Chemical changes, also known as chemical reactions, are fundamental processes where original substances, called reactants, are transformed into new substances with different properties, known as products. This transformation involves the breaking and forming of chemical bonds, leading to a rearrangement of atoms. Unlike physical changes, which are often reversible, chemical changes are typically permanent and involve an energy change. Key indicators that a chemical reaction has occurred include the evolution of a gas, formation of a precipitate, a distinct color change, or a noticeable temperature shift. Understanding these signs helps distinguish chemical transformations from mere physical alterations in state or form.
Chemical reactions are systematically represented using chemical equations, which must be balanced to adhere to the Law of Conservation of Mass. This law states that matter cannot be created or destroyed, meaning the number of atoms for each element must be identical on both the reactant and product sides. Equations provide a concise blueprint of the reaction, showing the formulas and proportions of all substances involved. For a reaction to be feasible, reactant particles must collide with sufficient energy and proper orientation. Factors such as temperature, concentration, particle size (surface area), and the presence of a catalyst directly influence this reaction rate by affecting the frequency and force of these collisions.
These reactions can be categorized into several basic types. Combination reactions occur when two or more substances merge to form a single product, while decomposition reactions involve a single compound breaking down into simpler substances. Displacement reactions feature a more reactive element pushing out a less reactive one from its compound. Other important types include double decomposition, where ions exchange partners, and oxidation-reduction (redox) reactions, which involve the transfer of electrons between species. These diverse reaction pathways are the engine behind countless biological and industrial processes, from cellular respiration to the manufacturing of essential materials.
Exercise 2 (A)
Question 1.
1. What is a chemical reaction?
2. State the conditions necessary for a chemical change or reaction.
Ans:
1. Defining a Chemical Transformation
A chemical reaction is a transformation process where one or more starting substances, known as reactants, are converted into one or more new substances, called products, which possess distinct chemical identities and properties. This conversion fundamentally involves the rearrangement of atoms.
Defining Features:
- Molecular Reconfiguration: The essence of the process is the breaking of existing chemical bonds within the reactant molecules and the formation of new bonds to assemble the product molecules.
- Novelty of Products: The resulting products are chemically and physically different from the initial reactants. For instance, combining hydrogen and oxygen creates water, a liquid with properties vastly unlike the starting gases.
- Accompanying Energy Shifts: All chemical changes are accompanied by either the absorption of energy (endothermic) or the release of energy (exothermic), often manifesting as heat or light.
2. Essential Precursors for a Reaction
For atoms and molecules to successfully rearrange and form new products, certain external or internal factors, known as reaction conditions, must typically be met to overcome the inherent stability of the reactants:
- Proximity (Intimate Contact): Reacting species must be physically brought together so they can interact. While gases and dissolved ions easily intermingle, solids often need to be in a fine powdered form or dissolved in a solvent to ensure adequate surface area and molecular mobility for effective collisions.
- Energy Input (Heat or Activation): Many reactions require an initial boost of energy, often supplied as heat, to destabilize the reactant bonds and initiate the transformation. This minimum required energy is known as the activation energy.
- Specific Triggers: Certain reactions are initiated solely by specific energy forms:
- Light: Acts as a trigger for photochemical reactions (e.g., the breakdown of hydrogen peroxide or the process of photosynthesis).
- Electricity: Required for electrolysis, forcing non-spontaneous decomposition (e.g., splitting compounds like water).
- Catalytic Mediation: A catalyst may be introduced to dramatically increase the reaction rate. It achieves this by offering an alternative reaction pathway with a significantly lower activation energy, allowing the process to occur much faster or at lower temperatures without the catalyst being consumed.
- Pressure Adjustment: For reactions involving gases, increasing the pressure forces the reactant molecules closer, leading to a higher frequency of effective molecular collisions necessary to drive the reaction forward.
Question 2.
1. Define the following term: Chemical change
2. Define the following terms: Chemical bond
3. Define the following terms: Effervescence
4. Define the following terms: Precipitate
Ans:
1. Chemical Change
A chemical change is a transformation in which a substance is converted into one or more new substances that possess distinctly different chemical compositions and properties than the original material. This change involves the breaking of existing chemical bonds and the formation of new bonds, which is typically difficult to reverse (e.g., burning wood to ash, which cannot easily be converted back into wood).
2. Chemical Bond
A chemical bond is the strong electrostatic force of attraction that holds atoms, ions, or molecules together in a compound or crystalline structure. This force results from the rearrangement of valence electrons between atoms—either through the complete transfer of electrons (forming ionic bonds) or the sharing of electrons (forming covalent bonds)—to achieve a more stable, lower-energy electron configuration.
3. Effervescence
Effervescence is the visual process of gas liberation from a liquid, often seen as bubbling or frothing, that occurs during a chemical reaction or when a gas is released from a supersaturated solution. It is commonly observed when carbonates or bicarbonates react with an acid to produce carbon dioxide gas (e.g., adding Alka-Seltzer to water).
4. Precipitate
A precipitate is the insoluble solid that separates out of a liquid solution during a chemical reaction. When two solutions containing dissolved ions are mixed, they may react to form an ionic compound that has extremely low solubility in the solvent, causing it to fall out of the solution and settle at the bottom. This process is known as precipitation.
Question 3.
1. Give an example of a reaction where the following are involved:Heat
2. Give an example of a reaction where the following are involved: Light
3. Give an example of a reaction where the following are involved: Electricity
4. Give an example of a reaction where the following are involved: Close contact
5. Give an example of a reaction where the following are involved: Solution
6. Give an example of a reaction where the following are involved:Pressure
7. Give an example of a reaction where the following are involved: Catalyst
Ans:
- Heat: A classic example is the thermal decomposition of limestone in a lime kiln. When solid calcium carbonate (limestone) is heated strongly, it breaks down to form quicklime (calcium oxide) and releases carbon dioxide gas. This process is fundamental in the manufacturing of cement and steel.
- Light: The reaction between hydrogen and chlorine gas is a photochemical process. When these two colorless gases are mixed in a sealed vessel and exposed to direct sunlight, they combine explosively to form hydrogen chloride gas. This demonstrates how light energy can provide the activation energy needed to initiate a vigorous chemical change.
- Electricity: The electrolysis of acidified water is a direct application of electrical energy. When a direct current is passed through water containing a little sulphuric acid, the water molecules decompose. This results in the liberation of hydrogen gas at the cathode and oxygen gas at the anode, showing the conversion of electrical energy into chemical change.
- Close Contact: The rusting of iron is a reaction that proceeds through intimate contact. When iron is exposed to moist air, a layer of water on its surface allows oxygen from the air to make close and continuous contact with the metal. This leads to the formation of a flaky, brown substance known as hydrated ferric oxide, commonly called rust.
- Solution: The precipitation of silver chloride occurs when two clear solutions are mixed. If a solution of silver nitrate is added to a solution of sodium chloride, an immediate white, curdy precipitate of silver chloride forms and settles out of the solution. This reaction is a common test for the presence of chloride ions.
- Pressure: The industrial synthesis of ammonia via the Haber process relies on high pressure. Nitrogen gas from the air and hydrogen gas, typically from natural gas, are forced to react at a pressure of around 200 atmospheres in the presence of a catalyst. This high pressure favors the formation of ammonia gas from its constituent elements.
- Catalyst: The decomposition of hydrogen peroxide into water and oxygen is dramatically sped up by a catalyst. While hydrogen peroxide decomposes slowly on its own, adding a small amount of manganese dioxide powder causes a rapid effervescence as oxygen gas is released briskly. The manganese dioxide itself remains unchanged and can be recovered after the reaction.
Question 4.
1. Define Photochemical reaction Give an example
2. Define Electrochemical reaction Give an example
Ans:
1. Photochemical Reaction
A photochemical reaction is a chemical process that is initiated when the atoms or molecules of a substance absorb light energy, typically from the sun or an artificial ultraviolet (UV) source. The key principle is that the energy from a photon of light is transferred to the substance, causing electrons to jump to a higher energy state. This new, energized state makes the substance highly reactive, allowing it to undergo chemical changes that would not occur in the dark.
Example: The Fading of Colored Fabric in Sunlight
A common example you can observe is the fading of a brightly colored curtain or t-shirt left in direct sunlight for a long time. The dyes used in the fabric contain molecules that are stable in normal indoor light. However, when these dye molecules absorb high-energy photons from sunlight, they become electronically excited. This excess energy causes chemical bonds within the dye molecule to break or rearrange, transforming it into a different, colorless molecule. The gradual breakdown of countless dye molecules across the fabric is what we perceive as fading.
2. Electrochemical Reaction
An electrochemical reaction is a chemical process where the energy from a spontaneous redox reaction is directly converted into electrical energy, or conversely, where electrical energy is used to drive a non-spontaneous chemical reaction. These reactions always involve the transfer of electrons between a chemical substance and an electrode. They are characterized by the separation of the oxidation (loss of electrons) and reduction (gain of electrons) processes into two half-reactions connected by an external electrical circuit.
Example: The Corrosion of Iron (Rusting)
The formation of rust on an iron nail is a widespread and destructive electrochemical reaction. It requires water and oxygen, but it proceeds via a miniature electrical circuit on the metal’s surface.
On one area of the iron (acting as an anode), oxidation occurs: Iron atoms lose electrons to form iron ions (Fe → Fe²⁺ + 2e⁻).
These freed electrons travel through the moist iron to another area (acting as a cathode), where reduction occurs: Oxygen in the water gains the electrons to form hydroxide ions (O₂ + 2H₂O + 4e⁻ → 4OH⁻).
The iron ions and hydroxide ions then combine to form iron(III) hydroxide, which dehydrates to become the familiar reddish-brown rust (Fe₂O₃·xH₂O). This process naturally converts the chemical energy of the reacting iron and oxygen into a small electrical current, which facilitates the corrosion.
Question 5.
(a) 1. Give an example of the following chemical changes. A photochemical reaction involving silver salt.
2. Give an example of the following chemical changes. A photochemical reaction involving water.
(b) 1. Give an example of the following chemical changes.A reaction involving blue solution.
2. Give an example of the following chemical changes. A reaction involving formation of dirty green precipitate.
3. Give an example of the following chemical change. Two gases combine to form white solid
4. Give an example of the following chemical change. Two solids combine to form a liquid
5. Give an example of the following chemical change. A reaction where colour change is noticed
Ans:
(a) 1. A photochemical reaction involving a silver salt.
Example: The fading of silver-based tattoo pigments when exposed to intense sunlight over time. The pigments, often made from silver salts, undergo a photochemical reduction. Light energy causes the silver ions to gain electrons and form dark colloidal silver particles within the skin, which can lead to the tattoo darkening or changing color.
Chemical Reaction: Ag⁺ (in pigment) + e⁻ → Ag⁰ (colloidal) (driven by light energy)
(a) 2. A photochemical reaction involving water.
Example: The degradation of hydrogen peroxide in sunlight. A bottle of hydrogen peroxide left in a bright window will slowly decompose into water and oxygen gas. The energy from sunlight provides the activation energy needed to break the chemical bonds.
Chemical Reaction: 2H₂O₂ (l) → 2H₂O (l) + O₂ (g) (catalyzed by light)
(b) 1. A reaction involving a blue solution.
Example: The reaction between a bright copper penny and concentrated nitric acid produces a deep blue solution as the copper metal oxidizes. This occurs after initial brown fumes of nitrogen dioxide are released, and the solution turns blue due to the formation of hydrated copper(II) ions.
Chemical Reaction: Cu (s) + 4HNO₃ (aq) → Cu(NO₃)₂ (aq) + 2NO₂ (g) + 2H₂O (l)
(The resulting Cu(NO₃)₂ solution is blue.)
(b) 2. A reaction involving the formation of a dirty green precipitate.
Example: The reaction between iron(II) chloride solution and ammonium hydroxide results in the immediate formation of a dirty green, gelatinous precipitate of iron(II) hydroxide. This precipitate is unstable and will slowly turn brown as it reacts with oxygen in the air.
Chemical Reaction: FeCl₂ (aq) + 2NH₄OH (aq) → Fe(OH)₂ (s) + 2NH₄Cl (aq)
(Dirty Green Precipitate)
3. Two gases combine to form a white solid.
Example: The reaction between ammonia gas and hydrogen chloride gas is a classic demonstration. When the vapors from concentrated hydrochloric acid and ammonium hydroxide solution meet in the air, they instantly react to form a dense, white smoke of fine ammonium chloride crystals.
Chemical Reaction: NH₃ (g) + HCl (g) → NH₄Cl (s)
(White Solid Smoke)
4. Two solids combine to form a liquid.
Example: The reaction between solid calcium chloride and solid potassium hydroxide. When these two dry, white solids are ground together in a mortar, they begin to react, absorbing water vapor from the air to form a syrupy, liquid mixture of calcium hydroxide and potassium chloride.
Observation: Two initial dry solids gradually deliquesce and combine to form a liquid paste.
5. A reaction where a colour change is noticed.
Example: The oxidation of colorless iodide ions to brown iodine. When a colorless solution of potassium iodide is added to an acidified pale orange solution of potassium dichromate, the mixture turns a deep brown color due to the formation of elemental iodine. This is a common test for an oxidizing agent.
Chemical Reaction: Cr₂O₇²⁻ (aq) + 6I⁻ (aq) + 14H⁺ (aq) → 2Cr³⁺ (aq) + 3I₂ (aq) + 7H₂O (l)
(Orange) (Colorless) (Green) (Brown)
Question 6.
1. Write the chemical reaction where the following changes are observed. Gas is evolved
2. Write the chemical reaction where the following changes are observed. Colour change is noticed
3. Write the chemical reaction where the following changes are observed. Precipitate is formed
4. Write the chemical reaction where the following changes are observed. Physical state is changed
Ans:
1. Gas is evolved
When a piece of zinc metal is dropped into a solution of dilute sulfuric acid, a vigorous reaction occurs. This interaction produces zinc sulfate, which remains dissolved in the solution, and hydrogen gas, which is released as visible bubbles. The chemical equation for this reaction is:
Zn(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂(g)↑
2. Colour change is noticed
In a reaction between copper sulfate and iron, a striking colour shift is observed. A shiny iron nail, when placed in a blue copper sulfate solution, slowly develops a reddish-brown coating. This coating is copper metal, while the solution’s colour fades as blue copper ions are replaced by pale green iron ions. The reaction is:
Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)
3. Precipitate is formed
Mixing clear solutions of sodium sulfate and barium chloride results in the immediate formation of a white, insoluble solid. This solid, called a precipitate, is barium sulfate and settles at the bottom of the container. The reaction is represented as:
Na₂SO₄(aq) + BaCl₂(aq) → BaSO₄(s)↓ + 2NaCl(aq)
4. Physical state is changed
The burning of a candle is a classic example where the physical state changes, even though it is also a chemical process. Solid wax, when heated by the flame, melts into a liquid. This liquid wax is then drawn up the wick and vaporizes into a gaseous fuel before combusting. The initial melting is a physical change in state, represented for a typical hydrocarbon wax as:
Wax(s) → Heat → Wax(l) → Wax(g)
Question 7.
1. Give reason for the following: Silver nitrate solution is kept in coloured bottles.
2. Give reason for the following: Molybdenum is used in the manufacture of ammonia.
3. Give reason for the following: The blue solution of copper sulphate changes to green when a piece of iron is added to this solution.
4. Give reason for the following:Colourless concentrated sulphuric acid in a test tube changes to blue on adding a small piece of copper to it.
Ans:
1. Reason for Silver Nitrate Solution being Kept in Coloured Bottles
Silver nitrate is highly photosensitive, meaning it decomposes chemically when exposed to light. The coloured bottles, typically amber or brown, act as a filter that blocks the specific wavelengths of light (primarily in the ultraviolet and blue spectrum) that initiate this decomposition reaction. When exposed to light, the silver ions (Ag⁺) are reduced to metallic silver (Ag), which appears as a blackish-brown precipitate, rendering the solution impure and unfit for quantitative analysis. The coloured glass effectively prevents this photochemical reduction, ensuring the stability and purity of the solution.
2. Reason for Molybdenum being Used in the Manufacture of Ammonia
In the industrial synthesis of ammonia via the Haber-Bosch process, nitrogen and hydrogen gases combine over a solid catalyst. While iron is the primary catalyst, it is “promoted” with other substances to drastically enhance its efficiency and longevity. Molybdenum acts as a structural promoter in some modern catalyst formulations. It helps to stabilize the porous, high-surface-area structure of the iron catalyst against a process called “sintering,” where particles fuse together and lose active surface area at high temperatures and pressures. By maintaining the physical integrity of the catalyst, molybdenum ensures a consistently high reaction rate for ammonia production over a longer operational period.
3. Reason for the Blue Copper Sulphate Solution changing to Green when Iron is added
This colour change is a visual demonstration of a single displacement reaction, driven by the relative positions of metals in the reactivity series. Iron is more reactive than copper. When an iron piece is introduced into the copper sulphate solution, the more reactive iron atoms (Fe) readily lose electrons and displace the less reactive copper ions (Cu²⁺) from the solution. The copper ions are reduced to metallic copper, which deposits on the iron surface, while the iron is oxidized to ferrous ions (Fe²⁺). The blue colour of the solution, characteristic of hydrated copper ions, fades and is replaced by a pale green colour, which is the characteristic colour of hydrated ferrous ions in the solution.
4. Reason for Colourless Concentrated Sulphic Acid turning Blue on adding Copper
Concentrated sulphuric acid is a powerful oxidizing agent, especially at high temperatures. When copper is added, it does not produce hydrogen gas as a dilute acid would with a more reactive metal. Instead, the acid itself gets reduced. The copper metal (Cu) is oxidized to copper(II) ions (Cu²⁺). Simultaneously, the sulphuric acid is reduced to sulphur dioxide (SO₂) gas. The resulting solution contains dissolved copper(II) ions, which in the presence of a small amount of water (often present in the acid or formed during the reaction) form the complex ion [Cu(H₂O)₆]²⁺. This hydrated copper ion is responsible for the characteristic blue colour that appears in the test tube.
Exercise 2 (B)
Question 1.
1. Complete the following statement. The chemical change involving iron and hydrochloric acid illustrates a _________________ reaction.
2. Complete the following statement. In the type of reaction called, two compounds exchange their positive and negative radicals.
3. Complete the following statement. A catalyst either ______ or _____________ the rate of a chemical change but itself remains ______________ at the end of the reaction.
4. Complete the following statement. On heating, hydrated copper sulphate changes its colour from ________ to __________.
Ans:
- The chemical change involving iron and hydrochloric acid illustrates a displacement reaction.
- In the type of reaction called double decomposition, two compounds exchange their positive and negative radicals.
- A catalyst either increases or decreases the rate of a chemical change but itself remains unchanged in mass and composition at the end of the reaction.
- On heating, hydrated copper sulphate changes its colour from blue to white.
Question 2.
When hydrogen burns in oxygen, water is formed; when electricity is passed through water, hydrogen and oxygen are given out. Name the type of chemical change involved in the two cases.
Ans:
When hydrogen gas ignites and reacts with oxygen, the result is the formation of water. This is a classic example of a combination reaction, a process characterized by multiple reactants uniting to create one distinct product. The fundamental nature of this change is the merging of separate elements into a single, more complex compound.
Conversely, if an electric current is passed through water, it can be split back into its core components of hydrogen and oxygen gas. This process demonstrates a decomposition reaction, which acts as the direct opposite of combination. In this scenario, a single compound is broken apart into simpler substances. The driving force for this breakdown is electrical energy, leading to the specific term for this reaction: electrolysis. Electrolysis, therefore, is the precise name for the decomposition of a substance induced by an electric current.
Question 3.
1. Explain, giving one example of the following chemical changes: Double decomposition
2. Explain, giving one example of the following chemical change: Thermal dissociation
3. Explain,giving one example of the following chemical changes: Reversible reaction
4. Explain, giving one example for each of the following chemical changes: Displacement
Ans:
1. Double Decomposition
Explanation:
Double decomposition is a chemical reaction where two compounds react by exchanging their ions or radicals to form two new compounds. It typically occurs in solutions, and one of the new products is often a precipitate (an insoluble solid), a gas, or a weak electrolyte like water.
Example: The Reaction Between Silver Nitrate and Sodium Chloride
When a solution of silver nitrate (AgNO₃) is mixed with a solution of sodium chloride (NaCl), a double decomposition reaction occurs. The positive ions (Ag⁺ and Na⁺) swap their negative partners (NO₃⁻ and Cl⁻).
- Word Equation: Silver Nitrate + Sodium Chloride → Silver Chloride + Sodium Nitrate
- Chemical Equation: AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq)
In this reaction, the silver ions (Ag⁺) from the first compound combine with the chloride ions (Cl⁻) from the second to form a white, insoluble precipitate of silver chloride (AgCl). The sodium ions (Na⁺) and nitrate ions (NO₃⁻) remain in the solution as soluble sodium nitrate (NaNO₃). The formation of the solid precipitate is the driving force behind this reaction.
2. Thermal Dissociation
Explanation:
Thermal dissociation is the process where a single compound breaks down into two or more simpler substances when heated. The key feature is that this decomposition is reversible; when the products are cooled, they recombine to form the original compound.
Example: The Heating of Ammonium Chloride
When solid ammonium chloride (NH₄Cl) is heated, it does not melt but instead breaks down directly into two gases: ammonia and hydrogen chloride.
- Word Equation: Ammonium Chloride ⇌ Ammonia + Hydrogen Chloride
- Chemical Equation: NH₄Cl(s) ⇌ NH₃(g) + HCl(g)
If you heat ammonium chloride in one end of a long tube, you will see a white solid (the recombined ammonium chloride) forming in the cooler part of the tube. This is because the gaseous ammonia and hydrogen chloride diffuse and, upon cooling, readily recombine to form the original solid ammonium chloride. This reversibility is the hallmark of thermal dissociation.
3. Reversible Reaction
Explanation:
A reversible reaction is one where the products can react with each other to re-form the original reactants. The reaction can proceed in both the forward and backward directions simultaneously under the same conditions. It is denoted by a double arrow (⇌) in a chemical equation.
Example: The Reaction for the Haber Process (Formation of Ammonia)
The industrial production of ammonia from nitrogen and hydrogen is a classic reversible reaction. The forward reaction (forming ammonia) is exothermic (releases heat), while the reverse reaction (decomposing ammonia) is endothermic (absorbs heat).
- Word Equation: Nitrogen + Hydrogen ⇌ Ammonia
- Chemical Equation: N₂(g) + 3H₂(g) ⇌ 2NH₃(g) + Heat
In a closed container, nitrogen and hydrogen gases will react to form ammonia. However, as soon as some ammonia is formed, it starts to decompose back into nitrogen and hydrogen. The reaction never goes to completion; instead, it reaches a state of balance called equilibrium, where the rates of the forward and reverse reactions are equal.
4. Displacement
Explanation:
Displacement is a chemical reaction where a more reactive element (a better “pusher” of its own electrons) kicks out a less reactive element from its compound. The more reactive element takes the place of the less reactive one.
Example: Iron Nail in Copper Sulfate Solution
When a clean iron nail is placed into a blue copper sulfate (CuSO₄) solution, a displacement reaction occurs. Iron is more reactive than copper, so iron atoms displace copper ions from the solution.
- Word Equation: Iron + Copper Sulfate → Iron Sulfate + Copper
- Chemical Equation: Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)
As the reaction proceeds, the blue color of the copper sulfate solution fades (as Cu²⁺ ions are removed), and a brownish coating of metallic copper forms on the iron nail. The new solution contains greenish iron(II) sulfate. This reaction visually demonstrates the reactivity series of metals.
Question 4.
(a) What is synthesis?
(b) What kind of chemical reaction is synthesis? Support your answer by an example.
Ans:
(a) What is synthesis?
Synthesis is a fundamental chemical process where two or more simple substances, known as reactants, combine to form a single, more complex product. The core idea is one of building and unification, where multiple components come together to create a new compound with properties entirely different from the original ingredients. This process is a direct application of the principle that matter can be rearranged to create new materials, which is central to both chemistry and the formation of complex molecules in biological systems.
(b) What kind of chemical reaction is synthesis? Support your answer by an example.
Synthesis is classified as a combination reaction. This classification is based on its defining characteristic of multiple reactants merging into one product. A classic and clear example is the formation of iron sulfide. In this reaction, the element iron (Fe) chemically unites with the element sulfur (S) when heat is supplied. The result is an entirely new substance, iron sulfide (FeS), which is a dark, brittle solid completely unlike the shiny, malleable iron or the yellow, powdery sulfur it was made from. The chemical equation, Fe + S → FeS, perfectly illustrates the essence of a combination reaction.
Question 5.
Decomposition brought about by heat is known as thermal decomposition. What is the difference between thermal dissociation and thermal decomposition?
Ans:
The terms thermal decomposition and thermal dissociation both refer to chemical breakdowns caused by heat, but they are distinguished by the reversibility of the reaction.
Thermal Decomposition
Thermal decomposition is a chemical reaction in which a single compound breaks down into two or more simpler, different substances upon heating, and the reaction is typically irreversible.
| Aspect | Description |
| Reversibility | Irreversible. The products formed usually cannot readily recombine to form the original substance simply by cooling. |
| Conditions | Requires a continuous supply of heat (activation energy) to maintain the breakdown of stable bonds. |
| Example | The thermal decomposition of Calcium Carbonate (CaCO3): CaCO3HeatCaO+CO2 |
Thermal Dissociation
Thermal dissociation is a specific type of thermal decomposition where the breakdown is a reversible process. The products can readily recombine to form the original reactant simply by changing the conditions, such as cooling the system or increasing the pressure. This process often reaches a state of chemical equilibrium.
| Aspect | Description |
| Reversibility | Reversible. The reaction can proceed in both the forward (dissociation) and backward (recombination) directions. |
| Conditions | Is an equilibrium reaction that is highly dependent on temperature and pressure. |
| Example | The thermal dissociation of Ammonium Chloride (NH4Cl): NH4ClHeatingCoolingNH3+HCl |
Key Difference Summary
The core difference is that thermal dissociation is reversible and reaches equilibrium, while general thermal decomposition is irreversible.
Question 6.
(a) Define the neutralization reaction with an example.
(b) Give a balanced equation for this reaction.
(c) Give three applications of neutralization reactions.
Ans:
(a) Definition and Example
A neutralization reaction is a specific type of chemical process where an acid and a base, when combined, effectively cancel out each other’s characteristic properties. The reaction consistently produces two new substances: a salt and water. The driving force behind this reaction is the combination of hydrogen ions (H⁺) from the acid and hydroxide ions (OH⁻) from the base to form neutral water molecules (H₂O).
For example, when hydrochloric acid, which is found in our stomach gastric juices, reacts with sodium hydroxide, a common base, the result is sodium chloride, which is ordinary table salt, and water.
(b) Balanced Chemical Equation
The balanced equation for the reaction between hydrochloric acid and sodium hydroxide is:
HCl + NaOH → NaCl + H₂O
This equation demonstrates a one-to-one ratio where one molecule of acid reacts with one molecule of base to yield one molecule of salt and one molecule of water.
(c) Three Applications of Neutralization Reactions
- Treating Soil Acidity for Agriculture: Excessive rainfall can leach away basic nutrients, leaving farmland too acidic for optimal crop growth. To remedy this, farmers often spread slaked lime (calcium hydroxide), a base, over their fields. This treatment neutralizes the excess acidity in the soil, creating a more favorable chemical environment for healthy plant development.
- Relieving Acidity in the Human Stomach: Overproduction of hydrochloric acid in the stomach leads to indigestion and discomfort. To alleviate this, people take antacid tablets, which contain mild bases like magnesium hydroxide or sodium bicarbonate. These bases react with and neutralize the excess stomach acid, providing relief from the burning sensation.
- Reducing the Effects of Acidic Factory Waste: The wastewater from industrial factories and mines is often highly acidic. If released directly into rivers and soil, it would cause severe environmental damage. To prevent this, the waste is first treated with a mild base, such as calcium carbonate (limestone), to neutralize the acid before the water is discharged, thereby protecting the ecosystem.
Question 7.
What do you understand about the precipitation reaction? Explain with examples.
Ans:
At its heart, a precipitation reaction is a type of chemical “showdown” in a liquid solution that results in the formation of a solid. This solid, which often looks like a cloudy powder or tiny crystals, is what we call the precipitate. It’s the star of the show.
Think of it like mixing salt and sand in water. The salt dissolves and disappears, but the sand sinks to the bottom. In a precipitation reaction, the “sand” is created during the reaction itself.
The Core Principle: The Swap and Drop
These reactions almost always involve two ionic compounds dissolved in water. Ions are atoms or molecules that have an electrical charge. When dissolved, these compounds break apart into their positive and negative ions, moving around freely.
When we mix these two solutions, the positive ions from one solution can meet and partner up with the negative ions from the other solution. If this new partnership forms a compound that is insoluble (meaning it cannot dissolve in water), it can no longer stay in the liquid. It “falls out” of the solution as a solid precipitate.
A handy way to predict this is by using a solubility table, which tells us which ionic combinations are soluble (dissolve) and which are not.
The Classic Example: The Cloudy Mixture
One of the most famous examples is the reaction between silver nitrate and sodium chloride (common table salt).
- The Solutions:
- Silver Nitrate Solution (AgNO₃): This contains free-moving Ag⁺ (silver) ions and NO₃⁻ (nitrate) ions.
- Sodium Chloride Solution (NaCl): This contains free-moving Na⁺ (sodium) ions and Cl⁻ (chloride) ions.
- The Mixing: When we pour these two clear, colorless solutions together, all four types of ions (Ag⁺, NO₃⁻, Na⁺, Cl⁻) are now swimming around in the same beaker.
- The “Swap”: The ions swap partners. The silver ions (Ag⁺) are strongly attracted to the chloride ions (Cl⁻) to form silver chloride (AgCl). Meanwhile, the sodium ions (Na⁺) pair with the nitrate ions (NO₃⁻) to form sodium nitrate (NaNO₃).
- The “Drop”: Now, we check the solubility:
- Sodium Nitrate (NaNO₃): According to solubility rules, all nitrates are soluble. So, NaNO₃ stays dissolved in the water.
- Silver Chloride (AgCl): Most chloride salts are soluble, but a key exception is silver chloride. AgCl is insoluble.
Because AgCl cannot dissolve, its molecules clump together to form tiny solid particles. This is what makes the clear solution instantly turn cloudy with a white, chalky-looking precipitate.
We can write the reaction like this:
AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq)
The (aq) means “aqueous” (dissolved in water), and the (s) means “solid” – our precipitate!
Another Common Example: The Kidney Stone Connection
Another great example is the reaction between calcium chloride and sodium carbonate, which mimics how some kidney stones form.
- The Solutions:
- Calcium Chloride (CaCl₂): Contains Ca²⁺ and Cl⁻ ions.
- Sodium Carbonate (Na₂CO₃): Contains Na⁺ and CO₃²⁻ ions.
- The Mixing and Swap: When mixed, the calcium ions (Ca²⁺) meet the carbonate ions (CO₃²⁻).
- The “Drop”: Checking solubility, we find that calcium carbonate (CaCO₃) is insoluble. It is the same compound that makes up limestone, eggshells, and some types of kidney stones.
A white, gritty precipitate of calcium carbonate forms immediately.
CaCl₂(aq) + Na₂CO₃(aq) → CaCO₃(s) + 2NaCl(aq)
Why Are These Reactions Important?
Precipitation reactions are not just classroom experiments; they are crucial in many real-world applications:
- Water Purification: Undesirable ions are removed from water by causing them to form insoluble precipitates that can be filtered out.
- Analytical Chemistry: Scientists use these reactions to identify the presence of specific ions in an unknown sample. The color and texture of the precipitate act as a fingerprint.
- Wastewater Treatment: Toxic heavy metals like lead or mercury are often removed by converting them into insoluble precipitates.
Question 8.
1. What is a double displacement reaction?
2. Give an example of a double displacement reaction, where a gas is evolved.
Ans:
- A double displacement reaction is a type of chemical process where two ionic compounds in an aqueous solution exchange their constituent ions. This swap results in the formation of two new compounds. For a reaction to be noticeable, one of these new products must be insoluble (forming a solid precipitate), a gas that bubbles out, or a stable molecule like water.
- A common example where a gas is evolved is the reaction between hydrochloric acid and sodium sulfide. When these two compounds are mixed, they exchange ions. This leads to the formation of sodium chloride and hydrogen sulfide gas. The hydrogen sulfide gas is immediately recognizable by its distinct smell, similar to that of rotten eggs. The reaction can be represented as follows:
2HCl (aq) + Na₂S (aq) → 2NaCl (aq) + H₂S (g)↑
Question 9.
1. What is a decomposition reaction?
2. Decomposition reactions can occur by heat. Give two balanced reactions for each.
3. Decomposition reactions can occur by electricity. Give balanced reactions.
4. Decomposition reactions can occur by sunlight. Give two balanced reactions.
Ans:
1. What is a decomposition reaction?
A decomposition reaction is a chemical process where a single compound breaks down into two or more simpler substances. These simpler products can be individual elements or new compounds. This type of reaction is essentially the opposite of a combination or synthesis reaction. For the breakdown to occur, energy is always required, which can be supplied in the form of heat, light, or electricity.
2. Decomposition reactions can occur by heat. Give two balanced reactions for each.
Decomposition triggered by heat is known as a thermal decomposition reaction.
Reaction 1: Decomposition of Limestone (Calcium Carbonate)
When limestone is heated strongly, it decomposes into quicklime (calcium oxide) and carbon dioxide gas.
Word Equation: Calcium Carbonate → Calcium Oxide + Carbon Dioxide
Balanced Chemical Equation:
CaCO₃(s) → CaO(s) + CO₂(g)
Reaction 2: Decomposition of Potassium Chlorate
Heating potassium chlorate in the presence of a catalyst like manganese dioxide leads to its decomposition, producing potassium chloride and oxygen gas.
Word Equation: Potassium Chlorate → Potassium Chloride + Oxygen
Balanced Chemical Equation:
2KClO₃(s) → 2KCl(s) + 3O₂(g)
3. Decomposition reactions can occur by electricity. Give balanced reactions.
Decomposition caused by electric current is called electrolysis.
Reaction: Electrolysis of Water
When an electric current is passed through acidified water, it decomposes into hydrogen and oxygen gases.
Word Equation: Water → Hydrogen + Oxygen
Balanced Chemical Equation:
2H₂O(l) → 2H₂(g) + O₂(g)
Reaction: Electrolysis of Molten Sodium Chloride
Passing electricity through molten (liquid) sodium chloride decomposes it into elemental sodium metal and chlorine gas.
Word Equation: Sodium Chloride → Sodium + Chlorine
Balanced Chemical Equation:
2NaCl(l) → 2Na(l) + Cl₂(g)
4. Decomposition reactions can occur by sunlight. Give two balanced reactions.
Decomposition initiated by sunlight is known as a photochemical decomposition reaction.
Reaction 1: Decomposition of Silver Chloride
Silver chloride, a white solid, is unstable in sunlight. It decomposes to form silver metal, which appears greyish-purple or black, and chlorine gas.
Word Equation: Silver Chloride → Silver + Chlorine
Balanced Chemical Equation:
2AgCl(s) → 2Ag(s) + Cl₂(g)
Reaction 2: Decomposition of Hydrogen Peroxide
While hydrogen peroxide decomposes slowly on its own, the reaction is significantly accelerated by exposure to light, breaking it down into water and oxygen.
Word Equation: Hydrogen Peroxide → Water + Oxygen
Balanced Chemical Equation:
2H₂O₂(l) → 2H₂O(l) + O₂(g)
Question 10.
1. State the type of reactions for the following represent and balance the ones that are not balanced. Cl2 + 2KBr → 2KCl + Br2
2. State the type of reactions for the following represent and balance the ones that are not balanced. NaOH + HCl → NaCl + H2O
3. State the type of reactions for the following represent and balance the ones that are not balanced. 2HgO → 2Hg + O2
4. State the type of reactions for the following represent and balance the ones that are not balanced. Fe + CuSO4→ FeSO4 + Cu
5. State the type of reactions for the following represent and balance the ones that are not balanced. PbO2 + SO2→ PbSO4
6. State the type of reactions for the following represent and balance the ones that are not balanced. 2KClO3→ 2KCl + 3O2
7. State the type of reactions for the following represent and balance the ones that are not balanced. 2H2O2→ 2H2O + O2
8. State the type of reactions for the following represent and balance the ones that are not balanced. KNO3 + H2SO4 → HNO3 + KHSO4
9. State the type of reactions for the following represent and balance the ones that are not balanced.CuO+H2→ Cu+ H2O
10. State the type of reactions for the following represent and balance the ones that are not balanced.CaCO3→ CaO+ CO2
11. State the type of reactions for the following represent and balance the ones that are not balanced. NH4Cl → NH3 + HCl
12. State the type of reactions for the following represent and balance the ones that are not balanced. PbO + 2HNO3→ Pb(NO3) + 2H2O
13. State the type of reactions for the following represent and balance the ones that are not balanced.AgNO3 + NaCl → AgCl + NaNO3
Ans:
- Cl₂ + 2KBr → 2KCl + Br₂
- Reaction Type: Displacement Reaction (Chlorine is more reactive than bromine and displaces it from its salt).
- NaOH + HCl → NaCl + H₂O
- Reaction Type: Neutralization Reaction (a specific type of double decomposition where an acid and a base react to form salt and water).
- 2HgO → 2Hg + O₂
- Reaction Type: Decomposition Reaction (a single compound breaks down into two simpler substances).
- Fe + CuSO₄ → FeSO₄ + Cu
- Reaction Type: Displacement Reaction (Iron, being more reactive than copper, pushes it out of its compound).
- PbO₂ + 2SO₂ → 2PbSO₄
- Reaction Type: This is a combination reaction where lead dioxide and sulfur dioxide combine to form a single product, lead sulfate. It also involves a change in oxidation states, making it a Redox reaction.
- Balanced Equation: PbO₂ + 2SO₂ → PbSO₄ + SO₃? (The given product is incorrect for a simple combination. A more accurate representation for forming lead sulfate from these reactants requires a redox process: PbO₂ + 2SO₂ → PbS₂O₆ or it occurs in multiple steps. However, if the intended product is PbSO₄, the reaction is not straightforward and requires specific conditions.)
- 2KClO₃ → 2KCl + 3O₂
- Reaction Type: Decomposition Reaction (a complex substance breaks down into two simpler substances upon heating).
- 2H₂O₂ → 2H₂O + O₂
- Reaction Type: Decomposition Reaction (this breakdown is often sped up by a catalyst like manganese dioxide).
- KNO₃ + H₂SO₄ → HNO₃ + KHSO₄
- Reaction Type: Double Decomposition Reaction (also a precipitation reaction if concentrated acids are used, leading to the evolution of HNO₃ vapors). The equation is already balanced.
- CuO + H₂ → Cu + H₂O
- Reaction Type: Redox Reaction (specifically, a reduction-oxidation reaction where copper oxide is reduced, and hydrogen is oxidized). The equation is already balanced.
- CaCO₃ → CaO + CO₂
- Reaction Type: Decomposition Reaction (this breakdown occurs when limestone is strongly heated). The equation is already balanced.
- NH₄Cl → NH₃ + HCl
- Reaction Type: Decomposition Reaction (ammonium chloride breaks down upon heating into ammonia and hydrogen chloride gases). The equation is already balanced.
- PbO + 2HNO₃ → Pb(NO₃)₂ + H₂O
- Reaction Type: Neutralization Reaction (a base, lead oxide, reacts with an acid to form salt and water).
- Balanced Equation: The product should be lead nitrate, which has the formula Pb(NO₃)₂. The correct balanced equation is PbO + 2HNO₃ → Pb(NO₃)₂ + H₂O.
- AgNO₃ + NaCl → AgCl + NaNO₃
- Reaction Type: Double Decomposition Reaction (the ions exchange partners, and a white precipitate of silver chloride is formed, making it a precipitation reaction). The equation is already balanced.
Question 11.
Match the following:
| a. Zn(s) + H2SO4(aq) → ZnSO4(aq) + H2(g) | i. Photochemical decomposition |
| b. 2AgCI(s) Sunlight 2Ag(s)+CI2(g) | ii. Thermal decomposition |
| c. 2KCI Electricity 2K+CI 2 | iii. Displacement reaction |
| d. 2HgO(s) Δ 2Hg(s)+O(g) | iv. Electrolytic decomposition |
Ans:
| Reaction | Type of Reaction |
| a. Zn(s)+H2SO4(aq)→ZnSO4(aq)+H2(g) | iii. Displacement reaction |
| b. 2AgCl(s)Sunlight2Ag(s)+Cl2(g) | i. Photochemical decomposition |
| c. 2KClElectricity2K+Cl2 | iv. Electrolytic decomposition |
| d. 2HgO(s)Δ2Hg(s)+O2(g) | ii. Thermal decomposition |
Question 12.
1. Multiple choice: Which of the following is not a characteristic of a chemical change?
- It is irreversible.
- No net energy change is involved.
- New substances are formed.
- Involves absorption or liberation of energy.
Question 2.
Multiple choice: A reaction of a type: AB + CD → AD + CD, involves
- No chemical change
- Decomposition of AB and CD
- Exchange of ions of AB and CD
- Combination of AB and CD
Question 3.
Multiple choice: The reaction BaCl2(aq) + H2SO4(aq) → BaSO4(s) + 2HCl(aq) is
- Displacement reaction
- Neutralisation reaction
- Decomposition reaction
- Double displacement reaction
Question 4.
Multiple choice: Thermal decomposition of sodium carbonate will produce
- Carbon dioxide
- Oxygen
- Sodium hydroxide
- No other product
Exercise 2 (C)
Question 1.
What is a chemical change? Give two examples of chemical change?
Ans:
A chemical change is a fundamental process where one or more substances are transformed into entirely new materials with different chemical properties and compositions. This transformation, also called a chemical reaction, is not superficial; it involves the breaking of chemical bonds within the original substances (reactants) and the formation of new bonds to create the resulting substances (products). Because a new chemical identity is formed, the change is typically permanent and cannot be reversed by simple physical means. Key indicators that a chemical change has occurred include the production of a gas, the formation of a solid precipitate, a distinct color change, or a noticeable release or absorption of energy as heat or light.
Two clear examples of chemical change are:
- The Rusting of Iron: When a piece of iron is left exposed to damp air, it slowly reacts with oxygen and water vapor. This process transforms the hard, shiny metallic iron into a flaky, reddish-brown substance called hydrated iron(III) oxide, commonly known as rust. The rust has a completely different composition and is weaker and more brittle than the original iron, demonstrating a permanent chemical transformation.
- The Burning of Wood: When wood is set on fire, it undergoes a rapid chemical reaction with oxygen in the air called combustion. The complex organic molecules in the wood, like cellulose, break down and recombine with oxygen to form new substances, primarily carbon dioxide gas and water vapor, along with a residue of ash. The wood is permanently converted into these new substances, releasing heat and light in the process, and cannot be turned back into its original form.
Question 2.
Why is energy involved in a chemical change?
Ans:
At its heart, every chemical change is a dramatic rearrangement of atoms. They break away from their existing partners and form new, more stable connections. Energy is the absolute requirement for this entire process to begin and to conclude. It is the currency spent to break old bonds and the dividend paid when new ones form.
To understand this, imagine a chemical reaction as a journey over a hill. The substances you start with, the reactants, are in one valley. The substances you end up with, the products, are in another, lower valley. Energy is involved in two critical stages of this journey.
First, energy must be invested. Even if the final products are more stable, the initial reactants are comfortably settled. The bonds holding their atoms together are strong. To break these bonds and allow the atoms to be available for new connections, a surge of energy is required. This initial, uphill part of the journey is known as the activation energy. It’s the push needed to get the whole process started, like the effort required to roll a boulder to the top of a hill before it can roll down the other side. A common example is the heat from a match needed to initiate the burning of paper. The paper will burn and release vast energy, but it needs that initial spark to overcome its stable, unburned state.
Second, energy is released. Once the old bonds are broken and the atoms are free, they immediately seek new, more stable arrangements. As they form these new bonds, they release energy. This is because the new bonds are stronger and exist at a lower, more stable energy level than the old ones. The energy released as the system “rolls downhill” into the new valley is often greater than the initial energy investment. This net release is what we feel as heat from a fire or the explosive force from a firework.
The overall character of a reaction—whether it feels hot or cold—depends on the balance between this energy investment and dividend.
In an exothermic reaction, the energy released from forming the new, strong bonds is greater than the energy needed to break the old ones. There is a net release of energy, usually as heat or light. The surroundings get warmer. A burning log is a classic example; it releases far more thermal energy than the small amount from the match used to light it.
In an endothermic reaction, the opposite is true. The new bonds in the products are weaker, or less stable, than the original bonds. The energy required to break the old bonds is greater than the energy released when the new ones form. This creates a net absorption of energy from the surroundings, which we feel as cooling. Photosynthesis is a prime example, where plants absorb solar energy to convert carbon dioxide and water into sugar.
Question 3.
What do you understand about ‘chemical reaction’?
Ans:
At its heart, a chemical reaction is a process of profound transformation, much like a caterpillar becoming a butterfly. It begins with a set of substances known as reactants, each with its own unique identity and properties. During the reaction, the bonds that hold the atoms within these reactants together are broken. This breakdown then allows the atoms to rearrange themselves, forming new connections and bonds. The outcome of this atomic reshuffling is an entirely new set of substances called products, which possess chemical and physical characteristics distinct from the original reactants. This fundamental reorganization of atoms is the true essence of a chemical change.
We can witness this invisible dance of atoms through observable clues in the macroscopic world. Tell-tale signs often signal that a chemical reaction is underway. These include a noticeable change in colour, the formation of a solid precipitate within a liquid, the release of gas bubbles, or a significant shift in temperature, either releasing warmth or drawing heat from the surroundings. For instance, when iron is exposed to moist air, it reacts with oxygen to form rust, a brittle, reddish-brown compound that is entirely different from the original, strong metal. This formation of rust is a classic, visible manifestation of the atoms of iron and oxygen rearranging to create a new chemical substance.
To describe these transformations clearly and concisely, scientists use a symbolic shorthand known as a chemical equation. This equation must be balanced, meaning the number of atoms for each element is identical on both sides. This balancing act is not just a mathematical exercise; it is a direct representation of the Law of Conservation of Mass, which states that matter cannot be created or destroyed in a chemical reaction. The atoms themselves are merely the players being rearranged on a molecular stage, and the chemical equation is the script that documents their new roles and relationships.
Question 4.
1. Give an example of a reaction where the following is involved . Evolution of heat
2. Give an example of a reaction where the following is involved . Absorption of heat
3. Give an example of a reaction where the following is involved . High pressure is required
Ans:
1. Evolution of Heat (Exothermic Reaction)
A classic and easily observable example is the reaction involved in a single-use hand warmer. The rusting of iron is the core process, but it is dramatically sped up.
- Reaction: The rapid oxidation of iron powder in the presence of air (oxygen), water, and a catalyst like salt.
- Chemical Change: Iron + Oxygen → Iron Oxide (Rust)
- Why Heat is Evolved: Breaking the bonds in the iron and oxygen molecules requires energy, but forming the new, stronger bonds in the iron oxide releases a significantly larger amount of energy. This net release of energy surges into the surroundings as palpable heat, warming your hands on a cold day.
2. Absorption of Heat (Endothermic Reaction)
A common kitchen experiment demonstrates this perfectly. The reaction feels cold to the touch because it pulls heat from its environment, including your skin.
- Reaction: Mixing baking soda (sodium bicarbonate) with household vinegar (acetic acid).
- Chemical Change: Acetic Acid + Sodium Bicarbonate → Sodium Acetate + Water + Carbon Dioxide
- Why Heat is Absorbed: While the fizzing from the carbon dioxide is obvious, the reaction has a hidden energy cost. The energy required to break the bonds in the vinegar and baking soda is greater than the energy released from forming the new bonds in the products. This energy deficit is covered by absorbing thermal energy from the immediate surroundings, causing a noticeable drop in temperature.
3. High Pressure is Required
This is essential for the industrial-scale production of ammonia, a critical component in agricultural fertilizers, through the Haber-Bosch process.
- Reaction: The direct combination of nitrogen gas from the air and hydrogen gas (often from natural gas).
- Chemical Change: Nitrogen + Hydrogen → Ammonia
- Why High Pressure is Required: The reaction takes four molecules of reactants (1 N₂ + 3 H₂) and turns them into two molecules of product (2 NH₃). This is a decrease in the number of gas molecules. According to fundamental principles, applying a very high pressure (around 200 atmospheres) favors the side of the reaction with fewer gas molecules. It literally squeezes the nitrogen and hydrogen together, making them much more likely to collide and react to form ammonia, making an otherwise slow and inefficient process commercially viable.
Question 5.
State the main characteristics of chemical reactions. Give at least one example in each case.
Ans:
A tell-tale signal that a chemical reaction is underway is the sudden appearance of a solid where none existed before, known as a precipitate. This happens as the newly formed substances possess their own unique physical characteristics, which are often entirely distinct from the reactants that created them. A clear illustration of this can be seen by passing hydrogen sulfide gas, recognizable by its foul, sulfurous odor, into a solution of copper sulfate. The result is not a mixture but an immediate chemical transformation, yielding a stark black solid of copper sulfide that was not present initially.
Another reliable marker is the release of a gas, a process that frequently partners with a measurable shift in temperature. Reactions that draw in heat from their environment, leaving it feeling cooler, are classified as endothermic. A simple experiment using household items demonstrates this: combining citric acid from a lemon with sodium bicarbonate (baking soda) in water triggers a fizzing reaction as carbon dioxide gas bubbles out. Simultaneously, the container feels cold to the touch because the reaction is pulling thermal energy from the water and the surrounding air to fuel the process.
Perhaps one of the most visually straightforward clues is a lasting alteration in color. This is particularly convincing when the change is permanent and cannot be undone by simple physical means like heating or filtering. A familiar example from daily life is seen when a fresh apple slice is left on a counter. The slow transition of the exposed flesh from a pale white to a rusty brown is not spoilage in the conventional sense, but a specific chemical process called oxidation, where compounds within the fruit react with atmospheric oxygen to generate new molecules with a brown pigment.
Question 6.
1. Give an example of the following chemical changes. A reaction involving
(i) Change of state
(ii) formation of precipitate
2. Give an example of the following chemical changes.An exothermic and an endothermic reaction involving carbon as one of the reactants.
3. Give an example of the following chemical changes.A reaction where colour change is noticed.
Ans:
1. Examples of Reactions Involving:
(i) Change of state: Burning of Wax (a candle)
When a wax candle burns, the wax (a hydrocarbon) melts from a solid to a liquid, and then the liquid wax is drawn up the wick. Here, it undergoes a chemical change (combustion) with oxygen in the air to produce two new substances that are gases: carbon dioxide and water vapour. This is a clear example of a chemical reaction producing a change in state, as solid reactants are converted into gaseous products.
Word Equation: Wax + Oxygen → Carbon Dioxide + Water Vapour + Heat and Light
(ii) Formation of a precipitate: Reaction of Washing Soda and Epsom Salt
A common and easy-to-observe precipitation reaction occurs between solutions of washing soda (sodium carbonate) and Epsom salt (magnesium sulfate). When their solutions are mixed, a double displacement reaction takes place, forming a white, cloudy solid of magnesium carbonate that is insoluble in water.
Word Equation: Sodium Carbonate + Magnesium Sulfate → Sodium Sulfate + Magnesium Carbonate (white precipitate)
2. Examples of Exothermic and Endothermic Reactions involving Carbon:
(i) Exothermic Reaction: Combustion of Charcoal
This is a classic and highly exothermic reaction. When charcoal (which is primarily carbon) is burned in air, it reacts with oxygen to form carbon dioxide gas. This reaction releases a significant amount of energy in the form of heat and light, which is why we use it for barbecuing and heating.
Word Equation: Carbon + Oxygen → Carbon Dioxide + Heat
(ii) Endothermic Reaction: Production of Water Gas
When carbon in the form of coke (a coal derivative) is passed over red-hot steam, an endothermic reaction occurs. The carbon reacts with water vapour, absorbing a large amount of thermal energy from the surroundings to produce a mixture of carbon monoxide and hydrogen gases, known as “water gas.” This reaction requires continuous heat input to proceed.
Word Equation: Carbon + Water Vapour → Carbon Monoxide + Hydrogen (absorbs heat)
3. Example of a Reaction where Colour Change is Noticed: The Blue Bottle Experiment
A fascinating and visually striking reaction is the “Blue Bottle Experiment.” In this demonstration, a colorless solution of glucose and sodium hydroxide containing a small amount of methylene blue indicator is prepared in a sealed bottle. Upon shaking the bottle, the solution instantly turns a deep blue color. When left to stand, the blue color slowly fades back to colorless. This cycle can be repeated many times.
The chemistry involves the oxidation of glucose by atmospheric oxygen. Shaking the bottle dissolves oxygen into the solution, which oxidizes the colorless methylene blue (its reduced form) into its blue form. Upon standing, the glucose reduces the methylene blue back to its colorless form. The constant change from colorless to blue and back again provides a very clear and dramatic example of a colour change in a chemical reaction.
Process Description: Shaking a colorless alkaline glucose solution with methylene blue introduces oxygen, causing a chemical oxidation that turns the solution blue. On standing, a reduction reaction occurs, and the solution becomes colorless again.
Question 7.
Define exothermic and endothermic changes. Give two example in each case.
Ans:
While many everyday processes give off heat, others actively absorb it, and this fundamental difference classifies them as exothermic or endothermic. Exothermic changes are characterized by a net release of energy, typically as heat or light, into the environment. This occurs because the energy required to break the bonds in the reactants is less than the energy released when new bonds form in the products. The result is a rise in temperature in the immediate surroundings. Common illustrations include the combustion of fuels like wood in a fireplace, where the chemical reaction produces noticeable warmth and light. Similarly, the chemical process of respiration within our cells is exothermic, constantly releasing the thermal energy that helps maintain our body temperature.
Conversely, endothermic changes proceed by drawing energy in from their surroundings, which usually causes a detectable drop in temperature. For these reactions to occur, the energy needed to break existing chemical bonds is greater than the energy released from the new bonds being formed, creating an overall energy deficit. This necessitates a continuous inflow of external energy for the reaction to be sustained. A prime biological example is photosynthesis, where plants harness solar energy to power the transformation of carbon dioxide and water into complex sugars. On a more industrial scale, the operation of an instant cold pack is based on an endothermic reaction; mixing certain chemicals, such as ammonium nitrate with water, rapidly absorbs heat, providing a portable cooling effect for injuries.
Question 8.
State the effects of endothermic and exothermic reactions on the surroundings.
Ans:
Many everyday chemical processes are defined by their thermal interaction with the environment. Exothermic reactions are characterized by a net release of energy, often felt as heat. During these reactions, the energy required to break the bonds in the reactants is less than the energy liberated when new bonds form in the products. This energy surplus manifests as a noticeable increase in temperature in the immediate surroundings. Common examples that demonstrate this principle include the combustion of fuels like wood or propane, and the process of respiration in living organisms, where both situations result in a distinct and measurable release of thermal energy.
Conversely, endothermic reactions operate in an opposite manner, requiring a continuous absorption of energy from their surroundings to proceed. In these cases, the energy needed to dismantle the existing bonds in the reactants surpasses the energy given off during the creation of new product bonds. This creates an energy deficit, leading the reaction to draw heat from its environment, which results in a perceptible cooling effect. Practical instances of this are seen in the dissolution of ammonium nitrate in water or the chemical process within instant cold packs used for sports injuries, where a distinct drop in temperature is the primary observable outcome.
Question 9.
1. Define: Photochemical reaction Give one example
2. Define: Electrochemical reaction Give one example
Ans:
1. Photochemical Reaction
Core Concept: This is a chemical transformation that begins or speeds up when a substance captures energy from light, most often from the visible or ultraviolet part of the spectrum. The absorbed light energy provides a “push” that allows molecules to reorganize, break apart, or form new bonds in ways that would not happen on their own without light.
Illustrative Example: A classic example is the fading of dyes and pigments in fabrics, posters, or artwork when they are exposed to direct sunlight over time. The vibrant colors in these materials are often from complex organic molecules. When photons of light are absorbed, they provide enough energy to break specific chemical bonds within these dye molecules. This breakage alters the molecule’s structure, changing the wavelengths of light it absorbs and reflects, which we perceive as the color fading or bleaching.
2. Electrochemical Reaction
Core Concept: This describes a process where chemical energy and electrical energy are directly interchanged. It is fundamentally governed by the movement of electrons. In essence, a chemical reaction can be set up in a way that forces electrons to travel through an external path, creating an electrical current. Conversely, electrical energy can be forced into a chemical system to provoke a chemical change that would not occur naturally.
Illustrative Example: The process of chrome plating (electroplating) a car part or a bathroom fixture is a direct application. In this setup, the object to be plated is submerged in a solution containing chromium ions and connected to the negative terminal of a power source. A piece of chromium metal is connected to the positive terminal. When electric current is applied, it drives a non-spontaneous reaction: chromium atoms from the metal source lose electrons (oxidize), becoming ions in the solution, while chromium ions in the solution gain electrons (reduce) onto the surface of the object, forming a thin, shiny, and protective metallic layer.
Question 10.
1. Complete and balance the following reaction: NaCl(aq) + AgNO3(aq) →
2. Complete and balance the following reaction: Pb(NO3)2 + KI →
3. Complete and balance the following reaction: CuCO3Δ →
4. Complete and balance the following reaction: Pb(NO3)2Δ →
5. Complete and balance the following reaction: NH3+O2Δ →
Ans:
1. NaCl(aq) + AgNO₃(aq) →
This is a classic double displacement reaction, also known as a precipitation reaction. The silver ion (Ag⁺) and the chloride ion (Cl⁻) combine to form an insoluble solid, silver chloride, which appears as a white precipitate.
The complete and balanced reaction is:
NaCl(aq) + AgNO₃(aq) → AgCl(s) + NaNO₃(aq)
2. Pb(NO₃)₂ + KI →
This is another double displacement precipitation reaction. The lead ion (Pb²⁺) reacts with the iodide ion (I⁻) to form a bright yellow precipitate of lead iodide.
The balanced reaction requires two molecules of potassium iodide to provide the necessary iodide ions for one molecule of lead nitrate.
Pb(NO₃)₂(aq) + 2KI(aq) → PbI₂(s) + 2KNO₃(aq)
3. CuCO₃ → (with heat, Δ)
This is a decomposition reaction driven by heat. Many metal carbonates break down upon strong heating to yield the metal oxide and carbon dioxide gas.
The complete and balanced reaction is:
CuCO₃(s) → CuO(s) + CO₂(g)
4. Pb(NO₃)₂ → (with heat, Δ)
This is the thermal decomposition of lead nitrate. It breaks down to form lead monoxide, nitrogen dioxide (a reddish-brown gas), and oxygen gas.
The balanced reaction is:
2Pb(NO₃)₂(s) → 2PbO(s) + 4NO₂(g) + O₂(g)
5. NH₃ + O₂ → (with heat, Δ)
This represents the combustion of ammonia. However, without a specific catalyst, the primary reaction in excess oxygen is the formation of nitrogen monoxide and water vapor. This is a key step in the Ostwald process for manufacturing nitric acid.
The balanced reaction for this oxidation is:
4NH₃(g) + 5O₂(g) → 4NO(g) + 6H₂O(g)
Question 11.
1. What do you observe? When Lead nitrate is heated.
2. What do you observe? When Silver chloride is exposed to sunlight.
3. What do you observe? When Hydrogen peroxide is exposed to sunlight.
4. What do you observe? When H2S gas is passed through copper sulphate solution.
5. What do you observe? When Barium chloride is added to sodium sulphate solution
6. What do you observe? When water is added to the quick lime.
7. What do you observe? When Sodium chloride solution is added to silver nitrate solution.
Ans:
1. When Lead nitrate is heated.
A crackling sound is often heard initially. The white crystalline solid changes color, becoming yellow while hot due to the formation of lead (II) oxide. Upon cooling, the residue turns to a pale yellow or buff color. A reddish-brown gas (nitrogen dioxide) is evolved, which can be identified by its characteristic choking odor, and a colorless gas (oxygen) that can rekindle a glowing splinter.
2. When Silver chloride is exposed to sunlight.
The white or off-white solid gradually darkens, taking on a greyish-violet or purplish-grey hue. This color change starts in patches and eventually spreads throughout the entire sample. The compound appears to decompose and lose its pristine white form.
3. When Hydrogen peroxide is exposed to sunlight.
Tiny, invisible gas bubbles form steadily within the solution. There is no dramatic color change or precipitation; the primary observation is the slow, silent decomposition of the liquid, releasing oxygen gas. If a glowing splinter is brought near the mouth of the container, it will be reignited, confirming the gas is oxygen.
4. When H₂S gas is passed through copper sulphate solution.
The beautiful transparent blue solution becomes cloudy and then transforms into a precipitate with a very distinctive black coloration. The final mixture appears as an opaque, inky black suspension, completely obscuring the original blue color of the copper ions.
5. When Barium chloride is added to sodium sulphate solution.
Upon mixing the two clear, colorless solutions, an immediate, milky white precipitate forms, making the mixture opaque and resembling diluted milk. This precipitate is insoluble and settles slowly to the bottom if left undisturbed.
6. When water is added to the quick lime.
A vigorous hissing and sputtering sound is produced, accompanied by a significant release of heat, making the container hot to the touch. The solid quicklime (calcium oxide) slakes and crumbles, swelling up to form a fine, white, dry powder (calcium hydroxide). The process is highly exothermic and can sometimes produce steam.
7. The reaction between the two clear, colorless solutions is instantaneous, resulting in the formation of a thick, curdy white precipitate. This precipitate swirls through the liquid before gradually settling at the bottom of the container as a white mass.
Question 12.
1. Name a carbonate which does not decompose on heating.
2. Name a nitrate which produces oxygen as the only gas.
3. Name a compound which produces carbon dioxide on heating.
4. Name a nitrate which produces brown gas on heating.
Ans:
- A carbonate which does not decompose on heating: Sodium Carbonate (Na₂CO₃). Unlike most other carbonates, it is thermally stable and requires an extremely high temperature to decompose, which is not achieved by normal laboratory heating.
- A nitrate which produces oxygen as the only gas: Sodium Nitrate (NaNO₃). When heated strongly, it decomposes to sodium nitrite and releases only oxygen gas (2NaNO₃ → 2NaNO₂ + O₂).
- A compound which produces carbon dioxide on heating: Calcium Carbonate (CaCO₃). This is a common compound found in limestone and marble. Upon heating, it breaks down into calcium oxide and carbon dioxide gas (CaCO₃ → CaO + CO₂).
- A nitrate which produces brown gas on heating: Lead Nitrate (Pb(NO₃)₂). Its decomposition produces nitrogen dioxide (NO₂), which is a distinctive brown-colored gas (2Pb(NO₃)₂ → 2PbO + 4NO₂ + O₂).
