Revision Guide
Explore the fundamentals of chemical reactions, types, and balancing equations in this chapter, essential for understanding chemistry basics.
Chemical Reactions and Equations - Quick Look Revision Guide
Your 1-page summary of the most exam-relevant takeaways from Science.
This compact guide covers 20 must-know concepts from Chemical Reactions and Equations aligned with Class X preparation for Science. Ideal for last-minute revision or daily review.
Complete study summary
Essential formulas, key terms, and important concepts for quick reference and revision.
Key Points
Define chemical reaction with an example.
A chemical reaction involves the transformation of reactants into products through chemical changes. Example: Burning of magnesium ribbon in air to form magnesium oxide.
State the law of conservation of mass.
Mass can neither be created nor destroyed in a chemical reaction. The total mass of reactants equals the total mass of products.
Explain balanced chemical equations.
Balanced equations have equal numbers of each type of atom on both sides. Example: 2H2 + O2 → 2H2O.
Describe combination reaction.
Two or more substances combine to form a single product. Example: CaO + H2O → Ca(OH)2.
Explain decomposition reaction.
A single compound breaks down into two or more simpler substances. Example: 2FeSO4 → Fe2O3 + SO2 + SO3.
Define exothermic reaction.
Reactions that release heat energy. Example: Respiration releases energy used by cells.
Define endothermic reaction.
Reactions that absorb heat energy. Example: Decomposition of calcium carbonate requires heat.
Explain displacement reaction.
A more reactive element displaces a less reactive one from its compound. Example: Fe + CuSO4 → FeSO4 + Cu.
Describe double displacement reaction.
Exchange of ions between two reactants to form new compounds. Example: Na2SO4 + BaCl2 → BaSO4 + 2NaCl.
Define redox reaction.
Reactions involving oxidation and reduction. Example: CuO + H2 → Cu + H2O.
Explain oxidation.
Gain of oxygen or loss of hydrogen. Example: 2Cu + O2 → 2CuO.
Explain reduction.
Loss of oxygen or gain of hydrogen. Example: CuO + H2 → Cu + H2O.
Define corrosion.
Deterioration of metals due to reaction with environment. Example: Rusting of iron.
Explain rancidity.
Oxidation of fats/oils leading to bad smell/taste. Example: Spoilage of butter.
State the importance of balancing equations.
Ensures the law of conservation of mass is followed and provides correct stoichiometry.
Describe thermal decomposition.
Decomposition caused by heating. Example: 2Pb(NO3)2 → 2PbO + 4NO2 + O2.
Explain electrolysis of water.
Decomposition of water into hydrogen and oxygen using electricity. Example: 2H2O → 2H2 + O2.
Define precipitation reaction.
Formation of an insoluble solid (precipitate) during a reaction. Example: BaCl2 + Na2SO4 → BaSO4 + 2NaCl.
Explain the effect of light on silver chloride.
Silver chloride decomposes in sunlight to form silver and chlorine. Example: 2AgCl → 2Ag + Cl2.
Describe the reaction of zinc with sulphuric acid.
Zinc reacts with dilute sulphuric acid to form zinc sulphate and hydrogen gas. Example: Zn + H2SO4 → ZnSO4 + H2.
Explore the properties, reactions, and uses of acids, bases, and salts in everyday life and their importance in chemistry.
Explore the properties, reactions, and uses of metals and non-metals, understanding their role in daily life and industrial applications.
Explore the versatile world of carbon, its allotropes, and the vast array of compounds it forms, including hydrocarbons and their derivatives, in this comprehensive chapter.
Life Processes explores the essential functions that sustain living organisms, including nutrition, respiration, transportation, and excretion.
