Chemical Reactions and Equations | Class 10 Chemical Reactions and Equations | Class 10 Chapter 1 Science | Chemical Reactions and Equations Short Notes | Chemical equations | Balancing Chemical Equations | Balancing Equations

Chemical Reactions and Equations | Class 10 Chemical Reactions and Equations | Class 10 Chapter 1 Science | Chemical Reactions and Equations Short Notes | Chemical equations | Balancing Chemical Equations | Balancing Equations

What is a Chemical Reaction?

A chemical reaction refers to the process in which one or more substances are transformed into different substances with distinct properties. This transformation occurs due to a chemical change where atoms are rearranged. In a chemical reaction, new substances are formed, and they are completely different from the original substances in terms of their properties.

Chemical Reactions and Equations | Class 10 Chemical Reactions and Equations | Class 10 Chapter 1 Science | Chemical Reactions and Equations Short Notes | Chemical equations | Balancing Chemical Equations | Balancing Equations

Examples of Chemical Reactions:

• Rusting of iron: Iron reacts with oxygen from the air, resulting in the formation of rust (iron oxide).
• Milk turning into curd: The process of fermentation where bacteria act on the milk to form curd.
• Digestion of food: Food undergoes chemical changes in the digestive system to break down complex molecules into simpler ones.
• Respiration: The process by which organisms break down glucose with oxygen to release energy.

Key Terms in Chemical Reactions and Equations :

1. Reactants: The substances that participate in a chemical reaction and undergo change are known as reactants.

   • Example: In the reaction of magnesium with oxygen, the reactants are magnesium (Mg) and oxygen (O₂).

2. Products: The substances produced after a chemical reaction are called products.

   • Example: In the reaction where magnesium burns in oxygen, the product is magnesium oxide (MgO).

3. Chemical Change: A chemical reaction results in a change in the chemical composition of the substances, producing new substances. This change is permanent and cannot be easily reversed.

4. Rearrangement of Atoms: During a chemical reaction, the atoms of the reactants rearrange to form the products. The total number of atoms remains the same, but they are rearranged to create different substances.

Example of a Chemical Reaction:

The reaction of magnesium with oxygen in the air: $$2Mg(s)+O2(g)→2MgO(s)

• Before the reaction: Magnesium ribbon (Mg) is cleaned by rubbing with sandpaper to remove any protective layer of magnesium carbonate. This ensures the magnesium reacts with oxygen effectively.
• Reactants: Magnesium (Mg) and Oxygen (O₂)
• Product: Magnesium oxide (MgO)

Balanced and Unbalanced Chemical Equations:

• Balanced Equation: A chemical equation is said to be balanced when the number of atoms of each element is the same on both sides of the equation.

• Unbalanced Equation: If the number of atoms of any element is not equal on both sides of the equation, it is an unbalanced equation and needs to be balanced.

Balancing a Chemical Equation:

To balance a chemical equation, adjust the coefficients (the numbers in front of the chemical formulas) to ensure that the number of atoms of each element is the same on both sides of the equation.

Example of Balancing:

For the reaction of hydrogen and oxygen to form water: Unbalanced equation:
${\text{H}}_2(g) + {\text{O}}_2(g) \to \; {\text{H}}_2\text{O}(l)$

Balanced equation:
$2{\text{H}}_2(g) + {\text{O}}_2(g) \to 2{\text{H}}_2\text{O}(l)$

In the balanced equation, the number of atoms of each element is the same on both sides.

Conclusion:

Chemical reactions are essential in both nature and industry. They involve the transformation of substances through the rearrangement of atoms to form new products. Understanding these reactions and knowing how to balance chemical equations is crucial for studying and applying chemistry in real-life scenarios.

Product in Chemical Reactions:

A product is a substance that is formed as a result of a chemical reaction.

Example: Magnesium oxide (MgO) is a product formed when magnesium reacts with oxygen.

Characteristics of Chemical Reactions:

1. Gas Evolution: Some reactions result in the production of gas. For instance, when zinc reacts with dilute sulfuric acid, hydrogen gas is produced.

   • Reaction Example: $\text{Zn}(s)+ {\text{H}}_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}{\text{SO}}_{\mathrm{<span\; style="font-size:\; xx-small;">4</span>}}(aq)\to {\text{ZnSO}}_{\mathrm{<span\; style="font-size:\; xx-small;">4</span>}}(aq)+ {\text{H}}_2(g)\text{↑}$

2. Colour Change: Certain reactions cause a visible color change. For example, when citric acid reacts with potassium permanganate, the color of the solution changes from purple to colorless. Another example is when sulfur dioxide reacts with acidified potassium dichromate, causing a color change from orange to green.

3. Change in State: Chemical reactions can involve a change in the state of substances. For example, in the combustion of candle wax, the wax changes from solid to liquid and gas, while the water produced is liquid and the carbon dioxide is a gas.

4. Temperature Change: Some chemical reactions are characterized by a change in temperature. For example, the reaction between quicklime and water releases heat, while the reaction of zinc with sulfuric acid also results in an increase in temperature.

5. Precipitate Formation: Some reactions result in the formation of a solid product, known as a precipitate. For instance, when sulfuric acid reacts with barium chloride solution, a white precipitate of barium sulfate is formed.

   • Reaction Example: ${\text{BaCl}}_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}(aq) + {\text{H}}_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}{\text{SO}}_{\mathrm{<span\; style="font-size:\; xx-small;">4</span>}}(aq)\to {\text{BaSO}}_{\mathrm{<span\; style="font-size:\; xx-small;">4</span>}}(s)(\text{ppt})+ 2\text{HCl}(aq)$

What is a Chemical Equation?

A chemical equation is a symbolic representation of a chemical reaction where the reactants and products are shown using their chemical symbols and formulas.

Example:
In a simple chemical reaction between hydrogen and oxygen to form water, the word equation can be represented as:
Hydrogen + Oxygen → Water
This can be simplified using symbols as:
${\text{H}}_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}+{\text{O}}_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}\to {\text{H}}_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}\text{O}$

Here, H₂ and O₂ are reactants, while H₂O is the product. The arrow indicates the direction of the reaction.

Types of Chemical Equations:

Chemical equations can be classified into two main categories: Balanced and Unbalanced.

1. Balanced Chemical Equation: A chemical equation is balanced when the number of atoms of each element is the same on both sides of the equation. This ensures that mass is conserved during the reaction.

   • Example: $\text{Zn}+{\text{H}}_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}{\text{SO}}_{\mathrm{<span\; style="font-size:\; xx-small;">4</span>}}\to {\text{ZnSO}}_{\mathrm{<span\; style="font-size:\; xx-small;">4</span>}}+{\text{H}}_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}$
   • This equation is balanced because the number of zinc (Zn), hydrogen (H), and sulfate (SO₄) ions are equal on both sides.

   According to the Law of Conservation of Mass, the total mass of the reactants must be equal to the total mass of the products. This principle is upheld in a balanced equation.

2. Unbalanced Chemical Equation: In an unbalanced equation, the number of atoms of each element is not the same on both sides, meaning it needs to be balanced by adjusting the coefficients.

Unbalanced Chemical Equation:

An unbalanced chemical equation occurs when the number of atoms of each element on the reactant side does not match the number of atoms of those elements on the product side. In such an equation, the law of conservation of mass is not satisfied, as the mass is not conserved between the reactants and the products.

Example:
Consider the reaction: ${\mathrm{ }}_{}$Fe+H2O→Fe3O4+H2↑

In this equation:

• On the left-hand side, there is only one iron (Fe) atom.
• On the right-hand side, there are three iron (Fe) atoms in Fe₃O₄.

Since the number of iron atoms is not the same on both sides, this equation is unbalanced.

Balancing a Chemical Equation:

To balance a chemical equation, follow these steps:

1. Write the unbalanced equation: Start with the chemical equation as it is given. Fe+H2O→Fe3O4+H2↑

2. Count the number of atoms of each element in both the reactants and products.

   • Reactants:
     • Fe = 1
     • H = 2 (from H₂O)
     • O = 1 (from H₂O)
   • Products:
     • Fe = 3 (from Fe₃O₄)
     • H = 2 (from H₂)
     • O = 4 (from Fe₃O₄)

3. Balance the elements:

   • Begin by balancing iron (Fe). Since there are 3 Fe atoms in Fe₃O₄, place a coefficient of 3 in front of Fe on the left: $$(3Fe) + 4H2O → Fe3O4 + 4H2

   • Next, balance oxygen (O). There are 4 oxygen atoms in Fe₃O₄, so place a coefficient of 4 in front of H₂O: ${\mathrm{ }}_{}$3Fe+4H2O=Fe3O4+H2

   • Now, balance hydrogen (H). On the left, there are 8 hydrogen atoms (from 4 H₂O), so place a coefficient of 4 in front of H₂ on the right: 3Fe + 4H2O → Fe3O4 + 4H2

4. Double-check the equation to ensure all elements are balanced:

• Iron (Fe): 3 Fe on both sides
• Hydrogen (H): 8 H on both sides
• Oxygen (O): 4 O on both sides

The balanced equation is now: $3\text{Fe}+4{\text{H}}_2\text{O}\to {\text{Fe}}_3{\text{O}}_4+4{\text{H}}_2$

This equation now satisfies the law of conservation of mass, where the number of atoms of each element is the same on both sides of the equation.

Making Chemical Equations More Informative:

To make chemical equations more informative, it is useful to indicate the physical states of the substances involved in the reaction and the conditions under which the reaction takes place.

• Physical States:

  • Solid: Represented by (s).
  • Liquid: Represented by (l).
  • Gas: Represented by (g).
  • Aqueous Solution: Represented by (aq), indicating that the substance is dissolved in water.

• Conditions for the Reaction: Conditions like temperature, pressure, or the presence of a catalyst are typically written above or below the arrow in a chemical equation.

Including the physical states and conditions in a chemical equation provides a clearer understanding of how the reaction occurs.

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Types of Chemical Reactions:

Chemical reactions are categorized into several types based on how reactants transform into products. Here are the main types:

1. Combination Reaction:

   • A combination reaction occurs when two or more reactants combine to form a single product.

   • The general form of a combination reaction is: $A+B\to AB$

   • Examples:

     • When magnesium reacts with oxygen, it forms magnesium oxide: $\text{Mg}(s)+{\text{O}}_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}(g)\to 2\text{MgO}(s)$
     • When carbon burns in oxygen, carbon dioxide is produced: $\text{C}(s)+{\text{O}}_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}(g)\to {\text{CO}}_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}(g)$

2. Decomposition Reaction:

   • A decomposition reaction occurs when a compound breaks down into two or more simpler compounds or elements.

   • The general form of a decomposition reaction is: $AB\to A+B$

   • Examples:

     • When calcium carbonate is heated, it breaks down into calcium oxide and carbon dioxide: ${\text{CaCO}}_{\mathrm{<span\; style="font-size:\; xx-small;">3</span>}}(s) \; (heat)\stackrel{}{\to }\text{CaO}(s)+{\text{CO}}_2(g)$
     • When ferric hydroxide is heated, it decomposes into ferric oxide and water: $2{\text{Fe(OH)}}_{\mathrm{<span\; style="font-size:\; xx-small;">3</span>}}(s)\stackrel{\mathrm{\Delta }}{\to }{\text{Fe}}_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}{\text{O}}_{\mathrm{<span\; style="font-size:\; xx-small;">3</span>}}(s)+3{\text{H}}_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}\text{O}(l)$

• a) Thermal Decomposition:

This occurs when a substance breaks down due to heating.
Example:
• When lead nitrate is heated, it decomposes into lead oxide, nitrogen dioxide, and oxygen: $2{\text{Pb(NO}}_{\mathrm{<span\; style="font-size:\; xx-small;">3</span>}}{)}_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}(s)\stackrel{\text{heat}}{\to }2\text{PbO}(s)+4{\text{NO}}_2(g)+{\text{O}}_2(g)$

b) Electrolytic Decomposition (Electrolysis):

• In electrolytic decomposition, a compound decomposes into simpler substances when an electric current passes through it.
• Example:
• When water is electrolyzed, it decomposes into hydrogen and oxygen: $2{\text{H}}_2\text{O}(l)\stackrel{}{\to }$ 2H2(g)+O2(g)

c) Photolysis (Photo Decomposition):

Photolysis occurs when a substance breaks down into simpler substances due to the action of light (usually sunlight).
• Example:
  • Silver chloride decomposes into silver and chlorine when exposed to sunlight: $2\text{AgCl}(s)\; (\stackrel{}{sunlight)\; \to }2\text{Ag}(s)+{\text{Cl}}_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}(g)$

Photographic paper, which contains silver chloride, turns grey when exposed to sunlight due to the decomposition of silver chloride into silver (which is grey) and chlorine.]

3. Displacement Reaction:

A displacement reaction occurs when a more reactive element displaces a less reactive element from a compound. These reactions are also referred to as substitution reactions or single displacement/replacement reactions.

The general form of a displacement reaction is: $A + BC$

In such reactions, A must be more reactive than B for the displacement to occur. If B is more reactive than A, no reaction will take place, and A will not displace C from BC.

Examples:

• When zinc reacts with hydrochloric acid, hydrogen gas and zinc chloride are formed: $\text{Zn}(s)+2\text{HCl}(aq)\to {\text{ZnCl}}_2(aq)+{\text{H}}_2(g)$

• When zinc reacts with copper sulfate, zinc displaces copper from copper sulfate, forming zinc sulfate and copper metal: $\text{Zn}(s)+{\text{CuSO}}_4(aq)\to {\text{ZnSO}}_4(aq)+\text{Cu}(s)$

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4.  Double Displacement Reaction:

A double displacement reaction (also called a metathesis reaction) occurs when two ionic compounds exchange ions to form two new compounds. This reaction often results in the formation of a precipitate, gas, or water.

The general form of a double displacement reaction is: $AB+CD\to AC+BD$

Examples:

• When barium chloride reacts with sodium sulfate, a white precipitate of barium sulfate is formed, along with sodium chloride: $$BaCl2(aq)+Na2SO4(aq)→BaSO4(s)(precipitate)+2NaCl(aq)

• When sodium hydroxide (a base) reacts with hydrochloric acid, sodium chloride and water are produced: $$NaOH(aq)+HCl(aq)→NaCl(aq)+H2O(l)

5. Precipitation Reaction:

A precipitation reaction is a type of double displacement reaction where two aqueous solutions of salts combine to form a solid precipitate. This occurs when one of the products is insoluble in water.

Example: When solutions of barium chloride and sodium sulfate are mixed, barium sulfate precipitates out of the solution: ${\text{BaCl}}_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}(aq)+{\text{Na}}_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}{\text{SO}}_{\mathrm{<span\; style="font-size:\; xx-small;">4</span>}}(aq)\to {\text{BaSO}}_{\mathrm{<span\; style="font-size:\; xx-small;">4</span>}}(s)(\text{precipitate})+2\text{NaCl}(aq)$

This type of reaction is commonly observed in analytical chemistry to identify ions in solutions.

6. Neutralization Reaction:

A neutralization reaction occurs when an acid reacts with a base, resulting in the formation of salt and water. This happens due to the exchange of ions between the acid and the base.

Example:
When hydrochloric acid reacts with sodium hydroxide, sodium chloride (a salt) and water are formed:

$$\text{HCl}(aq)+\text{NaOH}(aq)\to \text{NaCl}(aq)+{\text{H}}_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}\text{O}(l)$$

In this reaction, the hydrogen ions (H⁺) from the acid and the hydroxide ions (OH⁻) from the base combine to form water, while the remaining ions form a salt.

7. Oxidation and Reduction Reactions:

Oxidation:

Oxidation refers to the process in which oxygen or a non-metallic element is added to a compound, or hydrogen or a metallic element is removed from it. When a substance undergoes oxidation, it is said to have been oxidized.

Reduction:

Reduction is the process where hydrogen or a metallic element is added to a compound, or oxygen or a non-metallic element is removed from it. The substance that undergoes reduction is said to have been reduced.

Oxidation and reduction always occur together, meaning that whenever one substance is oxidized, another is reduced. These reactions are often referred to as redox reactions.

Oxidizing Agent:

The substance that provides oxygen or removes hydrogen is called an oxidizing agent. It causes oxidation in another substance.

Reducing Agent:

The substance that provides hydrogen or removes oxygen is called a reducing agent. It causes reduction in another substance.

Example of Redox Reaction: When copper oxide (CuO) is heated with hydrogen (H₂), copper (Cu) and water (H₂O) are formed:

$$\text{CuO}+{\text{H}}_2\to \text{Cu}+{\text{H}}_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}\text{O}$$

• In this reaction:
  • CuO is reduced to Cu because oxygen is removed from copper oxide.
  • H₂ is oxidized to H₂O because hydrogen is combined with oxygen.

Here, the substance CuO is the oxidizing agent (because it takes oxygen from hydrogen), and H₂ is the reducing agent (because it donates hydrogen).

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8. Exothermic and Endothermic Reactions:

Exothermic Reaction:
An exothermic reaction is a type of chemical reaction that releases energy, often in the form of heat or light. These reactions generally increase the temperature of the surroundings.

Example:
Respiration is a biological decomposition reaction in which energy is released:

C6H12O6 + 6O2 → 6CO2 + 6H2O + Energy

In respiration, glucose reacts with oxygen, producing carbon dioxide, water, and energy, which is released as heat. This is an example of an exothermic reaction.

When calcium oxide (CaO) is mixed with water, it produces heat as a result of the reaction.

     An example of this reaction is when quick lime (CaO) is added to water, it forms slaked lime (Ca(OH)₂) and releases heat. The chemical equation for this process is:

$$\text{CaO}(s)+{\text{H}}_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}\text{O}(l)\to {\text{Ca(OH)}}_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}(aq)+\text{Heat}$$

Endothermic Reaction:
An endothermic reaction is a type of chemical reaction in which heat energy is absorbed from the surroundings.

Example:
One example is the decomposition of calcium carbonate. When calcium carbonate (CaCO₃) is heated, it breaks down into calcium oxide (CaO) and carbon dioxide (CO₂), absorbing heat in the process.

The chemical equation for this reaction is:

$${\text{CaCO}}_{\mathrm{<span\; style="font-size:\; xx-small;">3</span>}}(s)\text{heat}\to \text{CaO}(s)+{\text{CO}}_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}(g)$$

                                                

This reaction requires heat energy to proceed, making it endothermic.

Corrosion:

Corrosion is the gradual transformation of metals into unwanted compounds as a result of their interaction with elements like oxygen, water, acids, or gases present in the environment.

Example:
One common example of corrosion is the rusting of iron.

Rusting:

When iron reacts with oxygen in the presence of moisture, it forms a reddish-brown substance known as rust. This process is a form of corrosion where iron combines with oxygen and water over time.

                                    4Fe(s)+3O2​(g)+6H2​O(l)→2Fe2​O3​⋅3H2​O(s)

The rusting of iron is an example of a redox reaction. It involves the oxidation of iron and the reduction of oxygen. Over time, this process weakens iron and steel materials, leading to the deterioration of structures like railings, car bodies, bridges, and ships, significantly reducing their lifespan.

Methods to Prevent Rusting:

1. Painting: Applying a protective layer of paint to prevent exposure to air and moisture.
2. Greasing and Oiling: Coating the metal surfaces with grease or oil to shield them from water and oxygen.
3. Galvanization: Coating the metal with a layer of zinc to protect it from rusting.

Corrosion of Copper:

Copper objects gradually lose their shine as they interact with air. Over time, the surface of copper develops a greenish coating of basic copper carbonate (CuCO₃·Cu(OH)₂), a result of exposure to oxygen and moisture.

                                                                   2Cu(s)+O2​(g)+2H2​O(l)→CuCO3​⋅Cu(OH)2​(s)

In this reaction:

• Copper (Cu) reacts with oxygen (O₂) from the air and water (H₂O).
• The result is the formation of basic copper carbonate (CuCO₃·Cu(OH)₂), which gives copper its characteristic green patina.

This process occurs slowly over time as copper is exposed to moisture and air.

When silver is exposed to air, it gradually loses its shine and becomes dull. This happens because of the formation of a black coating of silver sulfide (Ag₂S) on its surface. This tarnishing occurs due to the reaction between silver and hydrogen sulfide (H₂S) gas present in the air. Over time, this chemical interaction leads to the dull appearance of silver objects.

                                                              The tarnishing of silver due to the reaction with hydrogen sulfide (H₂S) is represented by the following chemical equation:

$$2Ag(s)+H_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}S(g)\to A{g}_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}S(s)+H_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}(g)$$

In this reaction:

• Silver (Ag) reacts with hydrogen sulfide (H₂S) gas from the air.
• The result is the formation of silver sulfide (Ag₂S), which appears as a black tarnish on the silver surface.
• Hydrogen gas (H₂) is also produced as a by-product.

Rancidity:
Rancidity refers to the unpleasant change in taste and smell of food that contains fats and oils when they are exposed to air for an extended period. This is caused by the oxidation of the fats and oils in these foods.

Methods to Prevent Rancidity:

1. Addition of Anti-Oxidants: These compounds slow down the oxidation process.
2. Vacuum Packing: This method removes air from the packaging, preventing the exposure of food to oxygen.
3. Replacing Air with Nitrogen: Nitrogen is used in packaging to replace oxygen and inhibit oxidation.
4. Refrigeration: Cooling food can slow down the process of rancidity.

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Chemical Reactions:
Chemical reactions involve changes in the composition of substances, leading to the formation of new substances.

Chemical Equations:
Chemical reactions are often written in equation form, and these equations should always be balanced to follow the law of conservation of mass.

Types of Chemical Reactions:

1. Combination Reaction: Two or more reactants combine to form a single product.

   • Example:
     $2Mg+O_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}\to 2MgO$

2. Decomposition Reaction: A single reactant breaks down into two or more products.

   • Thermal Decomposition:
     $2Pb(N{O}_{\mathrm{<span\; style="font-size:\; xx-small;">3</span>}}{)}_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}\to 2PbO+4N{O}_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}+O_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}$
   • Electrolysis:
     $2H_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}O\to 2H_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}+O_2$
   • Photochemical Reaction:
     $2AgBr\to 2Ag+B{r}_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}$

3. Displacement Reaction: One element replaces another element in a compound.

   • Example:
     $Zn+CuS{O}_{\mathrm{<span\; style="font-size:\; xx-small;">4</span>}}\to ZnS{O}_{\mathrm{<span\; style="font-size:\; xx-small;">4</span>}}+Cu$

4. Double Displacement Reaction: Ions are exchanged between two compounds.

   • Example:
     $AgN{O}_{\mathrm{<span\; style="font-size:\; xx-small;">3</span>}}+NaCl\to AgCl+NaN{O}_{\mathrm{<span\; style="font-size:\; xx-small;">3</span>}}$

5. Redox Reaction: A reaction where both oxidation and reduction occur simultaneously.

   • Example:
     $CuO+H_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}\to Cu+H_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}O$

6. Exothermic Reaction: A reaction in which energy is released, usually as heat.

   • Example:
     $C+O_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}\to C{O}_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}+\text{heat}$

7. Endothermic Reaction: A reaction that absorbs heat energy.

   • Example:
     $ZnC{O}_{\mathrm{<span\; style="font-size:\; xx-small;">3</span>}}+\text{heat}\to ZnO+C{O}_2$

Oxidation and Reduction Reactions:

1. Oxidation: The process in which a substance gains oxygen or loses hydrogen.

   • Example:
     $ZnO+C\to Zn+CO$
   • In this example, ZnO is reduced to Zn (reduction), and C is oxidized to CO (oxidation).

2. Reduction: The process in which a substance loses oxygen or gains hydrogen.

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Effects of Oxidation Reactions in Daily Life:

1. Corrosion: The gradual deterioration of metals due to reactions with moisture, air, or acids.

   • Example: Rusting of Iron:
     $F{e}_{\mathrm{<span\; style="font-size:\; xx-small;">2</span>}}{O}_{\mathrm{<span\; style="font-size:\; xx-small;">3</span>}}\cdot n{H}_{2}O$ (Hydrated iron oxide)

2. Rancidity: The undesirable changes in the taste and odor of food items containing oils due to the oxidation of fatty acids.

• Preventive Methods for Rancidity:
• Adding antioxidants to food.
• Storing food in airtight containers.
• Replacing air with nitrogen in packaging.
• Refrigerating food to slow down the oxidation process.

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