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Chemical Reactions and Equations
Introduction:
- Chemical reactions occur in daily life situations, altering the nature of substances.
- These changes indicate chemical reactions, which we’ll explore further.
Examples of Everyday Chemical Reactions
- Milk left at room temperature during summers:
- Spoils due to the growth of bacteria, changing its nature.
- Iron tawa/pan/nail exposed to humid atmosphere:
- Develops rust, altering its composition.
- Grapes getting fermented:
- Turns into wine, indicating a chemical change.
- Food cooked:
- Ingredients transform, releasing aroma and changing taste, a result of a chemical reaction.
- Food digested in our body:
- Nutrients broken down into simpler substances during digestion, facilitating absorption.
- Respiration:
- Oxygen is used to convert food into energy, a vital chemical process in our body.
Features of Chemical Reactions
- Change in state, color, or temperature.
- Evolution of gas indicates chemical changes.
- Various types of chemical reactions occur, each with distinct characteristics.
By observing these changes, we can understand the diverse nature of chemical reactions happening around us.
Physical Change: Change in physical properties or Change in State or Appearance of substance like
– a change in state (solid, liquid, gas) or appearance (size, shape, texture) of the substance.
– Reversible as no new substance is formed.
Activities to understand Chemical Reactions
1. Burning Magnesium Ribbon (Activity 1.1):
Magnesium burns with a dazzling flame, forming magnesium oxide.
2. Mixing Lead Nitrate and Potassium Iodide (Activity 1.2):
Results in a color change, indicating a chemical reaction.
3. Reacting Zinc with Hydrochloric Acid (Activity 1.3):
Generates bubbles and a change in temperature, signs of a chemical reaction.
Chemical Equations
A chemical equation is a symbolic representation of a chemical reaction, showing the reactants transforming into products.
Chemical Equation Features:
- Word-Equations:
- Describe chemical reactions using words.
- Example: Magnesium + Oxygen → Magnesium oxide.
- Reactants and Products:
- Reactants are substances undergoing change.
- Products are new substances formed.
- Represented on the left and right sides of the equation, respectively.
- Chemical Formulae:
- Chemical equations can be written using chemical formulae.
- Example: Mg + O2 → MgO.
- Balanced Chemical Equations:
- Total mass of reactants equals total mass of products.
- Number of atoms of each element is the same on both sides.
- Achieved through balancing coefficients.
- Balancing Steps:
- Step I: Draw boxes around formulas.
Step II: List atoms on both sides.
Step III: Balance compound with maximum atoms.
Step IV: Balance remaining elements.
Step V: Check correctness of equation.
- we count atoms of each element on both sides of the equation.
Step VI: Mentioning of Physical States of reactants and products solid(s), gas(g), liquid(l), aqueous(aq).
Example: 3Fe(s) + 4H2O(g) → Fe3O4(s) + 4H2(g).
Step VII: Mentioning of Reaction Conditions like temperature, pressure, catalyst, etc. Sometimes indicated above or below the arrow.
By following these steps, chemical equations can be effectively balanced, representing chemical reactions accurately.
Skeletal Chemical Equation
– A simplified representation of a chemical reaction using chemical formulas to indicate reactants and products, without specifying the exact quantities of each substance. example: Fe + H2O → Fe3O4 + H2 “tap here to learn about balanced chemical equation”.
Balanced Chemical Equation
– A representation of a chemical reaction where the number of atoms of each element is the same on both sides, indicating the conservation of mass. example: 3Fe + 4H2O → Fe3O4 + 4H2
Types of Chemical Reactions
Some plants lack chlorophyll and cannot make their own food. They depend on other plants for nutrition, using the heterotrophic mode of nutrition.
1. Combination Reaction
- Definition: A combination reaction is when two or more substances combine to form a single product.
Example Activity:
- Take calcium oxide (quick lime) in a beaker.
- Add water slowly.
- The beaker gets warm due to the reaction.
- Reaction: CaO(s) + H2O(l) → Ca(OH)2(aq) + Heat
- Other Examples:
- Burning of Coal: C(s) + O2(g) → CO2(g)
- Formation of Water: 2H2(g) + O2(g) → 2H2O(l)
- Exothermic Reactions:
- Definition: Reactions that release heat.
- Examples:
- Burning Natural Gas: CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)
- Respiration: C6H12O6(aq) + 6O2(aq) → 6CO2(aq) + 6H2O(l) + energy
- We get energy from food.
- Food breaks down into simpler substances like glucose.
- Glucose combines with oxygen to provide energy.
- Decomposition of Vegetable Matter: Turns into compost and releases heat.
Do You Know?
Whitewashing Walls:
Reaction: Ca(OH)2(aq) + CO2(g) → CaCO3(s) + H2O(l)
– Produces calcium carbonate, giving a shiny finish to walls.
– The chemical formula for marble is also CaCO3.
Decomposition Reaction
Activity 1.5: Heating Ferrous Sulphate
- Steps:
- Take 2 g of ferrous sulphate crystals in a dry boiling tube.
- Note the green color of the crystals.
- Heat the boiling tube over a flame.
- Observe the color change and smell.
- Observation:
- Green crystals turn a different color.
- Smell of burning sulfur.
- Reaction: 2FeSO4(s) → Fe2O3(s) + SO2(g) + SO3(g)
- Explanation:
- A single reactant (ferrous sulphate) breaks down into simpler products (ferric oxide, sulfur dioxide, sulfur trioxide).
- This is a decomposition reaction.
Example: Decomposition of Calcium Carbonate
- Reaction: CaCO3(s) → CaO(s) + CO2(g)
- Explanation:
- Heating calcium carbonate produces calcium oxide (quick lime) and carbon dioxide.
- This process is called thermal decomposition.
Activity 1.6: Heating Lead Nitrate
- Steps:
- Take 2 g of lead nitrate powder in a boiling tube.
- Heat the tube over a flame.
- Observe the changes.
- Observation:
- Emission of brown fumes.
- Reaction: 2Pb(NO3)2(s) → 2PbO(s) + 4NO2(g) + O2(g)
Activity 1.7: Electrolysis of Water
- Steps:
- Drill two holes in a plastic mug and insert carbon electrodes.
- Connect electrodes to a 6-volt battery.
- Fill the mug with water and add dilute sulfuric acid.
- Invert test tubes filled with water over the electrodes.
- Switch on the current and observe bubbles.
- Observation:
- Bubbles form at both electrodes, displacing water in the test tubes.
- Test gases with a burning candle to identify them.
- Reaction: 2H2O(l) → 2H2(g) + O2(g)
Activity 1.8: Decomposition of Silver Chloride
- Steps:
- Take 2 g of silver chloride in a china dish.
- Observe its white color.
- Place the dish in sunlight.
- Observe the color change.
- Observation:
- White silver chloride turns grey.
- Reaction: 2AgCl(s) → 2Ag(s) + Cl2(g)
- Similar Reaction: 2AgBr(s) → 2Ag(s) + Br2(g)
- Explanation:
- Light causes silver chloride to decompose into silver and chlorine.
- Used in black and white photography.
Endothermic Reactions
- Definition: Reactions that absorb energy.
Activity:
- Mix 2 g of barium hydroxide with 1 g of ammonium chloride.
- Touch the bottom of the test tube to feel the temperature change.
- Observation: If the tube feels cold, it’s an endothermic reaction.
Double Displacement Reaction
Double Displacement Reactions
Activity 1.10: Mixing Sodium Sulphate and Barium Chloride
Steps:
- Take about 3 mL of sodium sulphate solution in a test tube.
- In another test tube, take about 3 mL of barium chloride solution.
- Mix the two solutions.
Observation:
- A white substance forms, which does not dissolve in water.
- This white substance is called a precipitate.
Explanation:
- The white precipitate formed is barium sulphate (BaSO4).
- The reaction can be written as:
- Na2SO4(aq) + BaCl2(aq) → BaSO4(s) + 2NaCl(aq)
- Barium sulphate (BaSO4) is the white precipitate.
- Sodium chloride (NaCl) stays dissolved in the solution.
Definition:
- Precipitation Reaction: A reaction that produces an insoluble substance (precipitate).
- Double Displacement Reaction: A reaction where ions are exchanged between the reactants.
Oxidation and Reduction
Activity 1.11: Heating Copper Powder
Steps:
- Heat a china dish containing about 1 gram of copper powder.
- Observe what happens.
Observation:
- The surface of the copper powder turns black.
- The black substance is copper(II) oxide (CuO), formed because oxygen is added to copper.
Reaction:
- 2Cu + O2 → 2CuO
Reversing the Reaction:
- If hydrogen gas is passed over the heated copper(II) oxide, the black coating turns brown as copper is obtained again.
- CuO + H2 → Cu + H2O
Concepts:
- Oxidation: When a substance gains oxygen, it is said to be oxidized.
- Reduction: When a substance loses oxygen, it is said to be reduced.
Redox Reaction:
- CuO + H2 → Cu + H2O
- Copper(II) oxide (CuO) is reduced to copper (Cu) by losing oxygen.
- Hydrogen (H2) is oxidized to water (H2O) by gaining oxygen.
Other Examples of Redox Reactions:
- ZnO + C → Zn + CO
- Carbon (C) is oxidized to carbon monoxide (CO).
- Zinc oxide (ZnO) is reduced to zinc (Zn).
- MnO2 + 4HCl → MnCl2 + 2H2O + Cl2
- Hydrogen chloride (HCl) is oxidized to chlorine (Cl2).
- Manganese dioxide (MnO2) is reduced to manganese chloride (MnCl2).
Key Points:
- Oxidation: Gain of oxygen or loss of hydrogen.
- Reduction: Loss of oxygen or gain of hydrogen.
Example Recall:
- In Activity 1.1, magnesium burns in oxygen and forms magnesium oxide. Magnesium is oxidized in this reaction.
Effects of oxidation reaction in daily life
Corrosion and Rancidity
1. Corrosion
- Observation: Iron articles are shiny when new but develop a reddish-brown coating over time, known as rust.
- Definition: Corrosion occurs when metals are attacked by substances like moisture and acids, leading to damage.
- Examples:
- Iron: Rusting, forming reddish-brown powder.
- Silver: Forms a black coating.
- Copper: Develops a green coating.
- Impact: Corrosion damages metal objects like car bodies, bridges, iron railings, and ships, especially those made of iron.
- Problem: Significant amounts of money are spent annually to replace damaged iron.
2. Rancidity
- Observation: Fats and oils in food become rancid, changing taste and smell when left for a long time.
- Cause: Oxidation of fats and oils.
- Prevention:
- Antioxidants: Added to foods to prevent oxidation.
- Airtight Containers: Slows down oxidation.
- Nitrogen Flushing: Chips manufacturers use nitrogen gas in bags to prevent oxidation.
Chapter Summary:
- A complete chemical equation represents the reactants, products, and their physical states symbolically.
- A chemical equation is balanced to have the same number of each type of atom on both sides.
- Equations must always be balanced.
- In a combination reaction, two or more substances combine to form a single new substance.
- Decomposition reactions are the opposite of combination reactions. One substance breaks down into two or more substances.
- Reactions that release heat along with the products are called exothermic reactions.
- Reactions that absorb energy are known as endothermic reactions.
- A displacement reaction occurs when one element displaces another element from its compound.
- In double displacement reactions, two different atoms or groups of atoms (ions) are exchanged.
- Precipitation reactions produce insoluble salts.
- Reactions can involve the gain or loss of oxygen or hydrogen.
- Oxidation is the gain of oxygen or loss of hydrogen.
- Reduction is the loss of oxygen or gain of hydrogen.
