Chemical Reactions and Equations Class 10 Science Notes

Chemical Equations

A chemical reaction is a process that involves the rearrangement of the structure of molecules or the ions of a chemical compound. In simpler terms, it's when substances change into new substances. We can observe that a chemical reaction has occurred if we notice a change in state, a change in colour, the evolution of a gas, or a change in temperature.

For example, when magnesium ribbon burns in air, it reacts with oxygen to form magnesium oxide. This is a chemical reaction.

Writing a Chemical Equation

Describing a chemical reaction using words can be lengthy. To make it more concise, we use a word-equation.

Word-Equation: Reactants $\to$ Products

Reactants are the substances that undergo a chemical change.
Products are the new substances formed during the reaction.

The reactants are written on the left-hand side (LHS) of an arrow, and the products are written on the right-hand side (RHS). A plus sign (+) is used to separate multiple reactants or products. The arrowhead points from reactants to products, indicating the direction of the reaction.

Example: Magnesium + Oxygen $\to$ Magnesium oxide

For a more compact representation, we can use chemical formulae instead of words. This is called a chemical equation.

Example:

M g + O_{2} \to M g O

This is a skeletal chemical equation. It represents the reaction but may not show the exact number of atoms involved.

Balanced Chemical Equations

The Law of Conservation of Mass states that mass can neither be created nor destroyed in a chemical reaction. This means the total mass of the elements present in the reactants must be equal to the total mass of the elements present in the products. Consequently, the number of atoms of each element must remain the same before and after a chemical reaction.

A chemical equation that satisfies this condition (i.e., the number of atoms of each element is the same on both sides of the equation) is called a balanced chemical equation. Skeletal chemical equations are often unbalanced and need to be balanced.

How to Balance a Chemical Equation (Hit-and-Trial Method):

Draw Boxes: Draw a box around each chemical formula in the equation. Do not change anything inside the boxes.

Example: $F e + H_{2} O \to F e_{3} O_{4} + H_{2}$ becomes $[F e] + [H_{2} O] \to [F e_{3} O_{4}] + [H_{2}]$

List Atoms: Count the number of atoms of each element on both the reactant and product sides.
Balance Oxygen: Start by balancing the element that appears in the maximum number of atoms in one of the compounds (often oxygen or a polyatomic ion). Use coefficients (numbers placed before the chemical formula) to equalize the number of atoms. You cannot change the formula itself.

In $F e + H_{2} O \to F e_{3} O_{4} + H_{2}$ , Oxygen has 1 atom on LHS and 4 on RHS. To balance, place '4' before $H_{2} O$ : $F e + 4 H_{2} O \to F e_{3} O_{4} + H_{2}$ . Now, Oxygen has 4 atoms on LHS and 4 on RHS.

Balance Other Elements: Next, balance other elements like hydrogen.

In $F e + 4 H_{2} O \to F e_{3} O_{4} + H_{2}$ , Hydrogen has 8 atoms on LHS ( $4 \times 2$ ) and 2 on RHS. To balance, place '4' before $H_{2}$ : $F e + 4 H_{2} O \to F e_{3} O_{4} + 4 H_{2}$ . Now, Hydrogen has 8 atoms on LHS and 8 on RHS.

Balance Remaining Elements: Finally, balance any remaining elements, typically metals.

In $F e + 4 H_{2} O \to F e_{3} O_{4} + 4 H_{2}$ , Iron has 1 atom on LHS and 3 on RHS. To balance, place '3' before $F e$ : $3 F e + 4 H_{2} O \to F e_{3} O_{4} + 4 H_{2}$ .

Verify: Count the atoms of each element on both sides to ensure the equation is balanced.

LHS: Fe = 3, H = 8, O = 4
RHS: Fe = 3, H = 8, O = 4
The balanced equation is: $3 F e + 4 H_{2} O \to F e_{3} O_{4} + 4 H_{2}$

Including Physical States:

To make a chemical equation more informative, the physical states of reactants and products are indicated using symbols:

(g) for gaseous state
(l) for liquid state
(aq) for aqueous state (solute dissolved in water)
(s) for solid state

Example:

3 F e (s) + 4 H_{2} O (g) \to F e_{3} O_{4} (s) + 4 H_{2} (g)

(Note: $H_{2} O (g)$ indicates water is used as steam).

Sometimes, reaction conditions like temperature, pressure, or a catalyst (a substance that speeds up a reaction without being consumed) are written above or below the arrow.

Example:

C O (g) + 2 H_{2} (g) \to_{340 atm}^{C a t a l y s t} C H_{3} O H (l)

Remember
When balancing chemical equations, only change the coefficients, not the subscripts within the chemical formulas, as changing subscripts alters the chemical identity of the substance.

Exam-smart
You might be given a word description of a reaction and asked to write and balance its chemical equation, or given an unbalanced equation and asked to balance it. You could also be presented with a balanced equation and asked to identify the reactants, products, or their physical states, or infer the conditions under which the reaction occurs.

Chemical Reactions and Equations

A chemical reaction is a process that leads to the transformation of one set of chemical substances to another. It involves the rearrangement of atoms and the breaking and forming of chemical bonds. For example, when iron rusts, a chemical reaction occurs where iron atoms react with oxygen in the air to form iron oxide.

Describing Chemical Reactions

We can describe chemical reactions in two ways: using word equations or chemical equations.

Word Equations: These are simple descriptions of a reaction using words. They state the reactants (substances that react) and the products (substances formed).
For instance, the rusting of iron can be written as: Iron + Oxygen $\to$ Iron Oxide
The substances on the left side of the arrow are the reactants, and those on the right side are the products. The arrow ( $\to$ ) indicates that a reaction has occurred, transforming the reactants into products.
Chemical Equations: These are more precise and use chemical formulas to represent reactants and products. They are a shorthand way of writing word equations.
For example, the rusting of iron can be written using chemical formulas:

F e + O_{2} \to F e_{2} O_{3}

This equation shows that iron (Fe) reacts with oxygen ( $O_{2}$ ) to produce iron(III) oxide ( $F e_{2} O_{3}$ ).

Balanced Chemical Equations

A balanced chemical equation is one where the number of atoms of each element is the same on both sides of the equation. This adheres to the Law of Conservation of Mass, which states that matter cannot be created or destroyed in a chemical reaction; it can only be rearranged.

Why Balancing is Important: In an unbalanced equation, the number of atoms of a particular element might differ between reactants and products. This is not physically possible. Balancing ensures that we account for all the atoms involved.
How to Balance:

Write the unbalanced chemical equation with the correct chemical formulas for reactants and products.
Count the number of atoms of each element on both sides of the arrow.
Use coefficients (numbers placed in front of chemical formulas) to balance the atoms. Coefficients multiply the entire molecule. Never change the subscripts within a chemical formula, as this would change the identity of the substance.
Balance one element at a time. It's often helpful to balance elements that appear in only one reactant and one product first.
Check your final equation to ensure that the number of atoms of each element is equal on both sides.

Example: Let's balance the equation for the formation of water from hydrogen and oxygen: Unbalanced: $H_{2} + O_{2} \to H_{2} O$

Count atoms:

Reactants: H = 2, O = 2
Products: H = 2, O = 1

Balance Oxygen: There are 2 oxygen atoms on the left and 1 on the right. Place a coefficient of 2 in front of $H_{2} O$ on the right side:

H_{2} + O_{2} \to 2 H_{2} O

Recount atoms:

Reactants: H = 2, O = 2
Products: H = 4, O = 2

Balance Hydrogen: Now there are 2 hydrogen atoms on the left and 4 on the right. Place a coefficient of 2 in front of $H_{2}$ on the left side:

2 H_{2} + O_{2} \to 2 H_{2} O

Final Check:

Reactants: H = $2 \times 2 = 4$ , O = 2
Products: H = $2 \times 2 = 4$ , O = $2 \times 1 = 2$ The equation is now balanced.

State Symbols

Chemical equations can also include state symbols to indicate the physical state of the reactants and products.

(g) for gaseous state
(l) for liquid state
(aq) for aqueous solution (substance dissolved in water)
(s) for solid state
Example: The balanced equation for the reaction between sodium hydroxide and hydrochloric acid, with state symbols, is:

N a O H (a q) + H C l (a q) \to N a C l (a q) + H_{2} O (l)

Reaction Conditions

Sometimes, the conditions under which a reaction takes place are crucial and are indicated above or below the arrow in the chemical equation. These conditions can include:

Temperature: For example, a specific temperature might be required.
Pressure: Some reactions need high pressure.
Catalyst: A catalyst is a substance that increases the rate of a chemical reaction without itself undergoing any permanent chemical change.
Example: The synthesis of methanol from carbon monoxide and hydrogen gas requires specific conditions:

C O (g) + 2 H_{2} (g) \overset{300 atm, catalyst}{\to} C H_{3} O H (l)

Here, 300 atm is the pressure and a catalyst is used.

Remember
Balancing chemical equations is essential to obey the Law of Conservation of Mass. Always use coefficients, never change subscripts.

Exam-smart
Questions might ask you to write a balanced chemical equation from a word equation, or to identify the state symbols of reactants and products given a description of a reaction. You might also be given conditions like temperature or a catalyst and asked to represent them correctly in the equation.

Types of Chemical Reactions

Chemical reactions involve the breaking and making of bonds between atoms to form new substances. These reactions can be classified based on how reactants combine or break down to form products, or based on the transfer of electrons.

Combination Reaction

A combination reaction is a reaction where two or more substances (elements or compounds) combine to form a single product.

How it works: Simpler substances merge to create a more complex substance.
Example: The reaction between calcium oxide (quick lime) and water to form calcium hydroxide (slaked lime).

C a O (s) + H_{2} O (l) \to C a (O H)_{2} (a q)

Exothermic Nature: Many combination reactions release heat. These are called exothermic reactions.
Burning of coal:

C (s) + O_{2} (g) \to C O_{2} (g)

Formation of water:

2 H_{2} (g) + O_{2} (g) \to 2 H_{2} O (l)

Respiration is also an exothermic process where glucose breaks down to provide energy.

C_{6} H_{12} O_{6} (a q) + 6 O_{2} (a q) \to 6 C O_{2} (a q) + 6 H_{2} O (l) + energy

Application: Slaked lime is used for whitewashing walls. It reacts with carbon dioxide in the air to form a thin layer of calcium carbonate, giving a shiny finish.

C a (O H)_{2} (a q) + C O_{2} (g) \to C a C O_{3} (s) + H_{2} O (l)

Remember
The key indicator of a combination reaction is the formation of only one product from multiple reactants.

Exam-smart
A question might describe a reaction where two gases combine to form a liquid, or a solid reacts with a gas to form a new solid. You need to identify that only one new substance is formed.

Decomposition Reaction

A decomposition reaction is a reaction where a single reactant breaks down to form two or more simpler products. These reactions require energy input.

How it works: A complex compound splits into simpler substances. This energy can be in the form of heat, light, or electricity.
Types of Decomposition Reactions:
Thermal Decomposition: Decomposition caused by heating.
Ferrous sulphate decomposes into ferric oxide, sulphur dioxide, and sulphur trioxide.

2 F e S O_{4} (s) \overset{Heat}{\to} F e_{2} O_{3} (s) + S O_{2} (g) + S O_{3} (g)

(Green ferrous sulphate crystals turn into reddish-brown ferric oxide)

Calcium carbonate (limestone) decomposes into calcium oxide (quick lime) and carbon dioxide. This is important in cement manufacturing.

C a C O_{3} (s) \overset{Heat}{\to} C a O (s) + C O_{2} (g)

Lead nitrate decomposes to form lead oxide, nitrogen dioxide, and oxygen.

2 P b (N O_{3})_{2} (s) \overset{Heat}{\to} 2 P b O (s) + 4 N O_{2} (g) + O_{2} (g)

(Brown fumes of $N O_{2}$ are observed)

Electrolytic Decomposition: Decomposition caused by passing electricity.
Electrolysis of water produces hydrogen gas at the cathode and oxygen gas at the anode. The volume of hydrogen collected is double that of oxygen.

2 H_{2} O (l) \overset{Electricity}{\to} 2 H_{2} (g) + O_{2} (g)

Photodecomposition: Decomposition caused by light.
Silver chloride decomposes into silver and chlorine gas when exposed to sunlight.

2 A g C l (s) \overset{Sunlight}{\to} 2 A g (s) + C l_{2} (g)

(White silver chloride turns grey)

Silver bromide also decomposes similarly.

2 A g B r (s) \overset{Sunlight}{\to} 2 A g (s) + B r_{2} (g)

These reactions are used in black and white photography.
Endothermic Nature: Reactions that absorb energy are called endothermic reactions. Decomposition reactions are typically endothermic.

Remember
Decomposition reactions require energy input. The type of energy used (heat, light, electricity) determines the sub-type of decomposition.

Exam-smart
You might be given the colour change of a substance upon heating or exposure to light (e.g., green to brown for ferrous sulphate, white to grey for silver chloride) and asked to identify the type of reaction or the products.

Displacement Reaction

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

How it works: A more reactive metal or non-metal "pushes out" a less reactive one from its solution. Reactivity is determined by a reactivity series.
Example: Iron displaces copper from copper sulphate solution.

F e (s) + C u S O_{4} (a q) \to F e S O_{4} (a q) + C u (s)

The blue colour of copper sulphate solution fades, and the iron nail gets coated with reddish-brown copper.
Other Examples:
Zinc displaces copper:

Z n (s) + C u S O_{4} (a q) \to Z n S O_{4} (a q) + C u (s)

Lead displaces copper:

P b (s) + C u C l_{2} (a q) \to P b C l_{2} (a q) + C u (s)

Reactivity Series: Elements higher in the reactivity series can displace elements lower in the series.
Metals: Potassium > Sodium > Calcium > Magnesium > Aluminium > Carbon > Zinc > Iron > Lead > Hydrogen > Copper > Mercury > Silver > Gold
Non-metals: Fluorine > Chlorine > Bromine > Iodine

Remember
For a displacement reaction to occur, the reactant element must be more reactive than the element it is trying to displace.

Exam-smart
Questions will present reactions like "Metal A reacts with the salt of Metal B to form the salt of Metal A and Metal B". You need to use your knowledge of the reactivity series to confirm if the reaction will happen and identify A and B.

Double Displacement Reaction

A double displacement reaction is a reaction in which ions are exchanged between two reactants, forming two new compounds.

How it works: The cations and anions of the two reacting compounds switch partners.
Precipitation Reaction: Often, one of the new compounds formed is insoluble in water and settles down as a precipitate. Such reactions are called precipitation reactions.
Example: The reaction between sodium sulphate and barium chloride.

N a_{2} S O_{4} (a q) + B a C l_{2} (a q) \to B a S O_{4} (s) + 2 N a C l (a q)

Barium sulphate ( $B a S O_{4}$ ) is a white precipitate. Sodium chloride ( $N a C l$ ) remains dissolved in the solution.
Other Examples:
Reaction between lead(II) nitrate and potassium iodide, forming a yellow precipitate of lead(II) iodide.

P b (N O_{3})_{2} (a q) + 2 K I (a q) \to P b I_{2} (s) + 2 K N O_{3} (a q)

Remember
Look for reactions where two compounds react to form two new compounds, and often a precipitate is formed.

Exam-smart
You might be given a description of mixing two solutions that results in a cloudy appearance or a solid settling down. You'll need to identify the reactants and the insoluble product (precipitate) and write the balanced chemical equation.

Oxidation and Reduction (Redox Reactions)

Oxidation is the gain of oxygen or loss of hydrogen by a substance. Reduction is the loss of oxygen or gain of hydrogen by a substance.

How it works: In a chemical reaction, one substance might lose electrons (get oxidized) while another gains electrons (gets reduced).
Redox Reactions: Reactions where both oxidation and reduction occur simultaneously are called oxidation-reduction reactions or redox reactions.
Examples:
When copper is heated in air, it gains oxygen to form copper(II) oxide. Copper is oxidized.

2 C u (s) + O_{2} (g) \overset{Heat}{\to} 2 C u O (s)

When copper(II) oxide is heated with hydrogen, it loses oxygen to form copper. Copper(II) oxide is reduced. Hydrogen gains oxygen to form water. Hydrogen is oxidized.

C u O (s) + H_{2} (g) \overset{Heat}{\to} C u (s) + H_{2} O (l)

In the reaction:

Z n O (s) + C (s) \to Z n (s) + C O (g)

ZnO is reduced to Zn (loses oxygen).
C is oxidized to CO (gains oxygen).
In the reaction:

M n O_{2} (s) + 4 H C l (a q) \to M n C l_{2} (a q) + 2 H_{2} O (l) + C l_{2} (g)

MnO $_{2}$ is reduced to MnCl $_{2}$ (loses oxygen).
HCl is oxidized to Cl $_{2}$ (loses hydrogen).

Remember
If a substance is being oxidized, another substance in the same reaction must be reduced. Electrons are conserved.

Exam-smart
You will be given a chemical equation and asked to identify which substance is oxidized and which is reduced, or to identify the oxidizing and reducing agents.

HAVE YOU OBSERVED THE EFFECTS OF OXIDATION REACTIONS IN EVERYDAY LIFE?

Oxidation has noticeable effects in our daily lives, particularly on metals.

Corrosion

Corrosion is the process where a metal is attacked by substances around it (like moisture, acids, or oxygen) and gets degraded.

Rusting of Iron: This is the most common example. When iron is exposed to moisture and air, it forms a reddish-brown coating of iron oxides/hydroxides, known as rust.
The process is complex, involving electrochemical steps.
Tarnishing of Silver: Silver articles turn black when left exposed to air due to the formation of silver sulphide.
Green Coating on Copper: Copper articles develop a green coating when exposed to moist air. This coating is a mixture of copper carbonate and copper hydroxide.
Prevention: Metals can be protected from corrosion by:
Painting: Creates a barrier to prevent contact with air and moisture.
Oiling/Greasing: Similar to painting, prevents contact.
Galvanization: Coating with a more reactive metal like zinc. Zinc corrodes preferentially, protecting the iron.
Alloying: Mixing the metal with other elements to form an alloy (e.g., stainless steel is an alloy of iron, chromium, and nickel, which is resistant to rust).
Anodizing: Forming a protective oxide layer on the surface (e.g., anodized aluminium).

Remember
Corrosion is a slow process of degradation of metals due to chemical reactions with their environment.

Exam-smart
Questions might describe an iron object turning reddish-brown, or silver jewellery becoming dull. You'll need to identify this as corrosion (specifically rusting or tarnishing) and suggest methods to prevent it.

Corrosion

Corrosion is the process by which a metal is attacked by substances in its surroundings, such as moisture and acids. This attack causes the metal to degrade or wear away. This is a natural process that cannot be prevented entirely, but it can be slowed down.

Examples of Corrosion:
Rusting of Iron: This is the most common and well-known example of corrosion. When iron is exposed to moist air, it reacts to form a reddish-brown flaky coating called rust.
Tarnishing of Silver: Silver objects develop a black coating when exposed to air for a period. This is due to the formation of silver sulfide ( $A g_{2} S$ ).
Green Coating on Copper: Copper objects, like statues or utensils, turn green when exposed to moist air. This green layer is a mixture of copper carbonate ( $C u C O_{3}$ ) and copper hydroxide ( $C u (O H)_{2}$ ).
Why Corrosion Matters:
Corrosion causes significant damage to iron structures like car bodies, bridges, iron railings, and ships.
It leads to a substantial economic loss as damaged iron articles need frequent replacement.
Chemical Process of Rusting: Rusting is an electrochemical process. It requires both moisture and oxygen.

Iron acts as the anode and gets oxidized to iron(II) ions:

F e (s) \to F e^{2 +} (a q) + 2 e^{-}

At the cathode (regions where oxygen is available), oxygen is reduced in the presence of hydrogen ions (from dissolved $C O_{2}$ in water) to form water:

O_{2} (g) + 4 H^{+} (a q) + 4 e^{-} \to 2 H_{2} O (l)

The $F e^{2 +}$ ions formed are further oxidized by oxygen to $F e^{3 +}$ ions:

4 F e^{2 +} (a q) + O_{2} (g) + 4 H^{+} (a q) \to 4 F e^{3 +} (a q) + 2 H_{2} O (l)

Finally, the $F e^{3 +}$ ions react with water to form hydrated iron(III) oxide, which is rust:

2 F e^{3 +} (a q) + 3 H_{2} O (l) \to F e_{2} O_{3} \cdot x H_{2} O (s) (Rust)

The 'x' indicates that rust is a hydrated compound, meaning it contains water molecules.

Methods to Prevent Corrosion:
Painting or Oiling: Applying a coat of paint or oil on the surface of iron objects prevents them from coming into contact with moisture and air.
Greasing: Similar to painting, greasing prevents contact with the environment.
Galvanization: This involves coating iron with a thin layer of zinc. Zinc is more reactive than iron and corrodes preferentially, protecting the iron underneath. This is called a sacrificial protection.
Alloying: Making stainless steel, an alloy of iron with nickel and chromium, makes it resistant to corrosion.
Electroplating: Coating iron with other less reactive metals like tin or chromium.

Remember
Rusting is an electrochemical process that requires both oxygen and water. The presence of acids (like in rainwater) and salts (like in seawater) accelerates the rate of rusting.

Exam-smart
A question might describe an object (e.g., a bridge, a ship) and ask why it is painted or how its lifespan can be extended. The answer would involve understanding that painting/oiling prevents contact with corrosive agents, or discussing sacrificial protection like galvanization where a more reactive metal corrodes instead of iron.

Rancidity

Rancidity is the condition where food substances containing fats and oils spoil over time, leading to an unpleasant smell and taste. This happens because the fats and oils get oxidised when exposed to air.

The Process: When fats and oils are exposed to oxygen in the air, they undergo oxidation. This chemical reaction breaks down the fats and oils into smaller molecules, such as aldehydes and fatty acids, which have a stale, sour smell and taste. This process is slow at room temperature but can be accelerated by heat and light.
Why it Matters: Rancid food is not only unpleasant to eat but can also be harmful to health. Therefore, it's crucial to prevent food from becoming rancid.
Preventing Rancidity: Several methods are used to prevent or slow down the oxidation of fats and oils in food:
Adding Antioxidants: Antioxidants are substances that prevent or slow down oxidation. They work by reacting with oxygen or by preventing the chain reactions that lead to oxidation. Common examples include BHA (Butylated Hydroxyanisole) and BHT (Butylated Hydroxytoluene), which are often added to packaged food items.
Storing in Airtight Containers: Airtight containers limit the exposure of food to oxygen, thereby slowing down the oxidation process.
Storing in Refrigerators: Low temperatures slow down the rate of chemical reactions, including oxidation. Refrigeration helps to preserve food for longer periods.
Storing away from Light: Light can also catalyze oxidation reactions, so storing food away from direct light is beneficial.
Flushing with Nitrogen Gas: Many food manufacturers, especially for chips and other snacks, flush the packaging with nitrogen gas ( $N_{2}$ ). Nitrogen is an inert gas, meaning it does not react with the fats and oils. By filling the bag with nitrogen, the amount of oxygen present is significantly reduced, thus preventing oxidation and keeping the chips fresh and crisp for a longer time.

Remember
Rancidity is the spoilage of fats and oils due to oxidation. Nitrogen gas is used in chip packets because it displaces oxygen, preventing oxidation.

Exam-smart
A question might ask why manufacturers add nitrogen gas to chip packets or why you should store oily foods in airtight containers. The answer relates to preventing the oxidation of fats and oils, which leads to rancidity and an unpleasant taste/smell.

Chemical Reactions and Equations — Notes

Read the notes

Chemical Equations

Writing a Chemical Equation

Balanced Chemical Equations

Chemical Reactions and Equations

Describing Chemical Reactions

Balanced Chemical Equations

State Symbols

Reaction Conditions

Types of Chemical Reactions

Combination Reaction

Decomposition Reaction

Displacement Reaction

Double Displacement Reaction

Oxidation and Reduction (Redox Reactions)

HAVE YOU OBSERVED THE EFFECTS OF OXIDATION REACTIONS IN EVERYDAY LIFE?

Corrosion

Corrosion

Rancidity

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