Chemical Reactions and Equations Notes

Chemical Reactions and Equations

Consider daily life situations where changes occur:

  • Milk left at room temperature during summer changes.

  • Iron exposed to humid atmosphere rusts.

  • Grapes ferment.

  • Food is cooked and digested.

  • Respiration occurs.

In these situations, the initial substance's nature and identity change, indicating a chemical reaction.

Activity 1.1: Burning Magnesium Ribbon

  • Clean a 3-4 cm long magnesium ribbon with sandpaper.

  • Hold it with tongs and burn it using a spirit lamp or burner.

  • Collect the ash (magnesium oxide) in a watch-glass.

  • Observation: Magnesium ribbon burns with a dazzling white flame and transforms into a white powder (magnesium oxide).

  • The reaction is 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 sulfuric acid.

  • Observe what happens around the zinc granules (evolution of gas).

  • Touch the flask or test tube to feel any temperature change.

Activity 1.3: Reaction of Lead Nitrate with Potassium Iodide

  • Take lead nitrate solution in a test tube.

  • Add potassium iodide solution.

  • Observe the reaction (formation of a precipitate).

Observations Indicating Chemical Reactions

From the above activities, these observations indicate a chemical reaction:

  • Change in state

  • Change in color

  • Evolution of a gas

  • Change in temperature

Chemical Equations

Word Equations

Activity 1.1 can be described as: Magnesium + Oxygen → Magnesium oxide

  • Reactants: Substances undergoing the chemical change (magnesium and oxygen).

  • Product: New substance formed (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 points towards the products, indicating the direction of the reaction.

Writing Chemical Equations

Using chemical formulas makes equations more concise. For example:

Mg + O_2 → MgO

  • If the number of atoms of each element is the same on both sides, the equation is balanced.

  • If not, the equation is unbalanced (a skeletal chemical equation).

Balanced Chemical Equations

Law of conservation of mass: Mass is neither created nor destroyed in a chemical reaction. 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. Therefore, we balance skeletal equations.

Example:

Zn + H2SO4 → ZnSO4 + H2

This equation is balanced because the number of atoms of each element is the same on both sides.

Balancing Chemical Equations: Step-by-Step

Consider: Fe + H2O → Fe3O4 + H2

Step I: Draw boxes around each formula (do not change anything inside the boxes).

Fe + H2O → Fe3O4 + H2

Step II: List the number of atoms of each element on both sides.

Element

Reactants (LHS)

Products (RHS)

Fe

1

3

H

2

2

O

1

4

Step III: Start with the compound containing the maximum number of atoms (e.g., Fe3O4) and select the element with the maximum number of atoms (oxygen).

Step IV: Balance oxygen atoms by adding a coefficient: Fe + 4H2O → Fe3O4 + H2

Step V: Balance hydrogen atoms: Fe + 4H2O → Fe3O4 + 4H2

Step VI: Balance iron atoms: 3Fe + 4H2O → Fe3O4 + 4H2

Step VII: Check the correctness of the balanced equation.

This method is called the hit-and-trial method.

Writing Symbols of Physical States

Indicate the physical states of reactants and products using (g) for gas, (l) for liquid, (aq) for aqueous solution, and (s) for solid. For example:

3Fe(s) + 4H2O(g) → Fe3O4(s) + 4H2(g)

Reaction conditions (temperature, pressure, catalyst) can be indicated above and/or below the arrow:

CO(g) + 2H2(g) \xrightarrow{340 atm} CH3OH(l)

Photosynthesis:

6CO2(aq) + 12H2O(l) \xrightarrow[Chlorophyll]{Sunlight} C6H{12}O6(aq) + 6O2(aq) + 6H_2O(l)

Types of Chemical Reactions

During a chemical reaction, atoms of one element do not change into those of another. Chemical reactions involve breaking and making bonds between atoms.

Combination Reaction

Activity 1.4: Reaction of Calcium Oxide with Water

  • Take calcium oxide (quick lime) in a beaker.

  • Slowly add water.

  • Touch the beaker; observe temperature change.

CaO(s) + H2O(l) → Ca(OH)2(aq) + Heat

Calcium oxide reacts vigorously with water to produce slaked lime (calcium hydroxide), releasing heat. A combination reaction is when a single product is formed from two or more reactants.

Other Examples:

  • Burning of coal: C(s) + O2(g) → CO2(g)

  • Formation of water: 2H2(g) + O2(g) → 2H_2O(l)

Reactions in which heat is released are called exothermic reactions.

Examples of Exothermic Reactions:

  • Burning of natural gas: CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)

  • Respiration: C6H{12}O6(aq) + 6O2(aq) → 6CO2(aq) + 6H2O(l) + energy

  • Decomposition of vegetable matter into compost.

Whitewashing:

Slaked lime (calcium hydroxide) reacts with carbon dioxide in air to form calcium carbonate:

Ca(OH)2(aq) + CO2(g) → CaCO3(s) + H2O(l)

The chemical formula for marble is also CaCO_3.

Decomposition Reaction

Activity 1.5: Heating Ferrous Sulphate Crystals

  • Take ferrous sulphate crystals in a dry boiling tube.

  • Note the color of the crystals.

  • Heat the boiling tube.

  • Observe the color change and smell the odor.

2FeSO4(s) \xrightarrow{Heat} Fe2O3(s) + SO2(g) + SO_3(g)

Ferrous sulphate crystals lose water when heated and decompose into ferric oxide, sulphur dioxide, and sulphur trioxide. A single reactant breaks down to give simpler products.

Thermal Decomposition: Decomposition by heating.

CaCO3(s) \xrightarrow{Heat} CaO(s) + CO2(g)

Calcium oxide (lime or quick lime) is used in the manufacture of cement.

Activity 1.6: Heating Lead Nitrate Powder

  • Take lead nitrate powder in a boiling tube.

  • Heat it over a flame.

  • Observe the emission of brown fumes (nitrogen dioxide).

2Pb(NO3)2(s) \xrightarrow{Heat} 2PbO(s) + 4NO2(g) + O2(g)

Activity 1.7: Silver Chloride in Sunlight

  • Take silver chloride in a china dish; note the color.

  • Place the dish in sunlight.

  • Observe the color change.

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

White silver chloride turns grey due to decomposition into silver and chlorine.

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

This reaction is used in black and white photography.

Activity 1.8: Electrolysis of Water

  • Set up electrolysis apparatus with carbon electrodes.

  • Add a few drops of dilute sulfuric acid to water.

  • Invert test tubes over the electrodes.

  • Switch on the current.

  • Observe bubble formation.

  • Test the gases with a burning candle.

Decomposition reactions require energy in the form of heat, light, or electricity. Reactions in which energy is absorbed are known as endothermic reactions.

Displacement Reaction

Activity 1.9: Iron Nails in Copper Sulphate Solution

  • Take three iron nails and clean them.

  • Take two test tubes with copper sulphate solution.

  • Immerse two iron nails in one test tube for 20 minutes.

  • Compare the color of the solutions and the nails.

Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s)

Iron displaces copper from copper sulphate solution. A displacement reaction involves one element displacing another from a compound.

Other Examples:

Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s)

Pb(s) + CuCl2(aq) → PbCl2(aq) + Cu(s)

Zinc and lead are more reactive than copper.

Double Displacement Reaction

Activity 1.10: Sodium Sulphate and Barium Chloride

  • Take sodium sulphate solution in a test tube.

  • Take barium chloride solution in another test tube.

  • Mix the two solutions.

  • Observe the formation of a white precipitate.

Na2SO4(aq) + BaCl2(aq) → BaSO4(s) + 2NaCl(aq)

A white precipitate of BaSO_4 is formed. A precipitation reaction produces an insoluble salt. Double displacement reactions involve the exchange of ions between reactants.

Oxidation and Reduction

Activity 1.11: Heating Copper Powder

  • Heat copper powder in a china dish.

  • Observe the formation of black copper(II) oxide.

2Cu + O_2 \xrightarrow{Heat} 2CuO

If hydrogen gas is passed over the heated material, the black coating turns brown, and copper is obtained.

CuO + H2 \xrightarrow{Heat} Cu + H2O

Oxidation: A substance gains oxygen.

Reduction: A substance loses oxygen.

In the above reaction, copper(II) oxide is reduced, and hydrogen is oxidized. Such reactions are called oxidation-reduction or redox reactions.

Other Redox Examples:

ZnO + C → Zn + CO

MnO2 + 4HCl → MnCl2 + 2H2O + Cl2

Oxidation: Gain of oxygen or loss of hydrogen.

Reduction: Loss of oxygen or gain of hydrogen.

Effects of Oxidation Reactions in Everyday Life

Corrosion

Iron articles get coated with reddish-brown powder (rust). Other metals also tarnish. Corrosion is when a metal is attacked by substances around it (moisture, acids, etc.). Examples include black coating on silver and green coating on copper. Corrosion damages car bodies, bridges, and iron railings.

Rancidity

Fats and oils in food become rancid when oxidized, changing their smell and taste. Antioxidants are added to foods to prevent oxidation. Keeping food in airtight containers slows down oxidation. Chips manufacturers flush bags with nitrogen gas to prevent oxidation.