MP Board Class 10 Science Chapter 1: Chemical Reactions and Equations Notes

MP Board Class 10 Science Chapter 1: Chemical Reactions and Equations — This chapter is the foundation of chemistry in Class 10. It introduces how substances transform into new substances through chemical reactions, how to write and balance chemical equations, and the different types of reactions. With 12-15 marks weightage in the MP Board exam, this is one of the most scoring chapters. Let’s master it with clear explanations, examples, and previous year questions.

🔬 1. Chemical Reactions — Basic Concepts

A chemical reaction is a process in which one or more substances (reactants) are converted into one or more different substances (products). The reaction involves rearrangement of atoms — bonds between atoms break in the reactants and new bonds form to create the products.

Signs of a Chemical Reaction

How do you know a chemical reaction has occurred? Watch for these observable changes:

📝 2. Writing and Balancing Chemical Equations

Word Equations vs Chemical Equations

A word equation describes the reaction in words: “Magnesium burns in oxygen to form magnesium oxide.” A chemical equation uses chemical formulas: 2Mg + O₂ → 2MgO

The Law of Conservation of Mass

This is the most important principle in chemistry: Mass can neither be created nor destroyed in a chemical reaction. Therefore, the total mass of reactants must equal the total mass of products. This is why we balance chemical equations — the number of atoms of each element must be the same on both sides.

Step-by-Step: How to Balance a Chemical Equation

Let’s balance the reaction of hydrogen with oxygen to form water:

1. Write the word equation: Hydrogen + Oxygen → Water
2. Write the skeletal equation: H₂ + O₂ → H₂O
3. Count atoms on each side: Left: H=2, O=2 | Right: H=2, O=1
4. Balance O first: Put 2 before H₂O → H₂ + O₂ → 2H₂O (now O: 2 on both sides, H: 2 on left, 4 on right)
5. Balance H next: Put 2 before H₂ → 2H₂ + O₂ → 2H₂O
6. Verify: Left: H=4, O=2 | Right: H=4, O=2 ✅ Balanced!

🧪 3. Types of Chemical Reactions

3.1 Combination Reaction

Two or more substances combine to form a single product.

General form: A + B → AB

More examples:
2H₂(g) + O₂(g) → 2H₂O(l) | C(s) + O₂(g) → CO₂(g) | 2Mg(s) + O₂(g) → 2MgO(s)

3.2 Decomposition Reaction

A single compound breaks down into two or more simpler substances. This usually requires energy in the form of heat, light, or electricity.

General form: AB → A + B

3.3 Displacement Reaction

A more reactive element displaces a less reactive element from its compound.

General form: A + BC → AC + B

More examples:
Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s) | Cu(s) + 2AgNO₃(aq) → Cu(NO₃)₂(aq) + 2Ag(s)
The Reactivity Series determines which metal displaces which: K > Na > Ca > Mg > Al > Zn > Fe > Pb > H > Cu > Ag > Au

3.4 Double Displacement Reaction

Two compounds exchange their ions to form two new compounds. Usually results in precipitate formation.

General form: AB + CD → AD + CB

3.5 Oxidation and Reduction (Redox Reactions)

Oxidation: Addition of oxygen or removal of hydrogen.
Reduction: Addition of hydrogen or removal of oxygen.
Redox Reaction: Both oxidation and reduction occur simultaneously.

🔥 3.6 Exothermic and Endothermic Reactions

Exothermic Reactions

Exothermic reactions release heat energy to the surroundings. The temperature of the reaction mixture increases. Energy is released because the total energy of the products is less than the total energy of the reactants. The excess energy is given out as heat, light, or sound.

Real-life examples of exothermic reactions:

• Respiration: C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O + energy. This is how our bodies produce energy from food. The energy released is used for all bodily functions — movement, thinking, growth, and repair.
• Burning of fuels: CH₄ + 2O₂ → CO₂ + 2H₂O + heat and light. This is why LPG, petrol, and coal are used as fuels — they release large amounts of heat when they burn.
• Neutralization reaction: HCl + NaOH → NaCl + H₂O + heat. When an acid reacts with a base, heat is always released.
• Slaking of lime: CaO + H₂O → Ca(OH)₂ + heat. Large amounts of heat are produced — enough to cook food! This is why quicklime is sometimes used in self-heating food packages.

Endothermic Reactions

Endothermic reactions absorb heat energy from the surroundings. The temperature of the reaction mixture decreases. Energy is absorbed because the products have more energy than the reactants. The extra energy comes from the surroundings, making the mixture feel cold.

Real-life examples of endothermic reactions:

• Photosynthesis: 6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂ (requires sunlight energy). Plants absorb sunlight to convert carbon dioxide and water into glucose and oxygen. This is the most important endothermic reaction on Earth — it produces food for almost all life forms.
• Thermal decomposition: CaCO₃ → CaO + CO₂ (requires heat). Limestone is heated in kilns to produce quicklime for construction.
• Photosynthesis in plants: Plants absorb sunlight energy to convert CO₂ and H₂O into glucose and oxygen — the reverse of respiration, and it requires continuous energy input.
• Melting of ice: H₂O(s) + heat → H₂O(l). While this is a physical change, it demonstrates how endothermic processes feel cold because they absorb heat from their surroundings.

⚡ 4. Oxidation and Reduction in Detail

Oxidising and Reducing Agents

🔥 5. Effects of Oxidation in Daily Life

5.1 Corrosion

Corrosion is the gradual destruction of metals by chemical reaction with environmental oxygen, moisture, or acids.

• Rusting of iron: 4Fe + 3O₂ + 2xH₂O → 2Fe₂O₃·xH₂O (rust). Rust is a hydrated iron(III) oxide — reddish-brown flaky substance.
• Silver tarnishing: 2Ag + H₂S → Ag₂S + H₂. Silver reacts with hydrogen sulphide in air, forming a black layer of silver sulphide.
• Copper corrosion: 2Cu + CO₂ + H₂O + O₂ → Cu₂(OH)₂CO₃. Copper develops a green coating of basic copper carbonate (verdigris).

Methods to Prevent Corrosion

• Galvanization: Coating iron/zinc with a layer of zinc. Zinc is more reactive, so it corrodes first (sacrificial protection). Used in iron gates, water pipes.
• Electroplating: Coating with a layer of chromium, nickel, or tin. Bicycle handles, car bumpers.
• Oil/Grease/Paint: Creates a barrier between metal and air/moisture.
• Alloying: Mixing iron with chromium and nickel makes stainless steel (resists rust).
• Anodizing: Creating a protective oxide layer on aluminium.

5.2 Rancidity

Rancidity is the oxidation of oils and fats in food, causing unpleasant smell and taste. Foods containing oils and fats (chips, butter, fried items) become rancid when exposed to air for too long.

Prevention methods:

• Packaging in nitrogen gas: Nitrogen is an inert gas that displaces oxygen. This is why chips packets are flushed with nitrogen before sealing.
• Adding antioxidants: Substances like BHA (butylated hydroxyanisole) and BHT are added to foods to prevent oxidation.
• Refrigeration: Low temperature slows down the oxidation reaction.
• Airtight containers: Minimize contact with oxygen.

📖 6. Important Chemical Equations for MP Board Exams

Memorize these equations — they are frequently asked in MP Board 10th Science papers:

⚠️ 7. Common Mistakes to Avoid

• ❌ Changing subscripts to balance equations: NEVER change the subscript (small number) in a chemical formula to balance the equation. Example: Changing H₂O to H₂O₂ is wrong! Only change coefficients (the big numbers before formulas).
• ❌ Forgetting state symbols: Always include (s), (l), (g), (aq) in balanced equations. These carry 1 mark in 3-mark questions.
• ❌ Confusing displacement and double displacement: Displacement = element + compound. Double displacement = compound + compound → exchange of ions.
• ❌ Not recognizing redox pairs: In any reaction given in the exam, identify which substance is oxidized (gained O / lost H) and which is reduced (lost O / gained H), then name the oxidizing and reducing agents.
• ❌ Writing incorrect chemical formulas: MgO not MgO₂. CaCO₃ not CaCO₂. H₂O not H₂O₂. Always double-check valencies.
• ❌ Confusing rancidity with corrosion: Rancidity = oil/food spoilage. Corrosion = metal deterioration. Both are caused by oxidation but affect different materials.
• ❌ Missing the colour changes: Always note the colour changes in displacement reactions — blue CuSO₄ → green FeSO₄; colourless AgNO₃ → blue Cu(NO₃)₂.

📋 Previous Year Questions (2017–2026)

❓ Frequently Asked Questions

Q1: What is the law of conservation of mass?

The law states that mass can neither be created nor destroyed in a chemical reaction. This means the total mass of reactants equals the total mass of products. This is why chemical equations must be balanced.

Q2: Why do we balance chemical equations?

Chemical equations are balanced to satisfy the law of conservation of mass. A balanced equation has the same number of atoms of each element on both sides, ensuring mass is conserved.

Q3: What is the difference between a combination and a decomposition reaction?

In a combination reaction, two or more substances combine to form a single product (A + B → AB). In a decomposition reaction, a single compound breaks down into two or more simpler substances (AB → A + B). They are opposite processes.

Q4: What is the use of photolytic decomposition in photography?

Silver chloride (AgCl) and silver bromide (AgBr) undergo photolytic decomposition in sunlight: 2AgCl → 2Ag + Cl₂. The grey-coloured silver metal deposited on the film creates the image. That’s why photographic films are stored in dark containers.

Q5: How does the reactivity series help predict displacement reactions?

A more reactive metal (higher in the reactivity series) can displace a less reactive metal from its salt solution. For example, iron displaces copper from CuSO₄ because iron is above copper in the reactivity series: K > Na > Ca > Mg > Al > Zn > Fe > Pb > H > Cu > Ag > Au.

Q6: What is the difference between an exothermic and endothermic reaction?

Exothermic reactions release heat energy (temperature increases), e.g., respiration, burning of fuels, CaO + H₂O. Endothermic reactions absorb heat energy (temperature decreases), e.g., photosynthesis, thermal decomposition of CaCO₃.

Q7: Why is nitrogen gas used in food packaging?

Nitrogen is an inert gas that does not react with food. When food packets are flushed with nitrogen, oxygen is removed from the packet. Since oxygen causes rancidity (oxidation of oils and fats), packaging in nitrogen prevents spoilage and keeps the food fresh for longer.

Q8: What happens when zinc granules are added to dilute sulphuric acid?

Zinc reacts with dilute sulphuric acid to form zinc sulphate and hydrogen gas: Zn + H₂SO₄ → ZnSO₄ + H₂↑. The hydrogen gas can be tested by bringing a burning matchstick near it — the gas burns with a pop sound. This is a displacement reaction.

Q9: What is the difference between rusting of iron and corrosion of copper?

Rusting: Iron reacts with oxygen and water to form hydrated iron(III) oxide, a reddish-brown flaky substance. Copper corrosion: Copper reacts with CO₂, moisture, and oxygen to form a green coating of basic copper carbonate, Cu₂(OH)₂CO₃. Both are oxidation processes but affect different metals.

Q10: How can you identify the type of chemical reaction from a given equation?

Look at the pattern: (i) A + B → AB = Combination; (ii) AB → A + B = Decomposition; (iii) A + BC → AC + B = Displacement; (iv) AB + CD → AD + CB = Double Displacement; (v) If both oxidation and reduction occur simultaneously = Redox. Also check if heat/light/electricity is required for decomposition reactions.

Q11: What is galvanization? Why is it effective?

Galvanization is the process of coating iron or steel with a layer of zinc. Zinc is more reactive than iron, so when exposed to air and moisture, zinc corrodes first instead of iron (sacrificial protection). This protects the underlying iron from rusting. It is used for iron gates, water pipes, and roofing sheets.

Q12: Why does the colour of copper sulphate solution change when an iron nail is dipped in it?

Copper sulphate solution is blue due to Cu²⁺ ions. When iron nail is dipped, iron displaces copper (Fe + CuSO₄ → FeSO₄ + Cu). The Fe²⁺ ions (ferrous) give a light green colour to the solution. Copper metal deposits on the nail as a reddish-brown layer. This colour change is a key indicator of a displacement reaction.

Q13: What is the difference between a physical change and a chemical change?

A physical change involves a change in physical properties like shape, size, or state but no new substance is formed (e.g., melting of wax, boiling of water). A chemical change involves the formation of one or more new substances with different chemical properties (e.g., burning of wood, rusting of iron). Chemical changes are usually irreversible, while physical changes are often reversible.

Q14: Why is the reaction between quicklime and water considered a combination reaction?

In the reaction CaO + H₂O → Ca(OH)₂, two substances (calcium oxide and water) combine to form a single product (calcium hydroxide). This fits the definition of a combination reaction: A + B → AB. It is also exothermic because large amounts of heat are released — you can feel the container getting hot.

Q15: What happens when lead nitrate is heated? Write the balanced equation.

When lead nitrate [Pb(NO₃)₂] is heated strongly, it undergoes thermal decomposition to produce lead oxide (PbO), nitrogen dioxide (NO₂), and oxygen (O₂). The balanced equation is: 2Pb(NO₃)₂ → 2PbO + 4NO₂ + O₂. The brown fumes of nitrogen dioxide (NO₂) are clearly visible. The colour of the residue changes from white to yellow (lead oxide is yellow when hot and turns pale yellow on cooling). This is one of the most commonly asked practical-based questions in MP Board exams.

Q16: How does the process of electrolysis of water demonstrate a decomposition reaction?

When an electric current is passed through water (with a few drops of acid added to make it conducting), water decomposes into its constituent elements: 2H₂O → 2H₂ + O₂. At the cathode (negative electrode), hydrogen gas is produced. At the anode (positive electrode), oxygen gas is produced. The volume of hydrogen collected is double that of oxygen — proving water has the formula H₂O. This is an example of electrolytic decomposition where electrical energy is used to break the compound.

Q17: What is a precipitation reaction? Give an example.

A precipitation reaction is a reaction in which an insoluble solid (called precipitate) is formed when two solutions are mixed. For example, when sodium sulphate solution (Na₂SO₄) is mixed with barium chloride solution (BaCl₂), a white precipitate of barium sulphate (BaSO₄) is formed: Na₂SO₄ + BaCl₂ → BaSO₄↓ + 2NaCl. The arrow pointing downward (↓) indicates the precipitate. Precipitation reactions are always double displacement reactions. These are important in water purification and qualitative analysis in chemistry labs.

Q18: What are the conditions required for rusting? How can it be prevented?

Rusting requires two things: (1) oxygen and (2) water or moisture. In the absence of either, rusting does not occur. That’s why an iron nail in boiled water (which has no dissolved oxygen) does not rust, and an iron nail kept with a drying agent (calcium chloride) in a sealed tube also does not rust. Prevention methods include: galvanization (zinc coating), painting, oiling/greasing, making alloys (stainless steel), and electroplating. In MP Board exams, a 3-mark question often asks: “Describe an experiment to show that both air and water are necessary for rusting.”

Q19: Why does silver jewellery turn black over time?

Silver reacts with hydrogen sulphide (H₂S) present in the air to form a black layer of silver sulphide (Ag₂S): 2Ag + H₂S → Ag₂S + H₂. This is a form of corrosion specific to silver. It is accelerated in polluted air and in areas with high humidity. This tarnishing can be removed by cleaning the jewellery with a mild abrasive or by using a silver polishing cloth. This is an example of oxidation of silver metal — the silver loses electrons (gets oxidized) and sulphur gains electrons (gets reduced).

Q20: What is the importance of the reactivity series in predicting chemical reactions?

The reactivity series ranks metals in order of their reactivity, from most reactive (potassium) to least reactive (gold). Its importance includes: (1) Predicting displacement reactions — a more reactive metal displaces a less reactive metal from its salt solution. (2) Determining extraction methods — highly reactive metals (K, Na, Ca) need electrolytic reduction, moderately reactive metals (Zn, Fe, Pb) can be extracted by reduction with carbon, and less reactive metals (Cu, Ag, Au) occur natively. (3) Predicting corrosion resistance — less reactive metals like gold and platinum do not corrode easily. (4) Understanding battery chemistry — more reactive metals make better anodes in electrochemical cells.