Chemical Reactions and Equations - Comprehensive Study Notes Chemical Reactions and Equations Daily Life Observations Milk left at room temperature during summers changes. Iron tawa/pan/nail exposed to humid atmosphere corrodes. Grapes get fermented. Food is cooked and digested. Respiration occurs. In these situations, the initial substance's nature and identity change, indicating a chemical reaction. Identifying Chemical Reactions Activity 1.1: Burning Magnesium Ribbon Clean a magnesium ribbon (3-4 cm long) with sandpaper. Hold it with tongs and burn it using a spirit lamp or burner. Collect the ash (magnesium oxide) in a watch-glass. Magnesium ribbon burns with a dazzling white flame and transforms into a white powder (magnesium oxide). This is due to the reaction between magnesium and oxygen in the air. Activity 1.2: Reaction of Zinc with Acid Take zinc granules in a conical flask or test tube. Add dilute hydrochloric acid or sulphuric acid. Observe any reactions happening around the zinc granules (evolution of gas). Touch the flask or test tube to check for temperature change. Hydrogen gas is formed by the action of dilute sulphuric acid on zinc. Activity 1.3: Reaction of Lead Nitrate with Potassium Iodide Take lead nitrate solution in a test tube. Add potassium iodide solution. Observe any changes (formation of precipitate). Indicators of a Chemical Reaction Change in state Change in color Evolution of a gas Change in temperature Chemical Equations Word Equations Represent a chemical reaction using words. Example: Magnesium + Oxygen → Magnesium oxide Reactants: Substances undergoing change (e.g., magnesium, oxygen). Product: New substance formed (e.g., magnesium oxide). Reactants are written on the left-hand side (LHS) with a plus sign (+) between them. Products are written on the right-hand side (RHS) with a plus sign (+) between them. An arrow (→) indicates the direction of the reaction. Use chemical formulas instead of words for a more concise representation. Example: M g + O 2 → M g O Mg + O_2 → MgO M g + O 2 → M g O Check if the number of atoms of each element is the same on both sides (balanced equation). If not, the equation is unbalanced (skeletal chemical equation). Balanced Chemical Equations Follow the law of conservation of mass: mass is neither created nor destroyed. The total mass of elements in the products must equal the total mass of elements in the reactants. The number of atoms of each element remains the same before and after the reaction. Balancing is done step by step. Example: Balancing F e + H < e m > 2 O → F e < / e m > 3 O < e m > 4 + H < / e m > 2 Fe + H<em>2O → Fe</em>3O<em>4 + H</em>2 F e + H < e m > 2 O → F e < / e m > 3 O < e m > 4 + H < / e m > 2 Draw boxes around each formula: F e + H < e m > 2 O → F e < / e m > 3 O < e m > 4 + H < / e m > 2 Fe + H<em>2O → Fe</em>3O<em>4 + H</em>2 F e + H < e m > 2 O → F e < / e m > 3 O < e m > 4 + H < / e m > 2 List the number of atoms of each element on both sides. Balance the element with the maximum number of atoms (e.g., oxygen in F e < e m > 3 O < / e m > 4 Fe<em>3O</em>4 F e < e m > 3 O < / e m > 4 F e + 4 H < e m > 2 O → F e < / e m > 3 O < e m > 4 + H < / e m > 2 Fe + 4H<em>2O → Fe</em>3O<em>4 + H</em>2 F e + 4 H < e m > 2 O → F e < / e m > 3 O < e m > 4 + H < / e m > 2 Balance hydrogen atoms.F e + 4 H < e m > 2 O → F e < / e m > 3 O < e m > 4 + 4 H < / e m > 2 Fe + 4H<em>2O → Fe</em>3O<em>4 + 4H</em>2 F e + 4 H < e m > 2 O → F e < / e m > 3 O < e m > 4 + 4 H < / e m > 2 Balance iron atoms.3 F e + 4 H < e m > 2 O → F e < / e m > 3 O < e m > 4 + 4 H < / e m > 2 3Fe + 4H<em>2O → Fe</em>3O<em>4 + 4H</em>2 3 F e + 4 H < e m > 2 O → F e < / e m > 3 O < e m > 4 + 4 H < / e m > 2 Check that the number of atoms of each element is equal on both sides. This method is called the hit-and-trial method. Symbols of Physical States Indicate the physical state of reactants and products. (g) for gaseous, (l) for liquid, (aq) for aqueous (solution in water), (s) for solid. Example: 3 F e ( s ) + 4 H < e m > 2 O ( g ) → F e < / e m > 3 O < e m > 4 ( s ) + 4 H < / e m > 2 ( g ) 3Fe(s) + 4H<em>2O(g) → Fe</em>3O<em>4(s) + 4H</em>2(g) 3 F e ( s ) + 4 H < e m > 2 O ( g ) → F e < / e m > 3 O < e m > 4 ( s ) + 4 H < / e m > 2 ( g ) Reaction conditions (temperature, pressure, catalyst) may be indicated above/below the arrow. Example: C O ( g ) + 2 H < e m > 2 ( g ) → 340 a t m C H < / e m > 3 O H ( l ) CO(g) + 2H<em>2(g) \xrightarrow[340 atm]{} CH</em>3OH(l) CO ( g ) + 2 H < e m > 2 ( g ) 340 a t m C H < / e m > 3 O H ( l ) Types of Chemical Reactions Combination Reaction Two or more reactants combine to form a single product. Activity 1.4: Calcium oxide (quick lime) reacts with water to produce slaked lime (calcium hydroxide) and heat.C a O ( s ) + H < e m > 2 O ( l ) → C a ( O H ) < / e m > 2 ( a q ) + H e a t CaO(s) + H<em>2O(l) → Ca(OH)</em>2(aq) + Heat C a O ( s ) + H < e m > 2 O ( l ) → C a ( O H ) < / e m > 2 ( a q ) + He a t Other examples:Burning of coal: C ( s ) + O < e m > 2 ( g ) → C O < / e m > 2 ( g ) C(s) + O<em>2(g) → CO</em>2(g) C ( s ) + O < e m > 2 ( g ) → CO < / e m > 2 ( g ) Formation of water: 2 H < e m > 2 ( g ) + O < / e m > 2 ( g ) → 2 H 2 O ( l ) 2H<em>2(g) + O</em>2(g) → 2H_2O(l) 2 H < e m > 2 ( g ) + O < / e m > 2 ( g ) → 2 H 2 O ( l ) Exothermic reactions : Reactions that release heat.Burning of natural gas: C H < e m > 4 ( g ) + 2 O < / e m > 2 ( g ) → C O < e m > 2 ( g ) + 2 H < / e m > 2 O ( g ) CH<em>4(g) + 2O</em>2(g) → CO<em>2(g) + 2H</em>2O(g) C H < e m > 4 ( g ) + 2 O < / e m > 2 ( g ) → CO < e m > 2 ( g ) + 2 H < / e m > 2 O ( g ) Respiration: C < e m > 6 H < / e m > 12 O < e m > 6 ( a q ) + 6 O < / e m > 2 ( a q ) → 6 C O < e m > 2 ( a q ) + 6 H < / e m > 2 O ( l ) + e n e r g y C<em>6H</em>{12}O<em>6(aq) + 6O</em>2(aq) → 6CO<em>2(aq) + 6H</em>2O(l) + energy C < e m > 6 H < / e m > 12 O < e m > 6 ( a q ) + 6 O < / e m > 2 ( a q ) → 6 CO < e m > 2 ( a q ) + 6 H < / e m > 2 O ( l ) + e n er g y Decomposition of vegetable matter into compost. Calcium hydroxide reacts with carbon dioxide in air to form calcium carbonate (used in whitewashing).C a ( O H ) < e m > 2 ( a q ) + C O < / e m > 2 ( g ) → C a C O < e m > 3 ( s ) + H < / e m > 2 O ( l ) Ca(OH)<em>2(aq) + CO</em>2(g) → CaCO<em>3(s) + H</em>2O(l) C a ( O H ) < e m > 2 ( a q ) + CO < / e m > 2 ( g ) → C a CO < e m > 3 ( s ) + H < / e m > 2 O ( l ) Decomposition Reaction A single reactant breaks down into simpler products. Activity 1.5: Ferrous sulphate crystals (F e S O < e m > 4 , 7 H < / e m > 2 O FeSO<em>4, 7H</em>2O F e SO < e m > 4 , 7 H < / e m > 2 O F e < e m > 2 O < / e m > 3 Fe<em>2O</em>3 F e < e m > 2 O < / e m > 3 S O < e m > 2 SO<em>2 SO < e m > 2 S O < / e m > 3 SO</em>3 SO < / e m > 3 2 F e S O < e m > 4 ( s ) → H e a t F e < / e m > 2 O < e m > 3 ( s ) + S O < / e m > 2 ( g ) + S O 3 ( g ) 2FeSO<em>4(s) \xrightarrow[]{Heat} Fe</em>2O<em>3(s) + SO</em>2(g) + SO_3(g) 2 F e SO < e m > 4 ( s ) He a t F e < / e m > 2 O < e m > 3 ( s ) + SO < / e m > 2 ( g ) + S O 3 ( g ) Decomposition of calcium carbonate (limestone):C a C O < e m > 3 ( s ) → H e a t C a O ( s ) + C O < / e m > 2 ( g ) CaCO<em>3(s) \xrightarrow[]{Heat} CaO(s) + CO</em>2(g) C a CO < e m > 3 ( s ) He a t C a O ( s ) + CO < / e m > 2 ( g ) Calcium oxide (lime or quick lime) is used in the manufacture of cement. Activity 1.6: Heating lead nitrate powder emits brown fumes of nitrogen dioxide (N O 2 NO_2 N O 2 2 P b ( N O < e m > 3 ) < / e m > 2 ( s ) → H e a t 2 P b O ( s ) + 4 N O < e m > 2 ( g ) + O < / e m > 2 ( g ) 2Pb(NO<em>3)</em>2(s) \xrightarrow[]{Heat} 2PbO(s) + 4NO<em>2(g) + O</em>2(g) 2 P b ( NO < e m > 3 ) < / e m > 2 ( s ) He a t 2 P b O ( s ) + 4 NO < e m > 2 ( g ) + O < / e m > 2 ( g ) Thermal decomposition : Decomposition by heating.Activity 1.7: Silver chloride turns grey in sunlight due to decomposition into silver and chlorine.2 A g C l ( s ) → S u n l i g h t 2 A g ( s ) + C l 2 ( g ) 2AgCl(s) \xrightarrow[]{Sunlight} 2Ag(s) + Cl_2(g) 2 A g Cl ( s ) S u n l i g h t 2 A g ( s ) + C l 2 ( g ) Activity 1.8: Electrolysis of water. Silver bromide also decomposes in sunlight:2 A g B r ( s ) → S u n l i g h t 2 A g ( s ) + B r 2 ( g ) 2AgBr(s) \xrightarrow[]{Sunlight} 2Ag(s) + Br_2(g) 2 A g B r ( s ) S u n l i g h t 2 A g ( s ) + B r 2 ( g ) Used in black and white photography. Endothermic reactions : Reactions that absorb energy (in the form of heat, light, or electricity).Displacement Reaction An element displaces another element from its compound. Activity 1.9: Iron nails dipped in copper sulphate solution become brownish, and the blue color of the solution fades.F e ( s ) + C u S O < e m > 4 ( a q ) → F e S O < / e m > 4 ( a q ) + C u ( s ) Fe(s) + CuSO<em>4(aq) → FeSO</em>4(aq) + Cu(s) F e ( s ) + C u SO < e m > 4 ( a q ) → F e SO < / e m > 4 ( a q ) + C u ( s ) Other examples:Z n ( s ) + C u S O < e m > 4 ( a q ) → Z n S O < / e m > 4 ( a q ) + C u ( s ) Zn(s) + CuSO<em>4(aq) → ZnSO</em>4(aq) + Cu(s) Z n ( s ) + C u SO < e m > 4 ( a q ) → Z n SO < / e m > 4 ( a q ) + C u ( s ) P b ( s ) + C u C l < e m > 2 ( a q ) → P b C l < / e m > 2 ( a q ) + C u ( s ) Pb(s) + CuCl<em>2(aq) → PbCl</em>2(aq) + Cu(s) P b ( s ) + C u Cl < e m > 2 ( a q ) → P b Cl < / e m > 2 ( a q ) + C u ( s ) Zinc and lead are more reactive than copper. Double Displacement Reaction Exchange of ions between reactants. Activity 1.10: Sodium sulphate reacts with barium chloride to form a white precipitate of barium sulphate.N a < e m > 2 S O < / e m > 4 ( a q ) + B a C l < e m > 2 ( a q ) → B a S O < / e m > 4 ( s ) + 2 N a C l ( a q ) Na<em>2SO</em>4(aq) + BaCl<em>2(aq) → BaSO</em>4(s) + 2NaCl(aq) N a < e m > 2 SO < / e m > 4 ( a q ) + B a Cl < e m > 2 ( a q ) → B a SO < / e m > 4 ( s ) + 2 N a Cl ( a q ) Precipitation reaction: Any reaction that produces a precipitate. Oxidation and Reduction Activity 1.11: Heating copper powder in a china dish causes the surface to become coated with black copper(II) oxide.2 C u + O 2 → H e a t 2 C u O 2Cu + O_2 \xrightarrow[]{Heat} 2CuO 2 C u + O 2 He a t 2 C u O Passing hydrogen gas over heated copper oxide turns the black coating brown, forming copper.C u O + H < e m > 2 → H e a t C u + H < / e m > 2 O CuO + H<em>2 \xrightarrow[]{Heat} Cu + H</em>2O C u O + H < e m > 2 He a t C u + H < / e m > 2 O Oxidation : Gain of oxygen.Reduction : Loss of oxygen.Redox reactions : Reactions where one reactant gets oxidized while the other gets reduced.Examples:Z n O + C → Z n + C O ZnO + C → Zn + CO Z n O + C → Z n + CO M n O < e m > 2 + 4 H C l → M n C l < / e m > 2 + H < e m > 2 O + C l < / e m > 2 MnO<em>2 + 4HCl → MnCl</em>2 + H<em>2O + Cl</em>2 M n O < e m > 2 + 4 H Cl → M n Cl < / e m > 2 + H < e m > 2 O + Cl < / e m > 2 Oxidation: A substance gains oxygen or loses hydrogen. Reduction: A substance loses oxygen or gains hydrogen. Effects of Oxidation Reactions in Everyday Life Corrosion Metals are attacked by substances around them (moisture, acids, etc.). Rusting of iron: Iron articles get coated with reddish-brown powder. Black coating on silver and green coating on copper are examples of corrosion. Corrosion damages car bodies, bridges, iron railings, ships, etc. Rancidity Fats and oils are oxidized, changing their smell and taste. Substances that prevent oxidation (antioxidants) are added to foods. Keeping food in airtight containers slows down oxidation. Chips manufacturers flush bags with nitrogen gas to prevent oxidation. Key Concepts Chemical Equation : Symbolic representation of reactants, products, and their physical states.Balanced Equation : Number of atoms of each element is the same on both sides.Combination Reaction : Two or more substances combine to form a single substance.Decomposition Reaction : A single substance breaks down into two or more substances.Exothermic Reaction : Heat is released.Endothermic Reaction : Energy is absorbed.Displacement Reaction : An element displaces another from its compound.Double Displacement Reaction : Exchange of ions between reactants.Precipitation Reaction : Produces insoluble salts.Oxidation : Gain of oxygen or loss of hydrogen.Reduction : Loss of oxygen or gain of hydrogen.Corrosion : Metal is attacked by substances around it.Rancidity : Oxidation of fats and oils, leading to changes in smell and taste