}=1$.

Multiplying $\text{HNO}_3$ by 2 and $\text{H}_2\text{O}$ by 2:

$$2\text{HNO}_3\text{(aq)} + \text{Ca(OH)}_2\text{(aq)} \rightarrow \text{Ca(NO}_3)_2\text{(aq)} + 2\text{H}_2\text{O(l)}$$

(ii) Left: $\text{Fe}=1, \text{H}=2, \text{O}=1$. Right: $\text{Fe}=3, \text{H}=2, \text{O}=4$.

Multiplying $\text{Fe}$ by 3, $\text{H}_2\text{O}$ by 4, and $\text{H}_2$ by 4:

$$3\text{Fe(s)} + 4\text{H}_2\text{O(g)} \rightarrow \text{Fe}_3\text{O}_4\text{(s)} + 4\text{H}_2\text{(g)}$$

(b) **Four critical pieces of information convey:**

1. Physical state of each participant: solid (s), liquid (l), gas (g), aqueous solution (aq).

2. Exact molar and molecular ratios of reactants consumed and products generated.

3. Reaction conditions above/below arrow: temperature, pressure, heat ($\Delta$), sunlight, or catalyst.

4. Energy changes: whether heat is released ($+\text{Heat}$ exothermic) or absorbed (endothermic).

(a) **Redox Reaction:**

- **Classical Concept:** Oxidation is the addition of oxygen or removal of hydrogen. Reduction is the addition of hydrogen or removal of oxygen.

- **Electronic Concept:** Oxidation is the loss of electrons (OIL: Oxidation Is Loss). Reduction is the gain of electrons (RIG: Reduction Is Gain).

- In a redox reaction, oxidation and reduction always occur simultaneously.

(b) **For $\text{MnO}_2 + 4\text{HCl} \rightarrow \text{MnCl}_2 + 2\text{H}_2\text{O} + \text{Cl}_2$:**

- (i) **Substance oxidised:** $\text{HCl}$ (Hydrogen is removed / electrons lost by chloride to form $\text{Cl}_2$).

- (ii) **Substance reduced:** $\text{MnO}_2$ (Oxygen is removed to form $\text{MnCl}_2$).

- (iii) **Oxidising agent:** $\text{MnO}_2$ (It provides oxygen and oxidises $\text{HCl}$).

- (iv) **Reducing agent:** $\text{HCl}$ (It reduces $\text{MnO}_2$ to $\text{MnCl}_2$).

(i) **Product:** Slaked Lime (Calcium hydroxide), formula: $\text{Ca(OH)}_2$.

(ii) **Type of Reaction:** Combination reaction and highly **Exothermic reaction** ($\Delta H < 0$).

(iii) **Chemical Equation:**

$$\text{CaO(s)} + \text{H}_2\text{O(l)} \rightarrow \text{Ca(OH)}_2\text{(aq)} + \text{Heat}$$

(iv) **Bubbling $\text{CO}_2$:** The solution turns milky due to the formation of insoluble white precipitate of calcium carbonate:

$$\text{Ca(OH)}_2\text{(aq)} + \text{CO}_2\text{(g)} \rightarrow \text{CaCO}_3\downarrow\text{(s)} + \text{H}_2\text{O(l)}$$

(i) **Color Changes:** The blue color of copper sulphate solution fades and changes to light green (due to $\text{FeSO}_4$). The iron nail gets coated with a reddish-brown layer.

(ii) **Metal Deposited:** Copper metal ($\text{Cu}$).

(iii) **Chemical Equation:**

$$\text{Fe(s)} + \text{CuSO}_4\text{(aq)} \rightarrow \text{FeSO}_4\text{(aq)} + \text{Cu(s)}$$

(iv) **Reverse condition:** No reaction will occur because copper is less reactive than iron in the reactivity series and cannot displace iron from ferrous sulphate solution.