Chemical Reactions And Equations Class 10 Notes Science Chapter 1

Welcome to CBSE Wale, your ultimate destination for CBSE Class 10 Science Notes. In this blog post, we learn about Chemical Reactions and Equations Class 10 Notes Science Chapter 1. As budding scientists and learners, understanding the fundamentals of chemical reactions is like unlocking the hidden language of the universe. So, fasten your seatbelts as we embark on a journey to unravel the mysteries of transformations that shape our world!

Chemical Reactions and Equations Class 10 Notes Science Chapter 1

Chemical Reaction

A chemical reaction is a process in which one or more substances, known as reactants, are converted into one or more different substances, known as products. During the reaction, the atoms of the reactants rearrange to form the atoms of the products.

Daily Life Examples of Chemical Reactions:

Baking: When baking a cake, the mixture of flour, sugar, eggs, and other ingredients undergoes a chemical reaction during the baking process. The reactants react to produce carbon dioxide gas, causing the cake to rise and become fluffy.
Rusting: The process of rusting involves a chemical reaction between iron and oxygen in the presence of water or moisture. The iron reacts with oxygen to form iron oxide (rust), which weakens and corrodes metal objects.
Photography: In traditional film photography, light interacts with chemicals in the film to create a chemical reaction. This results in the formation of a latent image that becomes visible after developing the film.
Respiration: The process of respiration in living organisms involves a series of chemical reactions. In cellular respiration, glucose and oxygen react to produce carbon dioxide, water, and energy (in the form of ATP) that cells use for their activities.
Digestion: The process of digestion involves chemical reactions that break down food molecules into simpler substances. Enzymes in the digestive system catalyse these reactions to allow the body to absorb nutrients.
Photosynthesis: In plants, photosynthesis is a vital chemical reaction that converts carbon dioxide and water into glucose and oxygen, using sunlight as an energy source.
Combustion: Burning of fuels, such as wood, gasoline, or natural gas, involves a chemical reaction with oxygen to release heat and produce carbon dioxide and water vapour as products.

Burning Of Magnesium Ribbon In Oxygen

Magnesium is a highly reactive metal that burns rapidly in air.
When ignited, magnesium ribbon reacts with oxygen to form magnesium oxide, a white solid.
The reaction is exothermic and releases heat.
The heat from the reaction can sustain the burning even after the initial ignition source is removed.
The reaction is vigorous and should be carried out with caution.
Magnesium has a strong affinity for oxygen and forms a protective layer of magnesium oxide when exposed to air.
Scratching or breaking the magnesium ribbon exposes the metal, leading to a rapid exothermic reaction.
The reaction is self-sustaining as the heat vaporises magnesium oxide, exposing more metal to oxygen.
The reaction produces a bright white flame and generates a large amount of heat.
Safety precautions when working with magnesium ribbon include wearing safety glasses, working in a well-ventilated area, and avoiding ignition near flammable materials.
In case of ignition, do not use water to extinguish the flames; use sand or a fire blanket instead.

Reaction between lead nitrate solution and potassium iodide solution

The reaction is a precipitation reaction, forming a solid precipitate.
The precipitate formed is lead iodide, which is yellow in colour.
The reaction is exothermic and releases heat.
The reaction is rapid and can be violent.
It should be conducted in a well-ventilated area due to the heat and potential boiling of the solutions.
Eye protection and gloves are essential to prevent skin or eye contact with the solutions.
Lead nitrate and potassium iodide are both soluble salts that react to form lead iodide.
The balanced chemical equation for the reaction is: Pb(NO3)2(aq) + 2KI(aq) → PbI2(s) + 2KNO3(aq).
Lead iodide is insoluble in water but soluble in potassium iodide solution.
This reaction is commonly used to test for the presence of lead ions.

Reaction between zinc and dilute hydrochloric acid

The reaction is a single displacement reaction, with zinc displacing hydrogen from hydrochloric acid.
The products of the reaction are zinc chloride and hydrogen gas.
The reaction is exothermic and releases heat.
The reaction produces bubbles of hydrogen gas.
Hydrogen gas generated from this reaction can be collected and utilised for various purposes, such as fuel cells or balloons.
The balanced chemical equation for the reaction is: Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g).
Safety precautions include wearing safety glasses, working in a well-ventilated area, and avoiding excessive heating that could cause zinc to melt and splash.
In case of ignition of hydrogen gas, do not use water to extinguish the flames; use sand or a fire blanket to smother the flames.

Reaction between zinc and dilute sulphuric acid

The reaction is a single displacement reaction, with zinc displacing hydrogen from dilute sulfuric acid.
The products of the reaction are zinc sulphate and hydrogen gas.
The reaction is exothermic and releases heat.
The reaction produces bubbles of hydrogen gas.
Hydrogen gas generated from this reaction can be collected and utilised for various purposes, such as fuel cells or balloons.
The balanced chemical equation for the reaction is: Zn(s) + H2SO4(aq) → ZnSO4(aq) + H2(g).
Safety precautions include wearing safety glasses, working in a well-ventilated area, and avoiding excessive heating that could cause zinc to melt and splash.
In case of ignition of hydrogen gas, do not use water to extinguish the flames; use sand or a fire blanket to smother the flames.

Observations to determine whether a chemical reaction has taken place

Change in colour: The appearance of a new colour or a colour change in the reactants or products.
Evolution of gas: The formation of bubbles or a noticeable release of gas during the reaction.
Change in temperature: Detection of heat being released or absorbed during the reaction, leading to a change in temperature.
Formation of a precipitate: The appearance of a solid substance when two solutions are mixed together.
Change in state of matter: Conversion of a substance from one state (solid, liquid, gas) to another during the reaction.
Light emission: The production of light during the reaction, which can indicate a chemical change.
Formation of odour: The development of a new smell or a change in the smell of the reactants or products.
Effervescence: The release of gas bubbles in a liquid during the reaction.
Change in texture or appearance: Alteration in the physical properties, such as texture or appearance, of the substances involved in the reaction.
Release of sound: Audible noises or hissing during the reaction.

Chemical Equation

A chemical equation is a symbolic representation of a chemical reaction, describing the reactants, products, and their corresponding stoichiometric coefficients.

Parts of a Chemical Equation

Reactants: The reactants are the starting substances or chemicals present before the reaction takes place. They are written on the left side of the chemical equation.
Products: The products are the new substances formed as a result of the chemical reaction. They are written on the right side of the chemical equation.
Arrow: The arrow (→) indicates the direction of the reaction, showing the conversion of reactants into products.
Stoichiometric Coefficients: These are the numbers written in front of each chemical formula (reactants and products) in the chemical equation. The stoichiometric coefficients represent the relative number of molecules or moles of each substance involved in the reaction.

Example of a Chemical Equation:

Consider the reaction of hydrogen gas (H2) reacting with oxygen gas (O2) to form water (H2O): 2H2(g) + O2(g) → 2H2O(g). In this example:
The reactants are hydrogen gas (H2) and oxygen gas (O2).
The product is water (H2O).
The arrow indicates the direction of the reaction, from reactants to products.
The stoichiometric coefficients (2 and 1) indicate that two molecules of hydrogen gas react with one molecule of oxygen gas to form two molecules of water.

Steps involved in writing a chemical equation correctly:

Identify the reactants and products: In a chemical equation, the reactants are the substances that undergo a chemical change at the beginning of the reaction, while the products are the new substances formed as a result of the reaction.
Write the symbols of the reactants and products: Use the chemical symbols of the elements to represent the reactants and products. If a substance is already in its compound form, write its chemical formula.
Use stoichiometric coefficients: Balancing a chemical equation involves adjusting the stoichiometric coefficients (numbers in front of the chemical formulas) to ensure that the same number of atoms of each element appears on both sides of the equation. This maintains the law of conservation of mass, where the total mass of the reactants equals the total mass of the products.
Include the physical state of the reactants and products if necessary: In some cases, it is important to specify the physical state of the substances involved in the reaction. Common states include (s) for solid, (l) for liquid, (g) for gas, and (aq) for aqueous (dissolved in water).
Use the correct arrow to indicate the direction of the reaction: The arrow points from the reactants to the products, indicating the direction of the chemical change.

For example:

Chemical Reaction: 2H2(g) + O2(g) → 2H2O(l)
Reactants: The reactants in this equation are hydrogen gas (H2) and oxygen gas (O2).
Products: The products are water (H2O).
Stoichiometric Coefficients: The coefficients 2 and 1 in front of H2 and O2, respectively, indicate that two molecules of hydrogen gas react with one molecule of oxygen gas to form two molecules of water.
Physical State: The physical state of water is specified as (l), indicating that it is in the liquid state.
Arrow Direction: The arrow points from the reactants to the products, indicating the direction of the reaction.

Balanced Chemical Equations

Definition: A balanced chemical equation ensures that the same number of atoms of each element are represented on both the reactant and product sides of the equation.
Conservation of Mass: Balancing chemical equations reflects the law of conservation of mass, stating that matter cannot be created or destroyed during a chemical reaction.
Balancing Process: To balance a chemical equation, add coefficients to the formulas of reactants and products to ensure the same number of atoms of each element on both sides.
Methods for Balancing: Two common methods are trial and error, where coefficients are adjusted until balanced, and systematic approaches like the “crisscross” or “molecule-to-molecule” methods.
Importance: Balancing chemical equations is fundamental in chemistry. They represent chemical reactions, help calculate product amounts from reactants, and understand stoichiometry.
Example: The equation 2H2(g) + O2(g) → 2H2O(g) represents the reaction of hydrogen gas and oxygen gas to form water. It shows the stoichiometric coefficients and the number of molecules involved in the reaction.

Balance the chemical equation (Fe + H2O → Fe3O4 + H2) step by step

Step 1: Count the number of atoms of each element on both sides of the equation.

On the left side: Fe: 1 atom H: 2 atoms O: 1 atom
On the right side: Fe: 3 atoms H: 2 atoms O: 4 atoms

Step 2: Start by balancing the elements that appear in the fewest molecules first.

Let’s balance Fe atoms first.

Step 3: Add a coefficient of 3 to Fe on the left side to balance the Fe atoms.

3Fe + H2O → Fe3O4 + H2

Now, let’s count the atoms again:

On the left side: Fe: 3 atoms H: 2 atoms O: 1 atom
On the right side: Fe: 3 atoms H: 2 atoms O: 4 atoms

Step 4: Now, let’s balance the O atoms.

Step 5: Add a coefficient of 4 to H2O on the left side to balance the O atoms.

3Fe + 4H2O → Fe3O4 + H2

Now, let’s count the atoms again:

On the left side: Fe: 3 atoms H: 8 atoms O: 4 atoms
On the right side: Fe: 3 atoms H: 2 atoms O: 4 atoms

Step 6: Lastly, balance the H atoms.

Step 7: Add a coefficient of 4 to H2 on the right side to balance the H atoms.

3Fe + 4H2O → Fe3O4 + 4H2

Now, let’s count the atoms one last time:

On the left side: Fe: 3 atoms H: 8 atoms O: 4 atoms
On the right side: Fe: 3 atoms H: 8 atoms O: 4 atoms

The equation is now balanced, and the final balanced chemical equation is:

3Fe + 4H2O → Fe3O4 + 4H2

Symbols of Physical States in Chemical Equations:

Solid: (s) – Indicates that the substance is in a solid state. For example, carbon dioxide (CO2) is written as CO2 (s) when it is in a solid state, such as dry ice.
Liquid: (l) – Indicates that the substance is in a liquid state. For example, water (H2O) is written as H2O (l) when it is in a liquid state.
Gas: (g) – Indicates that the substance is in a gaseous state. For example, oxygen (O2) is written as O2 (g) when it is in a gaseous state.
Aqueous: (aq) – Indicates that the substance is dissolved in water. For example, sodium hydroxide (NaOH) is written as NaOH (aq) when it is dissolved in water.
Diluted: (dil.) – Indicates that the substance is a dilute solution. For example, hydrochloric acid (HCl) is written as HCl (dil.) when it is a dilute solution.
Concentrated: (conc.) – Indicates that the substance is a concentrated solution. For example, sulfuric acid (H2SO4) is written as H2SO4 (conc.) when it is a concentrated solution.

Methanol Synthesis Reaction:

Chemical Equation: The balanced chemical equation for the methanol synthesis reaction is CO(g) + 2H2(g) → CH3OH(g).
Exothermic Reaction: The reaction is exothermic, meaning it releases heat during the process.
Reaction Conditions: The methanol synthesis reaction is typically carried out at high temperatures (200-300°C) and pressures (50-100 atm). These conditions favor the formation of methanol.
Catalyst: A catalyst, such as zinc oxide or molybdenum, is used to speed up the reaction and enhance its efficiency.
Initiation: The reaction is initiated by the collision of a carbon monoxide molecule and a hydrogen molecule.
Mechanism: The collision breaks the bonds in the carbon monoxide and hydrogen molecules, leading to the recombination of atoms to form a methanol molecule.
Product Release: The methanol molecule is released from the catalyst’s surface and escapes from the reaction vessel.
Reversibility: The reaction is reversible, but at high temperatures and pressures, the equilibrium favors the formation of methanol.
Industrial Importance: The methanol synthesis reaction is a crucial industrial process for producing methanol, which serves as an essential chemical feedstock and fuel source.

Chemical Reaction for photosynthesis:

Chemical Equation: The balanced chemical equation for photosynthesis is 6CO2 + 12H2O → C6H12O6 + 6O2.
Process: Photosynthesis is a biological process driven by sunlight and catalysed by chlorophyll.
Location: The reaction occurs in the chloroplasts of plant cells, where chlorophyll is present.
Products: The products of photosynthesis are glucose (C6H12O6) and oxygen (O2).
Role of Glucose: Glucose is a sugar used by plants as an energy source for various metabolic processes.
Release of Oxygen: Oxygen is released into the atmosphere during photosynthesis and is vital for respiration in living organisms.
Redox Reaction: Photosynthesis is a redox reaction involving the transfer of electrons from one substance to another.
Oxidation and Reduction: In photosynthesis, carbon dioxide (CO2) is oxidised to form glucose (C6H12O6), and water (H2O) is reduced to form oxygen (O2).
Energy Source: The energy from sunlight is utilised to power the oxidation of carbon dioxide and the reduction of water.
Significance: Photosynthesis is essential for life on Earth as it is the primary process by which green plants and some other organisms produce their food and release oxygen, supporting the biosphere and ecosystems.

Different Types Of Chemical Reactions

Combination Reactions

Definition: Combination reactions are chemical reactions where two or more substances combine to form a single new substance.
Reactants: Typically involve elements or simple compounds as reactants.
Products: The products of combination reactions are usually more complex compounds.
Exothermic Nature: Combination reactions are often exothermic, meaning they release heat during the reaction.

Examples:

Hydrogen and oxygen combine to form water: 2H2 + O2 → 2H2O
Magnesium and oxygen combine to form magnesium oxide: Mg + O2 → 2MgO
Carbon and oxygen combine to form carbon dioxide: C + O2 → CO2
Calcium oxide and water combine to form calcium hydroxide: CaO + H2O → Ca(OH)2

Uses Of Combination Reactions:

Hydrogen and oxygen combination produces drinking water and hydrogen peroxide.
Magnesium and oxygen combination is used in fire retardants and brake pads.
Carbon and oxygen combination is used to produce soft drinks and dry ice.
Calcium oxide and water combination is used in making mortar and concrete.

Application

Combination reactions are commonly employed in the production of new materials and various industrial processes.

Chemical Reaction between Calcium Oxide and Water (Slaking of Lime):

Reaction Name: The reaction is known as the “slaking of lime.”
Balanced Chemical Equation: CaO + H2O → Ca(OH)2
Exothermic Nature: The reaction is exothermic, meaning it releases heat during the process.
Temperature: The reaction is typically carried out at room temperature, but heating calcium oxide can accelerate the reaction.
Product: The product of the reaction is calcium hydroxide, also known as “slaked lime.”
Properties of Calcium Hydroxide: Calcium hydroxide is a white, powdery substance that is soluble in water.

Applications of Calcium Hydroxide:

Making mortar and concrete.
Purifying water.
Neutralising acids.
Making paper.
Treating textiles.

Chemical Reaction between Calcium Hydroxide and Carbon Dioxide:

Reaction Name: The reaction is known as the “limewater reaction.”
Balanced Chemical Equation: Ca(OH)2 + CO2 → CaCO3 + H2O
Exothermic Nature: The reaction is exothermic, meaning it releases heat during the process.
Temperature: The reaction is typically carried out at room temperature, but it can be accelerated by bubbling carbon dioxide gas through a solution of calcium hydroxide.
Product: The product of the reaction is calcium carbonate, also known as “limestone.”
Properties of Calcium Carbonate: Calcium carbonate is a white, solid substance that is insoluble in water.

Applications of Calcium Carbonate:

Making cement.
Making whitewash.
Purifying water.
Neutralising acids.
Making paper.
Treating textiles.

Chemical Reaction of Burning Coal:

Combustion Reaction: The burning of coal is a combustion reaction where a fuel (coal) reacts with oxygen in the air to produce heat and light.
Fuel: Coal is a solid form of carbon and serves as the fuel in this reaction.
Oxygen Source: The oxygen required for the combustion comes from the surrounding air.
Products: The products of the reaction are carbon dioxide (CO2) and water vapor (H2O).
Balanced Chemical Equation: C(s) + O2(g) → CO2(g)
Exothermic Nature: The burning of coal is exothermic, meaning it releases heat.
Utilisation: The heat generated from burning coal is used to generate electricity, heat homes and businesses, and power industrial processes.
Environmental Impact: The burning of coal releases carbon dioxide into the atmosphere, contributing to climate change.

Chemical Reaction of Formation of Water:

Combination Reaction: The formation of water is a combination reaction, where hydrogen gas (H2) and oxygen gas (O2) combine to form a single new substance, water (H2O).
Reactants: Hydrogen gas (H2) and oxygen gas (O2) are the reactants in this reaction.
Product: The product of the reaction is water (H2O).
Balanced Chemical Equation: 2H2(g) + O2(g) → 2H2O(g)
Exothermic Nature: The formation of water is exothermic, meaning it releases heat during the process.
Heat Utilisation: The heat released during the formation of water is utilized in various applications, such as in boilers and steam engines.
Hydrogen Peroxide Production: The formation of water is also involved in the production of hydrogen peroxide (H2O2).
Reversibility: The formation of water is a reversible reaction. Water can be broken down into hydrogen gas and oxygen gas. However, at room temperature, the equilibrium favours the formation of water.
Catalyst Influence: The formation of water can be accelerated by using a catalyst, such as platinum or iron.

Chemical Reaction of Burning Natural Gas:

Fuel: Natural gas, primarily composed of methane (CH4), is a fossil fuel used in the burning process.
Combustion: Burning natural gas involves its reaction with oxygen (O2) to produce carbon dioxide (CO2) and water vapour (H2O).
Balanced Chemical Equation: CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)
Exothermic Nature: The burning of natural gas is exothermic, meaning it releases heat during the process.
Heat Utilisation: The heat generated from burning natural gas is utilised to generate electricity, heat homes and businesses, and power industrial processes.
Environmental Impact: The burning of natural gas releases carbon dioxide into the atmosphere, contributing to climate change as it is a greenhouse gas.
Relative Cleanliness: Compared to other fossil fuels like coal and oil, burning natural gas is relatively cleaner as it produces fewer pollutants like sulfur dioxide and nitrogen oxides.
Efficiency Enhancements: Various technologies, such as combustion turbines and heat exchangers, can be used to improve the efficiency of natural gas burning.
Cleaner Burning: Carbon capture and storage technology can be employed to make the burning of natural gas cleaner by capturing and storing carbon dioxide emissions.

Chemical Reaction of Respiration:

Process: Respiration is a vital biological process through which organisms break down food molecules to release energy.
Complexity: The chemical reaction of respiration is complex and occurs through a series of steps.
Overall Reaction (Aerobic Respiration): The balanced chemical equation for aerobic respiration is: C6H12O6 + 6O2 → 6CO2 + 6H2O + ATP
Reactants: Glucose (C6H12O6) is the food molecule that is broken down, and oxygen (O2) is the reactant used to break down glucose.
Products: The products of aerobic respiration are carbon dioxide (CO2), water (H2O), and ATP (adenosine triphosphate).
ATP: ATP is the energy molecule produced during respiration, which cells use to carry out various activities.
Exothermic Nature: Respiration is an exothermic reaction, meaning it releases heat.
Heat Utilisation: The heat released during respiration is used to maintain body temperature and power the cellular activities.
Vital for Life: Respiration is essential for life as it provides the energy required for various physiological processes.
Importance: Without respiration, organisms would be unable to produce energy, leading to their eventual death.

Chemical Reaction of Decomposition of Vegetable Matter into Compost:

Process: The decomposition of vegetable matter into compost is a complex process involving various microorganisms.
Organic Matter Breakdown: Microorganisms break down the organic matter in vegetable matter into simpler molecules, including carbon dioxide (CO2), water (H2O), nitrogen (N2), and other products.
Balanced Chemical Equation: The general chemical equation for the decomposition process is: C6H12O6 + 6O2 → 6CO2 + 6H2O + N2 + other products
Exothermic Nature: The decomposition process is exothermic, meaning it releases heat.
Heat Utilisation: The heat generated from decomposition helps break down organic matter and eliminates harmful bacteria.
Time Frame: Decomposition into compost is a gradual process, taking several months or even years to complete.
Soil Health: The decomposition of vegetable matter into compost is crucial for soil health.
Soil Improvement: Compost improves soil drainage and aeration, helping to retain moisture in the soil.
Nutrient Provision: Compost provides essential nutrients to plants, promoting their growth and overall health.

Decomposition Reactions

Definition: Decomposition reactions are chemical reactions in which a single compound breaks down into two or more simpler substances.
Reactant: The reactant in a decomposition reaction is known as the decomposing substance.
Products: The products of a decomposition reaction are called the decomposed substances.
Exothermic Nature: Decomposition reactions are often exothermic, meaning they release heat during the process.

Examples Of Decomposition Reactions:

Water decomposes into hydrogen and oxygen: 2H2O → 2H2 + O2
Calcium carbonate decomposes into calcium oxide and carbon dioxide: CaCO3 → CaO + CO2
Hydrogen peroxide decomposes into water and oxygen: 2H2O2 → 2H2O + O2
Silver chloride decomposes into silver and chlorine gas: 2AgCl → 2Ag + Cl2
Initiation: Decomposition reactions can be triggered by heat, light, or electricity.

Application Of Decomposition Reactions

Decomposition reactions are frequently employed in the production of various chemicals for industrial applications.
Water decomposition produces hydrogen gas used in fuel cells and hydrogen balloons.
Calcium carbonate decomposition produces lime used in construction and agriculture.
Hydrogen peroxide decomposition produces oxygen gas utilised in first-aid kits and oxygen tanks.
Silver chloride decomposition produces silver used in jewellery and photography.

Chemical Reaction of Decomposition of Ferrous Sulphate:

Process: Ferrous sulphate decomposes into ferric oxide, sulphur dioxide, and sulphur trioxide when heated.
Balanced Chemical Equation: The balanced chemical equation for the decomposition of ferrous sulphate is: 2FeSO4(s) → Fe2O3(s) + SO2(g) + SO3(g)
Endothermic Nature: The decomposition of ferrous sulphate is an endothermic reaction, meaning it absorbs heat during the process.
Initiation: The decomposition of ferrous sulphate can be initiated by heating, but it can also occur spontaneously at high temperatures.
Odour: The decomposition of ferrous sulphate produces a pungent odour.
Sulphur Dioxide Production: The decomposition of ferrous sulphate can be utilised to produce sulphur dioxide (SO2), which finds applications in the production of sulfuric acid and paper.
Reversibility: The decomposition of ferrous sulphate is a reversible reaction. Ferric oxide, sulphur dioxide, and sulphur trioxide can react to reform ferrous sulphate, but the equilibrium favours the decomposition at high temperatures.
Catalyst Influence: The decomposition of ferrous sulphate can be accelerated by the presence of a catalyst, such as manganese dioxide.

Chemical Reaction of Decomposition of Limestone:

Limestone Composition: Limestone is a calcium carbonate (CaCO3) sedimentary rock.
Decomposition: Limestone decomposes into calcium oxide (CaO) and carbon dioxide (CO2) when heated.
Balanced Chemical Equation: The balanced chemical equation for the decomposition of limestone is: CaCO3(s) → CaO(s) + CO2(g)
Endothermic Nature: The decomposition of limestone is an endothermic reaction, meaning it absorbs heat during the process.
Initiation: The decomposition of limestone can be initiated by heating, and it can also occur spontaneously at high temperatures.
Product: The decomposition of limestone produces calcium oxide, which is a white powder.
Lime Production: The decomposition of limestone can be utilized to produce lime (calcium oxide), which finds applications in construction and agriculture.
Reversibility: The decomposition of limestone is a reversible reaction. Calcium oxide and carbon dioxide can react to reform limestone, but the equilibrium favors the decomposition at high temperatures.
Catalyst Influence: The decomposition of limestone can be accelerated by the presence of a catalyst, such as manganese dioxide.

Chemical Reaction of Decomposition of Lead Nitrate:

Lead Nitrate Composition: Lead nitrate (Pb(NO3)2) is a chemical compound.
Decomposition: Lead nitrate decomposes into lead oxide (PbO), nitrogen dioxide (NO2), and oxygen gas (O2) when heated.
Balanced Chemical Equation: The balanced chemical equation for the decomposition of lead nitrate is: 2Pb(NO3)2(s) → 2PbO(s) + 4NO2(g) + O2(g)
Exothermic Nature: The decomposition of lead nitrate is an exothermic reaction, meaning it releases heat during the process.
Initiation: The decomposition of lead nitrate can be initiated by heating, and it can also occur spontaneously at high temperatures.
Product: The decomposition of lead nitrate produces a brown gas, which is nitrogen dioxide (NO2).
Nitrogen Dioxide Production: The decomposition of lead nitrate can be utilized to produce nitrogen dioxide, which finds applications in the production of nitric acid and explosives.
Reversibility: The decomposition of lead nitrate is a reversible reaction. Lead oxide, nitrogen dioxide, and oxygen gas can react to reform lead nitrate, but the equilibrium favours the decomposition at high temperatures.
Catalyst Influence: The decomposition of lead nitrate can be accelerated by the presence of a catalyst, such as manganese dioxide.

Chemical Reaction of Decomposition of Silver Chloride:

Silver Chloride Composition: Silver chloride (AgCl) is a chemical compound.
Decomposition: Silver chloride decomposes into silver metal (Ag) and chlorine gas (Cl2) when exposed to light.
Balanced Chemical Equation: The balanced chemical equation for the decomposition of silver chloride is: 2AgCl(s) → 2Ag(s) + Cl2(g)
Photolytic Reaction: The decomposition of silver chloride is a photolytic reaction, meaning it is initiated by light.
Initiation: The decomposition of silver chloride can also be accelerated by heat, and it can occur spontaneously at high temperatures.
Product: The decomposition of silver chloride produces silver metal, which appears as a greyish-black powder.
Silver Production: The decomposition of silver chloride can be utilized to produce silver, which finds applications in various fields, including jewelry and photography.
Reversibility: The decomposition of silver chloride is a reversible reaction. Silver metal and chlorine gas can react to reform silver chloride, but the equilibrium favors the decomposition in the presence of light.
Preventive Measures: The decomposition of silver chloride can be prevented by coating it with a protective layer, such as clear varnish, to block the light and inhibit the decomposition reaction.

Chemical Reaction of Decomposition of Silver Bromide:

Silver Bromide Composition: Silver bromide (AgBr) is a chemical compound.
Decomposition: Silver bromide decomposes into silver metal (Ag) and bromine gas (Br2) when exposed to light.
Balanced Chemical Equation: The balanced chemical equation for the decomposition of silver bromide is: 2AgBr(s) → 2Ag(s) + Br2(g)
Photolytic Reaction: The decomposition of silver bromide is a photolytic reaction, meaning it is initiated by light.
Initiation: The decomposition of silver bromide can also be accelerated by heat, and it can occur spontaneously at high temperatures.
Product: The decomposition of silver bromide produces silver metal, which appears as a greyish-black powder.
Silver Production: The decomposition of silver bromide can be utilized to produce silver, which finds applications in various fields, including jewelry and photography.
Reversibility: The decomposition of silver bromide is a reversible reaction. Silver metal and bromine gas can react to reform silver bromide, but the equilibrium favors the decomposition in the presence of light.
Preventive Measures: The decomposition of silver bromide can be prevented by coating it with a protective layer, such as clear varnish, to block the light and inhibit the decomposition reaction.

Displacement Reactions

Definition: Displacement reactions are chemical reactions in which one element displaces another element from a compound.
Reactivity Difference: In a displacement reaction, one element is more reactive than the other element present in the compound.
Reactive Element Replacement: The more reactive element replaces the less reactive element in the compound, forming a new compound and releasing the displaced element.

Examples Of Displacement Reactions:

Zinc displaces copper from copper sulphate: Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s)
Iron displaces silver from silver nitrate: Fe(s) + AgNO3(aq) → Fe(NO3)2(aq) + Ag(s)
Magnesium displaces hydrogen from hydrochloric acid: Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g)
Extraction of Metals: Displacement reactions are commonly used in the extraction of metals from their ores. For example, zinc is extracted from its ore, zinc sulphide, by reacting it with iron. Iron is more reactive than zinc, so it displaces the zinc from the zinc sulphide, and zinc is collected as a pure metal.
Metal Purification: Displacement reactions can also be used to purify metals. For instance, silver can be purified by reacting it with nitric acid. Nitric acid oxidises the impurities in the silver, leaving pure silver behind.

Chemical Reaction Between Iron and Copper sulphate:

Displacement Reaction: When iron is placed in a solution of copper sulphate, a displacement reaction occurs. Iron, being more reactive than copper, displaces copper from the copper sulphate solution.
Products: The products of the reaction are ferrous sulphate (FeSO4) and copper metal (Cu).
Balanced Chemical Equation: The balanced chemical equation for the reaction is: Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s)
Exothermic and Spontaneous: The reaction is exothermic, meaning it releases heat, and it is spontaneous, occurring without any external energy input.
Reaction Rate: The reaction rate can be increased by stirring the solution and raising the temperature.
Preventive Measures: The reaction can be prevented by coating the iron with a protective layer, such as clear varnish, to block the contact with the copper sulphate solution.
Single-Displacement Reaction: The reaction between iron and copper sulphate is an example of a single-displacement reaction, where one element displaces another element from a compound.
Redox Reaction: The reaction is also a redox reaction, involving the transfer of electrons between reactants. In this case, iron is oxidized, losing electrons, and copper is reduced, gaining electrons.

Chemical Reaction Between Zinc and Copper sulphate:

Single Displacement Reaction: When zinc metal is placed in a solution of copper sulphate, a single displacement reaction occurs. Zinc, being more reactive than copper, displaces copper from the copper sulphate solution.
Products: The products of the reaction are zinc sulphate (ZnSO4) and copper metal (Cu).
Balanced Chemical Equation: The balanced chemical equation for the reaction is: Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s)
Exothermic and Spontaneous: The reaction is exothermic, meaning it releases heat, and it is spontaneous, occurring without any external energy input.
Reaction Rate: The reaction rate can be increased by stirring the solution and raising the temperature.
Preventive Measures: The reaction can be prevented by coating the zinc with a protective layer, such as clear varnish, to block the contact with the copper sulphate solution.
Single Displacement and Redox Reaction: The reaction between zinc and copper sulphate is a classic example of a single-displacement reaction, where one element displaces another element from a compound. It is also a redox reaction, involving the transfer of electrons between reactants. In this case, zinc is oxidized, losing electrons, and copper is reduced, gaining electrons.

Chemical Reaction Between Lead and Copper Chloride:

Single Replacement Reaction: When lead metal is placed in a solution of copper chloride, a single-replacement reaction occurs.
Reactivity Difference: Lead is less reactive than copper, so it cannot displace copper from the copper chloride solution.
Formation of Lead Chloride: Instead, lead reacts with the chloride ions in the copper chloride solution to form lead chloride (PbCl2).
Products: The products of the reaction are lead chloride and copper metal (Cu).
Balanced Chemical Equation: The balanced chemical equation for the reaction is: Pb(s) + CuCl2(aq) → PbCl2(aq) + Cu(s)
Exothermic and Spontaneous: The reaction is exothermic, meaning it releases heat, and it is spontaneous, occurring without any external energy input.
Reaction Rate: The reaction rate can be increased by stirring the solution and raising the temperature.
Preventive Measures: The reaction can be prevented by coating the lead with a protective layer, such as clear varnish, to block the contact with the copper chloride solution.

Double Displacement Reactions

Double displacement reactions involve the exchange of ions between two compounds, leading to the formation of two new compounds.
Reactants in a double displacement reaction consist of cations and anions.
The cations and anions swap places to form new combinations in the products.
The resulting compounds in the products are different from the original reactants.

Examples of Double Displacement Reactions:

Sodium chloride and silver nitrate:

NaCl(aq) + AgNO3(aq) → NaNO3(aq) + AgCl(s)
Product formed: Sodium nitrate and silver chloride

Potassium hydroxide and hydrochloric acid:

KOH(aq) + HCl(aq) → KCl(aq) + H2O(l)
Product formed: Potassium chloride and water

Calcium carbonate and hydrochloric acid:

CaCO3(s) + 2HCl(aq) → CaCl2(aq) + CO2(g) + H2O(l)
Products formed: Calcium chloride, carbon dioxide, and water

Applications of Double Displacement Reactions:

Laboratory Synthesis: Double displacement reactions are employed to create new compounds in the laboratory. For example, sodium chloride and silver nitrate react to synthesize silver chloride, which finds application in photography.
Ion Testing: These reactions are used to test for the presence of specific ions in a solution. For instance, adding potassium hydroxide to a solution with chloride ions will produce a white precipitate if chloride ions are present.

Precipitation Reactions

Precipitation reactions result in the formation of an insoluble solid, known as a precipitate, when two soluble compounds react.
The precipitate separates from the solution as a solid due to its low solubility.
In a precipitation reaction, the products are two new compounds, with one of them being insoluble.

Examples of Precipitation Reactions:

Sodium chloride and silver nitrate:

NaCl(aq) + AgNO3(aq) → NaNO3(aq) + AgCl(s)
Precipitate formed: Silver chloride

Potassium hydroxide and barium chloride:

KOH(aq) + BaCl2(aq) → Ba(OH)2(s) + KCl(aq)
Precipitate formed: Barium hydroxide

Calcium hydroxide and sulfuric acid:

Ca(OH)2(aq) + H2SO4(aq) → CaSO4(s) + 2H2O(l)
Precipitate formed: Calcium sulphate

Applications of Precipitation Reactions:

Separation of Mixtures: Precipitation reactions are used in the laboratory to separate mixtures of substances. For instance, adding potassium hydroxide to a mixture of sodium chloride and silver nitrate will precipitate out the silver chloride, which can be easily separated from the solution.
Ion Testing: Precipitation reactions serve as a valuable method to test for the presence of specific ions in a solution. For example, adding barium chloride to a solution containing sulphate ions will cause a white precipitate to form if sulphate ions are present.

Chemical Reaction between Sodium sulphate and Barium Chloride:

When sodium sulphate (Na2SO4) and barium chloride (BaCl2) are mixed in a solution, a white precipitate of barium sulphate (BaSO4) forms.
The balanced chemical equation for the reaction is: Na2SO4(aq) + BaCl2(aq) → BaSO4(s) + 2NaCl(aq)
The reaction is classified as a precipitation reaction because it produces an insoluble solid, called a precipitate, which is barium sulphate in this case.
It is also a double displacement reaction since the ions of the two compounds exchange places.
The sodium ions from sodium sulphate react with the chloride ions from barium chloride to form sodium chloride (NaCl), which is soluble and remains in the solution.
The barium ions from barium chloride react with the sulphate ions from sodium sulphate to form barium sulphate, which is insoluble and precipitates out of the solution as a white solid.

Applications of the Reaction:

Testing for Barium Ions: The reaction between sodium sulphate and barium chloride is commonly used in the laboratory to test for the presence of barium ions in a solution. If barium ions are present, a white precipitate of barium sulphate will form.
Precipitation of Barium sulphate: The reaction can also be utilised to precipitate barium sulphate from solutions. Barium sulphate is a white solid that remains insoluble in water.
Barium sulphate Uses: Barium sulphate finds applications in various industries, such as in paints, plastics, and ceramics, due to its white colour and low solubility properties.

Oxidation Reactions

Oxidation reactions involve the loss of electrons from an atom or molecule.
The species that loses electrons is considered to be oxidised.
The species that gains electrons is considered to be reduced.
These reactions often release energy in the form of heat or light.

Examples of Oxidation Reactions:

Burning of wood:

C + O2 → CO2
Carbon (C) is oxidised as it loses electrons to oxygen (O2), which is reduced.

Rusting of iron:

Fe + O2 → Fe2O3
Iron (Fe) is oxidised as it loses electrons to oxygen (O2), which is reduced.

Decomposition of hydrogen peroxide:

2H2O2 → 2H2O + O2
Hydrogen peroxide (H2O2) is oxidised, breaking down into water (H2O) and oxygen (O2) molecules.

Importance of Oxidation Reactions:

Combustion: Oxidation reactions, such as burning wood, are crucial in the process of combustion, providing energy in the form of heat and light.
Rusting: Oxidation reactions play a significant role in the rusting of iron and other metals, causing corrosion over time.
Photosynthesis: In photosynthesis, plants use oxidation reactions to convert carbon dioxide and water into glucose and oxygen, capturing energy from sunlight.
Energy Generation: Oxidation reactions are utilised in fuel cells to generate energy by oxidising fuel substances.

Environmental Impact:

Harmful Effects: Some oxidation reactions can be harmful to the environment, like the production of smog, which involves the oxidation of pollutants in the presence of sunlight.
Environmental Balance: Oxidation and reduction reactions are essential for maintaining the balance of chemicals in ecosystems, as they are involved in various natural processes.

Reduction Reactions

Reduction reactions involve the gain of electrons by an atom or molecule.
The species that gains electrons is considered to be reduced.
The species that loses electrons is considered to be oxidised.
These reactions often absorb energy in the form of heat or light.

Examples of Reduction Reactions:

Reduction of iron oxide by carbon monoxide:

Fe2O3 + 3CO → 2Fe + 3CO2
Iron oxide (Fe2O3) is reduced by carbon monoxide (CO), which gains oxygen atoms.

Reduction of hydrogen peroxide by glucose:

2H2O2 + C6H12O6 → 2H2O + 2CO2 + 6H2O
Hydrogen peroxide (H2O2) is reduced by glucose (C6H12O6), which gains oxygen atoms.

Reduction of nitrobenzene to aniline by hydrogen:

C6H5NO2 + 3H2 → C6H5NH2 + 2H2O
Nitrobenzene (C6H5NO2) is reduced by hydrogen (H2), which gains hydrogen atoms.

Importance of Reduction Reactions:

Metal Extraction: Reduction reactions are crucial in the extraction of metals from their ores, where metal oxides are reduced to obtain pure metals.
Petroleum Refining: Reduction reactions are used in the refining of petroleum to convert unsaturated hydrocarbons to saturated ones.
Chemical Production: Reduction reactions are employed in the production of various chemicals, including pharmaceuticals and industrial compounds.
Energy Generation: Reduction reactions are used in fuel cells to generate energy by reducing reactants like hydrogen.

Environmental Impact:

Harmful Effects: Some reduction reactions can be harmful to the environment, such as the reduction of nitrogen oxides to form acid rain.

Terminologies Used in Reduction Reactions:

Reduction: Gain of electrons.
Oxidation: Loss of electrons.
Oxidising Agent: Substance that oxidises another substance by accepting electrons.
Reducing Agent: Substance that reduces another substance by donating electrons.

Corrosion

Corrosion is the deterioration of a material due to its reaction with the surrounding environment.
The most common result of corrosion is the formation of a metal oxide, but it can also involve other forms of material loss, like pits or cracks.
Corrosion poses a significant problem as it can lead to structural and equipment failure.

Types of Corrosion:

Galvanic Corrosion: Occurs when two dissimilar metals are in contact with each other and an electrolyte (e.g., water), causing a flow of electrons between them, leading to the corrosion of one of the metals.
Oxidation Corrosion: Results from the reaction of a metal with oxygen in the air or water, forming metal oxides.
Erosion Corrosion: Happens when a metal is exposed to a moving fluid (e.g., water or air), leading to accelerated corrosion due to the fluid’s mechanical action.
Stress Corrosion Cracking: Occurs when a metal is under stress and exposed to a corrosive environment, leading to cracks and fractures.

Methods to Prevent or Slow Down Corrosion:

Protective Coatings: Applying coatings like paint or plating can create a barrier between the metal and the environment, preventing direct contact and reducing corrosion.
Electrochemical Protection: Implementing sacrificial anodes, made of more reactive metals, can act as sacrificial protection, corroding in place of the main metal.
Corrosion Inhibitors: Using chemicals as inhibitors can be added to the environment to slow down the corrosion process by forming a protective layer on the metal surface.

Rancidity

Rancidity is the process in which fats and oils become oxidised, resulting in an unpleasant odour and taste.
Oxygen is a major contributor to rancidity as it breaks down the fatty acids in fats and oils into smaller, foul-smelling molecules.
Rancidity is more common in unsaturated fats, such as vegetable oils, compared to saturated fats.
Factors that accelerate rancidity include exposure to heat, light, and oxygen, as well as the presence of certain bacteria and enzymes.
Rancid fats and oils can be harmful to health due to the production of harmful compounds like free radicals.

Prevention and Disposal of Rancid Fats and Oils:

To prevent rancidity, store fats and oils in a cool, dark place, and keep them tightly sealed to limit their exposure to oxygen.
Rancid fats and oils should be discarded as they are not safe to eat and can have negative health effects.

Common Sources of Rancidity:

Oils: Vegetable oil, olive oil, and canola oil are susceptible to rancidity.
Nuts: Almonds, peanuts, and walnuts can become rancid due to their oil content.
Seeds: Sunflower seeds, sesame seeds, and flaxseeds are also prone to rancidity.
Dry Fruits: Raisins, figs, and dates contain fats that can undergo rancidity.
Cooked Foods: Fried foods, salad dressings, and mayonnaise, which contain fats and oils, can become rancid over time.

In conclusion, “Chemical Reactions and Equations” in Class 10 Science Chapter 1 offer a profound insight into the dynamic interplay of matter and energy, enabling us to comprehend the intricate dance of atoms and molecules. From the humble combustion of a candle to the complex processes occurring within our bodies, chemical reactions are the driving force of life as we know it. Armed with the knowledge gained from this chapter, you’re now equipped to see the world with a new perspective – one where every alteration is a symphony of particles rearranging to create something new. Stay curious, keep exploring, and remember that the universe is a vast laboratory of reactions waiting to be understood. For more enriching educational content, stay tuned to CBSE Wale. Happy learning!