Chemical Reactions and Equations

Chemical reactions occur when the nature and identity of initial substances change through a chemical change.
Daily Life Examples: Milk souring at room temperature in summer, rusting of iron tawa/pan/nail in humid air, fermentation of grapes, cooking food, digestion in the body, and respiration.
Observations determining a reaction: - Change in state - Change in colour - Evolution of a gas - Change in temperature

Word Equations: Represent reactions using names of substances, e.g., $\text{Magnesium} + ext{Oxygen} \rightarrow \text{Magnesium oxide}$ .
Skeletal Equations: Symbolic representations like $Mg + O_2 \rightarrow MgO$ are unbalanced (skeletal) if mass is not conserved.
Law of Conservation of Mass: Mass is neither created nor destroyed; hence, the number of atoms for each element must be equal on the reactant (LHS) and product (RHS) sides.
Balancing Technique: Uses the hit-and-trial method with the smallest whole number coefficients. The formulae of compounds must not be altered during this process.
Physical States: Specified using symbols $(s)$ for solid, $(l)$ for liquid, $(g)$ for gas, and $(aq)$ for aqueous (solution in water).
Reaction Conditions: Temperature, pressure ( $340\,atm$ ), or catalysts are indicated above or below the arrow.

Combination Reaction: Two or more substances combine to form a single product. - Example: Burning of coal ( $C(s) + O_2(g) \rightarrow CO_2(g)$ ). - Example: $CaO(s) + H_2O(l) \rightarrow Ca(OH)_2(aq) + \text{Heat}$ .
Whitewashing: Calcium hydroxide ( $Ca(OH)_2$ ) reacts with atmospheric $CO_2$ to form a thin layer of $CaCO_3$ (Calcium carbonate/marble) and water.
Exothermic Reactions: Reactions that release heat energy along with products. - Natural Gas: $CH_4(g) + 2O_2(g) \rightarrow CO_2(g) + 2H_2O(g)$ . - Respiration: $C_6H_{12}O_6(aq) + 6O_2(aq) \rightarrow 6CO_2(aq) + 6H_2O(l) + \text{energy}$ . - Decomposition of vegetable matter: Into compost is an exothermic process.

Decomposition Reaction: A single reactant breaks down into simpler products.
Thermal Decomposition: Done via heating. - Limestone decomposition: $CaCO_3(s) \xrightarrow{\text{Heat}} CaO(s) + CO_2(g)$ . - Ferrous sulphate: $2FeSO_4(s) \xrightarrow{\text{Heat}} Fe_2O_3(s) + SO_2(g) + SO_3(g)$ . - Lead nitrate: $2Pb(NO_3)_2(s) \xrightarrow{\text{Heat}} 2PbO(s) + 4NO_2(g) + O_2(g)$ (releases brown fumes of $NO_2$ ).
Electrolytic Decomposition: Electrolysis of water produces $H_2$ and $O_2$ gas.
Photolytic Decomposition: Silver chloride ( $AgCl$ ) or Silver bromide ( $AgBr$ ) decompose in sunlight, used in black and white photography: $2AgCl(s) \xrightarrow{\text{Sunlight}} 2Ag(s) + Cl_2(g)$ .
Endothermic Reactions: Reactions that absorb energy in the form of heat, light, or electricity.

Displacement Reaction: A more reactive element displaces a less reactive element from its compound. - Example: $Fe(s) + CuSO_4(aq) \rightarrow FeSO_4(aq) + Cu(s)$ (Iron nail turns brown, blue copper sulphate solution fades). - Zinc and lead are more reactive than copper.
Double Displacement Reaction: An exchange of ions occurs between the reactants. - Example: $Na_2SO_4(aq) + BaCl_2(aq) \rightarrow BaSO_4(s) + 2NaCl(aq)$ .
Precipitation Reaction: Any reaction that produces an insoluble solid known as a precipitate (e.g., $BaSO_4$ ).

Oxidation: Gain of oxygen or loss of hydrogen.
Reduction: Loss of oxygen or gain of hydrogen.
Redox Reactions: Both oxidation and reduction occur simultaneously. - Example: $CuO + H_2 \xrightarrow{\text{Heat}} Cu + H_2O$ ( $CuO$ is reduced to $Cu$ ; $H_2$ is oxidised to $H_2O$ ). - Other examples: $ZnO + C \rightarrow Zn + CO$ and $MnO_2 + 4HCl \rightarrow MnCl_2 + 2H_2O + Cl_2$ .

Corrosion: Deterioration of metals through moisture or acid attack. - Examples: Rusting of iron (reddish-brown powder), black coating on silver, and green coating on copper.
Rancidity: Oxidation of fats and oils in food leads to unpalatable changes in smell and taste. - Prevention: Use of antioxidants, airtight containers, or flushing food packages with nitrogen gas.

Q: Why should a magnesium ribbon be cleaned before burning?
A: To remove the oxide layer formed by atmospheric interaction so it burns effectively.
Q: Why is the amount of gas collected in one test tube in Activity 1.7 double the other?
A: The electrolysis of water ( $H_2O$ ) produces two volumes of Hydrogen for every one volume of Oxygen.
Q: Identify substance X used for whitewashing.
A: Substance X is Calcium oxide ( $CaO$ ), also known as quick lime. Its reaction with water is: $CaO(s) + H_2O(l) \rightarrow Ca(OH)_2(aq)$ .
Discussion Observations (Activities): - Activity 1.1: Magnesium ribbon burns with a dazzling white flame resulting in white powder ( $MgO$ ). - Activity 1.3: Zinc granules reacting with acid cause the conical flask temperature to rise (exothermic). - Activity 1.6: Heating lead nitrate produces brown fumes ( $NO_2$ ). - Activity 1.9: Iron nails in copper sulphate solution become brownish due to copper deposition.