Types of chemical reactions (Important Points 2023-24)

 Chemical reactions are phenomenon which we usually get to see around us in daily life also. Here we will get to read about them in detail. All chemical reactions occurs at atomic level, the outcome of which can be visible or measurable.

Important notes for Class 10 C.B.S.E. Board for 2023-24 Exam

Syllabus: Chemical reactions, Chemical equation, and Balanced chemical equation, implications of a balanced chemical equation, types of chemical reactions: combination, decomposition, displacement, double displacement, precipitation, endothermic exothermic reactions, oxidation and reduction.

Exploring Important Chemical Reactions

Important Chemical Reactions

Chemical reactions are fundamental to understanding the transformations of matter. Among various types of reactions, some key ones include:

1. Combination reactions,
2. Decomposition reactions,
3. Displacement reactions,
4. Double-displacement reactions,
5. Oxidation & Reduction reaction.

1. Combination Reactions:

Combination reaction

Definition: Combination reactions involve the merging of two or more substances to form a single compound.

A combination reaction, in simple terms, is a type of chemical reaction where two or more substances come together to form a single, new substance. It's like when you mix ingredients to create a new dish, but in the world of chemistry, where atoms and molecules combine to make new compound.

Key Points:

- Often exothermic, releasing energy.

- The reaction typically involves elements or compounds reacting to form a more complex compound.

- Examples often include synthesis reactions.

Common Examples:

- Formation of water from hydrogen and oxygen:

  \[ 2H_2 + O_2 \rightarrow 2H_2O \]

- Formation of magnesium oxide from magnesium and oxygen:

  \[ 2Mg + O_2 \rightarrow 2MgO \]

Some more examples of combination reactions are as follows:

1. Reaction between Hydrogen (H₂) and Oxygen (O₂) to form Water (2H₂O):         2H₂ + O₂ → 2H₂O 
2. Reaction between Iron (Fe) and Oxygen (O₂) to form Ferric Oxide or rust (Fe₂O₃):                                                                                                                        4Fe + 3O₂ → 2Fe₂O₃
3. Reaction between Calcium(Ca) and Oxygen (O₂) to form Calcium Oxide (CaO):   2Ca + O₂ → 2CaO
4. Reaction between Magnesium (Mg) and oxygen (O₂) to form Magnesium Oxide (MgO):                                                                                                                       2Mg + O₂ → 2MgO
5. Reaction between Carbon (C) and Oxygen (O₂) to form Carbon Dioxide (CO₂):   C + O₂ → CO₂

In all the above examples two substances are combining to form a new compound.

2. Decomposition Reactions:

Decomposition Reactions

Definition: Decomposition reactions involve the breakdown of a single compound into two or more simpler substances.

Key Points:

- Often endothermic, absorbing energy.

- The reactant usually breaks down into simpler components like elements or smaller compounds.

- Examples can include thermal decomposition or electrolysis.

Common Examples:

- Decomposition of hydrogen peroxide into water and oxygen:

  \[ 2H_2O_2 \rightarrow 2H_2O + O_2 \]

- Thermal decomposition of calcium carbonate to form calcium oxide and carbon dioxide:

  \[ CaCO_3 \rightarrow CaO + CO_2 \]

Some more examples of decomposition reactions are as follows:

      i. Decomposition of Ferrous Sulphate (FeSO4) into Ferric oxide (Fe₂O₃), Sulphur Dioxide (SO₂) and Sulphur Trioxide (SO₃):

2FeSO₄(s) → Fe₂O₃(s) + 2SO₂(g) + 2SO₃(g)

     ii. Decomposition of Calcium Carbonate (CaCO₃) into Calcium oxide (CaO) and Carbon Dioxide (CO₂):

CaCO₃(s) → CaO(s) + CO₂(g)

The above reactions are carried out by heating the compounds, which means decompositions have been carried out by the action of heat.

     iii. Decomposition of Silver Chloride (AgCl) into Silver (Ag) and Chloride (Cl₂):

2AgCl(s) → 2Ag(s) + Cl₂(g)

    iv. Decomposition of Silver Bromide (AgBr) into Silver (Ag) and Bromide (Br₂):

2AgBr(s) → 2Ag(s) + Br₂(g)

The above reactions are carried out by passing electric current through the molten form of these compounds, which means decomposition have been carried out by the action of electricity.

3. Displacement Reactions:

Displacement Reactions

Definition: Displacement reactions involve the replacement of one element in a compound by another element.

The important thing which has to be kept in mind here is that the most reactive metal “which has been referred as bully above” replaces the least reactive metal from its compound.

It is known which metal is most reactive and which element is least reactive and is given in the list provided in the next page.

Image 1: Metal Reactivity Series

Reactivity series of metals

Key Points:

- Can be single displacement or double displacement reactions.

- Reactivity of elements determines the displacement.

- Occurs when a more reactive element displaces a less reactive one.

Common Examples:

- Reaction of zinc with hydrochloric acid:

  \[ Zn + 2HCl \rightarrow ZnCl_2 + H_2 \]

- Displacement of copper from copper sulfate solution by iron:

  \[ CuSO_4 + Fe \rightarrow FeSO_4 + Cu \]

Some examples of displacement reactions are as follows:

      i. Reaction between Iron (Fe) and Copper Sulphate (CuSO₄) solution to form Ferrous Sulphate (FeSO₄) and Copper (Cu):

Fe(s) + CuSO₄(aq) → FeSO4(aq) + Cu(s)

     ii. Reaction between Zinc (Zn) and Copper Sulphate (CuSO₄) solution to form Zinc Sulphate (ZnSO₄) and Copper (Cu):

Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)

    iii. Reaction between Lead (Pb) and Copper Chloride (CuCl₂) solution to form Copper (Cu) and Lead Chloride (PbCl₂):

Pb(s) + CuCl₂(aq) → Cu(s) + PbCl₂(s)

In all the above reactions highest reactive metal displaces least reactive metal from another compound.

4. Double Displacement Reactions:

Double-displacement reactions

Definition: Double displacement reactions involve the exchange of ions between two compounds.

Key Points:

- Results in the formation of two different compounds.

- Often forms a precipitate, gas, or water.

- Usually occurs in aqueous solutions.

Common Examples:

- Reaction between silver nitrate and sodium chloride:

  \[ AgNO_₃ + NaCl \rightarrow AgCl + NaNO_₃ \]

- Formation of calcium carbonate precipitate in a reaction between calcium chloride and sodium carbonate:

  \[ CaCl_₂ + Na_₂CO_₃ \rightarrow CaCO_₃ + 2NaCl \]

Some examples of double displacement reactions are in following page:

      i. Reaction between Sodium Sulphate (Na₂SO₄) and Barium Chloride (BaCl₂) to form Sodium Chloride (NaCl) and Barium Sulphate (BaSO₄):

Na2SO₄(aq) + BaCl₂(aq) → 2NaCl(aq) + BaSO₄(s)

     ii. Reaction between Lead(II) Nitrate (Pb(NO₃)₂) and Potassium Iodide (KI) to form Potassium Nitrate (KNO₃)  and Lead Iodide (PbI₂) :

Pb(NO₃)₂(aq) + 2KI(aq) → 2KNO₃(aq) + PbI₂(s)

    iii. Reaction between Barium Chloride (BaCl₂) and Sodium Sulphate (Na₂SO₄) to form Barium Sulphate (BaSO₄) and Sodium Chloride (NaCl):

BaCl₂(aq) + Na₂SO₄(aq) → BaSO₄(s) + 2NaCl(aq)

    iv. Reaction between Sodium Hydroxide (NaOH) and Hydrochloric Acid (HCl) to form Sodium Chloride (NaCl) and Water (H₂O):

NaOH(aq) + HCl(aq) → NaCl(aq) + H₂O(l)

In all the above examples, it shows exchange of ions that took place between two compounds leading to the formation of two different compounds.

5. Oxidation and Reduction Reactions:

Oxidation and reduction reactions

Definition: Oxidation involves the loss of electrons, while reduction involves the gain of electrons.

Key Points:

- Oxidation and reduction always occur together (redox reactions).

- Oxidizing agents accept electrons, reducing agents donate electrons.

- Involve changes in oxidation states of elements.

Common Examples:

- Rusting of iron (oxidation of iron by oxygen):

  \[ 4Fe + 3O_₂ \rightarrow 2Fe_₂O_₃ \]

- Reaction of copper with silver nitrate (displacement):

  \[ 2AgNO_₃ + Cu \rightarrow Cu(NO_₃)_₂ + 2Ag \]

Some examples of oxidation and reduction reactions are as follows:

      i. Reaction between Copper (Cu) and Oxygen (O):

2Cu(s) + O₂(g) → 2CuO(s)

It is an example of oxidation reaction as here copper is oxidized and form copper oxide as here copper loses electron and undergoes oxidation; it can also be seen here that oxygen is added to copper. Here oxygen act as reducing agent.

    ii. Reaction between Zinc Oxide (ZnO) and Carbon (C):

ZnO(s) + C(s) → Zn(s) + CO(g)

It is an example of redox reaction in which where zinc oxide is reduced to form zinc, and carbon is oxidized to form carbon monoxide.

Few more examples of oxidation-reduction are as follows:

• Reaction between Iron and Water in presence of Oxygen also known as Rusting:

4Fe + 3O₂ + 6H₂O → 4Fe(OH)₃

In this reaction, iron (Fe) gets oxidized to form iron(III) hydroxide [Fe(OH)₃], and oxygen (O₂) gets reduced. Rusting is a common example of a redox process in everyday life.

In this reaction manganese dioxide is losing oxygen to form manganese dichloride, so manganese dioxide is being reduced to manganese dichloride. On the other hand, hydrochloric acid is losing hydrogen to form chloride, so hydrochloric is being oxidized to chlorine.

Here manganese dioxide is the oxidizing agent whereas hydrochloric acid is the reducing agent.

Oxidation reactions can be easily observed in everyday life. The two most common effects of oxidation reactions in everyday life are:

- Corrosion of metals-

• Definition: Corrosion is the gradual degradation of metals through chemical reactions with substances in their surroundings, resulting in the formation of metal compounds that weaken the material.
• Example of Corrosion: The corrosion of iron (Fe) in the presence of oxygen and moisture, forming iron oxide (Fe2O3), is a well-known example. This process is commonly observed in the rusting of iron.
• Difference from Rusting: Rusting is a specific type of corrosion that occurs in iron or steel when exposed to oxygen and moisture. It involves the formation of iron oxide (rust). The key difference between corrosion and rusting is that corrosion is a broader term that encompasses the degradation of various metals, whereas rusting specifically refers to the corrosion of iron or steel.

- Rancidity of food-

• Definition: Rancidity of food is the spoilage process in which fats and oils in food products break down due to exposure to oxygen, light, or heat, leading to the development of unpleasant odors and flavors.
• Example: One common example of rancidity is the development of a stale or off taste in nuts, such as almonds or peanuts, when they are exposed to air and become rancid. Another example is when cooking oil turns rancid after being used multiple times for frying.

Causes:

• Rancidity is more likely to occur in foods that contain high amounts of fats and oils, such as nuts, cooking oils, and processed snacks.
• Factors that accelerate rancidity include exposure to air (oxygen), exposure to light, and high temperatures.
• Antioxidants, like vitamin E, are often added to food products to slow down the oxidation process and prevent rancidity.
• Rancidity not only affects the taste and smell of food but can also lead to the loss of essential nutrients, making the food less nutritious.

Prevention:

• To prevent rancidity, it's important to store food products containing fats and oils in airtight containers in a cool, dark place.
• Additionally, using fresh cooking oil and properly sealing food packaging can help extend the shelf life of these products.

Understanding these important types of chemical reactions provides a foundation for comprehending complex chemical processes. They play a crucial role in various fields, from industry to everyday life, governing the transformations of matter around us.

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