Chapter -1

Chemical Reactions and Chemical Equations

1.1 Physical Change and Chemical Change

The change in which only the physical properties of the substances get changed, and no new substance is formed is called a physical change. For example, melting of ice, getting salt from sea water, etc.

The change in which the composition and chemical properties of the substances get changed and one or more new substances are formed is called chemical change. The properties of the new substances formed are different from those of the original substance. For example, when coal is burnt, carbon dioxide is formed which have completely different properties from those of carbon.

1.1.1 Difference between Physical change and Chemical change

1.2 Chemical reaction

A chemical reaction is a change in which one or more substances(s) or reactant(s) react(s) to form new substance(s) with entirely different properties.

• The substances which take part in a chemical reaction are called reactants.
• The substances formed in a chemical reaction are called products.

For example,

Notes:

1. During any chemical reaction, atoms retain their identity, i.e., the atoms remain unchanged. The transformation of the reactants into products takes place due to the rearrangement of atoms.
2. During a chemical reaction, chemical bonds in the reactant molecules get broken and new chemical bonds are formed between them to give products which have completely different properties than those of reactants.
3. In a chemical reaction, the amount of reactants decreases and that of products increases.
4. The rate of chemical reaction depends on
   • Physical state of the reactants
   • Temperature, pressure and concentration of the reactants
   • Catalyst

1.2.1 Identification of Chemical Reaction

The following observations help us to determine whether a chemical reaction has taken place-

1. Change in physical state
2. Change in colour
3. Evolution of a gas
4. Change in temperature
5. Formation of precipitate

For example,

(i) When a mixture of hydrogen and oxygen is ignited with an electric spark at room temperature, liquid water is formed.

(ii) When chlorine water is added to the KI solution, brown colour appears.

(iii) When a metal like Zn, Fe, or Mg reacts with H₂SO₄, H₂ gas is evolved.

(iv) When coal is burnt, heat and light are produced.

(iv) When silver nitrate solution is mixed with a solution of sodium chloride, white precipitate of silver chloride is formed.

1.3 Chemical Equation

A chemical equation is a shortened representation of a chemical reaction in terms of names, symbols, and formulae of the substances involved.

Consider a chemical reaction such as when a magnesium ribbon is burnt in oxygen, it gets converted to magnesium oxide. This reaction can by the following word equation:

Such long word equation can be shortened by using symbols and formulae of the substances involved in the reaction as

Other examples are:

1.3.1 Writing a chemical equation

We should follow the following steps:

(i) Identify the reactants and the products of the chemical reaction.

(ii) Write down the formulae or symbols of the reactants on the left-hand side and those of products on the right-hand side of an arrow with a sign of plus (+) between them.

(iii) The arrow head points towards the products and shows the direction of the reaction.

For example, combustion of methane gas in the presence of oxygen gas gives carbon dioxide gas and water vapour.

Here, reactants CH₄ and O₂ are written on the left of the arrow with plus (+) sign, and products CO₂ and H₂O on the right of the arrow with plus (+) sign.

1.3.2 Making a Chemical Equation More Informative

A chemical equation can be made more informative by adding additional information to the chemical equation. This is done as follows:

(i) Reaction conditions: The information regarding temperature, pressure and catalyst etc., is provided above the arrow separating the reactants and products. For example,

(ii) Physical states of reactants and products: The information regarding physical states of the reactants and products can be provided by using letters (s), (l), (g), and (aq) for solid, liquid, gas and a solution in water respectively.

(iii) Heat absorbed or evolved: Chemical reactions proceed with the absorption or evolution of heat. The information is provided by adding a ‘heat’ term on the product side of the chemical equation.

The reactions in which heat is absorbed are called endothermic reactions.

The reactions in which heat is released are called exothermic reactions.

(iv) Concentration of reactants and products: The information regarding the concentration is provided by adding the dil. (for dilute) or conc. (for concentrated) before the formulae of the reactants and products.

(v) Rate of reaction: This information is provided by writing the term ‘slow’ or ‘fast’ over arrow for corresponding reactions.

1.3.3 Unbalanced chemical equation

The chemical equation in which the number of atoms of one or more elements on the reactant and product sides are not equal is called an unbalanced chemical equation. For example, the equation for the reaction of hydrogen and oxygen is

On comparing both sides, we get

Here, the number of H atoms on both sides is equal but the number of oxygen atoms is unequal. Thus the equation written above is unbalanced equation.

1.3.4 Balanced chemical equation

The chemical equation in which the number of atoms each element on the reactant and product side is equal is called the balanced chemical equation.

The balanced chemical equation is based on the law of conservation of mass. According to the law of conservation of mass, ‘mass can neither be created nor be destroyed during a chemical reaction’. It means that the number of atoms of each element remains the same, before and after a chemical reaction. For example,

On comparing both sides, we get

Here, the number of atoms of all the elements involved in the reaction is equal on both the reactant and product sides.

1.4 Balancing of Chemical Equations

The method by which the number of atoms of each element on both sides of the arrow in a chemical reaction are made equal, is called balancing of chemical equation.

Balancing of a chemical equation is necessary because no matter i.e., no atom is lost or gained during a chemical reaction.

1.4.1 Trial and Error Method

There are several steps involved in balancing a chemical equation. These steps are-

Step 1. Writing unbalanced equation and enclosing the formulae in brackets.

Step 2. Making list of number of atoms of different elements as present in unbalanced equation.

Step 3. Balancing the first element. (To start balancing with the compound that contains the maximum number of atoms. In that compound, select the element which has the maximum number of atoms.)

Step 4. Balancing the second element and the third and so on.

Step 5. Checking the correctness of equation.

Let us try to balance the following chemical equation-

Step 1. Writing unbalanced equation and enclosing the formulae in brackets.

Step 2. Making list of number of atoms of different elements as present in unbalanced equation.

Step 3. Balancing oxygen (O) atom because it comes with the greatest number.

Thus, after balancing the oxygen atom on both sides, the equation becomes

Step 4. Balancing hydrogen (H) atom.

Thus after balancing the hydrogen atom, we get

Step 5. Balancing iron (Fe) atom.

Thus, after balancing the iron (Fe) atoms, we get

Step 6. Checking the correctness of the equation.

The numbers of atoms of elements on both sides of the above chemical equation are equal. This equation is now balanced.

1.4.2 Assuming Method

This balancing method of chemical equations includes assigning algebraic variables as stoichiometric coefficients to every species in the unbalanced chemical equation. And, these variables are used in mathematical equations and solved to obtain the values of each stoichiometric coefficient.

Let us write the unbalanced equation of a chemical reaction.

For the coefficients, we write a, b, c, d, etc. before each species of the above equation.

Write linear equations for each element taking arrow as equal sign as

For Ca: a = 3c ………….(i)

For C: a = d …………..(ii)

For O:3a + 4b = 8c + 3d ………….(iii)

For H: 3b = 2d ………….(iv)

For P: b = 2c ……….(v)

Now take one of the coefficients is equal to 1. Suppose we take, c = 1, then

from equation (i), we get a = 3 x 1 = 3, and

from equation (v), we get b = 2 x 1 = 2

since we have a = 3, from equation (ii), we get d = 3

So we now have values of all variables. So our balanced equation will be

1.5 Types of Chemical Reactions

The chemical reactions are classified into different types depending upon the type of chemical changes taking place. They are –

1. Combination reaction
2. Decomposition reaction
3. Exothermic and Endothermic reactions
4. Displacement reaction
5. Double-displacement reaction
6. Precipitation reaction
7. Oxidation and Reduction reactions

1.5.1 Combination reaction

A reaction in which two or more substances (elements or compounds) combine to form a single product, is called a combination reaction. Most of the addition reactions are exothermic reactions. For example, Quick lime (calcium oxide) reacts vigorously with water to produce slaked lime (calcium hydroxide) releasing a large amount of heat.

In this reaction, calcium oxide and water combine to form a single product, calcium hydroxide. So, it is a combination reaction.

Note: A solution of slaked line produced is used for whitewashing walls. Calcium hydroxide reacts slowly with the carbon dioxide in the air to form a thin layer of calcium carbonate on the walls. Calcium carbonate is formed after two or three days of whitewashing and gives a shiny finish to the walls. The reaction between calcium hydroxide and carbon to form calcium carbonate is taken place as

Other examples of combination reactions are –

(i) When coal is burnt in the excess oxygen carbon dioxide gas is formed.

(ii) When hydrogen gas reacts with oxygen gas water is formed.

(iii) Hydrogen gas and chlorine gas combine in the presence of sunlight to form hydrogen chloride gas.

(iv) Moist ammonia reacts with moist hydrogen chloride to form ammonium chloride.

(v) Phosphorus trichloride reacts with chlorine gas to form phosphorus pentachloride.

1.5.2 Decomposition reactions

A reaction in which a substance is broken down into two or more simpler substances is known as decomposition reaction.

Decomposition reaction is opposite of a combination reaction.

Decomposition reactions take place only when some energy in the form of heat, light or electricity is supplied to the substance.

Various types of decomposition reactions

(i) Thermal decomposition: The decomposition reaction caused by heating is called thermal decomposition. For example,

(a) Ferrous sulphate, the green colour crystals (FeSO₄.7H₂O) lose water of crystallisation when heated and forms dirty white dehydrated FeSO₄. It then decomposes to Ferric oxide (Fe₂O₃), Sulphur dioxide (SO₂) and Sulphur trioxide (SO₃). Ferric oxide is a solid, while other two are gases.

(b) When limestone (CaCO₃) is heated strongly, it decomposes to give quicklime (CaO) and carbon dioxide (CO₂).

(c) Lead nitrate, on heating, decomposes to give yellow lead monoxide, nitrogen dioxide and oxygen gas.

(ii) Photochemical decomposition or photolysis: The decomposition of a substance in the presence of light is called photochemical decomposition. For example,

(a) When silver chloride exposes to sunlight it decomposes to give silver metal and chlorine gas.

Similarly, when silver bromide exposes to sunlight it decomposes to silver metal and bromine gas.

These decomposition reactions of salts of Ag are used in black and white photography.

(b) Hydrogen iodide (HI) decomposes in sunlight to give hydrogen gas and iodine gas.

(iii) Electrolytic decomposition: When a substance decomposes when electricity is passed through it, it is called electrolytic decomposition. For example,

(a) When electric current is passed through pure water, it decomposes into to give oxygen gas and hydrogen gas.

(b) When electric current is passed through molten sodium chloride (common salt), it decomposes to give sodium metal and chlorine gas.

1.5.3 Exothermic and Endothermic reactions

(a) Exothermic reactions: Those reactions in which heat is released along with the formation of products are called exothermic reactions. For example, addition reactions are exothermic reactions.

(i) Addition of water on quick lime produces large amount of heat

(ii) Burning of coal in the presence of excess oxygen produces heat and light.

(iii) Formation of water from H₂(g) and O₂ (g)

Examples, other than addition reactions, are

(i) Burning of natural gas (methane) produces large amount of energy.

(ii) Respiration: During digestion, carbohydrates from different foods like potatoes, rice, bread, etc. are broken down to form glucose. This glucose combines with oxygen in the cell of our body and provides energy. Hence, this is an exothermic reaction.

(iii) The decomposition of vegetable matter into compost is also an example of exothermic reaction.

(b) Endothermic reactions: Those reactions which occur by the absorption of heat/energy (either in the form of light or electricity) are known as endothermic reactions. For example, all decomposition reactions are endothermic reactions. Such reactions requires energy either in the form of heat, light or electricity for breaking down the reactants. Examples other than decomposition reactions are

(i) Reaction between nitrogen and oxygen: Nitrogen reacts with oxygen at about 3000 ⁰C to form nitric oxide.

(ii) Preparation of water gas: When steam is passed over red hot coke, water gas is formed.

(iii) Photosynthesis: It is the process by which plants use sunlight, water, and carbon dioxide to create oxygen and energy in the form of sugar.

1.5.4 Displacement reaction

A reaction in which a more reactive element displaces less reactive element from its compound is called displacement reaction.

A + BC → AC + B

Reactivity series of metals in decreasing order:

Potassium (K) > Sodium (Na) > Calcium (Ca) > Magnesium (Mg) > Aluminium (Al) > Carbon (C) > Zinc (Zn) > Iron (Fe) > Tin (Sn) > Lead (Pb) Hydrogen (H) > Copper (Cu) > Silver (Ag) > Gold (Au)

In short memorable line, it is “Please Stop Calling Me A Carless Zebra Instead Try Learning How Copper Saves Gold”.

Note: This is not the complete metal reactivity series. It is just a collection of the 10th level because we have about 85 metals in the periodic table. In the above reactivity series, hydrogen is not metal but it distinguishes reactive metals and almost inactive metals. Highly reactive non-metals also replaces low reactive non-metals.

Here, we can understand that Fe is more reactive than Zn, so Fe can replace Zn but Zn can not replace Fe because a more reactive metal can only replace low reactive metal.

For example,

(i) Displacement of copper (Cu) from copper sulphate solution (CuSO₄) by iron, zinc, and lead.

Metals such as iron, zinc, and lead are more reactive than copper so, they can replace copper from its salt solution. But, copper can not replace those metals from their salt solutions.

Note: In metals, more electropositive means more reactive. Electropositivity means a tendency to lose electrons from the valence shells.

(ii) Displacement of less reactive halogen from its salt solution with a more reactive halogen.

Chlorine displaces bromine and iodine from bromides and iodides.

Note: In non-metals, the more electronegative means more reactive. Electronegativity means the ease to add electrons in valence shells.

(iii) Displacement of hydrogen from acids by active metals. Active metals like Zn, Fe, Pb, etc., which is more reactive than hydrogen. The low reactive metals like Cu, Ag, and Au can not replace hydrogen from acids.

In the experiment, the evolution of hydrogen gas is detected by bringing a burning matchstick near the mouth of the test tube. If the gas burns with a pop sound, it will confirm the presence of hydrogen gas.

1.5.5 Double displacement reaction

The reactions in which there is an exchange of ions between the reactants are called double displacement reactions. For example,

(i) When silver nitrate reacts with sodium chloride, sodium nitrate, and silver chloride are produced.

In this reaction Ag⁺ of AgNO₃ is replaced by Na⁺ of NaCl and Na⁺ of NaCl is replaced by Ag⁺ of AgNO₃. Thus the exchange of ions takes place between the reactants such as

(ii) Sodium hydroxide reacts with hydrochloric acid to form sodium chloride and water.

Note: The reaction in which an acid reacts with a base to give salt and water is called a neutralisation reaction. In above reaction, sodium hydroxide (base) reacts with hydrochloric acid to give sodium chloride (salt) and water. So, it is a neutralisation reaction. This is named so because acid and base neutralise each other.

1.5.6 Precipitation reaction

The reaction in which an insoluble substance (precipitate) is produced is called a precipitation reaction. The precipitate settles down to the bottom of the container. A double displacement reaction can be a precipitation reaction if it produces an insoluble product. For example,

(i) When a solution of sodium sulphate is added with a solution of barium chloride, a curdy white precipitate of barium sulphate, and a solution of sodium chloride are produced.

(ii) On adding silver nitrate solution to sodium bromide, a yellow precipitate of silver bromide and solution of sodium nitrate are formed.

(iii) On adding lead nitrate solution to potassium iodide, a yellow precipitate of lead iodide, and potassium nitrate are formed.

1.5.7 Oxidation and Reduction reactions

(a) Oxidation reaction: The process which involves the addition of oxygen or removal of hydrogen is called an oxidation reaction. For example,

Addition of oxygen

(i) Burning of carbon: Carbon burns in the presence of oxygen to form carbon dioxide.

Here, oxygen is added to C to give CO₂ so it is an oxidation reaction.

(ii) Burning of magnesium ribbon: Magnesium ribbon burns in oxygen to give magnesium oxide.

Here, oxygen is added to Mg to give MgO so it is an oxidation reaction.

(iii) Oxidation of copper: When copper powder is exposed to air and heated, its surface becomes coated with black copper(II) oxide.

Removal of hydrogen

(i) Oxidation of H₂S: Hydrogen sulphide is oxidised by bromine to sulphur.

Here, H₂S removes hydrogen to become S. So, it is an oxidation reaction.

(ii) Oxidation of HCl: When hydrochloric acid (HCl) is heated with manganese dioxide (MnO₂). it gets oxidises to chlorine.

(b) Reduction reaction: The process which involves addition of hydrogen or removal of oxygen is called a reduction reaction. For example,

Addition of hydrogen

(i) Reduction of sodium: Sodium (Na) reacts to hydrogen (H₂) to give sodium hydride (NaH).

Here, hydrogen is added to Na to give NaH, so it is a reduction reaction.

(ii) Reduction of Cl₂: When hydrogen sulphide (H₂S) reacts with chlorine (Cl₂), chlorine gets reduced to hydrochloric acid (HCl).

Removal of oxygen

(i) Reduction of CuO: When hydrogen gas is passed over heated cupric oxide (CuO), CuO gets reduced to copper (Cu).

(ii) Reduction of KClO₃: When potassium chlorate (KClO₃) is heated, it gets reduced to potassium chloride (KCl).

Note: Oxidation and reduction are reciprocal processes. In a reaction, if a substance oxidises, another must be reduced.

Oxidising agent: The substance which can bring about oxidation of other substances is called an oxidising agent.

In other words, an oxidising agent is a substance that causes the addition of oxygen or the removal of hydrogen from other substances. For example, oxygen (O₂), ozone (O₃), hydrogen peroxide (H₂O₂), Chlorine (Cl₂), Nitric acid (HNO₃), conc. sulphuric acid (H₂SO₄), potassium permanganate (KMnO₄), potassium dichromate (K₂Cr₂O₇) etc.

Consider a reaction

In this reaction, O₂ is an oxidising agent as it oxidises of Cu to CuO.

Reducing agent: The substance which can bring about reduction of other substances is called an reducing agent.

In other words, a reducing agent is a substance that causes the addition of hydrogen or the removal of oxygen from other substances. For example, hydrogen (H₂), sulphur oxide (SO₂), hydrogen sulphide (H₂S), carbon (C), etc.

Consider a reaction

Here, hydrogen (H₂) is a reducing agent as it reduces Na to NaH.

Note: Oxidising agent is reduced itself to oxidise other substances while the reducing agent oxidised itself to reduce other.

Redox reaction: A reaction in which oxidation and reduction take place simultaneously is called an oxidation-reduction or redox reaction.

Consider the reaction

In this reaction, H₂S is oxidised to S, and Cl₂ is reduced to HCl simultaneously. So, it is a redox reaction. Here, H₂S reduces Cl₂ so it is a reducing agent while Cl₂ oxidises H₂S so it is an oxidising agent.

1.6 Effect of oxidation reactions in everyday life

Oxidation reactions are significant reactions that have a variety of consequences in our daily lives. Some of its manifestations, such as the combustion of fuels and the digestion of food in our bodies, are boons to humanity and extremely beneficial to the continuation of life. Some of its side effects, on the other hand, are extremely harmful, such as air pollution from burning fuels, food rancidification, metal corrosion, and so on.

1.6.1 Corrosion

When a metal is attacked by substances around it such as moisture, acids, etc., it is said to corrode and this process is called corrosion.

In other words, the phenomenon due to which the open surface of the metals is slowly eaten away (or corroded) by the reaction of moisture, acids, etc., present in the atmosphere, is called corrosion. For example,

(i) Rusting of iron: Iron articles are shiny when new, but get coated with a reddish brown powder (rust) when left form some time.

(ii) Silver loses its shine when its surface is black-coated.

(iii) Copper and brass get a green coloured deposit on their surfaces.

Effects of corrosion

Corrosion causes damage to car bodies, bridges, iron railings, ships and to all objects made of metals, specially those of iron.

Corrosion is a very serious problem in industries. The chemical reactors and plants need replacement due to corrosion. This requires a lot of money and time.

Prevention from corrosion

Corrosion can be prevented by

(i) Coating/depositing a thin layer of any other metal that does not corrode. For example, coating iron with zinc, brass with chromium, etc.

(ii) Coating the surface with paint, oil, grease or varnish, etc.

1.6.2 Rancidity

It is the process of slow oxidation of oil and fats present in the food materials resulting in a change of smell and taste in them.

Methods to prevent rancidity

(i) By adding antioxidant materials to the food material. These materials slow down the action of oxygen on fats and oils present in the food and thus save them from getting rancid.

(ii) By keeping the food materials in air-tight containers.

(iii) By packing the food materials under oxygen-free nitrogen atmosphere.

(iv) By refrigeration of cooked foods at low temperatures.

Questions

Q. 1. Why should a magnesium ribbon be cleaned before burning in the air?

Ans: Magnesium exposed to air for a long time, gets covered with a layer of magnesium oxide. This layer is unreactive and prevents the ribbon from burning. So, this layer of magnesium oxide is removed before burning the magnesium ribbon.

Q.2. Why is the amount of gas collected in one of the test tubes during the electrolysis of water double of the amount collected in the other? Name this gas.

Ans: Each water molecule contains 2 atoms of hydrogen and 1 atom of oxygen. So, water contains hydrogen and oxygen in the ratio 2:1 by volume. The gas having double volume is hydrogen gas.

Q.3. What is a balanced chemical equation? Why should chemical equations be balanced?

Ans: The chemical equation in which the number of atoms of each element on both sides is equal is called the balanced equation.

A chemical equation should be balanced because there is no loss or gain of any matter during a chemical reaction, i.e., the law of conservation of matter must hold good for the reaction.

Q.4. Why is respiration considered an exothermic reaction? Explain.

Ans: During respiration, the digested food gets oxidised and the energy is released. That is why it is considered as an exothermic reaction.

Q.5. Why do we apply paint on iron articles?

Ans: We apply paint on iron articles to protect them from rusting because paints do not allow oxygen and water (moisture) to come in contact with the surface of the iron.

Q.6. Oil and fat containing food items are flushed with nitrogen. Why?

Ans: Nitrogen is unreactive gas as compared to oxygen. It prevents the oxidation of the oil and fat-containing food so it is flushed before packaging the food.

Q.7. Why do we store silver chloride in dark coloured bottles?

Ans: Silver chloride is stored in dark coloured bottles because it is light sensitive. It gets decomposed to produce fine particles of silver metal.

Q.8. Why do fireflies glow at night?

Ans: Fireflies have a protein called luciferin. This protein undergoes enzymatic oxidation. This reaction involves emission of visible light and makes the fireflies to glow at night.

Q.9. Why do gold and silver not corrode in moist air?

Ans: Gold and silver are amongst the least reactive metals. They do not react with oxygen and water (present as moisture in the air). Therefore, silver and gold do not corrode in moist air.

Previous Years Questions

Short Questions

Q.1. State one example each characterised by the following along with the chemical equation:

(i) Change in state,

(ii) Evolution of gas,

(iii) Change in temperature. [2016]

Q.2. Write the balanced chemical equations for the following and identify the type of chemical reactions.

(i) Hydrogen iodide on reacting with chlorine gas gives iodine and hydrochloric acid.

(ii) Methane gas burns in oxygen of air to form carbon dioxide and water.

(iii) On passing electric current through molten aluminium oxide, it decomposes to form aluminium metal and oxygen gas. [2015]

Q.3. Write chemical equations for the reactions taking place when

(i) Magnesium reacts with dilute HNO₃.

(ii) Sodium reacts with water.

(iii) Zinc reacts with dilute hydrochloric acid. [2016]

Q.4. A metal salt MX when exposed to light, split up to form metal M and a gas X₂. Metal M is used in making ornaments whereas gas X₂ is used in making bleaching powder. The salt MX is itself used in black and white photography.

(i) Identify metal M and gas X₂.

(ii) Mention the type of chemical reaction involved when salt MX is exposed to light. [2018]

Q.5. Decomposition reactions require energy either in the form of heat or light or electricity for breaking down the reactants. Write one equation each for decomposition reactions where energy is supplied in the form of heat, light, and electricity. [2018]

Q.6. State the type of chemical reactions with chemical equations that take place in the following:

(i) Magnesium wire is burnt in air.

(ii) Electric current is passed through water.

(iii) Ammonia and hydrogen chloride gases are mixed. [2016]

Q.7. When potassium iodide solution is added to a solution of lead (II) nitrate in a test tube, a precipitate is formed.

(i) What is the colour of this precipitate?

(ii) Write the balanced chemical equation for this reaction.

(iii) List two types of reactions in which this reaction can be placed. [2019]

Q.8. (i) Classify the following reactions into different types:

(a) AgNO₃(aq) + NaCl(aq)  → AgCl(s) + NaNO₃(aq)

(b) CaO(s) + H₂O(l)  → Ca(OH)₂(aq)

(c) 2KClO₃(s)  → 2KCl(aq) + 3O₂(g)

(d) Zn + CuSO₄  → ZnSO₄ + Cu

(ii) Translate the following statement into a balanced chemical equation. “Barium chloride reacts with aluminium sulphate to give aluminium chloride and barium sulphate.” [2019]

Q.9. A metal nitrate ‘A’ on heating gives a metal oxide along with evolution of a brown coloured gas ‘B’ and a colourless gas, which helps in burning. Aqueous solution of ‘A’ when reacted with potassium iodide forms a yellow precipitate.
(a) Identify ‘A’ and ‘B’.
(b) Name the types of both the reactions involved in the above statement. [2023]

Long Questions

Q.1. (i) Identify the type of reactions taking place in each of the following cases and write the balanced chemical equations for the reactions

(a) Barium chloride solution is mixed, with copper sulphate solution and a white precipitate is obtained.

(b) On heating copper powder in air, the surface of the copper powder turns black.

(ii) What happens when hydrogen gas is passed over the heated copper oxide? Write the chemical equation involved in this reaction. [2016]

Q.2. (i) A student mixes sodium sulphate powder in barium chloride. What change would the student observe on mixing the two powders? Justify your answer and explain how he can obtain the desired change.

(ii) List two observations you would record in your 30 minutes after adding iron fillings to copper sulphate solution. [2019]