Chemical Reactions

The “Chemical Reactions” chapter lays the groundwork for understanding how substances transform. It delves into the definition of a chemical reaction as a process involving the breaking and forming of bonds to create new products. Students learn to identify reactions by observing characteristics such as color changes, temperature shifts (exothermic or endothermic), gas formation, or precipitate formation.

The chapter then categorizes reactions into key types:

• Combination, where substances merge (A + B → AB).
• Decomposition, where a compound breaks down (AB → A + B).

Emphasis is also placed on representing these changes through chemical equations, including details like physical states and reaction conditions. A crucial skill taught is balancing chemical equations, ensuring that atomic quantities remain consistent on both sides, thus upholding the Law of Conservation of Mass. Finally, the chapter distinguishes between reactions that release energy (exothermic) and those that absorb it (endothermic).

Exercise – I

Question 1.
(a) Define a chemical reaction.
(b) What happens during a chemical reaction?
(c) What do you understand about a chemical bond?
Answer:(a) what’s a chemical reaction all about?

Put simply, a chemical reaction is just a process where certain substances are changed into completely new ones. Imagine it like playing with LEGOs – the initial structures are taken apart, and the individual pieces are then put back together in different ways to create entirely new structures.   

(b)what actually occurs when a chemical reaction takes place?

When a chemical reaction happens, a few key events unfold. First, the chemical “ties” that hold the atoms together in the starting materials (what we call reactants) need to be broken. Once these ties are undone, the individual atoms become free to move around. Then, they start forming new chemical “ties,” linking up in different configurations to produce the new substances (which we call products). Energy usually plays a role in this entire process – it might be taken in to break the initial ties or given off when the new ones are formed.  

(c) what’s the story with a chemical bond, then?

Well, a chemical bond is basically the pulling force that keeps two or more atoms connected to form a stable molecule or a crystal structure. It’s like an atomic-level “stickiness.” This attraction comes about from how the atoms’ tiny particles with negative charge interact with each other. These interactions can involve sharing these particles (in covalent bonds) or the electrical attraction between oppositely charged particles (in ionic bonds). Ultimately, these “sticky” connections are what give substances their unique characteristics.

Question 2.
Give one example each of which illustrates the following characteristics of a chemical reaction:
(a) evolution of a gas
(b) change of colour
(c) change in state
Answer:
The Release of Gas

You’ve likely seen this if you’ve ever mixed baking soda and vinegar. The immediate fizzing and bubbling you observe is carbon dioxide gas being rapidly produced and released. This effervescence is a strong signal that a chemical reaction is taking place.

A Change in Color

When substances undergo a chemical reaction, their color can change. Gradually, the blue color of the solution will fade, often turning a murky green, while a reddish-brown coating of copper will appear on the nail. These simultaneous color shifts in both the liquid and the solid indicate that a chemical reaction, specifically the displacement of copper by iron, has happened.

Formation of a Solid (Precipitate)

Another clear sign of a chemical reaction is the creation of a new substance in a different physical state, such as a solid appearing from two liquids. This sudden appearance of a solid from liquid reactants unmistakably shows a change in state, confirming a chemical reaction has occurred.

Question 3.
How do the following help in bringing about a chemical change?
(a) pressure (b) light
(c) catalyst (d) heat.
Answer:
Factors Affecting Reaction Rates

• Pressure: Increasing pressure forces molecules closer together, leading to more frequent and powerful collisions. This enhanced interaction helps break old bonds and form new ones, speeding up the reaction.
  
• Light: Molecules can absorb light energy, which makes them more reactive. This absorbed energy can weaken or even break chemical bonds, kicking off a chemical reaction (e.g., how sunlight fades colors).
  
• Catalyst: A truly remarkable factor is a catalyst. Imagine a shortcut on a long journey. That’s essentially what a catalyst does for a reaction. It doesn’t get consumed in the process, but it provides an alternative route for the reaction to occur that requires much less energy. By lowering the energy hurdle, the reaction can proceed much more quickly and easily.
  
• Heat: Adding heat gives molecules more kinetic energy, causing them to move faster and collide with greater force. This extra energy makes it easier to break existing bonds and create new ones, thereby speeding up the chemical reaction.

Question 4.
(a) Define catalyst.
(b) What are (i) positive catalysts and (ii) negative catalysts? Support your answer with one example for each of them.
(c) Name three biochemical catalysts found in the human body.

Ans :

(a) Define catalyst. A catalyst is a substance that alters (usually increases) the rate of a chemical reaction without itself being consumed in the reaction. It participates in the reaction but is regenerated at the end, allowing it to be used repeatedly.

(b) What are (i) positive catalysts and (ii) negative catalysts? Support your answer with one example for each of them.

(i) Positive Catalysts: It does this by lowering the activation energy barrier for the reaction, thereby allowing more reactant molecules to convert into products in a given time.

• Example: In the decomposition of potassium chlorate (KClO₃) to potassium chloride (KCl) and oxygen (O₂), manganese dioxide (MnO₂) acts as a positive catalyst. Without MnO₂, the decomposition requires a very high temperature; with MnO₂, it occurs rapidly at a much lower temperature. 2KClO3​(s)MnO2​​2KCl(s)+3O2​(g)

(ii) Negative Catalysts (Inhibitors): A negative catalyst, more commonly referred to as an inhibitor, is a substance that decreases or retards the rate of a chemical reaction. They work by increasing the activation energy, by poisoning a positive catalyst, or by removing a reactive intermediate.

• Example: In the decomposition of hydrogen peroxide (H₂O₂) to water (H₂O) and oxygen (O₂), glycerol acts as a negative catalyst (inhibitor). It slows down the decomposition, helping to preserve hydrogen peroxide for longer periods. 2H2​O2​(aq)Glycerol​2H2​O(l)+O2​(g) (Here, glycerol slows down the reaction significantly compared to its natural decomposition rate).

(c) Name three biochemical catalysts found in the human body. Biochemical catalysts found in the human body are called enzymes. Here are three examples:

1. Amylase: Breaks down complex carbohydrates (starch) into simpler sugars (e.g., in saliva, pancreatic juice).
2. Pepsin: Breaks down proteins into smaller peptides in the stomach.
3. Lipase: Breaks down fats (lipids) into fatty acids and glycerol (e.g., pancreatic lipase).

Question 5.
What do you observe when
(a) dilute sulphuric acid is added to granulated zinc?
(b) a few pieces of iron are dropped in a blue solution of copper sulphate?
(c) silver nitrate is added to a solution of sodium chloride?
(d) ferrous sulphate solution is added to an aqueous solution of sodium hydroxide.
(e) solid lead nitrate is heated?
(f) when dilute sulphuric acid is added to the barium chloride solution?
Answer:

(a) Zinc and Dilute Sulphuric Acid: When zinc reacts with dilute sulfuric acid, you observe fizzing (hydrogen gas), the test tube feels warm (exothermic), and the zinc eventually disappears, leaving a clear solution of zinc sulfate.

(b) Iron in Blue Copper Sulphate: Dropping iron into blue copper sulfate solution causes a reddish-brown copper coating to form on the iron. The blue color of the solution fades, turning pale green as iron displaces copper, forming iron(II) sulfate.

(c) Silver Nitrate and Sodium Chloride Solutions: Mixing clear silver nitrate and sodium chloride solutions instantly produces a white precipitate of silver chloride, with a clear sodium nitrate solution remaining.

(d) Ferrous Sulphate and Sodium Hydroxide Solutions: Adding ferrous sulfate to sodium hydroxide solution forms a murky green precipitate of iron(II) hydroxide. If exposed to air, this green solid slowly turns brownish due to oxidation to iron(III) hydroxide.

(e) Heating Solid Lead Nitrate:Heating solid lead nitrate crystals produces a dramatic visual effect. You’ll see reddish-brown fumes of nitrogen dioxide gas being released, along with invisible oxygen gas. As the reaction progresses, the solid lead nitrate will decompose, leaving behind a yellowish solid residue of lead(II) oxide in the test tube.

(f) Dilute Sulphuric Acid and Barium Chloride Solution: Combining dilute sulfuric acid and barium chloride solution immediately forms a thick, white precipitate of barium sulfate, which is highly insoluble, leaving behind a solution of hydrochloric acid.

Question 6.
Complete and balance the following chemical equations:

Answer:

Exercise – II

Question 1.
1. Fill in the blanks.
(a) A reaction in which two or more substances combine to form a single substance is called a _________reaction.

Ans : combination
(b) A _________ is a substance which changes the rate of a chemical reaction without undergoing a chemical change.

Ans : catalyst
(c) The formation of gas bubbles in a liquid during a reaction is called _________

Ans : effervescence
(d) The reaction between an acid and a base is called a ___________

Ans : neutralization reaction.

(e) Soluble bases are called __________.

Ans : alkalis
(f) The chemical change involving iron and hydrochloric acid illustrates a _________reaction.

Ans : displacement
(g) In the type of reaction called ____________ two compounds exchange their positive and negative radicals respectively.

Ans : double decomposition reaction, ions
(h) A catalyst either____________the rate of a chemical change but itself remains______ at the end of the reaction.

Ans : increases or decreases , unchanged
(i) The chemical reaction between hydrogen and chlorine is a__________reaction

Ans : combination
(j) When a piece of copper is added to a silver nitrate solution, it turns ________ in colour.

Ans : blue

Question 2.
Classify the following reactions as a combination, decomposition, displacement, precipitation, and neutralization. Also, balance the equations.


Answer:

Question 3.

Define:
(a) precipitation (b) neutralization (c) catalyst
Answer:
(a) Precipitation is when a solid substance, known as a precipitate, forms and separates out of a liquid solution during a chemical reaction. Think of it like when you mix two clear liquids and suddenly see cloudy bits appearing and settling at the bottom.
Acid + Base → Salt + Water
Example.

(b) Neutralization: A chemical reaction where an acid and a base react, losing their individual properties to form a salt and usually water.

(c) A catalyst is a substance that speeds up a chemical reaction without itself being permanently changed or consumed in the process. It’s like a facilitator that helps the reaction happen faster but remains unchanged at the end, ready to help more reactions along.

Question 4.
Explain the following types of chemical reactions giving two examples for each of them.
(a) combination reaction
(b) decomposition reaction
(c) displacement reaction
(d) double decomposition reaction
Answer:
(a) Combination Reaction:  It’s like combining flour and eggs to make a batter – the individual components are no longer distinct but have merged into something unified. For instance, when hydrogen gas and oxygen gas are ignited, they combine to produce water. Similarly, the rusting of iron is a combination reaction where iron reacts with oxygen and water to form iron oxide.

Example:

• When carbon undergoes combustion in the presence of oxygen, it forms the gaseous compound carbon dioxide.

(b) Decomposition Reaction:  Imagine taking a completed LEGO model and breaking it back down into its individual bricks. This often requires an input of energy, such as heat or electricity. For example, applying heat to calcium carbonate (limestone) causes it to decompose into calcium oxide (quicklime) and carbon dioxide gas.

Example:

• When mercuric oxide is heated, it decomposes to yield two distinct elements: mercury and oxygen.

(c) Displacement Reaction: Picture a dance floor where a more confident dancer cuts in and takes the spot of a less confident one. This usually happens when a more reactive element replaces a less reactive one in a compound. A classic example is when an iron nail is placed in a copper sulfate solution; the more reactive iron displaces the copper, leading to the formation of iron sulfate and the deposition of copper on the nail.

Examples:

• Zinc is more reactive than copper, so it displaces copper from copper sulfate solution, forming zinc sulfate and solid copper: Zn + CuSO$_4$ (aq) → ZnSO$_4$ (aq) + Cu
• Similarly, when an iron piece is added to a copper sulfate solution, iron displaces copper: Fe + CuSO$_4$ → FeSO$_4$ + Cu.

(d) Double Decomposition Reaction (also known as Metathesis or Double Displacement): This type of reaction is like a “partner swap” between two compounds. Often, one of the new compounds formed is insoluble and precipitates out of the solution as a solid. 

Examples:

• When silver nitrate reacts with hydrochloric acid, silver chloride (a precipitate) and nitric acid are formed: AgNO$_3$ + HCl → AgCl + HNO$_3$ (aq)

Question 5.
Write the missing reactants and products and balance the equations.

Answer:

Question 6.
How will you obtain it?
(a) Magnesium oxide from magnesium.
(b) Silver chloride from silver nitrate.
(c) Nitrogen dioxide from lead nitrate.
(d) Zinc chloride from zinc.
(e) Ammonia from nitrogen.
Also, give balanced equations for the reactions
Answer:
(a) Magnesium oxide from magnesium:
To make magnesium oxide, you just need to burn a strip of magnesium metal in air. Magnesium reacts quickly with oxygen when it’s heated, producing a bright white flame.

(b) Silver chloride from silver nitrate:
This is a simple double displacement reaction. If you mix a solution of silver nitrate with another solution that has chloride ions, like sodium chloride or even hydrochloric acid, you’ll see a white solid form.

(c) Nitrogen dioxide from lead nitrate:
When lead nitrate is heated, it decomposes. During this breakdown, it gives off nitrogen dioxide gas, which you can recognize by its reddish-brown color. At the same time, solid lead oxide and oxygen gas are also produced. The gas makes it very easy to spot this reaction.

(d) Zinc chloride from zinc:
If you put a piece of zinc metal into hydrochloric acid, it reacts right away. You’ll also notice bubbles of hydrogen gas escaping during the reaction.

(e) Ammonia from nitrogen:
To get ammonia from nitrogen, industries use something called the Haber-Bosch process. It involves reacting nitrogen gas with hydrogen gas under high pressure and high temperature, using an iron catalyst. This setup helps make ammonia on a large scale, which is mainly used in fertilizers.

Question 7.
What do you observe when
(a) Iron nail is kept in copper sulphate solution for some time.
(b) Phenolphthalein is added to sodium hydroxide solution.
(c) Blue litmus paper is dipped in dilute hydrochloric acid.
(d) Lead nitrate is heated.
(e) Magnesium ribbon is burnt in oxygen.
(f) Ammonia is brought in contact with hydrogen chloride. gas.
Answer:
(a) Imagine dropping a common iron nail into a bright blue copper sulphate solution. The nail itself starts to get a reddish-brown crust – that’s copper metal, freshly deposited. At the same time, the vibrant blue liquid gradually loses its intensity, becoming a pale, almost watery green. What’s happening here is that the iron, being a more assertive metal, is essentially pushing the copper out of its chemical partnership in the solution, forming a new compound: iron sulphate.

(b) If you were to add a few drops of phenolphthalein, which is a chemical indicator, to a clear solution of sodium hydroxide, you’d see an immediate and striking color change.This is because phenolphthalein reacts specifically with basic substances like sodium hydroxide, signaling their presence with that characteristic hue.

(c) The instant it touches the acid, you’d observe a quick and definite change: the blue paper would immediately transform into a bright red. 

(d) When you apply heat to lead nitrate, a white solid, you’ll notice a few things. First, there might be a subtle crackling sound as it begins to break down. Then, as the heating continues, the white powder will change, forming a new, yellow solid (which is lead oxide). Simultaneously, you’ll see distinctive reddish-brown fumes wafting away – these are nitrogen dioxide gases being released.

(e) If you hold a magnesium ribbon in a flame, especially in an oxygen-rich environment, it won’t just burn; it will ignite with an incredibly dazzling, bright white light. It’s truly brilliant to behold. As it burns, it transforms into a soft, fine, white powder. This powder is magnesium oxide, the result of magnesium combining rapidly with oxygen in a reaction that releases a significant amount of both heat and light.

(f) When you bring ammonia gas and hydrogen chloride gas together, even without a flame or other obvious catalyst, something remarkable happens. You’ll instantly see dense, white fumes forming where the two gases meet. These aren’t just misty vapors; they’re tiny solid particles of ammonium chloride, created as the two gases directly combine to form a new compound.

Question 8.
Give reason:
(a) A person suffering from acidity is advised to take an antacid.
(b) Acidic soil is treated with quick lime.
(c) Wasp sting is treated with vinegar.
Answer:
(a) So, when your stomach feels all churned up because of too much acid, antacids come to the rescue. These medicines are the opposite of acid – they’re basic, or alkaline. When you swallow an antacid, it’s like sending in a neutralizer to calm things down in your stomach by reacting with and canceling out that extra acid. It brings you relief by balancing things out.

(b) Imagine soil that’s too acidic – it’s not a happy place for most plants to grow strong and healthy. That’s where quick lime, which is calcium oxide, can help. It’s a basic substance, so when you sprinkle it on acidic soil, it gets to work by reacting with the acid and taming it. It’s all about getting the soil’s pH just right.

(c) Now, a wasp sting is usually alkaline, which is why it can be so irritating. But here’s a simple trick: vinegar is acidic. This neutralization can really ease the pain and that uncomfortable feeling. It’s like using the opposite chemical to soothe the sting.

Question 9.
What is meant by the metal reactivity series? State its importance, (any two points).
Answer:
The metal reactivity series arranges metals by their tendency to react, with the most reactive at the top, easily forming positive ions. Hydrogen is included to show its reactions with metals.

1.Predicting Displacement Reactions: It helps predict which metal can displace another from its compound; a more reactive metal will always displace a less reactive one.

2.Understanding Metal Extraction: It guides metal extraction methods. Highly reactive metals need energy-intensive processes like electrolysis, while less reactive ones can be found naturally or extracted simply.

Question 10.
What are oxides? Give two examples of each of the following oxides.
(a) Basic oxide (b) Acidic oxide
(c) Amphoteric oxide (d) Neutral oxide
Answer:
Oxides are compounds featuring oxygen and at least one other element, usually with oxygen in a -2 oxidation state. They are classified by their reactions:

(a) Basic Oxides: These are oxides that react with acids to form salt and water. They are typically formed by metals.

Sodium oxide (Na2​O)

Calcium oxide (CaO)

(b)Acidic oxides (non-metal-derived) react with bases to produce salt and water, frequently forming acids in aqueous solutions.

Carbon dioxide (CO₂)
Sulphur dioxide (SO₂)

(c)Amphoteric oxides exhibit dual nature, reacting as both acids and bases to form salt and water.

Aluminium oxide (Al₂O₃)
Zinc oxide (ZnO)

(d)Neutral oxides display no acidic or basic characteristics and remain unreactive with either acids or bases.

Carbon monoxide (CO)
Nitrous oxide (N₂O)

Question 11.
Define exothermic and endothermic reactions. Give two examples of each.
Answer:
Exothermic reactions are like tiny heaters, releasing energy (usually as heat) and warming their surroundings. Think of burning natural gas or mixing strong acids and bases – they both produce heat because the starting materials have more energy than the products.

Endothermic reactions are like little refrigerators, absorbing energy (often as heat) and cooling their surroundings. Melting ice or dissolving ammonium nitrate in water are examples where heat is taken in, making things feel colder because the starting materials have less energy than the products.

Question 12.
State the effect of:
(a) an endothermic reaction
(b) an exothermic reaction on the surroundings.
Answer:
a) Picture an endothermic reaction as a process where the system actively draws in heat from everything around it, almost as if it’s sucking up the warmth. This energy intake by the system inevitably results in the surroundings losing energy, which we perceive as them getting colder.

(b) This release of energy causes the temperature of the immediate surroundings to increase, making them feel warmer.

Question 13.
What do you observe when
(a) an acid is added to a basic solution.
(b) ammonium chloride is dissolved in water.
Answer:(a) When you mix an acid and a base, neutralization happens. The solution will warm up (it’s exothermic!), and if there’s an indicator, you’ll see a color change. Also, the base’s signature slippery texture will vanish.

(b) Dissolving ammonium chloride in water makes the solution feel significantly colder.