Chemical reactions & equations
Chemical reactions and equations- 10th CBSE- Physical science
The chapter on chemical reactions and equations seeks to familiarize students with the various chemical transformations that occur in the natural world. These reactions include diverse phenomena such as food digestion, photosynthesis, or the rusting of iron. Understanding these chemical processes occurring in their surroundings sparks curiosity about chemistry, prompting students to delve into the intricacies ( complexities) of balanced chemical equations that depict them.
What is a chemical reaction?
A chemical reaction refers to a transformation
in which reactants combine under specific conditions of temperature and
pressure, with or without the presence of catalysts, resulting in the creation
of products.
During a chemical reaction, the bonds between
the reactant molecules are disrupted, leading to the formation of new bonds
within the resulting product molecule.
When a chemical reaction takes place, several
noticeable changes occur:
- Alteration in the physical state of the chemical compounds involved
- Modification in the color of the substances participating in the reaction
- Release of gases or residues
- Variation in temperature throughout the reaction process.
Examples of common chemical reactions
encountered in daily life include the conversion of milk into curd, the
fermentation of grapes, the ionization of common salt, and the preparation of
crispy/spongy recipes using baking soda, among others.
Important question & answers on chemical reaction?
1. What is a chemical reaction? Describe one activity each to show that a chemical change has occurred in which (i) change of color and (ii) change in temperature has taken place
[2024, AISSE]
Answer:
A chemical reaction expresses the chemical change occurring as one or
more reactants transform into products with new properties.
When the gray-white colored magnesium
ribbon burns in air, it imparts a characteristic dazzling white color to the
flame along with the formation of white powder of Magnesium oxide.
2Mg (s) + O2 (g) ---------→ 2 MgO (s)
Magnesium oxide
The reaction of
magnesium with atmospheric oxygen to form magnesium oxide is exothermic,
releasing a significant amount of heat energy and emitting a dazzling white
light. Consequently, this chemical reaction involves a temperature change
during the process of chemical transformation.
What is a chemical equation?
A chemical equation serves as a concise
representation of a genuine chemical reaction through the utilization of
symbols and formulas.
Essential components of a chemical equation:
- All reactants and products involved in the chemical reaction are expressed using symbols and formulas.
- It must accurately depict a chemical reaction that can be conducted in a laboratory setting.
- The equation must be balanced, ensuring that the number of atoms of each element is equivalent on both sides of the equation.
Information conveyed by a chemical equation:
- It identifies the reactants participating in the chemical reaction and the resultant products formed after the chemical transformation.
- The equation specifies the quantity of molecules of the reactants required and those of the products generated.
Activity-1.1:
This activity aims to
observe the chemical change that occurs when a clean Magnesium ribbon undergoes
combustion in the air.
Experiment:
- Begin by cleaning a 4-5 cm long magnesium ribbon using sandpaper.
- Ignite the Bunsen burner or spirit lamp.
- Using a pair of tongs, hold the Magnesium ribbon over the flame of the burner to initiate burning.
- Carefully collect the resulting white ash on a watch glass.
Observation:
The Magnesium ribbon burns with a brilliant white flame, a distinctive characteristic confirming Magnesium in the flame test.
Additionally, the grey-white colored Magnesium
ribbon transforms into white powder upon heating.
Chemical reaction:
Magnesium + Oxygen → Magnesium oxide
2Mg + O2 → 2 MgO
Grey-white White powder
Conclusion:
When heated in the presence of atmospheric oxygen, Magnesium undergoes a reaction to form magnesium oxide.
This process involves oxidation, wherein Magnesium gains
oxygen.
Important questions
and answers on Activity-1.1:
1. Why is it necessary to
clean a Magnesium ribbon before burning it in air?
Answer:
Magnesium is a highly
reactive metal that forms a protective layer of Magnesium oxide when exposed to
atmospheric oxygen, which prevents further reaction. To ensure complete burning
and avoid interference from the protective layer, the Magnesium ribbon must be
cleaned by removing the magnesium oxide layer with sandpaper before burning.
2. What occurs when a
Magnesium ribbon burns in air?
Answer:
Burning a Magnesium
ribbon in air produces a dazzling white flame and results in the formation of
white powder, which is Magnesium oxide.
3. Provide a balanced
chemical equation for the burning of Magnesium in air and explain.
Answer:
The combustion of
Magnesium in air involves an oxidation reaction with atmospheric oxygen,
leading to the formation of Magnesium oxide, represented as a white-colored
ash.
2Mg + O2 → 2 MgO
Magnesium oxide
4. Which type of change occurs when magnesium metal reacts with oxygen?
Answer:
When magnesium
undergoes a reaction with oxygen, it initiates a chemical change that leads to
alterations in the physical state and color of the chemical compounds formed,
accompanied by a change in temperature.
5. What kind of reaction
occurs during the formation of magnesium oxide?
Answer:
A chemical combination
reaction occurs between magnesium and oxygen, leading to the formation of
magnesium oxide. Furthermore, as magnesium metal acquires oxygen while burning
in the presence of atmospheric oxygen, the reaction is also referred to as an oxidation
reaction.
6. Despite the fact that the burning of magnesium requires an initial input of heat energy, it is still classified as an exothermic reaction. Explain
Answer:
This is due to the significant release of heat and light that occurs during the chemical combination reaction of Mg & O2. Although a small amount of heat is absorbed to initiate the burning of magnesium with atmospheric oxygen, the quantity of heat absorbed is minimal as compared to the substantial heat released during the chemical transformation. Therefore, despite the initial heat absorption, the overall process is exothermic.
Activity-1.2:
This activity is
designed to explore the chemical reaction between lead nitrate and potassium
iodide solutions under room temperature conditions.
Experiment:
Dissolve 1 g of white
lead nitrate solid in 3-4 ml of distilled water in a clean test tube to prepare
its solution. The resulting Pb(NO₃)₂ solution remains white.
Similarly, dissolve 1
g of white potassium iodide in 3-4 ml of distilled water in another test tube
to prepare its solution. The resulting KI solution is colorless.
Combine both solutions
in a 50 ml clean glass beaker and stir for a minute to ensure thorough mixing.
Yellow-colored
crystalline solid will be observed floating in the solution.
Filter the contents of
the beaker to separate them.
Observation:
The appearance of a yellow-colored solid confirms the formation of PbI₂.
Chemical reaction:
Lead nitrate + Potassium iodide → Lead iodide + Potassium nitrate
Pb(NO₃)₂ + 2KI →PbI₂ + 2KNO3
White Yellow
Conclusion:
The reaction between
lead nitrate and potassium iodide is a double displacement reaction, resulting
in the formation of potassium nitrate and lead iodide.
KNO₃ is a white-gray solid that dissolves in water, rendering it water-soluble and appearing as a colorless liquid in the beaker.
However, PbI₂, being water-insoluble, appears as a yellow crystalline salt in the beaker.
Since lead iodide precipitates
during the reaction, it is also categorized as a precipitation reaction.
Important questions
and answers on Activity-1.2:
1. Write a balanced
chemical equation for the following word equation:
Lead nitrate +
Potassium iodide → Lead iodide + Potassium nitrate
[2024, AISSE]
Answer:
Step-1: Formulating
the formulas of the compounds engaged in the balanced chemical equation.
Lead nitrate +
Potassium iodide → Lead iodide + Potassium nitrate
Pb (NO3)2
+ KI → PbI2 + KNO3
Step-2: Equalizing the
count of oxygen atoms on both sides of the reaction.
Pb (NO3)2
+ KI → PbI2 + 2KNO3
Step-3: Adjusting the
number of iodine atoms present on each side of the reaction.
Pb (NO3)2
+2KI → PbI2 + 2KNO3
Step-4: Scrutinizing
the chemical equation to ensure the proper balance of every compound's atoms on
both sides.
Step-5: Notating the
symbols denoting the physical states of the compounds participating in the
chemical reaction.
Pb (NO3)2 (aq) +2KI (aq) → PbI2 (s) + 2KNO3 (aq)
(a) Is this a double
displacement reaction? Justify your answer.
Answer:
(a) Yes, it is a double displacement reaction because both K⁺ and Pb²⁺ cations exchange their anions (NO₃⁻ and I⁻) during the reaction.
(b) Name the compound
precipitated and write the balanced chemical equation for the reaction involved.
During the reaction,
lead nitrate breaks down into Pb²⁺ and NO₃⁻ ions upon dissolution in water.
Pb(NO₃)₂ + H₂O → Pb²⁺
+ 2NO₃⁻
Similarly, KI
dissociates to release K⁺ and I⁻ ions.
KI + H₂O → K⁺ + I⁻
Subsequently, one Pb²⁺
ion combines with two iodide ions, leading to the formation of solid PbI₂.
Pb²⁺ + 2I⁻ → PbI₂
Activity-1.3:
This activity focuses
on the reaction between metals and acids.
Experiment:
Place 2g of silvery
grey zinc granules into a 250 ml conical flask.
Add 5 mL of dilute
sulphuric acid to the flask.
Immediately seal the
mouth of the conical flask with a rubber cork containing a single hole.
Insert a glass tube
into the hole of the rubber cork to collect the released gas without any loss.
Gently shake the
conical flask to ensure thorough mixing of the added chemical compounds.
Observation:
Following the reaction, a colorless zinc sulphate solution forms, accompanied by the release of hydrogen gas. We can verify the presence of hydrogen gas using a burning candle. Placing a burning candle at the mouth of the glass tube, hydrogen gas burns with a pop-up sound.
Similarly, the chemical reaction between zinc granules and dilute hydrochloric acid results in the formation of zinc chloride, accompanied by the release of hydrogen gas.
Therefore, this activity
serves to confirm that the reaction of any metal with acids involves the
formation of the corresponding metal salt along with the release of hydrogen
gas.
Chemical reaction:
Zinc + dil. Sulphuric
acid ---------------→ Zinc sulphate + hydrogen
Zn (s) + H2SO4 (aq) ----------------→ ZnSO4 (aq) + H2(g) ↑
Zinc + dil. Hydrochloric
acid ---------------→ Zinc chloride + hydrogen
Zn (s) + 2 HCl
(aq) ----------------→ ZnCl2 (aq) + H2(g) ↑
Important question and answers on activity-1.3:
1. Can metals less reactive than hydrogen displace it from acids?
Answer:
According to the
displacement reaction, metals with higher reactivity can displace metals with
lower reactivity from their salt solutions.
Therefore, metals such
as potassium, sodium, and magnesium, which are more reactive than hydrogen in
the reactivity series, can react with acids and water by displacing hydrogen
from them.
However, metals that
are below hydrogen in the reactivity series, such as copper, silver, gold, and
platinum, cannot displace hydrogen from acids and water. Consequently, these
metals do not react with acids.
2. Some metal react with acids to produce salt and hydrogen gas. Illustrate it with an example. How will test the presence of this gas?
[2024, AISSE]
Answer:
Metals more reactive
than hydrogen undergo reactions with acids, yielding the respective metal salt
and hydrogen gas.
For instance, when
zinc reacts with dilute hydrochloric acid, zinc chloride and hydrogen gas are
produced. This reaction involves the liberation of hydrogen gas.
Zinc + dil. Hydrochloric acid ---------------→ Zinc chloride + hydrogen
Zn (s) + 2 HCl (aq) ----------------→ ZnCl2 (aq) + H2(g) ↑
To confirm the
presence of hydrogen gas, we employ a burning candle. When a burning candle is
brought close to the hydrogen gas, it burns with a distinctive popping sound.
Activity-1.4:
This activity
demonstrates the production of slaked lime through the reaction of calcium
oxide with water.
Experiment:
Place a small quantity
of calcium oxide in a glass beaker.
Slowly add water to
the calcium oxide.
After a few minutes,
observe the beaker becoming warm, indicating the formation of Ca (OH)2.
Observation:
Upon adding a small
amount of water to the white calcium oxide solid, it reacts vigorously,
producing a white calcium hydroxide solution.
The reaction is
exothermic proceeds with the release of heat.
Chemical reaction:
Calcium oxide +
water -----------------→ Calcium
hydroxide + heat
CaO (s) + H2O
(l) ----------------→ Ca(OH)2 (aq) + heat
Conclusion:
When quicklime (CaO) reacts with water, it undergoes a process resulting in the formation of calcium hydroxide, known as slaked lime.
This reaction releases heat, indicating its
exothermic nature. Furthermore, since calcium hydroxide is the sole product
formed from the combination of calcium oxide and water, the reaction is
classified as a combination reaction.
Important question and answers on activity-1.4:
1. A solution of a
substance ‘X’ is used for whitewashing.
(i) Name the substance ‘X’ and write its formula
Answer:
The substance used for
whitewashing is quick lime and its chemical formula is CaO.
(ii) Write the
reaction of the substance ‘X’ named in (i) above with water
Answer:
CaO (s) + H2O
(l) --------------→ Ca(OH)2 (aq) + heat
Quick lime Lime water
Ca(OH)2 is
also known as lime water as it is formed from lime (CaO) and water
2. A clear solution of slaked lime is made by dissolving Ca(OH)2 in excess of water. If this solution is left exposed to air the solution slowly goes milky. Why a faint white precipitate is formed? Support your answer with the help of chemical equation.
Answer:
When a solution of
calcium hydroxide is left exposed to air, it interacts with the carbon dioxide
present in the air, resulting in the formation of calcium carbonate. This
reaction causes the solution's color to change to milky. The faint white
precipitate formed is calcium carbonate, which settles down as it is insoluble
in water.
Ca (OH)2 (aq)
+ CO2 (g) ------------→ CaCO3 (s) + H2O (l)
3. What happens when
excess carbon dioxide is passed through the lime water?
Answer:
Lime water reacts with
atmospheric carbon dioxide and moisture, leading to the formation of calcium
bicarbonate. This compound is colorless and soluble in water, hence it remains
in the form of an aqueous solution.
Ca (OH)2 (aq)
+ CO2 (g) ------------→ CaCO3 (s) + H2O (l)
CaCO3 (s) +
CO2 (g) + H2O (l) -----------------→ Ca(HCO3)2 (aq)
4. Describe why calcium carbonate is employed in whitewashing and write any one of its application other than whitewashing
Answer:
Calcium carbonate is
utilized in whitewashing due to its ability to create a bright, glossy, and
durable coating on walls.
Apart from
whitewashing, it serves as an antacid.
It is also employed to
raise calcium levels in the body.
5. Write the method of preparation of calcium hydroxide. What happens when carbon dioxide is passed through it? Write a balanced chemical equation for the reaction involved.
[2024, AISSE]
Answer:
Calcium hydroxide is
prepared from lime and water.
CaO (s) + H2O (l) --------------→ Ca(OH)2 (aq) + heat
When carbon dioxide
gas is passed through it, it gives lime stone (CaCO3)
Ca (OH)2 (aq) + CO2 (g) ------------→ CaCO3 (s) + H2O (l)
Activity 1.5:
It entails noting the chemical alterations observed when heating
ferrous sulfate crystals.
Experiment:
Place approximately 2 grams of ferrous sulfate crystals into a dry
boiling tube.
Upon heating, the green hue of the ferrous sulfate crystals
transforms to brown.
Additionally, a distinctive odor of burning sulfur becomes
noticeable.
Observation:
Upon heating, the green ferrous sulfate crystals lose their water
content and transform into a white substance, a phenomenon known as dehydration
or water of crystallization.
Upon further heating, the FeSO4 molecules decompose, forming a
brown-colored solid mass of ferric oxide along with the release of sulfur
dioxide and sulfur trioxide gases.
The distinct smell of burning sulfur observed during the heating
of ferrous sulfate crystals is primarily due to the release of sulfur dioxide
gas.
Chemical reaction:
Heat
Ferrous sulphate ---------------------→ Ferric oxide + Sulphur
dioxide + Sulphur trioxide
Heat
2FeSO4 (s) ----------------→ Fe2O3
(s) + SO2 (g) + SO3 (g)
Green Brown
The breakdown of ferrous sulfate into ferric oxide, sulfur
dioxide, and sulfur trioxide represents a chemical decomposition reaction.
Ferric oxide forms a solid product, while sulfur dioxide and
sulfur trioxide are produced as gases.
Decomposition reactions typically involve the absorption of
energy, and this particular reaction, known as thermal decomposition.
Because the formation of products occurs with the absorption of
heat energy, it is classified as an endothermic reaction.
Important question and answers on activity 1.5:
When subjected to heat, the green-colored ferrous sulfate crystals
undergo a change where they lose the water molecules incorporated in their
structure, transforming into a white solid. This process is commonly referred
to as dehydration or water of crystallization.
Δ
FeSO4, 7H2O --------------→ FeSO4
+ 7H2O
Green White
2. Write one equation each for decomposition reactions where energy is supplied in the form of heat, light or electricity
Answer:
Decomposition reactions occur with the absorption of various types of energy such as electricity, light, or heat.
When electricity is the energy
source involved in a decomposition reaction, it's referred to as electrolytic
decomposition. Likewise, thermal and photolytic decomposition reactions involve
the absorption of heat and light energies, respectively.
Electric current
2H2O (l)
------------------------→ 2H2 (g) + O2 (g) [Electrolytic decomposition]
Heat
2Cu (NO3)2 (s)
-------------→ 2CuO (s) + 4NO2 (g) + O2 (g) [Thermal decomposition]
Sun light
2AgCl (s) ------------------→ 2Ag (s)
+ Cl2 (g) [Photolytic
decomposition]