Kerala Syllabus Class 9 Chemistry Chapter 5 Chemical Kinetics Notes Solutions

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Kerala SCERT Class 9 Chemistry Chapter 5 Solutions Chemical Kinetics

Kerala Syllabus Std 9 Chemistry Chapter 5 Chemical Kinetics Notes Solutions Questions and Answers

Class 9 Chemistry Chapter 5 Let Us Assess Answers Chemical Kinetics

Question 1.
Classify the following chemical reactions into combination reaction, decomposition reaction, double decomposition reaction and displacement reaction.
Zn + FeSCO₄ → ZnSO₄ + Fe
NaCl + AgNO₃ → NaNO₃ + AgCl
2Pb(NO₃)₂ → 2PbO + 4NO₂ + O₂
H₂ + I₂ → 2HI
Pb(NO₃)₂ + 2KI → PbI₂ + 2KNO₃
NH₄NO₂ → 2H₂O + N₂

Answer:

Question 2.
Write the balanced chemical equations of the chemical reactions given below and classify them into combination reactions, decomposition reactions and displacement reactions.
a. Magnesium + Nitrogen → Magnesium nitride
b. Zinc carbonate \(\xrightarrow{\text { Heat }}\) Zinc oxide + Carbon dioxide
c. Aluminum + Lead nitrate → Aluminum nitrate + Lead
Answer:
a. 3Mg + N₂ → Mg₃N₂ [Combination reaction]
b. 2ZnCO₃ → 2ZnO + 2CO₂ [ Decomposition Reaction]
c. 2Al + 3Pb(NO₃)₂ → 2Al(NO₃)₃ + 3Pb [Displacement Reaction]

Question 3.
Identify the types of chemical reactions given below.
i. Formation of black copper oxide on heating copper powder in a China dish
ii. Silver nitrate solution reacts with sodium chloride solution to form silver chloride and sodium nitrate solution.
iii. Formation of ferric oxide, Sulphur dioxide and Sulphur trioxide on heating ferrous sulphate granules in a test tube.
Answer:
i. Combination Reaction
Copper (Cu) combines with oxygen (O₂) from the air to form copper (II) oxide (CuO).

ii. Double Displacement Reaction
The silver ions (Ag⁺) and chloride ions (Cl^–) exchange take place, forming silver chloride (AgCl) and sodium nitrate (NaNO₃).

iii. Decomposition Reaction
Ferrous sulphate (FeSO₄) decomposes upon heating to form ferric oxide (Fe₂O₃), sulphur dioxide (SO₂), and sulphur trioxide (SO₃).

Question 4.
What are displacement reactions? Write an example
Answer:
Displacement reactions are chemical reactions where one element or ion is replaced by another in a compound.
Zn + CuSO₄ → ZnSO₄ + Cu
In this reaction, zinc (Zn) displaces copper (Cu) from copper sulphate (CuSO₄), resulting in the formation of zinc sulphate (ZnSO₄) and copper.

Question 5.
Write two examples each of chemical reactions that take days or months to complete and those that take just a few seconds or minutes.
Answer:
Reactions That Take Days or Months to Complete:

• Rusting of Iron:
  Iron reacts with oxygen and moisture oVer time to form iron oxide (rust). This process can take days to years, depending on environmental conditions.
• Fermentation:
  The fermentation of sugars by yeast to produce alcohol (ethanol) can take several days to weeks. For example, in winemaking, the process typically lasts from a few days to several weeks.

Reactions That Take Just a Few Seconds or Minutes:

• Combustion of a Fuel:
  The burning of gasoline in an engine occurs very quickly, usually in seconds. For example, igniting a lighter or a matchstick produces a flame almost instantaneously.
• Neutralisation Reaction:
  The reaction between hydrochloric acid (HCl) and sodium hydroxide (NaOH) to form sodium chloride (NaCl) and water (H₂O) is rapid, often completing within minutes.

Question 6.
How does the rate of a chemical reaction vary when a gas reacts with large pieces of a solid substance and its fine powder? Explain the reason.
Answer:
The rate of a chemical reaction depends on the surface area.

Large Pieces of Solid:
Only the outer surface reacts with the gas, which slows down the reaction because there’s less area for the gas to interact with.

Fine Powder:
The fine powder has a much larger surface area, allowing more gas molecules to collide with it at once. This speeds up the reaction.
The surface area of a reactant significantly affects the rate of a chemical reaction. A larger surface area leads to more frequent collisions between reactant molecules, which, in turn, increases the rate of the reaction.

Question 7.
The rate of a chemical reaction for a fixed amount of a reactant is slow at room temperature.
Is it possible to increase the rate of the chemical reaction without changing the temperature? Explain.
Answer:
Yes, it is possible to increase the rate of a chemical reaction without changing the temperature. Here are a few methods to achieve this:

• Increase Surface Area
• Increase Concentration
• Stirring of Agitation
• Use of Catalysts

These methods can enhance the reaction rate without raising the temperature.

Question 8.
Zinc reacts with hydrochloric acid to produce hydrogen gas.
a. Write the chemical equation of this reaction.
b. What type of chemical reaction is this?
c. Suggest any two methods for increasing the amount of hydrogen.
Answer:
a. Zn + 2HCl → ZnCl₂ + H₂

b. This is a displacement reaction because zinc displaces hydrogen from hydrochloric acid.

c. Increase the Concentration of Hydrochloric Acid:
Using a more concentrated solution of hydrochloric acid increases the availability of H⁺ ions, leading to more frequent collisions and a higher rate of hydrogen gas production.

Increase the Surface Area of Zinc:
Using zinc in the form of a fine powder instead of larger pieces will increase the surface area, allowing more collisions with hydrochloric acid and producing more hydrogen gas.

Question 9.
The chemical equation of the conversion of hydrogen peroxide to water and oxygen is given.
2H₂O₂ → 2H₂O + O₂
a. What type of chemical reaction is this?
b. Which substance is used to accelerate the decomposition of H₂O₂? By what name is it commonly known?
Answer:
a. This is a decomposition reaction. In this type of reaction, a single compound (hydrogen peroxide) breaks down into two or more products (water and oxygen).

b. The substance commonly used to accelerate the decomposition of hydrogen peroxide is manganese dioxide (MnO₂). It is often referred to simply as a catalyst. Catalysts speed up chemical reactions without being consumed in the process.

Question 10.
Solutions of ammonium chloride and sodium hydroxide are taken in a test tube and heated.
a. Which gas is produced?
b. Which factor helps to increase the rate of the chemical reaction?
c. What is threshold energy?
d. Why does the rate of a chemical reaction increase as temperature increases?
Answer:
a. When solutions of ammonium chloride (NH₄Cl) and sodium hydroxide (NaOH) are heated, ammonia gas (NH₃) is produced along with water and sodium chloride (NaCl). The reaction can be represented as follows:
NH₄Cl + NaOH → NH₃(g) + H₂O + NaCl

b. One factor that helps to increase the rate of this chemical reaction is increasing the temperature. Heating the reactants speeds up the movement of particles, leading to more frequent and energetic collisions.

c. Threshold energy, also known as activation energy, is the minimum amount of energy required for reactants to undergo a chemical reaction. Only particles with energy equal to or greater than this threshold can effectively collide and react to form products.

d. The rate of a chemical reaction increases as temperature increases because:

• Increased Kinetic Energy: Higher temperatures give reactant particles more kinetic energy, causing them to move faster.
• More Frequent Collisions: Faster-moving particles collide more often, which increases the rate of successful collisions.
• Greater Energy of Collisions: Increased temperature means that a greater proportion of the particles have energy equal to or greater than the threshold energy, leading to more reactions occurring.

Extended Activities

Question 1.
Carry out the reaction between dilute hydrochloric acid and magnesium, following the experimental procedure given below.

Take 25 mL of dilute HCl in a conical flask. Place a 15 cm long Mg ribbon in the flask and start a stopwatch. Record the volume of hydrogen gas produced every 30 seconds, Swirl the flask slowly at regular intervals to ensure continuous reaction of the reactants. After fhe experiment, plot a graph between the volume of hydrogen on the ‘Y’ axis and time on the ‘X’ axis.

i. When is the rate of the chemical reaction highest in this experiment?
ii. How can the variation in the rate of the chemical reaction be explained?
iii. Repeat the above experiment using 25 mL of concentrated HCl instead of dilute HCl. Draw the graph. What change is noticed in the rate of the chemical reaction?
iv. Repeat the experiment by placing a 5 cm long Mg ribbon in 25 mL hydrochloric acid. Then, repeat the experiment by replacing the Mg ribbon with magnesium powder of equal mass. Compare the rates of chemical reactions using the graph.
v. Heat the system to 50°C. Then, put a 5 cm long Mg ribbon and record the volume of hydrogen every 30 seconds. Explain how the change in temperature affects the rate of chemical reaction by plotting the graph.
Answer:
i. The rate of the reaction is highest at the beginning when magnesium is freshly exposed to the acid.
Initially, the surface area available for reaction is maximised, and both reactants are at their highest concentrations.

ii. The rate decreases over time as:

• The concentration of HCl decreases as it reacts with Mg.
• The magnesium surface may become coated with magnesium chloride, hindering further reaction.
• The reaction is exothermic, but the initial conditions and reactant concentrations have the most impact on the rate.

iii. When using concentrated HCl:

• The rate of hydrogen gas production will increase significantly.
• The graph will show a steeper initial slope compared to the dilute HCl graph due to a higher concentration of HCl leading to more frequent collisions between reactant molecules.

iv. 5 cm Mg ribbon: The surface area is less than that of the magnesium powder, resulting in a slower reaction rate. The graph will show a lower rate of hydrogen production.
Magnesium powder: This increases the surface area exposed to the acid, leading to a rapid reaction and a steeper slope in the graph compared to the ribbon. The volume of hydrogen gas produced over time will be greater with the powder.

Comparing two reactions – the greater the surface area, the greater the rate of reaction.
v. Heating the system to 50°C increases the kinetic energy of the molecules, resulting in more frequent and effective collisions.
You would observe an increased volume of hydrogen gas produced over time, reflected in a steeper slope on the graph compared to the reaction at room temperature.

Comparing two reactions – the higher the temperature, the greater the rate of reaction.

Question 2.
Take an equal volume of potassium persulphate (K₂S₂O₈) solution in two test tubes. Heat one of the test tubes. Pour an equal amount of potassium iodide solution into both the test tubes.
a. In which test tube does iodine precipitate rapidly and form a brown colour?
b. Write the chemical equation of this reaction.
c. What is the effect on the rate of reaction by adding a little manganese dioxide to the unheated test tube?
d. What is the role of manganese dioxide here?
Answer:
a. Iodine precipitates rapidly and forms a brown colour in the test tube that was heated. Heating increases the reaction rate due to higher kinetic energy, leading to more frequent and effective collisions between reactants.

b. The reaction can be represented as follows
2KI + K₂S₂O₈ → 2K₂SO₄ + I₂

c. Adding manganese dioxide (MnO₂) to the unheated test tube increases the rate of reaction. MnO₂ acts as a catalyst, speeding up the formation of iodine without being consumed in the reaction.

d. Manganese dioxide acts as a catalyst in this reaction. It provides an alternative pathway with a lower activation energy, allowing the reaction to proceed more quickly, even at lower temperatures. This results in a more rapid formation of iodine, contributing to a more pronounced brown colour in the solution.

Question 3.
Observe the rate of formation of hydrogen iodide by passing a fixed amount of hydrogen gas at room temperature into the iodine crystal taken in a closed container. Repeat this experiment by heating up to 130°C. What change occurs in the rate of the chemical reaction? Explain.
Answer:
At Room Temperature:
Rate of Reaction: The reaction will proceed but at a relatively slow rate due to the lower kinetic energy of the molecules. The collisions between hydrogen and iodine molecules will happen less frequently and with less energy, resulting in a gradual formation of hydrogen iodide.

At 130°C:
Rate of Reaction: When the temperature is raised to 130°C, the rate of formation of hydrogen iodide significantly increases. This is due to the following reasons:

• Increased Kinetic Energy: Heating the system increases the kinetic energy of the reactant molecules, leading to more frequent and more energetic collisions.
• Higher Reaction Rates: The increased energy helps overcome the activation energy barrier more effectively, resulting in a faster reaction rate and more rapid formation of HI Heating the system to 130°C accelerates the reaction between hydrogen and iodine, resulting in a faster formation of hydrogen iodide compared to room temperature conditions.

Chemical Kinetics Class 9 Notes Questions and Answers Kerala Syllabus

Question 1.
Take some sugar (C₁₂H₂₂O₁₁) in a spatula and heat it on a bunsen burner.
What is the colour of the product obtained?
Answer:
Black
The product formed is carbon. On strong heating, sugar decomposes to form carbon and water.
C₁₂H₂₂O₁₁ + Heat → 12C + 11H₂O

Question 2.
Take some sodium carbonate (Na₂CO₃) in a test tube. Add 2 – 3 mL of dilute hydrochloric acid to it.
What is your observation?
Answer:
Brisk effervescence is observed, indicating the presence of CO₂ gas.
Na₂CO₃ + 2HCl → 2NaCl + H₂O + CO₂

Question 3.
Add 2 – 3 mL of potassium iodide solution to 2 – 3 mL of lead acetate (Pb(CH₃COO₂) solution in a test tube.
a. In which part of the test tube is the product found?
Answer:
At the bottom as a precipitate.

b. What is its colour?
Answer:
Yellow
The product formed in yellow colour is lead iodide.
Pb(CH₃-COO)₂ + 2KI → PbI₂ + K₂(CH₃-COO)₂

Question 4.
Take some iron sulphide of a test tube. Add 2 – 3 mL of dilute hydrochloric acid to it and heat, Is the smell familiar?
Answer:
When dilute hydrochloric acid is heated with iron sulphide, a gas is produced that smells like rotten eggs. This gas is hydrogen sulphide.
FeS + 2HCl → FeCl₂ + H₂S
Inferences of the activities given above

Experiment	How the chemical change is identified
1	The colour of the sugar changed
2	A gas is liberated
3	A precipitate is formed
4	A pungent gas is produced

Question 5.
What do you observe?
Answer:
The outer side of the beaker feels hot.
Here, calcium oxide reacts with water to form calcium hydroxide (Ca(OH)₂). It is an exothermic reaction as well.

Question 6.
Write the chemical equation of the reaction.
Answer:
CaO + H₂O → Ca(OH)₂ + heat
An exothermic reaction is a chemical reaction that releases heat.
Eg: Burning of fuel
Endothermic reactions are chemical reactions that involve the absorption of heat.
Eg: CaCO₃ + Heat → CaO + CO₂

Question 7.
Which are the substances combined in the following reactions? What are the products obtained as a result of these reactions?
a. N₂ + 3H₂ → 2NH₃
b. 2SO₂ + O₂ → 2SO₃
Answer:

Reaction	Substances combined	Products formed
a	N2, H2	NH3
b	2SO2, O2	SO3

The reaction in which two or more simple substances (elements/compounds) combine to* form a compound is called combination reaction.
More examples of combination reaction

• 2H₂ + O₂ → 2H₂O
• 2Na + Cl₂2 → 2NaCl
• C + O₂ → CO₂
• H₂ + Cl₂ → 2HCl
• 2Mg + O₂ → 2MgO
• S + O₂ → SO₂

Question 8.
What change is observed?
Answer:
Ammonium dichromate bums in a purple (pink) colour.
When ammonium dichromate undergoes thermal decomposition, chromium trioxide (Cr₂O₃), water vapour and nitrogen are formed.
(NH₄)₂Cr₂O₇ → Cr₂O₃ + 4H₂O + N₂↑

Question 9.
Examine the following chemical reactions.
a. 2Pb(NO₃)₂ → 2PbO + 4NO₂↑ + O₂↑
b. CaCO₃ → CaO + CO₂↑

Which substances underwent decomposition? What are the products obtained?
Answer:

Reaction	Substances that underwent decomposition	Products obtained
a	Pb(NO3)2	PbO, NO2, O2
b	CaCO3	CaO, CO2

Decomposition reaction is the process by which a compound breaks down into two or more substances

More examples of decomposition reaction

• 2HgO → 2Hg + O₂
• 2H₂O → 2H₂ + O₂
• C₁₂H₂₂O₁₁ → 12C + 11H₂O
• Ca(OH)₂ → CaO + H₂O
• 2HI → H₂ + I₂
• 2Al(OH)₃ → Al₂O₃ + 3H₂O
• NH₄Cl → NH₃ + HCl

Question 10.
Heat a small amount of copper carbonate (CuCO₃) in a boiling tube. What change is observed? What are the products obtained? Write the balanced chemical equation of the chemical reaction.
Answer:
CO₂ gas is produced. Blue copper carbonate turns into black-coloured cupric oxide.
CuCO₃ + Heat → CuO + CO₃

Question 11.
What do you observe?
Answer:
A white curdy precipitate is formed.

Question 12.
Examine the chemical equation of this reaction given below.
NaCl + AgNO₃ → AgCl↓ + NaNO₃
Can you identify the product precipitated as a result of this reaction?
Answer:
Silver chloride is the insoluble precipitate formed.
(Sodium nitrate is soluble in solution)

Question 13.
In one of the reactants, sodium chloride, which ion is bonded to the sodium ion?
Answer:
Chloride ion.

Question 14.
Among the products formed, with which metal ion does the chloride ion combine?
Answer:
Silver ion.

Question 15.
Consider the nitrate ion in the second reactant silver nitrate. With which metal ion does the nitrate ion combine to form the product?
Answer:
Sodium ion.
In the reaction between sodium chloride and silver nitrate, ions are exchanged.

Double decomposition reaction is a reaction in which two compounds, reacting with each other, interchange their ions to form two new compounds.

Double decomposition reactions can be classified into three types.
1. Precipitation reactions – Precipitates are insoluble solid compounds which can be separated from the solution.
MgCl₂ + H₂SO₄ → MgSO₄ + 2HCl

2. Reactions that involve formation of gas.
CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂

3. Reactions that involve formation of products among which at least one compound does not dissociate into ions.
The neutralisation reaction between an acid and a metal hydroxide base to give water and salt is an example of such a chemical reaction.
HCl + NaOH → NaCl + H₂0

More examples of double decomposition reaction

• Na₂SO₄ + BaCl₂ → BaSO₄ + 2NaCl
• HCl + NaOH → NaCl + H₂O
• CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂
• MgCl₂ + H₂SO₄ → MgSO₄ + 2HCl
• 2NaCl + H₂SO₄ → Na₂SO₄ + 2HCl
• 2NaNO₃ + H₂SO₄ → Na₂SO₄ + 2HNO₃

Question 16.
What do you observe?
Answer:
The reaction bubbles vigorously as hydrogen gas is produced.

Question 17.
Which gas burns with a ‘pop’ sound?
Answer:
Hydrogen gas bums with a pop sound.

Question 18.
Notice the balanced chemical equation of this chemical reaction, given below.
Zn + 2HCl → ZnCl₂ + H₂↑
During this chemical reaction, which atom replaces hydrogen in the reactant, HCl, to form the product?
Answer:
The zinc atom replaces the hydrogen atom in hydrochloric acid.
Reactions in which an element in a compound is displaced by another element are called displacement reactions
More examples of displacement reactions

• Mg + 2HCl → MgCl₂ + H₂
• Zn + H₂SO₄ → ZnSO₄ + H₂
• CuSO₄ + Fe → FeSO₄ + Cu
• CuSO₄ + Zn → ZnSO₄ + Cu
• Fe₂O₃ + Al → 2Fe + Al₂O₃

Question 19.
Classify the following chemical reactions into combination reaction, decomposition reaction, double decomposition reaction and displacement reaction. Also, include more examples.
1. Zn + 2AgNO₃ → Zn(NO₃)₂ + 2Ag
2. CaCO₃ → CaO + CO₂
3. 2HI → H₂ + I₂
4. Cu(NO₃)₂ + Na₂S → CuS + 2NaNO₃
5. SO₂ + H₂O → H₂SO₃
Answer:
Combination reactions
SO₂ + H₂O → H₂SO₃
N₂ + 3H₂ → 2NH₃
2CO + O₂ → 2CO₂
NH₃ + HCl → NH₄Cl
2H₂ + O₂ → 2H₂O

Decomposition reactions
2HI → H₂ + I₂
CaCO₃ → CaO + CO₂
NH₄Cl → NH₃ + HCl
2H₂O → 2H₂ + O₂
CuCO₃ → CuO + CO₂

Double decomposition reactions
Cu(NO₃)₂ + Na₂S → CuS + 2NaNO₃
CuSO₄ + BaCl₂ → CuCl₂ + BaSO₄
Na₂SO₄ + BaCl₂ → 2NaCl + BaSO₄
Pb(NO₃)₂ + 2KI → PbI₂ + 2KNO₃

Displacement reactions
Zn + 2AgNO₃ → Zn(NO₃)₂ + 2Ag
CuSO₄ + Fe → FeSO₄ + Cu
Zn + 2HCl → ZnCl₂ + H₂
Mg + H₂SO₄ → MgSO₄ + H₂

Question 20.
What measures do we take to prevent the spoilage of food?
Answer:

• Preserve using salt.
• Use of preservatives.
• Keep refrigerated.
• Sterilisation
• Dry storage

All chemical reactions do not take place at the same speed. There are several factors that affect the speed of a chemical reaction.

Question 21.
Take an equal volume of dilute hydrochloric acid in three test tubes. Place pieces of copper (Cu), zinc (Zn), and magnesium (Mg) of the same size in each of the three test tubes.
Observation:
Test tube 1: -No reaction
Test tube 2: – A moderate reaction will be observed with zinc. Gas evolution will be slower than with magnesium, and the metal will dissolve at a slower rate.
Test tube 3: – The most vigorous reaction will occur with magnesium. You will observe rapid gas evolution, and the metal will dissolve quickly.

a. Which test tube shows the highest rate of chemical reaction?
Answer:
Test tube 3
Here, since the metals are of the same size and the volume of hydrochloric acid is equal, it is clear that the rate of chemical reaction is different due to the characteristic properties of metals like copper, zinc and magnesium.

b. Why does magnesium react very slowly with water while sodium reacts vigorously with water at room temperature?
Answer:
The rate of chemical reaction depends on the nature of the reactants. Sodium is a more reactive metal than magnesium. So sodium reacts strongly with water even at normal temperature.

Question 22.
Take an equal volume of concentrated hydrochloric acid and dilute hydrochloric acid in two test tubes.
a. If so, which test tube contains a greater number of HCl molecules per unit volume?
Answer:
The test tube containing concentrated HCl contains a greater number of HCl molecules Place magnesium ribbons of equal mass in both test tubes.
Record the observation of the chemical reaction taking place in each test tube.

Test tube 1: – Vigorous reaction: Immediate and intense bubbling occurs.
Test tube 2: – Slower reaction: Bubbling is observed, but it is less intense than with concentrated HCl.

b. Which test tube shows the higher rate of a chemical reaction?
Test tube 1
Answer:

• Because the number of reactant molecules per unit volume is greater in a test tube containing concentrated hydrochloric acid.
• As the concentration of reactant increases, the number of particles per unit volume increases, and
  consequently, the number of effective collisions increases. Due to this, the rate of chemical reaction also increases.

Question 23.
If steel wool is heated in air, it will become red hot and get oxidised slowly. However, if the heated steel wool is exposed to an oxygen-rich environment, it will burn strongly and produce the product at a faster rate. What is the reason for this?
Answer:
This is due to the concentration of oxygen which is one of the properties. As the density increases, the number of effective collisions increases, and the rate of chemical reaction increases.

a. Which are the reactants in this chemical reaction?
Answer:
Piece of marble, Hydrochloric acid (HCl)

b. What are their physical states?
Answer:
Piece of marble: Solid state
Hydrochloric acid (HCl): Liquid
Let us write the chemical equation of this reaction.
CaCO₃(s) + 2HCl(l) → CaCl₂(s) + H2O(1) + CO(g)
Experiment:

Question 24.
Now, take a piece of marble and powdered marble of equal mass. Treat them with dilute hydrochloric acid of equal volume and same concentration. Record the observations.

a. Is the rate of the reaction the same in both the beakers?
Answer:
No

b. Is the surface area of the marble equal?
Ans:
No

c. What change occurs in the rate of collision between the reactant molecules as the surface area increases?
Answer:
The surface area of a reactant significantly affects the rate of a chemical reaction. A larger surface area leads to more frequent collisions between reactant molecules, which, in turn, increases the rate of the reaction.

d. What is the change in the rate of reaction if the marble pieces are ground to fine powder?
Answer:
If marble pieces are ground to fine powder, the rate of reaction will increase significantly. This is
because grinding the marble into a fine powder increases its surface area. With a larger surface area, there are more points of contact between the marble and the hydrochloric acid, leading to more frequent collisions between the reactant molecules. This increased collision rate results in a faster reaction.

Example: Wood bums quickly when it is cut into small pieces. The rate of chemical reaction increases with an increase in surface area.

Physical state of reactants and rate of reaction
The rate of a chemical reaction is related to the physical state of the reactants. When zinc sulphate and barium nitrate in a solid state are mixed, no noticeable reaction takes place between them. But when their aqueous solution is mixed, a white precipitate of barium sulphate is immediately formed. This chemical reaction occurs due to the effective collision between the molecules. That is, the physical state of the reactants plays an important role in determining the rate of reaction. This explains why gasoline vapour ignites more explosively than liquid gasoline.

Question 25.
Take dilute hydrochloric acid in two test tubes. Put a large piece of eggshell in the first test tube. In the second test tube, add an equal amount of powdered eggshell.

a. In which test tube is the rate of chemical reaction found to be higher?
Answer:
The rate of chemical reaction will be higher in the test tube containing powdered eggshell.

b. What is the reason for this?
Answer:
The reason for the faster reaction with powdered eggshell is increased surface area. When the eggshell is powdered, its surface area is significantly larger than that of a large piece. This increased surface area allows for more contact between the eggshell and the hydrochloric acid, leading to ‘ more frequent collisions between the reactant molecules and a faster reaction rate.

c. Write the chemical equation of the chemical reaction taking place here.
Answer:
The chemical reaction between eggshell (calcium carbonate, CaCO₃) and hydrochloric acid (HCl) produces calcium chloride (CaCl₂), water (H₂O), and carbon dioxide gas (CO₂). The balanced chemical equation is:
CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g)

Question 26.
Firecrackers used to enrich celebrations contain magnesium. Magnesium gives off a rich glow when it burns. Magnesium powder is used here. What could be the reason?
Answer:
Magnesium powder is used in firecrackers because it bums very brightly and produces a dazzling white light.

Here’s why:

• High reactivity: Magnesium is a highly reactive metal, meaning it readily reacts with oxygen in the air.
• Exothermic reaction: The reaction between magnesium and oxygen is highly exothermic, releasing a large amount of heat and light.
• Bright flame: The intense heat causes the magnesium to bum with a brilliant white flame.
• Fine powder: Using magnesium powder increases the surface area, which leads to a faster and more intense reaction, resulting in a brighter and more dazzling light.

Question 27.
Temperature is the measure of the hotness or coldness of a matter.
Let us examine the reaction between ammonium chloride (NH₄Cl) and sodium hydroxide (NaOH).
a. What are the products obtained when these compounds react?
Answer:
Sodium chloride (NaCl), Water (H₂O) and Ammonia gas (NH₃)
NH₄Cl + NaOH → NaCl + H₂O + NH₃↑
Experiment:
Take 2 mL of ammonium chloride solution in a test tube and pour into it 2 mL of sodium hydroxide solution taken in another test tube.

b. Did you observe the formation of any product?
Answer:
Yes

c. Wave your hand so that air from the mouth of the test tube moves towards your nose.
Do you experience any smell?
Answer:
A slight odour of ammonia.

d. Heat the test tube and check the smell again. Write your observation.
Answer:
The pungent smell of ammonia is obtained when the test tube is heated. That is, the increase in the rate of reaction is due to the increase in temperature.

Question 28.
The procedure for the chemical reaction between sodium thiosulphate (Na₂S₂O₃) and hydrochloric acid (HCl) is given below. Identify the factor that affects the rate of a chemical reaction by doing the experiment.
Prepare a dilute solution of sodium thiosulphate in a beaker. Take an equal amount of this solution in two test tubes. Heat one of the test tubes for a short while. Pour an equal volume of dilute hydrochloric acid into both test tubes.
Record your observations.
Answer:
Appearance of the cloudy precipitate.

a. In which test tube did the precipitate form quickly?
Answer:
The faster appearance of the cloudy precipitate in the heated test tube indicates that the reaction rate is higher at higher temperatures. This demonstrates that temperature is a factor that affects the rate of a chemical reaction.

b. Which factor influenced the rate of the chemical reaction here and write the chemical equation for the reaction involved.
Answer:
Temperature influences the rate of the chemical reaction.
The experiment showed that the heated test tube produced a cloudy precipitate more quickly than the one at room temperature. This indicates that increasing the temperature accelerated the reaction.
Na₂S₂O₃ + HCl → 2NaCl + S + H₂O + SO₂

Let us recognise the effect of temperature on some of the reactions around us,

• The dough that we grind to make dosa and idli can be seen to ferment quickly at normal temperature, but if the dough is kept in the refrigerator, it will rise slowly.
• You know that light worms can emit light by the chemiluminescence effect. Light sticks mixed with certain chemicals can also emit light. Such light sticks can emit intense light at high temperatures.

Question 29.
Take some hydrogen peroxide solution in a boiling tube. Insert a burning incense stick in the boiling tube.
Let us write the chemical equation of this reaction.
2H₂O₂ → 2H₂O + O₂ ↑
a. Is there any difference in the burning of the incense stick?
Answer:
When you insert a burning incense stick into a boiling tube containing hydrogen peroxide, you’ll likely observe a more vigorous and brighter flame compared to its burning in air alone.
Now, add a little manganese dioxide (MnO₂) into the boiling tube and insert the burning incense stick in the tube again.

b. What is observed now?
Answer:
It can be seen that the speed of combustion of the incense stick has increased. This is because the
rate of decomposition of hydrogen peroxide increases, and more oxygen is released.
Manganese dioxide is used as a catalyst in the decomposition of hydrogen peroxide. Here, the decomposition of hydrogen peroxide is accelerated by the substance manganese dioxide.

c. Can you identify the catalysts in the following chemical reaction
i. Manufacture of ammonia by Haber process.

Answer:
The catalyst used in the Haber process for the manufacture of ammonia is iron (Fe)

ii. Manufacture of sulphuric acid by contact process

Answer:
The catalyst used in the Contact Process is vanadium pentoxide (V₂O₅).

Catalysis in nature

• Enzymes are protein molecules that act as catalysts in living cells.
• They accelerate chemical reactions at normal temperature and pressure.
• Amylase is an example of an enzyme found in saliva.
• It converts starch into maltose.

Catalyst and Magic
Catalyst and Sugar Candy

• Sugar candy melts when heated but doesn’t bum.
• Adding ash to the suger candy before heating can cause it to bum.
• Metal compounds in the ash act as catalysts for the combustion of the sugar candy.

Bombardier Beetle Defense

• Bombardier beetles use catalytic decomposition of hydrogen peroxide for self-defence.
• Enzymes produced by the beetle accelerate this reaction, releasing steam and other irritating chemicals.