Knowledge That Ignites
Original, Copyright-Free Notes aligned with Maharashtra Board syllabus
Complete conceptual coverage with examples and practice
A chemical reaction is a process in which substances undergo chemical change to form new substances with different properties. During a chemical reaction, old chemical bonds break and new chemical bonds form, resulting in the transformation of reactants into products.
Chemical reactions are fundamental processes that occur continuously in nature and in human-made systems. They are responsible for countless phenomena in our daily lives, from the food we digest to the fuels we burn.
A chemical reaction involves the transformation of reactants (starting materials) into products (newly formed substances). This transformation occurs through the breaking and forming of chemical bonds.
A solid may form from liquids, a gas may be produced, etc.
Example: Burning of candle wax (solid → liquid → gas)
Reactants and products have different colors.
Example: Rusting of iron (silver → reddish-brown)
Formation of bubbles or fumes indicates gas production.
Example: Vinegar + baking soda → carbon dioxide gas
An insoluble solid forms when solutions are mixed.
Example: Lead nitrate + potassium iodide → yellow precipitate
Heat is either released (exothermic) or absorbed (endothermic).
Example: Burning of fuel (releases heat), photosynthesis (absorbs heat)
New substances often have different odors than reactants.
Example: Spoilage of food produces foul smell
A chemical equation is a symbolic representation of a chemical reaction using chemical formulas, symbols, and coefficients. It shows the reactants on the left side and products on the right side, separated by an arrow.
A word equation uses the names of reactants and products to describe a chemical reaction.
This tells us that magnesium reacts with oxygen to form magnesium oxide.
A skeletal equation uses chemical formulas but is not balanced according to the law of conservation of mass.
This equation is unbalanced because oxygen atoms don't match on both sides.
A balanced chemical equation has an equal number of atoms of each element on both sides of the equation, satisfying the law of conservation of mass.
Now balanced: 2 Mg atoms and 2 O atoms on both sides.
This fundamental law states that mass is neither created nor destroyed in a chemical reaction. The total mass of reactants equals the total mass of products.
Implication for equations: Number of atoms of each element must be equal on both sides of the equation.
Balancing chemical equations ensures they obey the law of conservation of mass. Follow these systematic steps:
Write the correct chemical formulas for all reactants and products. Never change subscripts in formulas to balance equations.
Count the number of atoms of each element on both sides of the equation.
Start with metals, then non-metals, and balance hydrogen and oxygen last. Use coefficients (numbers in front of formulas).
Only change coefficients (numbers in front of formulas). Never change subscripts (numbers within formulas).
Recount all atoms to verify balance. Simplify coefficients to smallest whole numbers if possible.
Unbalanced equation: Fe + H₂O → Fe₃O₄ + H₂
Count atoms on both sides:
Balance iron atoms: Put 3 before Fe on left
Balance oxygen atoms: Put 4 before H₂O on left
Balance hydrogen atoms: Put 4 before H₂ on right
Recheck: Left: Fe=3, H=8, O=4 | Right: Fe=3, H=8, O=4 ✓
Balanced equation:
Including physical state symbols in chemical equations provides more information about the reaction conditions and the physical forms of reactants and products.
| State Symbol | Meaning | Example |
|---|---|---|
| (s) | Solid state | Iron(s), Sodium chloride(s) |
| (l) | Liquid state (pure liquid) | Water(l), Mercury(l) |
| (g) | Gaseous state | Oxygen(g), Hydrogen(g) |
| (aq) | Aqueous (dissolved in water) | Sodium chloride(aq), Hydrochloric acid(aq) |
This tells us: Solid zinc reacts with aqueous hydrochloric acid to produce aqueous zinc chloride and hydrogen gas.
Chemical reactions are classified into different types based on how reactants combine or break apart. Understanding these categories helps predict products and understand reaction mechanisms.
Two or more substances combine to form a single product. Also called synthesis reaction.
Explanation: Quicklime (calcium oxide) combines with water to form slaked lime (calcium hydroxide). This is an exothermic reaction (releases heat).
Explanation: Magnesium combines with oxygen to form magnesium oxide (burning of magnesium ribbon).
A single compound breaks down into two or more simpler substances. These reactions usually require energy input.
| Type | Energy Source | Example | Equation |
|---|---|---|---|
| Thermal Decomposition | Heat | Heating calcium carbonate | CaCO₃(s) → CaO(s) + CO₂(g) |
| Electrolytic Decomposition | Electricity | Electrolysis of water | 2H₂O(l) → 2H₂(g) + O₂(g) |
| Photochemical Decomposition | Light | Decomposition of silver chloride | 2AgCl(s) → 2Ag(s) + Cl₂(g) |
A more reactive element displaces a less reactive element from its compound. These reactions follow the reactivity series of metals.
Explanation: Zinc is more reactive than copper, so it displaces copper from copper sulfate solution. Blue color of copper sulfate fades as colorless zinc sulfate forms, and brown copper deposits on zinc.
Metals arranged in order of decreasing reactivity:
Rule: A metal can displace any metal below it in the series from its compound.
Ions of two compounds exchange places to form two new compounds. Often results in precipitate formation or gas evolution.
Forms an insoluble solid (precipitate) when solutions are mixed.
White precipitate of barium sulfate forms.
Acid reacts with base to form salt and water.
Hydrochloric acid neutralizes sodium hydroxide.
Explanation: Sodium carbonate reacts with hydrochloric acid to produce sodium chloride, water, and carbon dioxide gas (effervescence observed).
Redox reactions involve both oxidation and reduction occurring simultaneously. One substance gets oxidized while another gets reduced.
Memory aid: OIL - Oxidation Is Loss (of electrons)
Memory aid: RIG - Reduction Is Gain (of electrons)
Oxidation and reduction always occur together in the same reaction. If one substance is oxidized, another must be reduced.
Step-by-step analysis:
Oxidizing agent: Cu²⁺ (causes oxidation of Zn)
Reducing agent: Zn (causes reduction of Cu²⁺)
| Concept | Definition | Example |
|---|---|---|
| Oxidizing Agent | Substance that causes oxidation (gets reduced itself) | Oxygen, chlorine, potassium permanganate |
| Reducing Agent | Substance that causes reduction (gets oxidized itself) | Hydrogen, carbon, metals like zinc and sodium |
| Redox Reaction | Reaction involving both oxidation and reduction | Combustion, respiration, corrosion |
Chemical reactions involve energy changes, usually in the form of heat. Reactions are classified based on whether they release or absorb energy.
Combustion of carbon: Releases heat and light energy.
Other examples:
Decomposition of calcium carbonate: Requires continuous heating.
Other examples:
Oxidation reactions have both beneficial and harmful effects in our daily lives. Two important effects are corrosion and rancidity.
Corrosion is the slow destruction of metals due to chemical reactions with environmental substances like oxygen, moisture, acids, etc.
Chemical reaction: Iron reacts with oxygen and water (moisture) to form hydrated iron(III) oxide (rust).
Conditions required for rusting:
If any one condition is absent, rusting doesn't occur.
Forms protective coating that prevents contact with air and moisture.
Coating iron with zinc layer. Zinc is more reactive and corrodes first, protecting iron.
Mixing iron with other metals to form stainless steel, which is corrosion-resistant.
Electroplating with chromium for shiny, corrosion-resistant surface.
Attaching more reactive metal (like magnesium or zinc) to iron structure.
Using silica gel packs to absorb moisture in packaged products.
Rancidity is the oxidation of oils and fats present in food, resulting in unpleasant smell and taste. It makes food unfit for consumption.
Prevents exposure to oxygen which causes oxidation.
Low temperature slows down oxidation reactions.
Substances like BHA, BHT, vitamin E that prevent oxidation.
Packing food items in nitrogen atmosphere instead of air.
Removing air from food packages to prevent oxidation.
Storing oils in dark bottles to prevent photo-oxidation.
Identification of gases produced in chemical reactions is an important practical skill. Different gases have characteristic tests.
| Gas | Test | Observation | Example Reaction |
|---|---|---|---|
| Hydrogen (H₂) | Bring burning splinter near gas | Burns with a pop sound | Zn + 2HCl → ZnCl₂ + H₂ |
| Oxygen (O₂) | Bring glowing splinter near gas | Relights glowing splinter | 2KClO₃ → 2KCl + 3O₂ |
| Carbon Dioxide (CO₂) | Pass through lime water (Ca(OH)₂) | Turns lime water milky (due to CaCO₃) | CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂ |
| Ammonia (NH₃) | Bring moist red litmus paper near | Turns red litmus blue (alkaline gas) | NH₄Cl + NaOH → NaCl + H₂O + NH₃ |
| Chlorine (Cl₂) | Bring moist blue litmus paper near | Bleaches litmus paper (turns white) | MnO₂ + 4HCl → MnCl₂ + 2H₂O + Cl₂ |
| Sulfur Dioxide (SO₂) | Pass through acidified K₂Cr₂O₇ | Turns orange solution to green | Na₂SO₃ + H₂SO₄ → Na₂SO₄ + H₂O + SO₂ |
Combination:
Decomposition:
Displacement:
Double Displacement:
Neutralization:
Rusting: