Selina Concise Solutions for ICSE Class 9 Chemistry Chapter 5 Chemical Changes And Reactions

Page No: 22

Question 1.
Answer: 1. A chemical reaction is the process of breaking the chemical bonds of the reacting substances (reactants) and making new bonds to form new substances (products).
2. Conditions necessary for a chemical change or reaction are
• Evolution of gas
• Change of colour
• Formation of precipitate
• Change of state
In simple words: A chemical reaction is like taking apart LEGO structures (breaking bonds) and rebuilding them into something new (forming products). You can tell it happened if you see bubbles, a color change, or a new solid forming.

📝 Teacher's Note: Use the analogy of a kitchen where ingredients (reactants) are mixed and cooked (energy/conditions) to create a cake (product) that cannot be turned back into raw eggs. Emphasize that "breaking bonds" requires energy.

🎯 Exam Tip: When listing conditions for a chemical change, always include "Evolution of gas" and "Formation of precipitate" as these are the most observable physical evidences.

Question 2.
Answer: • A chemical bond is the force which holds the atoms of a molecule together as in a compound.
• Formation of gas bubbles in a liquid during a reaction is called effervescence.
• Chemical reactions which are characterised by the formation of insoluble solid substances are called precipitates.
In simple words: A bond is like "atomic glue" that keeps atoms stuck together. Effervescence is just the fancy word for fizzing or bubbling, and a precipitate is solid "dust" or "chunks" that appear in a clear liquid.

📝 Teacher's Note: Demonstrate effervescence in class using ENO or baking soda and vinegar. For precipitates, use the reaction between silver nitrate and common salt to show the white "cloud" forming.

🎯 Exam Tip: Define "precipitate" as an *insoluble* solid; the word "insoluble" is a key term examiners look for.

Question 3.
Answer:
(a) \( \text{CuCO}_{3(s)} \xrightarrow{\Delta} \text{CuO} + \text{CO}_{2(g)} \)
(b) \( 6\text{CO}_2 + 6\text{H}_2\text{O} \rightarrow \text{C}_6\text{H}_{12}\text{O}_6 + 6\text{O}_2 \)
(c) \( 2\text{H}_2\text{O} \xrightarrow{\text{Electricity}} 2\text{H}_2 + \text{O}_2 \)
(d) \( \text{PbNO}_{3(s)} + 2\text{KI}_{(s)} \rightarrow 2\text{KNO}_{3(s)} + \text{PbI}_{2(s)} \)
(e) \( \text{NaCl}_{(s)} + \text{AgNO}_{3(s)} \rightarrow \text{AgCl} \downarrow + \text{NaNO}_{3(aq)} \)
(f) \( \text{N}_2 + 3\text{H}_2 \rightarrow{\text{Above 200 atm}} 2\text{NH}_3 \)
(g) \( 2\text{KClO}_3 \xrightarrow{\text{MnO}_2} 2\text{KCl} + 3\text{O}_2 \)
In simple words: These are chemical recipes written in symbols. They show how things like heat, electricity, or pressure can change one group of chemicals into another.

📝 Teacher's Note: Explain the symbols over the arrows: \( \Delta \) means heating, and "MnO2" is a catalyst that helps the reaction without being used up. Ensure students know (s) means solid and (g) means gas.

🎯 Exam Tip: Always include state symbols like (s), (l), (g), or (aq) if you know them, as they provide a complete description of the reaction.

Question 4.
Answer: (a) It is a reaction which occurs with absorption of light energy.
\[ 6\text{CO}_2 + 6\text{H}_2\text{O} \rightarrow \text{C}_6\text{H}_{12}\text{O}_6 + 6\text{O}_2 \]
(b) It is a reaction which occurs with absorption of electrical energy.
\[ 2\text{H}_2\text{O} \xrightarrow{\text{Electricity}} 2\text{H}_2 + \text{O}_2 \]
In simple words: Some reactions need energy to start. Plants use light to make food (photosynthesis), and we can use electricity to split water into hydrogen and oxygen gases.

📝 Teacher's Note: Use the term "Photochemical reaction" for sub-part (a) and "Electrochemical reaction" for sub-part (b) to introduce technical vocabulary.

🎯 Exam Tip: For the electrolysis of water, remember that Hydrogen is collected at the cathode and Oxygen at the anode.

Question 5.
Answer: (a)
(i) \( 2\text{AgNO}_3 \xrightarrow{\text{Sunlight}} 2\text{Ag} + 2\text{NO}_2 + \text{O}_2 \)
(ii) \( 6\text{CO}_2 + 6\text{H}_2\text{O} \rightarrow \text{C}_6\text{H}_{12}\text{O}_6 + 6\text{O}_2 \)

(b)
(i) \( \text{Fe} + \text{CuSO}_{4(aq)} \rightarrow \text{FeSO}_4 + \text{Cu} \)
(Blue solution \( \rightarrow \) Green solution)
(ii) \( \text{FeSO}_4 + 2\text{NaOH} \rightarrow \text{Fe(OH)}_2 + \text{Na}_2\text{SO}_{4(aq)} \)
(Dirty green ppt.)

(c) \( \text{NH}_{3(g)} + \text{HCl}_{(g)} \rightleftharpoons \text{NH}_4\text{Cl}_{(s)} \)

(d) \( \text{Fe} + \text{CuSO}_{4(aq)} \rightarrow \text{FeSO}_4 + \text{Cu} \)
(Blue solution \( \rightarrow \) Green solution)
In simple words: These equations show how silver nitrate breaks down in light, how iron changes the color of copper solution, and how two gases can combine to make a solid white smoke.

📝 Teacher's Note: The "dirty green precipitate" in (b)(ii) is a classic test for the presence of Ferrous (\( \text{Fe}^{2+} \)) ions. Contrast this with the reddish-brown precipitate formed by Ferric (\( \text{Fe}^{3+} \)) ions.

🎯 Exam Tip: Note the double arrow in (c) which indicates a reversible reaction. This is important for reactions involving Ammonium Chloride.

Question 6.
Answer: (a) \( \text{Zn} + \text{H}_2\text{SO}_4 \rightarrow \text{ZnSO}_4 + \text{H}_2 \)
(b) \( \text{Fe} + \text{CuSO}_{4(aq)} \rightarrow \text{FeSO}_4 + \text{Cu} \)
(Blue solution \( \rightarrow \) Green solution)
(c) \( \text{NaCl}_{(s)} + \text{AgNO}_{3(s)} \rightarrow \text{AgCl} \downarrow + \text{NaNO}_{3(aq)} \)
(d) \( \text{NH}_{3(g)} + \text{HCl}_{(g)} \rightleftharpoons \text{NH}_4\text{Cl}_{(s)} \)
In simple words: These are examples of different reaction types like displacement (one metal pushing out another) and precipitation (making a solid).

📝 Teacher's Note: Use (a) to discuss the laboratory preparation of Hydrogen gas. Point out that Zinc is used because it reacts at a moderate, safe speed.

🎯 Exam Tip: In reaction (c), the downward arrow \( \downarrow \) is the symbol for a precipitate. Don't forget to write it!

Question 7.
Answer: • Silver nitrate solution is kept in brown bottles in the laboratory because it decomposes in the presence of light.
• Molybdenum increases the efficiency of the catalyst iron used in the manufacture of ammonia.
• This is because the blue colour of the copper sulphate solution fades and eventually turns into light green due to the formation of ferrous sulphate.
In simple words: Some chemicals are light-sensitive, so they need "sunglasses" (brown bottles). Molybdenum is like a "helper" for iron, making it work better as a catalyst.

📝 Teacher's Note: Explain that Molybdenum is a "promoter," which is a substance that enhances a catalyst's performance. Also, show how iron "displaces" copper to cause the color change from blue to green.

🎯 Exam Tip: The brown bottle question is a very common "reasoning" question. Ensure you mention that light causes "decomposition" of the chemical.

Page No: 27

Question 1.
Answer: • Displacement
• Double decomposition
• Accelerate, decelerate, unaffected
In simple words: These terms describe the types of chemical "swaps" that happen and how catalysts can speed up or slow down a reaction.

📝 Teacher's Note: Clarify that "Double decomposition" is often called "Double Displacement" in some textbooks. They mean the same thing: an exchange of ions.

🎯 Exam Tip: A positive catalyst *accelerates* a reaction, while a negative catalyst (inhibitor) *decelerates* it.

Question 2.
Answer: Combination
Decomposition
In simple words: In combination, things join together. In decomposition, things break apart.

📝 Teacher's Note: These are the two most fundamental types of reactions. Use simple math analogies: Addition (Combination) and Subtraction/Division (Decomposition).

🎯 Exam Tip: Always check if a reaction has one product (Combination) or one reactant (Decomposition) to identify its type quickly.

Question 3.
Answer: (a) Double decomposition: This is a type of chemical change in which two compounds in a solution react to form two new compounds by mutual exchange of radicals.
\( \text{NaCl}_{(s)} + \text{AgNO}_{3(s)} \rightarrow \text{AgCl} \downarrow + \text{NaNO}_{3(aq)} \)

(b) Thermal decomposition: A decomposition reaction brought about by heat is known as thermal decomposition.
\( 2\text{HgO}_{(s)} \xrightarrow{\Delta} 2\text{Hg}_{(s)} + \text{O}_{2(g)} \)

(c) Reversible reaction: A chemical reaction in which the direction of a chemical change can be reversed by changing the conditions under which the reaction is taking place is called a reversible reaction.
\( \text{CuSO}_4 \cdot 5\text{H}_2\text{O}_{(s)} \rightleftharpoons \text{CuSO}_{4(s)} + 5\text{H}_2\text{O}_{(g)} \)

(d) Displacement: It is a chemical change in which a more active element displaces a less active element from its salt solution.
\( \text{CuSO}_4 + \text{Zn} \rightarrow \text{ZnSO}_4 + \text{Cu} \)
In simple words: Double decomposition is like two couples swapping partners at a dance. Thermal decomposition is breaking something using heat. Reversible means it can go back and forth. Displacement is a stronger element "kicking out" a weaker one.

📝 Teacher's Note: Use the Activity Series of Metals to explain displacement. Zinc can kick out Copper because it is higher on the list (more "active").

🎯 Exam Tip: In the definition of double decomposition, the keyword is "mutual exchange of radicals." Make sure to include it.

Question 4.
Answer: A reaction in which two or more substances combine together to form a single substance is called a synthesis or combination reaction.
\( \text{A} + \text{B} \rightarrow \text{AB} \)
In the above reaction, substances A and B combine to give a molecule of a new substance, AB.
Carbon burns in oxygen to form a gaseous compound, carbon dioxide.
\( \text{C} + \text{O}_2 \rightarrow \text{CO}_2 \)
In simple words: When two things become one, it's a combination. For example, solid carbon and oxygen gas combine to make the carbon dioxide gas we breathe out.

📝 Teacher's Note: Contrast synthesis (A + B \( \rightarrow \) AB) with decomposition (AB \( \rightarrow \) A + B) to help students see they are opposites.

🎯 Exam Tip: Synthesis and Combination are synonyms. You can use either term unless the question specifically asks for one.

Question 5.
Answer: A decomposition reaction brought about by heat is known as thermal decomposition.
\( 2\text{HgO}_{(s)} \xrightarrow{\Delta} 2\text{Hg}_{(s)} + \text{O}_{2(g)} \)
A simultaneous reversible decomposition reaction brought about only by heat is thermal dissociation.
\( \text{NH}_4\text{Cl} \rightleftharpoons \text{NH}_3 + \text{HCl} \)
In simple words: Thermal decomposition breaks a substance permanently. Thermal dissociation is a "break-up" that can "make-up" (recombine) as soon as it cools down.

📝 Teacher's Note: The main difference between decomposition and dissociation is reversibility. Thermal dissociation always goes back to the original substance when the heat is removed.

🎯 Exam Tip: If you see a reversible arrow \( \rightleftharpoons \) with heat, it's likely thermal dissociation, not just decomposition.

Question 6.
Answer: The reaction between an acid and a base which forms salt and water only is referred to as reaction of neutralisation.
Applications of neutralisation reactions:
• When someone is stung by a bee, formic acid enters the skin and causes pain, which can be relieved by rubbing the spot with slaked lime or baking soda, both of which are bases.
• Acid which accidentally spills on to our clothes can be neutralised with ammonia solution.
• If the soil is somewhat acidic and thus unfavourable for growing of certain crops, then slaked lime is added to neutralise the excess acid.
In simple words: Neutralization is when an acid and a base cancel each other out to make salt and water. We use it to stop the sting of a bee or to fix soil that is too sour for plants.

📝 Teacher's Note: Neutralization is a very practical concept. Use the bee sting vs. wasp sting example: bee stings are acidic (neutralize with base like soda), while wasp stings are alkaline (neutralize with acid like vinegar).

🎯 Exam Tip: Remember the "Salt + Water only" part of the definition. If other products are formed, it might not be a pure neutralization.

Question 7.
Answer: Hydrolysis is the process in which a salt and water react to form an acidic or basic solution.
In the process of hydrolysis, only those salts hydrolyse which are formed by the reaction of:
• strong base and weak acid
• strong acid and weak base
This happens because a salt formed due to a strong base and a weak acid on dissolving in water will form a basic solution. A basic solution turns red litmus blue.
\( \text{Na}_2\text{CO}_3 + 2\text{H}_2\text{O} \rightarrow 2\text{NaOH} + \text{H}_2\text{CO}_3 \)
However, the salt formed due to a strong acid and a weak base on dissolving in water will make an acidic solution. Acidic solutions turn blue litmus red.
\( \text{FeCl}_3 + 3\text{H}_2\text{O} \rightarrow \text{Fe(OH)}_3 + 3\text{HCl} \)
In simple words: Hydrolysis is when a salt reacts with water to make the water slightly acidic or basic. It's like the "ghost" of the stronger parent (acid or base) coming back to change the pH.

📝 Teacher's Note: Explain that salts of strong acids and strong bases (like NaCl) do *not* undergo hydrolysis; they stay neutral. This is a common point of confusion.

🎯 Exam Tip: To remember litmus colors: Base turns it Blue. Acid stays (or turns) Red.

Question 8.
Answer: Iron(III) chloride is a salt prepared from strong acid \( \text{HCl} \) and a weak base \( \text{Fe(OH)}_3 \).
\( \text{Fe(OH)}_3 + 3\text{HCl} \rightarrow \text{FeCl}_3 + 3\text{H}_2\text{O} \)
On the other hand, sodium carbonate is a salt prepared from a strong base \( \text{NaOH} \) and a weak acid \( \text{H}_2\text{CO}_3 \).
In simple words: Not all salts are neutral like table salt. Some are born from strong acids (making them acidic salts) and others from strong bases (making them basic salts).

📝 Teacher's Note: This builds on the concept of hydrolysis. Use this to explain why \( \text{FeCl}_3 \) solution has a pH less than 7.

🎯 Exam Tip: If asked to identify the nature of a salt, look at the acid and base it came from. The "strong" parent always wins the pH battle.

Question 9.
Answer: Decomposition is the breaking up of a compound either into elements or simpler compounds such that these products do not combine to form the original compound. Decomposition may occur in the presence of heat or light or by the passage of an electric current.
Example: Mercuric oxide when heated decomposes to form two elements—mercury and oxygen.
\( 2\text{HgO}_{(s)} \xrightarrow{\Delta} 2\text{Hg}_{(s)} + \text{O}_{2(g)} \)
In simple words: Decomposition is like taking a toy apart. Once it's in pieces (like mercury and oxygen), it doesn't just jump back together on its own.

📝 Teacher's Note: Contrast this again with thermal dissociation. Use the Mercuric oxide experiment to show how a red powder turns into silver liquid (mercury) and a gas (oxygen) that supports combustion.

🎯 Exam Tip: When writing the example for decomposition, always include the heat symbol \( \Delta \) over the arrow.

Question 10.
Answer: • \( \text{Cl}_2 + 2\text{KBr} \rightarrow 2\text{KCl} + \text{Br}_2 \)
Displacement
• \( \text{Fe} + \text{CuSO}_4 \rightarrow \text{FeSO}_4 + \text{Cu} \)
Displacement
• \( 2\text{H}_2\text{O} \rightarrow 2\text{Hg} + \text{O}_2 \) (Note: Likely meant \( \text{HgO} \) or splitting of water)
Decomposition
• \( \text{PbO}_2 + \text{SO}_2 \rightarrow \text{PbSO}_4 \)
Combination
• \( \text{AgNO}_3 + \text{NaCl} \rightarrow \text{AgCl} + \text{NaNO}_3 \)
Double decomposition
• \( 2\text{KClO}_3 \rightarrow 2\text{KCl} + 3\text{O}_2 \)
Decomposition
• \( 2\text{H}_2\text{O}_2 \rightarrow 2\text{H}_2\text{O} + \text{O}_2 \)
Decomposition
• \( \text{KNO}_3 + \text{H}_2\text{SO}_4 \rightarrow \text{HNO}_3 + \text{KHSO}_4 \)
Double decomposition
• \( \text{CuO} + \text{H}_2 \rightarrow \text{Cu} + \text{H}_2\text{O} \)
Displacement
• \( \text{CaCO}_3 \rightarrow \text{CaO} + \text{CO}_2 \)
Decomposition
• \( \text{NH}_4\text{Cl} \rightarrow \text{NH}_3 + \text{HCl} \)
Decomposition
In simple words: This is a matching list. You look at the reaction and name it based on what the atoms are doing—swapping, joining, or breaking apart.

📝 Teacher's Note: Use this as a rapid-fire quiz. Have students identify the "type" of reaction based on the general formulas like \( \text{A} + \text{B} \rightarrow \text{AB} \) or \( \text{AB} + \text{CD} \rightarrow \text{AD} + \text{CB} \).

🎯 Exam Tip: Double decomposition reactions often involve ions in solution (\( aq \)) and result in a precipitate or a weak electrolyte like water.

Page No: 30

Question 1.
Answer: Main characteristics of chemical reactions are
(i) Evolution of gas: In many chemical reactions, one of the products is a gas.
\( \text{Zn} + \text{H}_2\text{SO}_4 \rightarrow \text{ZnSO}_4 + \text{H}_2 \uparrow \)
(ii) Change of colour:
Certain chemical reactions are characterised by a change in the colour of the reactants.
\( \text{Fe} + \text{CuSO}_{4(aq)} \rightarrow \text{FeSO}_4 + \text{Cu} \)
(iii) Formation of precipitates:
Certain chemical reactions are characterised by the formation of insoluble solid substances called precipitates.
\( \text{AgNO}_{3(aq)} + \text{NaCl}_{(aq)} \rightarrow \text{AgCl} \downarrow + \text{NaNO}_{3(aq)} \)
(iv) Change of state:
In many chemical reactions, a change of state is observed. The reaction might start with gaseous or liquid reactants and end with solid products and vice versa.
\( \text{NH}_{3(g)} + \text{HCl}_{(g)} \rightleftharpoons \text{NH}_4\text{Cl}_{(s)} \)
In simple words: Reactions give us "clues" that they happened. These clues are fizzing (gas), changing colors, making solid "clouds" (precipitate), or turning a gas into a solid.

📝 Teacher's Note: These four characteristics are the primary way we observe chemistry in the macro world. Always use these examples to ground abstract concepts in reality.

🎯 Exam Tip: If asked for "characteristics," always provide the name of the characteristic followed by a balanced chemical equation as an example.

Question 2.
Answer: Exothermic reaction:
A chemical reaction in which heat is given out is called an exothermic reaction.
Example: When carbon burns in oxygen to form carbon dioxide, a lot of heat is produced.
\( \text{C} + \text{O}_2 \rightarrow \text{CO}_2 + \text{Heat} \)
When hydrogen is burnt in oxygen, water is formed and heat is released.
\( 2\text{H}_2 + \text{O}_2 \xrightarrow{\Delta} 2\text{H}_2\text{O} + \text{Heat} \)
Endothermic reaction:
A reaction in which heat is absorbed is called endothermic reaction.
Example: When carbon is heated with sulphur at high temperature, liquid carbon disulphide is formed.
\( \text{C} + 2\text{S} \xrightarrow{\Delta} \text{CS}_2 \)
When nitrogen and oxygen are heated together to a temperature of about \( 3000^\circ\text{C} \), nitric oxide gas is formed.
\( \text{N}_2 + \text{O}_2 \xrightarrow{\text{Heat } 3000^\circ\text{C}} 2\text{NO} \)
In simple words: Exothermic reactions feel hot because they "exit" or throw out heat. Endothermic reactions feel cold because they "enter" or suck in heat from the surroundings.

📝 Teacher's Note: Use the prefixes "Exo" (exit) and "Endo" (into) to help students remember the direction of heat flow.

🎯 Exam Tip: In exothermic reactions, "+ Heat" is written on the products side. In endothermic reactions, "Heat" is written on the reactants side or above the arrow.

Question 3.
Answer: Exothermic reactions are spontaneous and warm their surroundings with the release of heat energy.
Endothermic reactions absorb heat energy from their surroundings and cause their surroundings to cool down.
In simple words: Exothermic is like a fireplace—it gives off warmth. Endothermic is like an ice pack—it takes away heat.

📝 Teacher's Note: While many exothermic reactions are spontaneous, remind students that some might still need a little "spark" (activation energy) to get started.

🎯 Exam Tip: If a reaction vessel feels hot to the touch, the reaction inside is exothermic. If it feels cold, it is endothermic.

Question 4.
Answer: (a) \( \text{C} + \text{O}_2 \rightarrow \text{CO}_2 + \text{Heat} \)
(b) \( \text{C} + 2\text{S} \xrightarrow{\Delta} \text{CS}_2 \)
(c) \( \text{N}_2 + 3\text{H}_2 \rightleftharpoons{\text{Above 200 atm}} 2\text{NH}_3 \)
In simple words: These are specific examples of energy-releasing (burning) and energy-absorbing reactions.

📝 Teacher's Note: Sub-part (c) refers to the Haber process. Note that while it is technically exothermic, it requires high pressure and some heat to overcome activation energy.

🎯 Exam Tip: Ammonia production is a classic example of a reversible reaction. Use the \( \rightleftharpoons \) symbol.

Question 5.
Answer: (a) It is a reaction which occurs with absorption of light energy.
Example: Photosynthesis
\( 6\text{CO}_2 + 6\text{H}_2\text{O} \xrightarrow{\text{light}} \text{C}_6\text{H}_{12}\text{O}_6 + 6\text{O}_2 \)

(b) It is a reaction which occurs with absorption of electrical energy.
Example:
Acidulated water breaks into hydrogen and oxygen.
\( 2\text{H}_2\text{O} \xrightarrow{\text{Electricity}} 2\text{H}_2 + \text{O}_2 \)
In simple words: Light and electricity are just different forms of "energy food" that reactions use to happen.

📝 Teacher's Note: Explain that "acidulated water" is just water with a few drops of acid added to make it conduct electricity better, as pure water is a poor conductor.

🎯 Exam Tip: Photosynthesis is the most famous photochemical reaction. Ensure you know the full balanced equation.

Question 6.
Answer: (a)
(i) Change of state
Ammonia gas reacts with \( \text{HCl} \) gas to give solid ammonium chloride.
\( \text{NH}_{3(g)} + \text{HCl}_{(g)} \rightleftharpoons \text{NH}_4\text{Cl}_{(s)} \)
(ii) Formation of precipitate
When a solution of silver nitrate is added to a solution of sodium chloride, a white insoluble substance, silver chloride, is formed.
\( \text{AgNO}_{3(aq)} + \text{NaCl}_{(aq)} \rightarrow \text{AgCl}_{(aq)} + \text{NaNO}_{3(aq)} \) (Note: \( \text{AgCl} \) is the precipitate)

(b)
Exothermic reaction:
When carbon burns in oxygen to form carbon dioxide, a lot of heat is produced.
\( \text{C} + \text{O}_2 \rightarrow \text{CO}_2 + \text{Heat} \)
Endothermic reaction:
When carbon is heated with sulphur at high temperature, liquid carbon disulphide is formed.
\( \text{C} + 2\text{S} \xrightarrow{\Delta} \text{CS}_2 \)

(c) Colour change
A few pieces of iron are added into a blue coloured copper sulphate solution; the blue colour of copper sulphate fades and eventually turns into light green due to the formation of ferrous sulphate.
\( \text{Fe} + \text{CuSO}_4 \rightarrow \text{FeSO}_4 + \text{Cu} \)
In simple words: These are all the visible "proofs" that a chemical reaction has taken place—color changes, temperature changes, and things turning from gas to solid.

📝 Teacher's Note: The reaction of \( \text{NH}_3 \) and \( \text{HCl} \) produces a thick white "smoke." It's a great visual for a change of state. The Iron/Copper experiment is the best way to show the Activity Series.

🎯 Exam Tip: For sub-part (c), always mention the specific colors (Blue to Green) to get full marks for the observation.

Question 7.
Answer: A chemical reaction is the process of breaking the chemical bonds of the reacting substances (reactants) and making new bonds to form new substances (products).
A chemical change or chemical reaction occurs when particles collide. Collisions occur when reactants are in close contact or by supply of energy.
In simple words: For a reaction to happen, the chemical bits have to bump into each other (collide) hard enough to break their old bonds and form new ones.

📝 Teacher's Note: This introduces "Collision Theory." The more the particles collide, the faster the reaction. This is why we stir solutions or heat them up.

🎯 Exam Tip: Remember that "Collision" and "Energy" are the two requirements for a reaction to occur according to this theory.

Question 8.
Answer: (a) \( \text{NaCl}_{(aq)} + \text{AgNO}_{3(aq)} \rightarrow \text{AgCl}_{(aq)} + \text{NaNO}_{3(aq)} \)
(b) \( \text{Pb(NO}_3)_2 + 2\text{KI} \rightarrow 2\text{KNO}_3 + \text{PbI}_2 \)
(c) \( \text{CuCO}_3 \xrightarrow{\Delta} \text{CuO}_{(s)} + \text{CO}_{2(g)} \)
(d) \( 2\text{Pb(NO}_3)_2 \xrightarrow{\Delta} 2\text{PbO} + 4\text{NO}_2 + \text{O}_2 \)
(e) \( 4\text{NH}_3 + 5\text{O}_2 \xrightarrow{\text{Pt}} 4\text{NO} + 6\text{H}_2\text{O} \)
In simple words: These are balanced equations showing different chemical changes, like the breakdown of lead nitrate or the making of nitrogen oxide using a catalyst.

📝 Teacher's Note: Reaction (d) is very important in the lab. Lead nitrate is a white solid that turns into yellow lead monoxide, releasing brown fumes of nitrogen dioxide. It's a very colorful decomposition!

🎯 Exam Tip: The catalyst "Pt" (Platinum) in reaction (e) must be written above the arrow.

Question 9.
Answer: • Lead nitrate decomposes on heating leaving a yellow residue lead monoxide, brown gas nitrogen dioxide and colourless gas oxygen.
• If chlorine water is exposed to sunlight, oxygen is evolved.
• Hydrogen peroxide breaks down to form water and oxygen gas along with heat energy.
• When hydrogen sulphide gas is passed through a blue solution of copper sulphate, a black precipitate of copper sulphide is obtained and the sulphuric acid so formed remains in the solution.
• A white insoluble precipitate of barium sulphate is formed.
In simple words: These are observations. If you do these experiments, you will see yellow powders, brown smoke, or black and white solids appearing.

📝 Teacher's Note: Observation-based questions are common. For the H2S/Copper sulphate reaction, emphasize that the black precipitate is Copper Sulphide (\( \text{CuS} \)).

🎯 Exam Tip: Always describe both the color and the state (gas/solid/liquid) when writing an observation.

Question 10.
Answer: • Sodium carbonate
• Sodium nitrate
• Zinc carbonate
• Lead nitrate
In simple words: These are common salts used in many of the chemical reactions described in this chapter.

📝 Teacher's Note: Most nitrates are soluble in water, but many carbonates (except for group 1 metals) are insoluble. This is a key pattern to help students predict precipitates.

🎯 Exam Tip: Know your solubility rules! They help you identify which product will be the precipitate in a double decomposition reaction.