1. Evidence of Chemical Reactions

2. Reactants and Products

3. Writing Chemical Equations

Introduction

Chemists around the world use chemical equations as a universal language to describe chemical reactions. Instead of writing the full names of substances every time, they use chemical symbols and chemical formulae to show exactly which substances react and which new substances are produced.

Chemical equations are much more than a shorthand notation—they show how atoms are rearranged during a reaction. Every equation represents billions upon billions of particles reacting in exactly the same way. Learning to read and write chemical equations is an essential step toward understanding chemical reactions, conservation of mass, and stoichiometry.

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What Is a Chemical Equation?

A chemical equation is a symbolic representation of a chemical reaction.

It uses:

• Chemical symbols.
• Chemical formulae.
• A reaction arrow.

to show:

• The reactants.
• The products.

General form:

Reactants → Products

Example:

2H₂ + O₂ → 2H₂O

This equation shows hydrogen reacting with oxygen to produce water.

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Figure 1. Chemical equations use symbols and formulae to represent chemical reactions.

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Word Equations vs Symbol Equations

There are two common ways to describe reactions.

Word Equation

Uses the names of substances.

Example:

Hydrogen + Oxygen → Water

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Symbol Equation

Uses chemical symbols and formulae.

Example:

H₂ + O₂ → H₂O

Symbol equations are shorter, more precise, and understood internationally.

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Comparing the Two

Scientists usually use symbol equations in laboratories and textbooks.

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Figure 2. Word equations use substance names, while symbol equations use chemical symbols and formulae.

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Using Chemical Symbols and Formulae

Every element has its own chemical symbol.

Examples:

Compounds are represented using chemical formulae.

Examples:

Correct symbols and formulae are essential for writing chemical equations accurately.

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Identifying the Substances

In every chemical equation:

• Reactants appear before the arrow.
• Products appear after the arrow.

Example:

CH₄ + 2O₂ → CO₂ + 2H₂O

Reactants:

• Methane (CH₄)
• Oxygen (O₂)

Products:

• Carbon dioxide (CO₂)
• Water (H₂O)

The arrow means "produces" or "forms."

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Reading Chemical Equations

Consider:

Mg + O₂ → MgO

This can be read as:

"Magnesium reacts with oxygen to produce magnesium oxide."

Chemical equations describe:

• What reacts.
• What is produced.

They do not usually explain how the reaction occurs.

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Figure 3. Reactants are written on the left of the arrow, while products are written on the right.

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What Happens at the Particle Level?

A chemical equation represents the behaviour of atoms and molecules.

During a reaction:

• Chemical bonds break.
• Atoms rearrange.
• New chemical bonds form.
• New substances are produced.

The atoms themselves are not created or destroyed.

Only the way they are connected changes.

For example:

Hydrogen molecules and oxygen molecules rearrange to form water molecules.

This particle-level view explains why the products have different properties from the reactants.

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Chemical Equations Represent Billions of Particles

Although a chemical equation appears small on paper, it represents an enormous number of particles.

For example:

2H₂ + O₂ → 2H₂O

Actually represents:

• Billions of hydrogen molecules.
• Billions of oxygen molecules.
• Billions of water molecules.

Chemists use equations because they apply equally well to a tiny laboratory sample or a large industrial process.

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Figure 4. Chemical equations represent atoms being rearranged into new molecules during a reaction.

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Examples of Chemical Equations

Notice that symbol equations use chemical symbols and formulae instead of names.

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Why Chemical Equations Are Important

Chemical equations allow scientists to:

• Describe reactions accurately.
• Predict products.
• Communicate internationally.
• Study industrial processes.
• Calculate quantities in reactions.

They are one of the most important tools in chemistry.

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Figure 5. Chemical equations are the universal language used by chemists to describe reactions.

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Worked Example

Question

Write the symbol equation for the following word equation.

Hydrogen + Chlorine → Hydrogen chloride

Solution

Step 1

Identify the symbols.

Hydrogen = H₂

Chlorine = Cl₂

Hydrogen chloride = HCl

Step 2

Write the equation.

H₂ + Cl₂ → 2HCl

This equation shows hydrogen molecules reacting with chlorine molecules to form hydrogen chloride molecules.

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Real-World Connection

Chemical equations are used everywhere in science and industry. Engineers use them when designing processes to manufacture fertilisers, medicines, plastics, and fuels. Environmental scientists use chemical equations to study pollution and climate change, while doctors and pharmacists use them to understand how medicines react inside the human body.

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Did You Know?

Every year, chemists write millions of chemical equations in scientific journals, laboratories, and factories. Because chemical symbols and equations are standardised worldwide, a chemist in Canada can read and understand a reaction written by a scientist in Brazil, India, or Germany without needing to translate it into another language.

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Key Terms

Chemical equation – A symbolic representation of a chemical reaction using chemical symbols and formulae.

Chemical formula – A combination of symbols and numbers showing the types and numbers of atoms in a substance.

Chemical symbol – A one- or two-letter abbreviation representing an element.

Particle – A general term for atoms, molecules, or ions that make up matter.

Product – A substance formed during a chemical reaction.

Reactant – A starting substance in a chemical reaction.

Reaction arrow (→) – A symbol meaning "produces" or "forms" in a chemical equation.

Symbol equation – A chemical equation written using chemical symbols and formulae.

Word equation – A chemical equation written using the names of substances.

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Key Takeaways

• A chemical equation uses symbols and formulae to represent a chemical reaction.
• Word equations use substance names, while symbol equations use chemical symbols and formulae.
• Reactants are written before the reaction arrow, and products are written after it.
• Chemical equations represent the rearrangement of atoms into new substances at the particle level.
• Chemical equations provide a universal way for scientists to describe chemical reactions.
• Understanding chemical equations prepares students for balancing equations and quantitative chemistry.

4. Conservation of Mass

Introduction

When a chemical reaction takes place, it may seem as though matter disappears or is created. For example, when wood burns, only a small amount of ash remains, while an effervescent tablet seems to disappear when dropped into water. However, careful scientific measurements show that matter is never lost or created during ordinary chemical reactions.

This idea is known as the Law of Conservation of Mass. It is one of the most important principles in chemistry because it explains why chemical equations must be balanced and why the same atoms are present before and after every reaction. Understanding conservation of mass helps scientists predict the amounts of substances produced during chemical reactions.

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The Law of Conservation of Mass

The Law of Conservation of Mass states:

Mass is neither created nor destroyed during a chemical reaction.

This means:

• The total mass of the reactants equals the total mass of the products.

In every ordinary chemical reaction:

Total mass before reaction = Total mass after reaction

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Figure 1. The total mass remains the same before and after a chemical reaction.

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Why Mass Is Conserved

Chemical reactions do not create or destroy atoms.

Instead:

• Existing chemical bonds break.
• Atoms are rearranged.
• New chemical bonds form.

The atoms are simply reorganised into different combinations.

Since the same atoms are present before and after the reaction, the total mass remains constant.

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The Particle Model

The particle model explains conservation of mass.

Before the reaction:

• Atoms are arranged in the reactant particles.

During the reaction:

• Bonds break.
• Atoms rearrange.

After the reaction:

• The same atoms are arranged into product particles.

No atoms disappear.

No new atoms appear.

Only their arrangement changes.

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Figure 2. The particle model shows that atoms are rearranged, not created or destroyed.

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Rearrangement of Atoms

Consider the reaction:

Hydrogen + Oxygen → Water

At the particle level:

• Hydrogen atoms remain hydrogen atoms.
• Oxygen atoms remain oxygen atoms.

The atoms simply combine differently to form water molecules.

Every hydrogen atom present before the reaction is still present afterwards.

The same is true for every oxygen atom.

This is why mass is conserved.

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Chemical Equations Show Conservation of Mass

Balanced chemical equations demonstrate conservation of mass.

Example:

2H₂ + O₂ → 2H₂O

Count the atoms.

Before:

• Hydrogen atoms = 4
• Oxygen atoms = 2

After:

• Hydrogen atoms = 4
• Oxygen atoms = 2

The number of each type of atom is the same on both sides.

Therefore, mass is conserved.

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Figure 3. Balanced chemical equations contain the same number of each type of atom on both sides.

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Why Chemical Equations Must Be Balanced

Chemical equations are balanced because:

• Atoms cannot be created.
• Atoms cannot be destroyed.

Only coefficients may be changed when balancing equations.

Subscripts must never be changed because they would change the identity of the substance.

Example:

Correct:

2H₂ + O₂ → 2H₂O

Incorrect:

H₂ + O₂ → H₂O₂

Changing the subscripts creates a different compound.

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Applying Conservation of Mass

The law can be used to solve simple problems.

Example

A reaction starts with:

• 12 g of carbon.
• 32 g of oxygen.

Total reactant mass:

44 g

If all the reactants form carbon dioxide,

Mass of carbon dioxide:

44 g

The total mass remains unchanged.

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Open and Closed Systems

Sometimes it appears that mass changes.

Example:

A burning candle loses mass.

Why?

Some products are gases that escape into the air.

If the reaction occurred in a closed container, the total mass would remain exactly the same.

The law always applies, provided all reactants and products are included.

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Figure 4. In a closed system, the total mass remains constant because no substances can enter or leave.

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Everyday Examples

Conservation of mass occurs in:

• Rusting iron.
• Burning fuels.
• Photosynthesis.
• Cellular respiration.
• Cooking.
• Digestion.

Although substances change, the total amount of matter remains the same.

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Why Conservation of Mass Is Important

The law helps scientists:

• Balance chemical equations.
• Predict product masses.
• Design industrial processes.
• Understand chemical reactions.
• Calculate quantities in chemistry.

It is one of the fundamental laws of science.

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Figure 5. Conservation of mass is the foundation of balancing chemical equations and quantitative chemistry.

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Worked Example

Question

A reaction begins with:

• 18 g of hydrogen.
• 144 g of oxygen.

Assuming all the reactants react completely, what is the total mass of water produced?

Solution

Total mass of reactants:

18 g + 144 g = 162 g

According to the Law of Conservation of Mass:

Total mass of products = Total mass of reactants

Answer: 162 g of water

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Real-World Connection

Chemical engineers rely on the Law of Conservation of Mass when designing factories that manufacture medicines, fertilisers, plastics, and fuels. By knowing the masses of the reactants, they can predict how much product will be formed and how much waste may be produced. This helps reduce costs, improve efficiency, and minimise environmental impacts.

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Did You Know?

The Law of Conservation of Mass was first clearly stated in the late 1700s by the French chemist Antoine Lavoisier, often called the "Father of Modern Chemistry." His careful experiments using sealed containers showed that the total mass of substances remains constant during chemical reactions, helping transform chemistry into a precise quantitative science.

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Key Terms

Balanced chemical equation – A chemical equation with the same number of each type of atom on both sides.

Closed system – A system in which no matter can enter or leave.

Coefficient – A number placed before a chemical formula indicating the number of particles involved in a reaction.

Conservation of mass – The principle that mass is neither created nor destroyed during ordinary chemical reactions.

Particle model – A model describing matter as tiny particles that can rearrange during chemical reactions.

Product – A substance formed during a chemical reaction.

Reactant – A starting substance in a chemical reaction.

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Key Takeaways

• The Law of Conservation of Mass states that mass is neither created nor destroyed during ordinary chemical reactions.
• During a reaction, atoms are rearranged, but no atoms are created or destroyed.
• The particle model explains conservation of mass because the same atoms are present before and after the reaction.
• Balanced chemical equations demonstrate conservation of mass by showing equal numbers of each type of atom on both sides.
• In a closed system, the total mass of the reactants always equals the total mass of the products.
• Conservation of mass is a fundamental principle used throughout chemistry to predict reaction outcomes and balance chemical equations.

5. Balancing Chemical Equations

Introduction

Chemical equations describe what happens during a chemical reaction, but they must also obey one of the most important laws in science: the Law of Conservation of Mass. Because atoms cannot be created or destroyed during ordinary chemical reactions, the same number of each type of atom must appear on both sides of every chemical equation.

Balancing chemical equations ensures that equations accurately represent what happens during a reaction. By adjusting coefficients (the numbers placed in front of chemical formulae), chemists show that the number of atoms remains the same before and after the reaction. Balancing equations is an essential skill that prepares students for quantitative chemistry and stoichiometry.

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Why Must Chemical Equations Be Balanced?

Chemical equations must be balanced because:

• Atoms are not created during reactions.
• Atoms are not destroyed during reactions.
• Matter is conserved.

A balanced equation has:

• The same number of each type of atom on both sides.

This demonstrates the Law of Conservation of Mass.

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Figure 1. Balanced chemical equations contain the same number of each type of atom before and after the reaction.

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Counting Atoms

Before balancing an equation, count the atoms on each side.

Example:

H₂ + O₂ → H₂O

Reactants

Hydrogen = 2 atoms

Oxygen = 2 atoms

Products

Hydrogen = 2 atoms

Oxygen = 1 atom

The oxygen atoms are not balanced.

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Using Coefficients

A coefficient is a whole number placed before a chemical formula.

Coefficients multiply every atom in the formula.

Example:

2H₂O

Contains:

• Hydrogen: 2 × 2 = 4 atoms
• Oxygen: 2 × 1 = 2 atoms

Only coefficients are changed when balancing equations.

Subscripts must never be changed, because changing a subscript changes the identity of the substance.

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Figure 2. Coefficients change the number of molecules, while subscripts define the substance itself.

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Balancing a Simple Equation

Example:

H₂ + O₂ → H₂O

Step 1

Count atoms.

Hydrogen:

2 → 2

Oxygen:

2 → 1

Not balanced.

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Step 2

Balance oxygen first.

Place a coefficient of 2 before H₂O.

H₂ + O₂ → 2H₂O

Now count again.

Hydrogen:

2 → 4

Oxygen:

2 → 2

Hydrogen is now unbalanced.

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Step 3

Balance hydrogen.

Place a coefficient of 2 before H₂.

2H₂ + O₂ → 2H₂O

Count again.

Hydrogen:

4 → 4

Oxygen:

2 → 2

The equation is now balanced.

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A Step-by-Step Method

To balance any equation:

1. Write the correct chemical formulae.
2. Count the atoms on each side.
3. Balance one element at a time using coefficients.
4. Recount all atoms after each change.
5. Continue until every element is balanced.
6. Check that the coefficients are in the simplest whole-number ratio.

Following the same method each time helps avoid mistakes.

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Figure 3. A systematic approach makes balancing chemical equations much easier.

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More Examples

Example 1

Unbalanced:

Mg + O₂ → MgO

Balanced:

2Mg + O₂ → 2MgO

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Example 2

Unbalanced:

Fe + O₂ → Fe₂O₃

Balanced:

4Fe + 3O₂ → 2Fe₂O₃

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Example 3

Unbalanced:

N₂ + H₂ → NH₃

Balanced:

N₂ + 3H₂ → 2NH₃

Each equation has the same number of every type of atom on both sides.

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Checking Conservation of Mass

Balanced equations obey the Law of Conservation of Mass.

Example:

2Na + Cl₂ → 2NaCl

Count the atoms.

Before:

• Sodium = 2
• Chlorine = 2

After:

• Sodium = 2
• Chlorine = 2

Because the number of each atom is unchanged, mass is conserved.

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Figure 4. Counting atoms confirms that balanced equations obey the Law of Conservation of Mass.

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Balancing More Complex Equations

As equations become more complex:

• More elements are involved.
• More coefficients are needed.

Tips:

• Leave hydrogen and oxygen until last if possible.
• Balance elements that appear in only one reactant and one product first.
• Never change subscripts.
• Recount atoms frequently.

Practice is the key to becoming confident.

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Common Mistakes

Avoid these errors:

Changing Subscripts

Incorrect:

H₂O → H₂O₂

This creates a different substance.

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Forgetting to Recount

Always count the atoms after changing a coefficient.

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Using Fractions

Final balanced equations should use the smallest whole-number coefficients.

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Figure 5. When balancing equations, only coefficients should be changed—never subscripts.

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Why Balancing Equations Is Important

Balanced equations allow chemists to:

• Predict reaction products.
• Calculate masses.
• Design industrial processes.
• Manufacture chemicals safely.
• Understand chemical reactions accurately.

Every area of chemistry depends on balanced equations.

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Worked Example

Question

Balance the equation:

Al + O₂ → Al₂O₃

Solution

Step 1

Count atoms.

Al:

1 → 2

O:

2 → 3

Not balanced.

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Step 2

Balance oxygen.

Use:

3O₂ and 2Al₂O₃

Now oxygen:

6 → 6

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Step 3

Balance aluminium.

Products contain:

4 aluminium atoms.

Place a coefficient of 4 before Al.

Balanced equation:

4Al + 3O₂ → 2Al₂O₃

Check:

Aluminium:

4 → 4

Oxygen:

6 → 6

The equation is balanced.

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Real-World Connection

Every chemical manufacturing plant uses balanced chemical equations to calculate exactly how much of each reactant is needed to produce the desired amount of product. Whether producing medicines, fertilisers, plastics, or fuels, balanced equations help reduce waste, lower costs, improve safety, and ensure efficient production.

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Did You Know?

A modern fertiliser factory may use millions of balanced chemical equations every year in computer-controlled systems. These calculations ensure that the correct amounts of raw materials are mixed together, helping produce fertilisers efficiently while minimising waste and energy use.

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Key Terms

Balanced chemical equation – A chemical equation with equal numbers of each type of atom on both sides.

Coefficient – A whole number placed before a chemical formula that multiplies the number of molecules or formula units.

Chemical equation – A symbolic representation of a chemical reaction.

Conservation of mass – The law stating that mass is neither created nor destroyed during ordinary chemical reactions.

Product – A substance formed during a chemical reaction.

Reactant – A starting substance in a chemical reaction.

Subscript – A small number in a chemical formula indicating the number of atoms of an element in one molecule or formula unit.

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Key Takeaways

• Chemical equations must be balanced because atoms are neither created nor destroyed during chemical reactions.
• Balancing ensures that the same number of each type of atom appears on both sides of the equation.
• Only coefficients should be changed when balancing equations; subscripts must never be altered.
• Counting atoms carefully is the key to balancing equations correctly.
• Balanced equations demonstrate the Law of Conservation of Mass.
• Balancing chemical equations is an essential skill for understanding chemical reactions and performing quantitative chemistry.