Chemical Reactions (Chemistry) PDF

Chemical Reactions and Equations The word-equation (Balanced equation) Magnesium + Oxygen → Magnesium oxide 2Mg + O₂ → 2MgO (Reactants) (Product) 3Fe + 4H₂O → Fe₃O₄ + 4H₂ Sometimes the reaction conditions...

Chemical Reactions and Equations The word-equation (Balanced equation) Magnesium + Oxygen → Magnesium oxide 2Mg + O₂ → 2MgO (Reactants) (Product) 3Fe + 4H₂O → Fe₃O₄ + 4H₂ Sometimes the reaction conditions, such as temperature, pressure, catalyst, etc., for the reaction are indicated above and/or below the arrow in the equation. For example – Combination Reaction: The reaction in which two or more reactants combine to form a single product. (i) Burning of coal C(s) + O₂(g) → CO₂(g) (ii) Formation of water 2H₂(g) + O₂(g) → 2H₂O(l) (iii) CaO(s) + H₂O(l) → Ca(OH)₂ (aq) + Heat (Quick lime) (Slaked lime) Calcium hydroxide reacts slowly with the carbon dioxide in air to form a thin layer of calcium carbonate on the walls. Ca(OH)₂ (aq) + CO₂ (g) → CaCO₃ (s) + H₂O(l) (Calcium hydroxide) (Calcium carbonate) Exothermic Reactions: Reaction in which heat is released along with formation of products. (i) Burning of natural gas. CH4(g) + 2O₂(g) → CO₂(g) + 2H₂O(g) + Heat (ii) Respiration is also an exothermic reaction. C₆H₁₂O₆(aq) + 6O₂(g) → 6CO₂(aq) + 6H₂O(l) + energy (Glucose) Decomposition Reaction: The reaction in which a compound splits into two or more simpler substances is called decomposition reaction. A→B+C (a) Thermal decomposition: When decomposition is carried out by heating. (iii) Heating of lead nitrate and emission of nitrogen dioxide (b) Electrolytic Decomposition: When decomposition is carried out by passing electricity. (c) Photolytic Decomposition When decomposition is carried out in presence of sunlight. Endothermic Reaction: The reactions which require energy in the form of heat, light or electricity to break reactants are called endothermic reactions. Displacement Reaction: The chemical reactions in which more reactive element displaces less reactive element from its salt solution. (iii) Pb(s) + CuCl₂ (aq) → PbCl₂ (aq) + Cu(s) (Copper chloride) (Lead chloride)  Double Displacement Reaction: A reaction in which new compounds are formed by mutual exchange of ions between two compounds. White precipitate of BaSO4 is formed, so it is also called precipitation reaction. Oxidation: Reduction It is a process of gaining oxygen It is the gain of electrons or a decrease in during a reaction by an atom, the oxidation state of an atom by another molecule or ion. atom, an ion or a molecule. In this reaction, CuO is reduced to Cu and H2 is oxidised to H2O. In other words, one reactant gets oxidised while the other gets reduced. Such reactions are called oxidation-reduction reactions or redox reactions.  Acid , Bases and Salts Reaction of Metals with: Acids: Acid + Metal → Salt + Hydrogen gas Bases: Base + Metal → Salt + Hydrogen gas 2HCl + Zn → ZnCl₂ + H₂ ↑ 2NaOH + Zn → Na₂ZnO₂ + H₂ ↑ (Zinc chloride) (Sodium zincate) Test for H2 gas: Hydrogen gas released can be tested by bringing a burning candle near gas bubbles, it bursts with pop sound. Reaction of Metal Carbonates / Metal Hydrogen Carbonates with: Acids: Acid + Metal Carbonate / Metal hydrogen Carbonate → Salt + CO₂ + H₂O 2HCl + Na₂CO₃ → 2NaCl + CO₂ + H₂O HCl + NaHCO₃ → NaCl + CO₂ + H₂O Bases: Base + Metal Carbonate / Metal Hydrogen Carbonate→ No Reaction 2NaOH + Zn → Na₂ZnO₂ + H₂ ↑ (Sodium zincate) Test for CO₂: CO₂ can be tested by passing it through lime water. Lime water turns milky. Ca(OH)₂ + CO₂ → CaCO₃ + H₂O (Lime water) (White precipitate) When excess CO₂ is passed, milkiness disappears. CaCO₃ + CO₂ + H₂O → Ca(HCO₃)₂ (Soluble in water) Reaction of Acids and Bases With Each Other: Acid + Base → Salt + Water NaOH(aq) + HCl(aq) → NaCl(aq) + H₂O(l) Common property between all acids and all bases When acids are dissolved in water they dissociate as H+ ions. HCl (aq) → H+(aq) + Cl– (aq) HNO₃(aq) → H+(aq) + NO₃ – (aq) CH₃COOH (aq) → H+ (aq) + CH₃COO− (aq) As H+ ion cannot exist alone so it combines with water molecules and forms H3O+ (hydronium) ions. Reaction of acids with metal oxides : Metal oxide + Acid → Salt + Water Copper oxide reacts with dil. hydrochloric acid to form copper chloride (salt) and water. CuO + 2HCl → CuCl₂ + H₂O Copper oxide Copper chloride Sodium Chloride (NaCl): Preparation of Sodium Hydroxide: Bleaching Powder: Ca(OH)₂(aq)+Cl₂(g)→CaOCl₂(aq)+H₂O(l) Baking Soda: Washing Soda: Plaster of Paris:  Metals and Non-metals Reaction of Metals with Air: Metals combine with oxygen to form metal oxide. Metals + O₂ → Metal oxide Examples: (i) 2Cu + O₂ → 2CuO Copper (II) oxide (black) (ii) 4Al + 3O₂ → 2Al₂O₃ Aluminium oxide (iii) 2Mg + O₂ → 2MgO Magnesium oxide Amphoteric Oxides: Metal oxides which react with both acids as well as bases to produce salt and water are called amphoteric oxides. Examples: Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O Aluminium chloride Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O Sodium aluminate Sodium oxide and potassium oxide dissolve in water to produce alkalis Na₂O(s) + H₂O(l) → 2NaOH(aq) K₂O(s) + H₂O(l) → 2KOH(aq) Reaction of Metals with Water: Metals react with water to produce metal hydroxide and hydrogen gas. Metal + Water → Metal oxide + Hydrogen Examples: Metal oxide + Water → Metal hydroxide 2Mg + 2H₂O → 2MgO + 2H₂ ↑ Magnesium oxide MgO + H₂O → Mg(OH)₂ Magnesium hydroxide Metals like potassium and sodium react violently with cold water. 2K(s) + 2H₂O(l) → 2KOH(aq) + H₂ (g) + heat energy 2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂ (g) + heat energy The reaction of calcium with water is less violent. Ca(s) + 2H₂O(l) → Ca(OH)₂ (aq) + H₂ (g) Metals like aluminium, iron and zinc do not react either with cold or hot water. But they react with steam to form the metal oxide and hydrogen. 2Al(s) + 3H₂O(g) → Al₂O₃ (s) + 3H₂ (g) 3Fe(s) + 4H₂O(g) → Fe₃ O₄ (s) + 4H₂ (g)  Reactions of Metals with Acid: Copper, mercury and silver don’t Metal + Dil. Acid → Salt + Hydrogen gas react with dilute e.g., Mg + H₂SO₄ → MgSO₄ + H₂ acids. Reaction of Metals with solutions of other Metal Salts: Metal A + Salt solution B → Salt solution A + Metal B Fe + CuSO₄ → FeSO₄ + Cu Reaction of Non-Metals: Reaction with oxygen: Non- metals react with oxygen to form acidic oxides. e.g., C + O₂ → CO₂ Reaction with chlorine: Non-metals react with chlorine to form their respective Chlorides. e.g., H₂ + Cl₂ → 2HCl Reaction with hydrogen: Non-metals react with hydrogen to form their respective hydrides. e.g., H₂ + S → H₂S Extracting Metals Low in the Activity Series: For example, cinnabar (HgS) is an ore of mercury. When it is heated in air, it is first converted into mercuric oxide (HgO). Mercuric oxide is then reduced to mercury on further heating. Similarly, copper which is found as Cu2S in nature can be obtained from its ore by just heating in air. Extracting Metals in the Middle of the Activity Series  when manganese dioxide is heated with aluminum powder, the following reaction takes place – 3MnO₂ (s) + 4Al(s) → 3Mn(l) + 2Al₂O₃ (s) + Heat Extracting Metals towards the Top of the Activity Series: These metals are obtained by electrolytic reduction. For example, sodium, magnesium and calcium are obtained by the electrolysis of their molten chlorides. The metals are deposited at the cathode , whereas, chlorine is liberated at the anode. The reactions are – At cathode Na+ + e– → Na At anode 2Cl– → Cl₂ + 2e– Thermit reaction: Fe₂O₃ (s) + 2Al(s) → 2Fe(l) + Al₂O₃ (s) + Heat