Chemical Reactions and Equations Class 10 Chemistry Chapter 1 Notes

Core Definitions & Glossary

Chemical Reaction

A process in which one or more substances (reactants) react to form new substances (products) with entirely different chemical properties.

Chemical Equation

A symbolic representation of a chemical reaction using chemical formulas, symbols, and physical state indicators (s, l, g, aq).

Balanced Chemical Equation

An equation where the total number of atoms of each element is equal on both the reactant side and the product side.

Exothermic Reaction

A chemical reaction accompanied by the evolution or release of heat into the surroundings (e.g., respiration, combustion).

Endothermic Reaction

A chemical reaction that absorbs energy from its surroundings in the form of heat, light, or electricity (e.g., photosynthesis, decomposition reactions).

Precipitate

An insoluble solid substance formed when two clear aqueous solutions react chemically.

Redox Reaction

A reaction in which oxidation (gain of oxygen/loss of electrons) and reduction (loss of oxygen/gain of electrons) take place simultaneously.

Rancidity

The oxidation of unsaturated fats and oils present in food items, resulting in an unpleasant smell and taste.

Concept Focus: Law of Conservation of Mass & Equation Balancing

According to the Law of Conservation of Mass, matter can neither be created nor destroyed in a chemical reaction. Therefore, the total mass of the reactants must equal the total mass of the products. In terms of atoms, this means that the absolute number of atoms of each element before the reaction (on the left-hand side, reactant side) must equal the number of atoms of those elements after the reaction (on the right-hand side, product side).

To balance an equation, we use the Hit-and-Trial method. The most critical rule in balancing is that you must never alter the subscripts in a chemical formula (e.g., changing H₂O to H₂O₂ to balance oxygen is chemically incorrect). Instead, you add coefficients in front of the formulas. Once balanced, always write the physical states in parentheses: (s) for solid, (l) for liquid, (g) for gas, and (aq) for aqueous (dissolved in water) to make the equation informative.

Steps to Balance a Chemical Equation

1. — Convert the word statement into a basic chemical equation. Write the correct formulas for reactants on the left and products on the right.
2. — Draw boxes around each chemical formula to remind yourself not to change anything inside the boxes. Count and note down the number of atoms of each element present on both sides.
3. — Start balancing with the compound or element that has the maximum number of atoms (often oxygen or hydrogen), regardless of whether it is a reactant or product.
4. — Multiply the chemical formulas by whole-number coefficients to equalize the atom counts on both sides.
5. — Progressively balance the rest of the atoms one by one. Check the counts of all elements again to ensure they are equal on both sides.
6. — Add indicators like (s), (l), (g), or (aq). Write conditions like temperature, pressure, or catalysts above or below the reaction arrow.

Comparison of Chemical Reaction Types

Must-Remember Key Revision Points

• Five indicators that a chemical reaction has taken place: change in state, change in color, evolution of gas, change in temperature, or formation of a precipitate.
• Magnesium ribbon is cleaned with sandpaper before burning to remove the protective layer of basic magnesium carbonate (MgCO₃) from its surface.
• Decomposition reactions are opposite to combination reactions and are classified into Thermal (using heat), Electrolytic (using electricity), and Photolytic (using sunlight).
• Photolytic decomposition of Silver Chloride (AgCl) and Silver Bromide (AgBr) turns them grey in sunlight; this reaction is used in black-and-white photography.
• The reactivity series must be memorized to identify single displacement reactions. If the standalone element is less reactive than the metal in the compound, no reaction occurs.
• In a Redox reaction, the substance that gains oxygen (or loses hydrogen) is Oxidized. The substance that loses oxygen (or gains hydrogen) is Reduced.
• An Oxidizing Agent gets reduced in a reaction, while a Reducing Agent gets oxidized.
• Corrosion is the damage caused to metals by moisture, acids, and air (e.g., rusting of iron, black coating on silver, green coating on copper).
• Rancidity of oily food is prevented by adding antioxidants (like BHA), flushing packages with Nitrogen gas, or keeping food in airtight containers.

Worked Mini-Examples

• {"title":"Example 1: Balancing a Complex Equation","description":"Balance the skeletal equation: Fe (s) + H₂O (g) → Fe₃O₄ (s) + H₂ (g)\n\n1. Count atoms: Fe (1 vs 3), H (2 vs 2), O (1 vs 4).\n2. Balance 'O' by putting coefficient 4 in front of H₂O: Fe + 4H₂O → Fe₃O₄ + H₂. Now hydrogen atoms on left = 8.\n3. Balance 'H' by placing coefficient 4 in front of H₂ on the right: Fe + 4H₂O → Fe₃O₄ + 4H₂.\n4. Balance 'Fe' by putting coefficient 3 on the left: 3Fe (s) + 4H₂O (g) → Fe₃O₄ (s) + 4H₂ (g).\n5. Verify counts: Fe (3 on both sides), H (8 on both sides), O (4 on both sides). The equation is balanced."}
• {"title":"Example 2: Identifying Redox Components","description":"Identify the substance oxidized, substance reduced, oxidizing agent, and reducing agent in: \nCuO (s) + H₂ (g) --[Heat]--> Cu (s) + H₂O (l)\n\n Substance Oxidized: H₂ (gained Oxygen to become H₂O)\n Substance Reduced: CuO (lost Oxygen to become Cu)\n Oxidizing Agent: CuO (provided the Oxygen for oxidation)\n Reducing Agent: H₂ (removed the Oxygen from CuO)"}

CBSE Board Exam Traps & Cues

1. Watch the Color Changes: Examiners love questions based on observations. Remember: Copper sulphate solution is blue; after iron nails are dipped, it becomes light green due to FeSO₄ formation. Lead nitrate and potassium iodide reaction yields a yellow precipitate of Lead Iodide (PbI₂).
2. The Gas Tests: If a gas is evolved, name the test! CO₂ turns lime water milky. H₂ burns with a characteristic 'pop' sound. SO₂ has a suffocating, burning sulfur smell and turns acidified potassium dichromate paper green.
3. Thermal Decomposition of Lead Nitrate: Note down this reaction as it is asked frequently:

2Pb(NO₃)₂ (s) --[Heat]--> 2PbO (s) [Yellow] + 4NO₂ (g) [Brown fumes] + O₂ (g). Memorizing the brown fumes of Nitrogen Dioxide is a guaranteed marks saver!

Quick Revision Check

• Why is respiration considered an exothermic reaction? During respiration, glucose obtained from food undergoes slow combustion by reacting with oxygen inside our cells, releasing energy (ATP) and heat. Hence, it is classified as an exothermic process.
• What happens chemically when quicklime (Calcium Oxide) is added to water? Calcium Oxide reacts vigorously with water to form slaked lime (Calcium Hydroxide) while releasing a large amount of heat. Equation: CaO (s) + H₂O (l) → Ca(OH)₂ (aq) + Heat.
• Why does the color of copper sulfate solution change when an iron nail is kept in it? Iron is more reactive than copper (according to the reactivity series). It displaces copper from copper sulfate solution, forming iron sulfate which is light green, while the blue color of copper sulfate fades away.
• Write the balanced chemical equation for the electrolysis of water. What is the ratio of gases collected at the anode and cathode? Equation: 2H₂O (l) --[Electricity]--> 2H₂ (g) + O₂ (g). The gas collected at the cathode is Hydrogen and at the anode is Oxygen. The volume ratio of Hydrogen to Oxygen collected is 2:1.

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