Chemical Reactions and Equations

Fundamentals of Chemical Reactions

Definition of Chemical Change: A chemical reaction takes place whenever the original identity and chemical composition of a substance undergo a transformation, resulting in the creation of new substances with distinct physical and chemical properties.
Daily Life Examples of Chemical Changes:
- Leaving milk at room temperature during warm summer months (causes curdling/fermentation).
- Exposing an iron tawa, pan, or nail to a humid, moist atmosphere (leads to corrosion/rusting).
- Fermentation of grapes to form alcohol.
- Cooking of food raw ingredients into digestible dishes.
- Digestion of food within the gastrointestinal tract of living organisms.
- Respiration in biological cells where nutrients are broken down to release energy.
Visual and Physical Indicators of a Chemical Reaction: The occurrence of a chemical reaction can be confirmed by observing one or more of the following key indicators:
- Change in state: Transformation between solid, liquid, or gas states.
- Change in colour: Alteration in the hue or shade of the reacting species.
- Evolution of a gas: Production and release of gas bubbles during the process.
- Change in temperature: Release of heat (exothermic) or absorption of heat (endothermic).

Practical Laboratory Experiments and Procedures

Combustion of Magnesium Ribbon:
- Procedure: Take a strip of magnesium ribbon measuring approximately $3-4\,\text{cm}$ in length. Clean its surface thoroughly by rubbing it with sandpaper to remove any protective oxide coating formed due to atmospheric oxygen. Using a pair of tongs, hold the ribbon over a spirit lamp or Bunsen burner flame. Collect the resulting white ash in a watch-glass placed below.
- Safety Precautions: This experiment requires teacher assistance. Eyeglasses or safety goggles should be worn, and the burning ribbon must be held as far away from the eyes as possible.
- Observations: The magnesium ribbon burns with a bright, dazzling white flame and converts into a fine white powder.
- Chemical Identity: The white powder formed is magnesium oxide ( $MgO$ ), created via the chemical reaction between metallic magnesium and atmospheric oxygen.

Burning of a magnesium ribbon in air and collection of magnesium oxide in a watch-glass

Reaction of Zinc Granules with Dilute Acids:
- Procedure: Place a few zinc granules inside a glass conical flask or test tube. Carefully add dilute hydrochloric acid ( $HCl$ ) or dilute sulphuric acid ( $H_2SO_4$ ).
- Observations: Gas bubbles evolve continuously around the surface of the zinc granules. Touching the bottom of the flask reveals a noticeable rise in temperature, signifying an exothermic chemical process.

Formation of hydrogen gas by the action of dilute sulphuric acid on zinc

Reaction between Lead Nitrate and Potassium Iodide:
- Procedure: Take an aqueous solution of lead nitrate in a clean test tube and add an aqueous solution of potassium iodide.
- Observations: A distinct yellow precipitate of lead(II) iodide ( $PbI_2$ ) forms instantly in the liquid mixture.

Chemical Equations and Representations

Word Equations: A word equation describes a chemical transformation using the names of the participating substances:

$\text{Magnesium} + \text{Oxygen} \rightarrow \text{Magnesium oxide}$

*   **Reactants:** The starting substances undergoing chemical transformation (written on the left-hand side, LHS, separated by a plus sign  $+$ ).
*   **Products:** The newly formed chemical substances resulting from the reaction (written on the right-hand side, RHS, separated by a plus sign  $+$ ).
*   **Arrowhead:** Connects reactants to products, pointing towards the product side to indicate the direction of the reaction.

Skeletal Chemical Equations: A skeletal chemical equation replaces substance names with chemical formulae without requiring balanced atom counts on both sides:

$Mg + O_2 \rightarrow MgO$

*   An equation is termed unbalanced or skeletal when the chemical masses on the LHS and RHS are unequal.

Law of Conservation of Mass: Mass can neither be created nor destroyed in a chemical reaction. The total mass of elements in the product side must exactly equal the total mass of elements in the reactant side. Consequently, the total number of atoms of each individual element remains constant before and after a chemical reaction occurs.

Step-by-Step Procedure for Balancing Chemical Equations

Balancing chemical equations uses the systematic Hit-and-Trial Method, adjusting coefficients to obtain the smallest whole numbers.

Unbalanced Chemical Reaction Example:

$Fe + H_2O \rightarrow Fe_3O_4 + H_2$

Step I: Formula Isolation: Draw protective boxes around every formula in the unbalanced reaction. Do not modify any subscripts or numbers inside the boxes.

$[Fe] + [H_2O] \rightarrow [Fe_3O_4] + [H_2]$

Step II: Atom Inventory: List and compare the count of atoms for each distinct element present on the LHS and RHS:
- Iron ( $Fe$ ): LHS = $1$ , RHS = $3$
- Hydrogen ( $H$ ): LHS = $2$ , RHS = $2$
- Oxygen ( $O$ ): LHS = $1$ , RHS = $4$
Step III: Balancing the Oxygen Atoms: Select the compound containing the maximum total atoms ( $Fe_3O_4$ ) and focus on its element with the highest count (oxygen). Oxygen has $4$ atoms on the RHS and $1$ atom in $H_2O$ on the LHS. Multiply $H_2O$ by a coefficient of $4$ on the LHS (never alter subscripts like $H_2O_4$ ):

$Fe + 4H_2O \rightarrow Fe_3O_4 + H_2$

Step IV: Balancing Hydrogen Atoms: Check the hydrogen count. LHS now contains $8$ atoms ( $4 \times 2$ ), while RHS contains $2$ atoms in $H_2$ . Place a coefficient of $4$ in front of $H_2$ on the RHS:

$Fe + 4H_2O \rightarrow Fe_3O_4 + 4H_2$

Step V: Balancing Iron Atoms: Examine the iron count. RHS contains $3$ atoms in $Fe_3O_4$ , while LHS contains $1$ atom. Place a coefficient of $3$ in front of $Fe$ on the LHS:

$3Fe + 4H_2O \rightarrow Fe_3O_4 + 4H_2$

Step VI: Equation Verification: Count all atoms on both sides to verify equality:
- Iron ( $Fe$ ): LHS = $3$ , RHS = $3$
- Hydrogen ( $H$ ): LHS = $8$ , RHS = $8$
- Oxygen ( $O$ ): LHS = $4$ , RHS = $4$
- Since atom counts are identical, the equation is balanced.
Step VII: Inclusion of Physical States and Reaction Conditions:
- Notations represent physical states: solid $(s)$ , liquid $(l)$ , gaseous $(g)$ , and aqueous solution $(aq)$ .
- The complete balanced equation is:

$3Fe(s) + 4H_2O(g) \rightarrow Fe_3O_4(s) + 4H_2(g)$

*   The symbol  $(g)$  attached to  $H_2O$  specifies that water reacts in the gaseous phase as steam.
*   Specific conditions such as reaction temperature, applied pressure, or catalysts are written above or below the arrow:

$CO(g) + 2H_2(g) \xrightarrow{340\,\text{atm}} CH_3OH(l)$

$6CO_2(aq) + 12H_2O(l) \xrightarrow[\text{Chlorophyll}]{\text{Sunlight}} C_6H_{12}O_6(aq) + 6O_2(aq) + 6H_2O(l)$

Major Types of Chemical Reactions

Chemical reactions break and form chemical bonds between atoms to yield new chemical substances.

Combination Reactions

Definition: A chemical reaction in which two or more starting substances (elements or compounds) combine together to produce a single final product.
Reaction of Calcium Oxide with Water:
- Calcium oxide (commonly called quick lime, $CaO$ ) reacts vigorously with liquid water to form calcium hydroxide (slaked lime, $Ca(OH)_2$ ), liberating a substantial amount of heat energy:

$CaO(s) + H_2O(l) \rightarrow Ca(OH)_2(aq) + \text{Heat}$

Formation of slaked lime by the reaction of calcium oxide with water

Application in Whitewashing:
- A solution of slaked lime ( $Ca(OH)_2$ ) is applied to walls during whitewashing. It reacts slowly with atmospheric carbon dioxide ( $CO_2$ ) over a period of $2-3\,\text{days}$ to form a thin, durable, shiny coating of calcium carbonate ( $CaCO_3$ ) along with water:

$Ca(OH)_2(aq) + CO_2(g) \rightarrow CaCO_3(s) + H_2O(l)$

*   The chemical formula for marble is also  $CaCO_3$ .

Additional Examples of Combination Reactions:
- Combustion of Coal:

$C(s) + O_2(g) \rightarrow CO_2(g)$

*   **Synthesis of Water:**

$2H_2(g) + O_2(g) \rightarrow 2H_2O(l)$

Exothermic Chemical Reactions:
- Reactions that release thermal energy alongside product formation.
- Combustion of Natural Gas:

$CH_4(g) + 2O_2(g) \rightarrow CO_2(g) + 2H_2O(g)$

*   **Cellular Respiration:** Carbohydrates consumed in food (e.g., from potatoes, rice, bread) are broken down into glucose ( $C_6H_{12}O_6$ ). Glucose reacts with dissolved oxygen inside body cells to provide life-sustaining energy:

$C_6H_{12}O_6(aq) + 6O_2(aq) \rightarrow 6CO_2(aq) + 6H_2O(l) + \text{Energy}$

*   **Composting:** The decomposition of organic vegetable matter into nutrient-rich compost is an exothermic process.

Decomposition Reactions

Definition: A chemical reaction in which a single reactant breaks down into two or more simpler component products.
Thermal Decomposition: Decomposition processes driven by the application of thermal energy (heating).
- Decomposition of Ferrous Sulphate:
  - Procedure: Place about $2\,\text{g}$ of green ferrous sulphate crystals ( $FeSO_4 \cdot 7H_2O$ ) in a dry boiling tube and heat over a burner flame.
  - Observation: The green crystals lose water of crystallization and change colour. Upon further intense heating, a characteristic smell of burning sulphur is emitted as the substance decomposes into solid ferric oxide ( $Fe_2O_3$ ), sulphur dioxide gas ( $SO_2$ ), and sulphur trioxide gas ( $SO_3$ ):

$2FeSO_4(s) \xrightarrow{\text{Heat}} Fe_2O_3(s) + SO_2(g) + SO_3(g)$

Heating of boiling tube containing ferrous sulphate crystals

*   **Decomposition of Calcium Carbonate:**
    *   Heating limestone ( $CaCO_3$ ) produces quick lime ( $CaO$ ) and carbon dioxide ( $CO_2$ ). Quick lime is widely used in industry, particularly in cement manufacturing:

$CaCO_3(s) \xrightarrow{\text{Heat}} CaO(s) + CO_2(g)$

*   **Decomposition of Lead Nitrate:**
    *   Procedure: Heat  $2\,\text{g}$  of white lead nitrate powder ( $Pb(NO_3)_2$ ) in a boiling tube using tongs.
    *   Observation: Dense brown fumes of nitrogen dioxide gas ( $NO_2$ ) evolve, leaving behind a yellow residue of lead oxide ( $PbO$ ):

$2Pb(NO_3)_2(s) \xrightarrow{\text{Heat}} 2PbO(s) + 4NO_2(g) + O_2(g)$

Heating of lead nitrate and emission of nitrogen dioxide

Electrolytic Decomposition (Electrolysis):
- Electrolysis of Water: Drill two holes at the base of a plastic mug and insert rubber stoppers fitted with graphite/carbon electrodes connected to a $6\,\text{V}$ DC battery. Fill the mug with water until electrodes are submerged, then add a few drops of dilute sulphuric acid to enhance conductivity. Invert two water-filled test tubes over the electrodes and pass an electric current.
- Observations: Gas bubbles evolve at both electrodes, displacing water in the test tubes.
- Gas Identification and Volume Ratio: Oxygen gas collects at the anode (positive electrode), while hydrogen gas collects at the cathode (negative electrode). The volume of hydrogen collected at the cathode is exactly double ( $2:1$ ratio) the volume of oxygen collected at the anode, matching the stoichiometric ratio of water ( $H_2O$ ).

Photolytic Decomposition (Photolysis):
- Decomposition of Silver Chloride: Place $2\,\text{g}$ of white silver chloride powder ( $AgCl$ ) in a china dish under direct sunlight.
- Observation: The white powder turns grey due to photolytic decomposition into silver metal and chlorine gas:

$2AgCl(s) \xrightarrow{\text{Sunlight}} 2Ag(s) + Cl_2(g)$

Silver chloride turns grey in sunlight to form silver metal

*   **Decomposition of Silver Bromide:**

$2AgBr(s) \xrightarrow{\text{Sunlight}} 2Ag(s) + Br_2(g)$

*   **Practical Application:** Photolytic reactions of silver salts are utilized in black-and-white photography.

Endothermic Chemical Reactions:
- Reactions that absorb energy from their surroundings in the form of heat, light, or electricity.
- Example: Mixing $2\,\text{g}$ of barium hydroxide ( $Ba(OH)_2$ ) with $1\,\text{g}$ of ammonium chloride ( $NH_4Cl$ ) in a test tube. Stirring causes the temperature of the vessel to drop noticeably, indicating an endothermic reaction.

Displacement Reactions

Definition: A chemical reaction in which a more reactive element displaces a less reactive element from its aqueous compound solution.
Displacement of Copper by Iron:
- Procedure: Clean three iron nails using sandpaper. Submerge two nails in a test tube containing $10\,\text{mL}$ of blue copper sulphate solution ( $CuSO_4$ ) for $20\,\text{minutes}$ , leaving one nail unreacted for reference.
- Observations: The immersed iron nails acquire a reddish-brown coating of displaced copper metal, while the deep blue color of the copper sulphate solution fades to a light green shade due to iron sulphate ( $FeSO_4$ ) formation:

$Fe(s) + CuSO_4(aq) \rightarrow FeSO_4(aq) + Cu(s)$

Iron nails dipped in copper sulphate solution

Iron nails and copper sulphate solutions compared before and after experiment

Additional Displacement Examples:
- Displacement by Zinc: Metallic zinc displaces copper from copper sulphate solution because zinc possesses higher chemical reactivity:

$Zn(s) + CuSO_4(aq) \rightarrow ZnSO_4(aq) + Cu(s)$

*   **Displacement by Lead:** Metallic lead displaces copper from copper chloride solution:

$Pb(s) + CuCl_2(aq) \rightarrow PbCl_2(aq) + Cu(s)$

Double Displacement Reactions

Definition: Chemical reactions featuring a mutual exchange of ions between two reacting compounds to yield two new compounds.
Precipitation Reactions: Any reaction that forms an insoluble solid product (called a precipitate) that separates from the liquid solution.
Reaction between Sodium Sulphate and Barium Chloride:
- Procedure: Combine $3\,\text{mL}$ of sodium sulphate solution ( $Na_2SO_4$ ) with $3\,\text{mL}$ of barium chloride solution ( $BaCl_2$ ) in a test tube.
- Observation: An insoluble white precipitate of barium sulphate ( $BaSO_4$ ) forms immediately:

$Na_2SO_4(aq) + BaCl_2(aq) \rightarrow BaSO_4(s) + 2NaCl(aq)$

*   **Ionic Mechanism:** The white precipitate forms when sulphate ions ( $SO_4^{2-}$ ) react with barium ions ( $Ba^{2+}$ ) to form solid  $BaSO_4$ , while soluble sodium chloride ( $NaCl$ ) remains dissolved in solution.

Formation of barium sulphate and sodium chloride

Reaction between Lead Nitrate and Potassium Iodide:

$Pb(NO_3)_2(aq) + 2KI(aq) \rightarrow PbI_2(s) + 2KNO_3(aq)$

*   This reaction forms a bright yellow precipitate of lead iodide ( $PbI_2$ ).

Oxidation and Reduction (Redox) Reactions

Definitions:
- Oxidation: Chemical gain of oxygen or chemical loss of hydrogen by a substance during a reaction.
- Reduction: Chemical loss of oxygen or chemical gain of hydrogen by a substance during a reaction.
- Redox Reaction: A chemical process where oxidation and reduction occur simultaneously—one reactant is oxidised while the other is reduced.
Oxidation of Copper Metal:
- Heating $1\,\text{g}$ of reddish-brown copper powder in an open china dish coats its surface with a black layer of copper(II) oxide ( $CuO$ ):

$2Cu(s) + O_2(g) \xrightarrow{\text{Heat}} 2CuO(s)$

Reduction of Copper(II) Oxide:
- Passing hydrogen gas ( $H_2$ ) over heated black copper oxide ( $CuO$ ) reverses the reaction, restoring the reddish-brown copper metal:

$CuO(s) + H_2(g) \xrightarrow{\text{Heat}} Cu(s) + H_2O(g)$

*   In this redox process,  $CuO$  loses oxygen (it is reduced to  $Cu$ ), while  $H_2$  gains oxygen (it is oxidised to  $H_2O$ ).

Additional Redox Reaction Examples:
- Reduction of Zinc Oxide by Carbon:

$ZnO + C \rightarrow Zn + CO$

    *   Carbon ( $C$ ) is oxidised to carbon monoxide ( $CO$ ), while zinc oxide ( $ZnO$ ) is reduced to elemental zinc ( $Zn$ ).
*   **Reaction of Manganese Dioxide with Hydrochloric Acid:**

$MnO_2 + 4HCl \rightarrow MnCl_2 + 2H_2O + Cl_2$

    *   Hydrochloric acid ( $HCl$ ) is oxidised to chlorine gas ( $Cl_2$ ), while manganese dioxide ( $MnO_2$ ) is reduced to manganese chloride ( $MnCl_2$ ).

Oxidation Effects in Everyday Life

Corrosion:
- Mechanism: Metallic surfaces deteriorate when exposed to atmospheric environmental agents such as oxygen, water moisture, acids, and salts over time.
- Rusting of Iron: Exposing iron objects to moist air forms a flaky, reddish-brown surface coating of hydrated iron oxide ( $Fe_2O_3 \cdot xH_2O$ ).
- Other Metal Examples:
  - Silver: Silver items react with atmospheric sulphur compounds to form a dark, black tarnished film of silver sulphide ( $Ag_2S$ ).
  - Copper: Copper objects react with atmospheric moisture and carbon dioxide over time to form a green coating of basic copper carbonate ( $CuCO_3 \cdot Cu(OH)_2$ ).
- Socioeconomic Impacts: Corrosion degrades structural integrity, severely damaging car bodies, bridges, iron railings, railway lines, and ships. Replacing corroded infrastructure requires substantial global financial expenditure annually.
Rancidity:
- Mechanism: Fats and oils present in food items undergo ambient oxidation when exposed to air, producing volatile breakdown compounds with unpleasant tastes and foul odours.
- Prevention Methods:
  - Incorporating antioxidants (chemical oxidation inhibitors) into lipid-containing foods.
  - Packaging perishable food items inside sealed, airtight containers to minimize exposure to atmospheric oxygen.
  - Flushing sealed snack bags (such as potato chip packets) with unreactive, inert gas like nitrogen ( $N_2$ ) to displace reactive oxygen gas.

In-Text Conceptual Questions and Solutions

Question 1: Why should a magnesium ribbon be cleaned before burning in air?
- Answer: Magnesium reacts with atmospheric oxygen at room temperature to form a thin, stable protective layer of magnesium oxide ( $MgO$ ) on its surface. This coating acts as a barrier that prevents further contact with oxygen. Cleaning the ribbon with sandpaper removes this oxide film, allowing the underlying metal to burn smoothly.
Question 2: Write balanced chemical equations for the following reactions:
1. Hydrogen + Chlorine $\rightarrow$ Hydrogen chloride
2. Barium chloride + Aluminium sulphate $\rightarrow$ Barium sulphate + Aluminium chloride
3. Sodium + Water $\rightarrow$ Sodium hydroxide + Hydrogen
- Answer:
1. $H_2(g) + Cl_2(g) \rightarrow 2HCl(g)$
2. $3BaCl_2(aq) + Al_2(SO_4)_3(aq) \rightarrow 3BaSO_4(s) + 2AlCl_3(aq)$
3. $2Na(s) + 2H_2O(l) \rightarrow 2NaOH(aq) + H_2(g)$
Question 3: Write a balanced chemical equation with state symbols for the following reactions:
1. Solutions of barium chloride and sodium sulphate in water react to give insoluble barium sulphate and the solution of sodium chloride.
2. Sodium hydroxide solution (in water) reacts with hydrochloric acid solution (in water) to produce sodium chloride solution and water.
- Answer:
1. $BaCl_2(aq) + Na_2SO_4(aq) \rightarrow BaSO_4(s) + 2NaCl(aq)$
2. $NaOH(aq) + HCl(aq) \rightarrow NaCl(aq) + H_2O(l)$
Question 4: A solution of a substance 'X' is used for whitewashing.
1. Name the substance 'X' and write its formula.
2. Write the reaction of the substance 'X' named in (1) above with water.
- Answer:
1. The substance 'X' is calcium oxide (quick lime). Its chemical formula is $CaO$ .
2. Reaction with water:

$CaO(s) + H_2O(l) \rightarrow Ca(OH)_2(aq)$

Question 5: Why is the amount of gas collected in one of the test tubes during the electrolysis of water double the amount collected in the other? Name this gas.
- Answer: Water ( $H_2O$ ) consists of hydrogen and oxygen chemically combined in a stoichiometric mole ratio of $2:1$ . During electrolysis, decomposition releases two volumes of hydrogen gas at the cathode for every one volume of oxygen gas produced at the anode:

$2H_2O(l) \rightarrow 2H_2(g) + O_2(g)$

The gas collected in double the quantity is hydrogen ( $H_2$ ).

Question 6: Why does the colour of copper sulphate solution change when an iron nail is dipped in it?
- Answer: Iron possesses higher chemical reactivity than copper. When immersed in blue copper sulphate solution ( $CuSO_4$ ), iron displaces copper to form iron(II) sulphate ( $FeSO_4$ ), which imparts a light green color to the aqueous solution:

$Fe(s) + CuSO_4(aq) \rightarrow FeSO_4(aq) + Cu(s)$

Question 7: Give an example of a double displacement reaction other than the reaction between sodium sulphate and barium chloride.
- Answer: Reaction between lead(II) nitrate and potassium iodide, which forms a yellow precipitate of lead iodide alongside aqueous potassium nitrate:

$Pb(NO_3)_2(aq) + 2KI(aq) \rightarrow PbI_2(s) + 2KNO_3(aq)$

Question 8: Identify the substances that are oxidised and the substances that are reduced in the following reactions:
1. $4Na(s) + O_2(g) \rightarrow 2Na_2O(s)$
2. $CuO(s) + H_2(g) \rightarrow Cu(s) + H_2O(l)$
- Answer:
1. Sodium ( $Na$ ) gains oxygen and is oxidised to $Na_2O$ ; oxygen gas ( $O_2$ ) is reduced.
2. Hydrogen ( $H_2$ ) gains oxygen and is oxidised to $H_2O$ ; copper(II) oxide ( $CuO$ ) loses oxygen and is reduced to metallic Cu$.\n\n# End-of-Chapter Exercises and Complete Solutions\n\n* **Exercise 1:** Which of the statements about the reaction below are incorrect?\n\n2PbO(s) + C(s) \rightarrow 2Pb(s) + CO_2(g)\n\n (a) Lead is getting reduced.\n (b) Carbon dioxide is getting oxidised.\n (c) Carbon is getting oxidised.\n (d) Lead oxide is getting reduced.\n Options:\n (i) (a) and (b)\n (ii) (a) and (c)\n (iii) (a), (b) and (c)\n (iv) All\n * **Answer:** **(i) (a) and (b)**\n * *Explanation:* Lead oxide (PbO $) loses oxygen and is reduced to lead metal ($ Pb $); carbon ($ C $) gains oxygen and is oxidised to carbon dioxide ($ CO_2). Statements (a) and (b) incorrectly state that the products are undergoing reduction/oxidation.\n* **Exercise 2:** Fe_2O_3 + 2Al \rightarrow Al_2O_3 + 2Fe\n The above reaction is an example of a:\n (a) Combination reaction\n (b) Double displacement reaction\n (c) Decomposition reaction\n (d) Displacement reaction\n * **Answer:** **(d) Displacement reaction**\n * *Explanation:* Aluminium (Al $) displaces iron ($ Fe $) from iron(III) oxide ($ Fe_2O_3) because aluminium is more reactive than iron.\n* **Exercise 3:** What happens when dilute hydrochloric acid is added to iron filings? Tick the correct answer:\n (a) Hydrogen gas and iron chloride are produced.\n (b) Chlorine gas and iron hydroxide are produced.\n (c) No reaction takes place.\n (d) Iron salt and water are produced.\n * **Answer:** **(a) Hydrogen gas and iron chloride are produced.**\n * *Equation:* Fe(s) + 2HCl(aq) \rightarrow FeCl_2(aq) + H_2(g)\n* **Exercise 4:** What is a balanced chemical equation? Why should chemical equations be balanced?\n * **Answer:** A balanced chemical equation has an equal count of atoms for each element on both reactant and product sides. Chemical equations must be balanced to satisfy the **Law of Conservation of Mass**, which states that matter cannot be created or destroyed in a chemical transformation.\n* **Exercise 5:** Translate the following statements into chemical equations and then balance them:\n * **(a) Hydrogen gas combines with nitrogen to form ammonia.**\n * Unbalanced: N_2 + H_2 \rightarrow NH_3\n * Balanced: N_2(g) + 3H_2(g) \rightarrow 2NH_3(g)\n * **(b) Hydrogen sulphide gas burns in air to give water and sulphur dioxide.**\n * Unbalanced: H_2S + O_2 \rightarrow H_2O + SO_2\n * Balanced: 2H_2S(g) + 3O_2(g) \rightarrow 2H_2O(l) + 2SO_2(g)\n * **(c) Barium chloride reacts with aluminium sulphate to give aluminium chloride and a precipitate of barium sulphate.**\n * Unbalanced: BaCl_2 + Al_2(SO_4)3 \rightarrow AlCl_3 + BaSO_4\n * Balanced: 3BaCl_2(aq) + Al_2(SO_4)_3(aq) \rightarrow 2AlCl_3(aq) + 3BaSO_4(s)\n * **(d) Potassium metal reacts with water to give potassium hydroxide and hydrogen gas.**\n * Unbalanced: K + H_2O \rightarrow KOH + H_2\n * Balanced: 2K(s) + 2H_2O(l) \rightarrow 2KOH(aq) + H_2(g)\n* **Exercise 6:** Balance the following chemical equations:\n * **(a) HNO_3 + Ca(OH)_2 \rightarrow Ca(NO_3)_2 + H_2O**\n * Balanced: 2HNO_3(aq) + Ca(OH)_2(aq) \rightarrow Ca(NO_3)_2(aq) + 2H_2O(l)\n * **(b) NaOH + H_2SO_4 \rightarrow Na_2SO_4 + H_2O**\n * Balanced: 2NaOH(aq) + H_2SO_4(aq) \rightarrow Na_2SO_4(aq) + 2H_2O(l)\n * **(c) NaCl + AgNO_3 \rightarrow AgCl + NaNO_3**\n * Balanced: NaCl(aq) + AgNO_3(aq) \rightarrow AgCl(s) + NaNO_3(aq)\n * **(d) BaCl_2 + H_2SO_4 \rightarrow BaSO_4 + HCl**\n * Balanced: BaCl_2(aq) + H_2SO_4(aq) \rightarrow BaSO_4(s) + 2HCl(aq)\n* **Exercise 7:** Write the balanced chemical equations for the following reactions:\n * **(a) Calcium hydroxide + Carbon dioxide \rightarrow Calcium carbonate + Water**\n * Equation: Ca(OH)_2(aq) + CO_2(g) \rightarrow CaCO_3(s) + H_2O(l)\n * **(b) Zinc + Silver nitrate \rightarrow Zinc nitrate + Silver**\n * Equation: Zn(s) + 2AgNO_3(aq) \rightarrow Zn(NO_3)_2(aq) + 2Ag(s)\n * **(c) Aluminium + Copper chloride \rightarrow Aluminium chloride + Copper**\n * Equation: 2Al(s) + 3CuCl_2(aq) \rightarrow 2AlCl_3(aq) + 3Cu(s)\n * **(d) Barium chloride + Potassium sulphate \rightarrow Barium sulphate + Potassium chloride**\n * Equation: BaCl_2(aq) + K_2SO_4(aq) \rightarrow BaSO_4(s) + 2KCl(aq)\n* **Exercise 8:** Write the balanced chemical equation for the following and identify the type of reaction in each case:\n * **(a) Potassium bromide(aq) + Barium iodide(aq) \rightarrow Potassium iodide(aq) + Barium bromide(s)**\n * Equation: 2KBr(aq) + BaI_2(aq) \rightarrow 2KI(aq) + BaBr_2(s)\n * Type: **Double Displacement (Precipitation) Reaction**\n * **(b) Zinc carbonate(s) \rightarrow Zinc oxide(s) + Carbon dioxide(g)**\n * Equation: ZnCO_3(s) \xrightarrow{\text{Heat}} ZnO(s) + CO_2(g)\n * Type: **Thermal Decomposition Reaction**\n * **(c) Hydrogen(g) + Chlorine(g) \rightarrow Hydrogen chloride(g)**\n * Equation: H_2(g) + Cl_2(g) \rightarrow 2HCl(g)\n * Type: **Combination Reaction**\n * **(d) Magnesium(s) + Hydrochloric acid(aq) \rightarrow Magnesium chloride(aq) + Hydrogen(g)**\n * Equation: Mg(s) + 2HCl(aq) \rightarrow MgCl_2(aq) + H_2(g)\n * Type: **Displacement Reaction**\n* **Exercise 9:** What does one mean by exothermic and endothermic reactions? Give examples.\n * **Exothermic Reaction:** A chemical process that releases thermal energy into its surrounding environment along with products.\n * *Example:* Combustion of methane:\n\nCH_4(g) + 2O_2(g) \rightarrow CO_2(g) + 2H_2O(g) + \text{Heat}\n\n * **Endothermic Reaction:** A chemical process that absorbs thermal or electrical energy from its surroundings to proceed.\n * *Example:* Thermal decomposition of limestone:\n\nCaCO_3(s) \xrightarrow{\text{Heat}} CaO(s) + CO_2(g)\n\n* **Exercise 10:** Why is respiration considered an exothermic reaction? Explain.\n * **Answer:** During digestion, complex dietary carbohydrates are broken down into glucose (C_6H{12}O_6). In cellular respiration, glucose reacts with dissolved oxygen inside body cells to yield carbon dioxide, water, and energy:\n\nC_6H_{12}O_6(aq) + 6O_2(aq) \rightarrow 6CO_2(aq) + 6H_2O(l) + \text{Energy}\n\n Because energy is released, cellular respiration is classified as an exothermic process.\n* **Exercise 11:** Why are decomposition reactions called the opposite of combination reactions? Write equations for these reactions.\n * **Answer:** In a combination reaction, two or more reactants combine to yield a single product. In a decomposition reaction, a single reactant breaks down into two or more simpler products.\n * *Combination Equation:* 2H_2(g) + O_2(g) \rightarrow 2H_2O(l)\n * *Decomposition Equation:* 2H_2O(l) \xrightarrow{\text{Electricity}} 2H_2(g) + O_2(g)\n* **Exercise 12:** Write one equation each for decomposition reactions where energy is supplied in the form of heat, light, or electricity.\n * **Decomposition by Heat (Thermal):**\n\n2Pb(NO_3)_2(s) \xrightarrow{\text{Heat}} 2PbO(s) + 4NO_2(g) + O_2(g)\n\n * **Decomposition by Light (Photolytic):**\n\n2AgCl(s) \xrightarrow{\text{Sunlight}} 2Ag(s) + Cl_2(g)\n\n * **Decomposition by Electricity (Electrolytic):**\n\n2H_2O(l) \xrightarrow{\text{Electricity}} 2H_2(g) + O_2(g)\n\n* **Exercise 13:** What is the difference between displacement and double displacement reactions? Write equations for these reactions.\n * **Displacement Reaction:** A single, more reactive element displaces a less reactive element from its compound solution.\n * *Equation:* Fe(s) + CuSO_4(aq) \rightarrow FeSO_4(aq) + Cu(s)\n * **Double Displacement Reaction:** Two distinct ionic compounds react by mutually exchanging their component ions to form two new chemical products.\n * *Equation:* Na_2SO_4(aq) + BaCl_2(aq) \rightarrow BaSO_4(s) + 2NaCl(aq)\n* **Exercise 14:** In the refining of silver, the recovery of silver from silver nitrate solution involved displacement by copper metal. Write down the reaction involved.\n * **Answer:** Copper metal displaces silver from aqueous silver nitrate solution because copper possesses higher chemical reactivity:\n\nCu(s) + 2AgNO_3(aq) \rightarrow Cu(NO_3)_2(aq) + 2Ag(s)\n\n* **Exercise 15:** What do you mean by a precipitation reaction? Explain by giving examples.\n * **Answer:** A precipitation reaction is a chemical process that forms an insoluble solid product (called a precipitate) upon mixing aqueous solutions of two soluble salts.\n * *Example:* Mixing potassium iodide solution with lead nitrate solution forms an insoluble yellow precipitate of lead iodide (PbI_2):\n\nPb(NO_3)_2(aq) + 2KI(aq) \rightarrow PbI_2(s) + 2KNO_3(aq)\n\n* **Exercise 16:** Explain the following in terms of gain or loss of oxygen with two examples each:\n * **(a) Oxidation:** Gain of oxygen atoms by a substance.\n * *Example 1:* 2Cu + O_2 \xrightarrow{\text{Heat}} 2CuO $(Copper gains oxygen to form$ CuO)\n * *Example 2:* C + O_2 \rightarrow CO_2 $(Carbon gains oxygen to form$ CO_2)\n * **(b) Reduction:** Loss of oxygen atoms by a substance.\n * *Example 1:* CuO + H_2 \xrightarrow{\text{Heat}} Cu + H_2O $($ CuO $loses oxygen to form$ Cu)\n * *Example 2:* ZnO + C \rightarrow Zn + CO $($ ZnO $loses oxygen to form$ Zn)\n* **Exercise 17:** A shiny brown coloured element 'X' on heating in air becomes black in colour. Name the element 'X' and the black coloured compound formed.\n * **Answer:** The shiny brown element 'X' is metallic **copper** (Cu $). Upon heating in the presence of atmospheric oxygen, it oxidises to form a black surface coating of **copper(II) oxide** ($ CuO):\n\n2Cu(s) + O_2(g) \xrightarrow{\text{Heat}} 2CuO(s)\n\n* **Exercise 18:** Why do we apply paint on iron articles?\n * **Answer:** Applying a layer of paint covers the metallic iron surface, preventing direct exposure to atmospheric moisture (H_2O $) and oxygen ($ O_2). This barrier blocks the electrochemical process of rusting.\n* **Exercise 19:** Oil and fat containing food items are flushed with nitrogen. Why?\n * **Answer:** Nitrogen (N_2) is an unreactive, inert gas. Flushing food packaging with nitrogen excludes atmospheric oxygen, preventing the oxidative degradation (rancidity) of fats and oils and extending shelf life.\n* **Exercise 20:** Explain the following terms with one example each:\n * **(a) Corrosion:** Deterioration of a metal caused by chemical reactions with environmental oxygen, moisture, or acids.\n * *Example:* Formation of a reddish-brown rusty coating (Fe_2O_3 \cdot xH_2O$$) on iron structures exposed to humid air.
- (b) Rancidity: Oxidation of fats and oils in stored food items that alters their taste and produces foul odors.
  - Example: Butter or potato chips left exposed to open air develop a sour taste and rancid smell.