Chapter 1: Chemical Reactions and Equations - Comprehensive Study Notes

Introduction to Chemical Reactions

• Chemical reactions are processes that occur in various daily life situations where the nature and identity of the initial substance undergo a change.
• Daily Life Examples of Chemical Changes:
  • Milk being left at room temperature during the summer season.
  • An iron tawa, pan, or nail being left exposed to a humid atmosphere.
  • Grapes undergoing the process of fermentation.
  • The cooking of food items.
  • The digestion of food within the human body.
  • The process of respiration.
• Whenever a chemical change occurs, it is an indication that a chemical reaction has taken place.
• "Facts are not science — as the dictionary is not literature." — Martin H. Fischer.

Characteristics of Chemical Reactions

• Based on investigative activities, various observations help determine if a chemical reaction has occurred:
  • Change in state: The physical form (solid, liquid, gas) of the substances changes.
  • Change in colour: The visual appearance or hue of the substance is altered.
  • Evolution of a gas: The production and release of gaseous bubbles or fumes.
  • Change in temperature: The reaction mixture either releases heat (becomes warm) or absorbs heat (becomes cool).

Activity-Based Observations

• Activity 1.1: Burning of Magnesium Ribbon:
  • Procedure: Clean a magnesium ribbon (roughly $3-4\,cm$ long) with sandpaper to remove impurities. Hold it with tongs and burn it using a spirit lamp. Collect the resulting ash in a watch-glass.
  • Caution: Use teacher assistance and wear suitable eyeglasses. Keep the burning ribbon as far from the eyes as possible.
  • Observation: The magnesium ribbon burns with a dazzling white flame and is converted into a white powder.
  • Explanation: This powder is magnesium oxide ($MgO$), formed by the reaction between magnesium and oxygen present in the air.
• Activity 1.2: Lead Nitrate and Potassium Iodide:
  • Procedure: Take lead nitrate solution in a test tube and add potassium iodide solution.
  • Observation: A chemical reaction occurs, resulting in the formation of a precipitate.
• Activity 1.3: Zinc Granules with Acid:
  • Procedure: Add a few zinc granules to a conical flask. Add dilute hydrochloric acid or sulphuric acid.
  • Caution: Handle acids with extreme care.
  • Observation: Gas bubbles (hydrogen) form around the zinc granules. The conical flask or test tube becomes hot to the touch, indicating a temperature change.

Chemical Equations

• A chemical reaction can be described in sentence form, but it is more concise to use a word-equation or a symbolic chemical equation.
• Word-Equation Example:
  • $Magnesium + Oxygen \rightarrow Magnesium\ oxide$
  • Reactants: Substances that undergo chemical change, written on the Left-Hand Side (LHS) with a plus sign ($+$).
  • Products: New substances formed during the reaction, written on the Right-Hand Side (RHS) with a plus sign ($+$).
  • Arrowhead: Points toward the products and indicates the direction of the reaction.
• Writing a Chemical Equation:
  • Using chemical formulae makes representations more concise.
  • $Mg + O_{2} \rightarrow MgO$
  • This is a skeletal chemical equation if the number of atoms of each element is not equal on both sides.

Balanced Chemical Equations

• Law of Conservation of Mass: Mass can neither be created nor destroyed in a chemical reaction. The total mass of elements in products must equal the total mass of elements in reactants.
• To satisfy this law, the number of atoms of each element must remain the same before and after the reaction.

Step-by-Step Balancing (Hit-and-Trial Method)

• Example Equation: $Fe + H_{2}O \rightarrow Fe_{3}O_{4} + H_{2}$
• Step I: Draw boxes around each formula. Do not alter anything inside the boxes.
• Step II: List the initial number of atoms of different elements.
  • $Fe$: LHS = $1$, RHS = $3$
  • $H$: LHS = $2$, RHS = $2$
  • $O$: LHS = $1$, RHS = $4$
• Step III: Start balancing with the compound containing the maximum number of atoms ($Fe_{3}O_{4}$). To balance oxygen, multiply LHS $H_{2}O$ by $4$.
  • Equation: $Fe + 4H_{2}O \rightarrow Fe_{3}O_{4} + H_{2}$
• Step IV: Balance Hydrogen. LHS now has $8$ atoms of $H$. Multiply RHS $H_{2}$ by $4$.
  • Equation: $Fe + 4H_{2}O \rightarrow Fe_{3}O_{4} + 4H_{2}$
• Step V: Balance Iron ($Fe$). Multiply LHS $Fe$ by $3$.
  • Final Balanced Equation: $3Fe + 4H_{2}O \rightarrow Fe_{3}O_{4} + 4H_{2}$
• Step VI: Verify the atom counts on both sides.

Writing Symbols of Physical States and Reaction Conditions

• To provide more information, physical states are indicated:
  • (s) for solid
  • (l) for liquid
  • (g) for gas
  • (aq) for aqueous (a substance dissolved in water)
• The equation for the reaction of iron with steam is represented as:
  • $3Fe(s) + 4H_{2}O(g) \rightarrow Fe_{3}O_{4}(s) + 4H_{2}(g)$
  • Note: $(g)$ is used for $H_{2}O$ because it is in the form of steam.
• Reaction Conditions: Temperature, pressure, or catalysts are written above or below the arrow.
  • Carbon Monoxide and Hydrogen: $CO(g) + 2H_{2}(g) \xrightarrow{340\,atm} CH_{3}OH(l)$
  • Photosynthesis: $6CO_{2}(aq) + 12H_{2}O(l) \xrightarrow[Chlorophyll]{Sunlight} C_{6}H_{12}O_{6}(aq) + 6O_{2}(aq) + 6H_{2}O(l)$

Types of Chemical Reactions

• Chemical reactions involve the breaking and making of bonds between atoms to produce new substances.

1. Combination Reaction

• A reaction where a single product is formed from two or more reactants.
• Activity 1.4: Calcium oxide (Quick lime) reacts vigorously with water to produce slaked lime (calcium hydroxide).
  • $CaO(s) + H_{2}O(l) \rightarrow Ca(OH)_{2}(aq) + Heat$
• Application of Slaked Lime: Used for whitewashing walls. It reacts with atmospheric $CO_{2}$ to form a thin, shiny layer of calcium carbonate ($CaCO_{3}$) over $2-3$ days.
  • $Ca(OH)_{2}(aq) + CO_{2}(g) \rightarrow CaCO_{3}(s) + H_{2}O(l)$
  • Note: Marble also has the chemical formula $CaCO_{3}$.
• Other Examples:
  • Burning of coal: $C(s) + O_{2}(g) \rightarrow CO_{2}(g)$
  • Formation of water: $2H_{2}(g) + O_{2}(g) \rightarrow 2H_{2}O(l)$

2. Decomposition Reaction

• A reaction in which a single reactant breaks down to give simpler products. It is the opposite of a combination reaction.
• Thermal Decomposition: Decomposition carried out by heating.
  • Ferrous Sulphate: $2FeSO_{4}(s) \xrightarrow{Heat} Fe_{2}O_{3}(s) + SO_{2}(g) + SO_{3}(g)$
  • Green ferrous sulphate crystals ($FeSO_{4} \cdot 7H_{2}O$) lose water on heating and change colour. It decomposes into solid ferric oxide and gaseous sulphur dioxide and trioxide.
  • Calcium Carbonate (Limestone): $CaCO_{3}(s) \xrightarrow{Heat} CaO(s) + CO_{2}(g)$
  • Quick lime ($CaO$) produced is used in the manufacture of cement.
  • Lead Nitrate: $2Pb(NO_{3})_{2}(s) \xrightarrow{Heat} 2PbO(s) + 4NO_{2}(g) + O_{2}(g)$
  • This reaction produces brown fumes of nitrogen dioxide ($NO_{2}$).
• Electrolytic Decomposition (Electrolysis):
  • Activity 1.7: Decomposition of water into hydrogen and oxygen using an electric current ($6V$ battery).
  • Observation: Bubbles form at electrodes. The volume of gas at the cathode (hydrogen) is double the volume at the anode (oxygen).
• Photolytic Decomposition: Decomposition using light energy.
  • Silver Chloride: $2AgCl(s) \xrightarrow{Sunlight} 2Ag(s) + Cl_{2}(g)$
  • White silver chloride turns grey as it decomposes into silver metal and chlorine gas.
  • Silver Bromide: $2AgBr(s) \xrightarrow{Sunlight} 2Ag(s) + Br_{2}(g)$
  • This reaction is used in black and white photography.

3. Displacement Reaction

• A reaction where a more reactive element displaces or removes a less reactive element from its compound.
• Activity 1.9: Iron nails in copper sulphate solution.
  • $Fe(s) + CuSO_{4}(aq) \rightarrow FeSO_{4}(aq) + Cu(s)$
  • Observation: The blue colour of $CuSO_{4}$ solution fades, and the iron nail becomes brownish due to the deposition of copper.
• Other Examples:
  • $Zn(s) + CuSO_{4}(aq) \rightarrow ZnSO_{4}(aq) + Cu(s)$
  • $Pb(s) + CuCl_{2}(aq) \rightarrow PbCl_{2}(aq) + Cu(s)$
  • Zinc and lead are more reactive than copper.

4. Double Displacement Reaction

• Reactions in which there is an exchange of ions between the reactants.
• Activity 1.10: Mixing sodium sulphate and barium chloride.
  • $Na_{2}SO_{4}(aq) + BaCl_{2}(aq) \rightarrow BaSO_{4}(s) + 2NaCl(aq)$
  • Observation: A white insoluble substance (precipitate) of barium sulphate ($BaSO_{4}$) is formed by the reaction of $SO_{4}^{2-}$ and $Ba^{2+}$.
• Precipitation Reaction: Any reaction that produces an insoluble solid called a precipitate.

Energy Changes in Reactions

• Exothermic Reactions: Reactions in which heat is released along with the formation of products.
  • Examples: Burning of natural gas ($CH_{4}$), respiration, and decomposition of vegetable matter into compost.
  • Respiration Equation: $C_{6}H_{12}O_{6}(aq) + 6O_{2}(aq) \rightarrow 6CO_{2}(aq) + 6H_{2}O(l) + energy$
• Endothermic Reactions: Reactions in which energy (heat, light, or electricity) is absorbed to break down reactants.
  • Example: Mixing barium hydroxide with ammonium chloride results in the test tube becoming cold.

Oxidation and Reduction (Redox Reactions)

• Oxidation: The gain of oxygen or the loss of hydrogen by a substance.
• Reduction: The loss of oxygen or the gain of hydrogen by a substance.
• Redox Reactions: Reactions where one reactant is oxidised while the other is reduced simultaneously.
• Example: Heating copper powder in air.
  • Step 1: $2Cu + O_{2} \xrightarrow{Heat} 2CuO$ (Copper is oxidised to black copper(II) oxide).
  • Step 2: $CuO + H_{2} \xrightarrow{Heat} Cu + H_{2}O$ (Copper(II) oxide is reduced to copper; hydrogen is oxidised to water).
• Other Examples:
  • $ZnO + C \rightarrow Zn + CO$ ($C$ is oxidised, $ZnO$ is reduced).
  • $MnO_{2} + 4HCl \rightarrow MnCl_{2} + 2H_{2}O + Cl_{2}$ ($HCl$ is oxidised to $Cl_{2}$, $MnO_{2}$ is reduced).

Effects of Oxidation in Everyday Life

• Corrosion: The process where a metal is attacked by substances in its environment such as moisture, acids, etc.
  • Iron: Rusting results in a reddish-brown powder coating. This causes significant structural damage to bridges, ships, and cars.
  • Silver: Formation of a black coating.
  • Copper: Formation of a green coating.
• Rancidity: The oxidation of fats and oils in food materials, leading to changes in smell and taste.
  • Prevention: Adding antioxidants to food, using airtight containers to slow down oxidation, or flushing food bags (like potato chips) with nitrogen gas to prevent oxidation.

Questions & Discussion

• Pre-burning cleaning: Why clean magnesium ribbon with sandpaper? (To remove the layer of magnesium oxide that might interfere with burning).
• Respiration as Exothermic: During digestion, carbohydrates break down into glucose, which reacts with oxygen in cells to provide energy.
• Whitewashing Substance 'X': Substance 'X' is calcium oxide ($CaO$ or Quick lime). Reaction: $CaO(s) + H_{2}O(l) \rightarrow Ca(OH)_{2}(aq)$.
• Gas Volume in Water Electrolysis: The amount of gas collected in one tube is double the other because water ($H_{2}O$) contains two parts hydrogen for every one part oxygen. The gas in larger volume is hydrogen.
• Identification of Redox Components:
  • In $4Na(s) + O_{2}(g) \rightarrow 2Na_{2}O(s)$, sodium is oxidised.
  • In $CuO(s) + H_{2}(g) \rightarrow Cu(s) + H_{2}O(l)$, copper oxide is reduced and hydrogen is oxidised.
• Group Activity - Heat changes:
  • Beaker A: Potassium sulphate ($K_{2}SO_{4}$).
  • Beaker B: Ammonium nitrate ($NH_{4}NO_{3}$).
  • Beaker C: Anhydrous copper sulphate ($CuSO_{4}$).
  • Beaker D: Iron filings added to copper sulphate solution.
  • Goal: Measure temperature changes to classify as endothermic or exothermic.