Introduction

Very few pure elements are found by themselves in nature. Instead, most elements combine with other elements to form compounds. Water, table salt, carbon dioxide, and sugar are all examples of compounds that are essential for life and are used every day.

Compounds are created when atoms of different elements become chemically bonded together in fixed proportions. Once bonded, the elements form a completely new substance with properties that are often very different from those of the elements that formed it. Understanding compounds is one of the key foundations of chemistry because nearly all materials around us are compounds.

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What Is a Compound?

A compound is a pure substance formed when two or more different elements are chemically combined in fixed proportions.

The atoms in a compound are joined together by chemical bonds.

Examples include:

• Water (H₂O)
• Carbon dioxide (CO₂)
• Sodium chloride (NaCl)
• Ammonia (NH₃)

Each compound has:

• A fixed chemical formula.
• A fixed composition.
• Unique physical and chemical properties.

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Figure 1. Compounds are formed when atoms of different elements become chemically bonded together.

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Compounds Are Pure Substances

Like elements, compounds are pure substances.

This means they:

• Have a fixed composition.
• Have consistent properties.
• Contain only one type of particle (molecules or formula units).

For example:

Every water molecule contains:

• Two hydrogen atoms.
• One oxygen atom.

This ratio is always the same.

Changing the ratio would produce a different substance.

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How Compounds Form

Compounds form during chemical reactions.

During these reactions:

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

Example:

Hydrogen gas reacts with oxygen gas.

Hydrogen + Oxygen → Water

Although hydrogen and oxygen are both gases, the compound they form is a liquid at room temperature.

This demonstrates that compounds often have very different properties from the elements that compose them.

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Figure 2. Chemical reactions join atoms together to form compounds with new properties.

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Elements, Compounds, and Mixtures

It is important to distinguish between these three types of matter.

Understanding these differences helps scientists classify matter correctly.

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Understanding Chemical Formulas

A chemical formula shows:

• Which elements are present.
• How many atoms of each element are present.

Examples:

The subscript tells us how many atoms of that element are present.

Example:

H₂O

contains:

• 2 hydrogen atoms.
• 1 oxygen atom.

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Figure 3. Chemical formulas show the types and numbers of atoms present in a compound.

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Why Compounds Have Different Properties

When elements combine chemically, their atoms are arranged in a completely new way.

As a result, the compound often has properties that are very different from those of its elements.

Example: Water

Hydrogen:

• Colourless gas.
• Highly flammable.

Oxygen:

• Colourless gas.
• Supports combustion.

Water:

• Liquid.
• Extinguishes fires.
• Safe to drink.

The compound has completely different properties from either element.

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Example: Sodium Chloride

Sodium:

• Soft, highly reactive metal.
• Reacts violently with water.

Chlorine:

• Poisonous green gas.

Sodium chloride:

• White crystalline solid.
• Common table salt.
• Safe to eat in small amounts.

Again, the compound's properties are completely different from those of its elements.

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Figure 4. Compounds often have properties that are very different from the elements that form them.

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Classifying Substances

Scientists use evidence to determine whether a substance is an element, compound, or mixture.

Elements

Examples:

• Gold
• Oxygen
• Copper
• Helium

Contain only one type of atom.

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Compounds

Examples:

• Water
• Carbon dioxide
• Glucose
• Sodium chloride

Contain different elements chemically bonded together.

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Mixtures

Examples:

• Air
• Seawater
• Soil
• Salad
• Brass

Contain substances that are physically mixed rather than chemically bonded.

Scientists classify matter by examining its composition and properties.

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Why Compounds Are Important

Compounds make up most of the materials around us.

Examples include:

• Water
• Rocks
• Plastics
• Medicines
• Foods
• Fuels
• Living organisms

Nearly every biological process depends on compounds.

Without compounds, life on Earth would not exist.

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Figure 5. Many everyday materials are compounds formed from combinations of different elements.

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Worked Example

Question

Classify each substance.

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Real-World Connection

Many of the materials we rely on every day are compounds. Medicines are carefully designed compounds that treat disease, fertilisers are compounds that supply nutrients to crops, and concrete is made using compounds that give buildings their strength. Chemists continually develop new compounds to improve technology, medicine, agriculture, and environmental protection.

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Did You Know?

The human body is composed almost entirely of compounds, not pure elements. Water makes up about 60% of an adult's body mass, while proteins, carbohydrates, lipids, DNA, and minerals are all compounds made from combinations of different elements. Only tiny amounts of a few elements, such as oxygen gas in the lungs or iron in the blood, exist in relatively pure forms.

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Key Terms

Chemical bond – The force that holds atoms together in a compound.

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

Chemical reaction – A process in which substances are transformed into new substances through the breaking and forming of chemical bonds.

Compound – A pure substance made of two or more different elements chemically bonded together in fixed proportions.

Element – A pure substance consisting of only one type of atom.

Mixture – Two or more substances physically combined but not chemically bonded.

Molecule – A group of atoms held together by chemical bonds.

Pure substance – A material with a fixed composition and consistent properties.

Subscript – A small number written below the line in a chemical formula indicating the number of atoms of an element.

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Key Takeaways

• A compound is a pure substance formed when two or more different elements chemically combine.
• Compounds have a fixed composition and are held together by chemical bonds.
• Chemical formulas show which elements are present and how many atoms of each are in a compound.
• Compounds often have properties that are very different from the elements that form them.
• Elements, compounds, and mixtures can be distinguished by examining how their particles are arranged and whether substances are chemically bonded.
• Most of the materials found in living organisms and in everyday life are compounds.

Introduction

Chemists use a special language to describe substances and chemical reactions. Instead of writing the full names of elements every time, they use chemical formulae to show exactly which elements are present in a substance and how many atoms of each element it contains. Chemical formulae are like mathematical shorthand that allows scientists around the world to communicate clearly and efficiently.

Understanding chemical formulae is an essential skill in chemistry. By learning how to read and write formulae, you can identify the elements in a compound, count the number of atoms present, and prepare for balancing chemical equations and studying chemical reactions.

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What Is a Chemical Formula?

A chemical formula is a combination of chemical symbols and numbers that shows:

• The elements present in a substance.
• The number of atoms of each element.

Examples:

Each formula represents one type of compound.

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Figure 1. Chemical formulae show the elements present in a compound and the number of atoms of each.

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Reading Chemical Formulae

To interpret a chemical formula:

Step 1

Identify the chemical symbols.

These tell you which elements are present.

Step 2

Read the subscripts.

These tell you how many atoms of each element are present.

If there is no subscript, the number of atoms is 1.

Example:

CO₂

Contains:

• 1 carbon atom
• 2 oxygen atoms

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Counting Atoms

Examples:

Remember:

No subscript means one atom.

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Figure 2. Subscripts indicate the number of atoms of each element in a compound.

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What Is a Subscript?

A subscript is the small number written after and below a chemical symbol.

It tells us how many atoms of that element are present.

Examples:

H₂O

• 2 hydrogen atoms
• 1 oxygen atom

CO₂

• 1 carbon atom
• 2 oxygen atoms

C₆H₁₂O₆

Contains:

• 6 carbon atoms
• 12 hydrogen atoms
• 6 oxygen atoms

Subscripts are part of the chemical formula and must not be changed unless the substance itself changes.

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What Is a Coefficient?

A coefficient is a whole number written in front of a chemical formula.

It tells us how many molecules or formula units are present.

Example:

3H₂O

Means:

Three water molecules.

Total atoms:

Hydrogen:

3 × 2 = 6

Oxygen:

3 × 1 = 3

Coefficients multiply every atom in the formula.

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Figure 3. Subscripts count atoms within one molecule, while coefficients count the number of molecules.

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Subscripts vs Coefficients

Understanding this difference is essential when balancing chemical equations.

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Writing Chemical Formulae

Chemists use chemical symbols to write the correct formula for compounds.

Some common examples include:

The order of the elements and the subscripts must always be written correctly.

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Representing Compounds Accurately

A chemical formula gives an exact description of a compound.

For example:

H₂O

always represents water.

Changing the formula changes the substance.

Example:

CO

Carbon monoxide

CO₂

Carbon dioxide

Although both compounds contain carbon and oxygen, they are different substances with different properties.

Accurate chemical formulae are essential for clear scientific communication.

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Figure 4. Even small changes in a chemical formula create a different compound with different properties.

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Why Chemical Formulae Matter

Chemical formulae allow scientists to:

• Describe compounds accurately.
• Count atoms.
• Predict chemical reactions.
• Balance equations.
• Communicate internationally.

Every area of chemistry depends on correctly interpreting and writing chemical formulae.

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Worked Example

Question

Determine the number of atoms of each element in 2CaCl₂.

Solution

One formula unit of CaCl₂ contains:

• 1 calcium atom
• 2 chlorine atoms

The coefficient is 2, so multiply each amount by 2.

Total atoms:

• Calcium: 2 × 1 = 2
• Chlorine: 2 × 2 = 4

Answer: 2 calcium atoms and 4 chlorine atoms.

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Real-World Connection

Chemical formulae appear in many everyday products. Fertiliser bags list compounds such as NH₄NO₃ (ammonium nitrate), food labels may include NaCl (table salt), and medicines often contain active ingredients represented by chemical formulae. Scientists, pharmacists, engineers, and doctors all rely on accurate chemical formulae to communicate safely and precisely.

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Did You Know?

The chemical formula for glucose, C₆H₁₂O₆, contains 24 atoms in a single molecule:

• 6 carbon atoms
• 12 hydrogen atoms
• 6 oxygen atoms

Plants produce millions of glucose molecules every second during photosynthesis, providing the energy that supports nearly all life on Earth.

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Key Terms

Atom – The smallest particle of an element that retains its chemical properties.

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

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

Coefficient – A number written before a chemical formula indicating the number of molecules or formula units.

Compound – A pure substance made of two or more different elements chemically bonded together.

Formula unit – The simplest ratio of ions in an ionic compound.

Molecule – A group of atoms chemically bonded together.

Subscript – A small number written after a chemical symbol indicating the number of atoms of that element in one molecule or formula unit.

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Key Takeaways

• A chemical formula shows the types and numbers of atoms in a compound.
• Chemical symbols identify the elements present, while subscripts show how many atoms of each element are present.
• A missing subscript means there is one atom of that element.
• Coefficients count the number of molecules or formula units and multiply every atom in the formula.
• Writing chemical formulae accurately is essential for describing compounds and chemical reactions.
• Understanding chemical formulae is a key foundation for studying chemical bonding, reactions, and equations.

Introduction

Many of the substances we use every day, such as table salt, ceramics, and some minerals, are made of ionic compounds. These compounds are formed when atoms of metals and non-metals combine through the transfer of electrons. This type of bonding is called ionic bonding.

Ionic bonding occurs because atoms tend to become more stable when their outer electron shell is full. Metals usually lose electrons, while non-metals usually gain electrons. The resulting positively and negatively charged ions attract one another, forming strong ionic bonds. Understanding ionic bonding helps explain the properties of many important materials and lays the foundation for studying chemical reactions.

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What Is Ionic Bonding?

An ionic bond is the electrostatic force of attraction between positively charged ions and negatively charged ions.

Ionic bonding forms when:

• One or more electrons are transferred.
• A metal loses electrons.
• A non-metal gains those electrons.

This process produces oppositely charged ions that attract each other.

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Figure 1. Ionic bonding occurs when electrons are transferred from a metal atom to a non-metal atom.

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Why Do Atoms Transfer Electrons?

Atoms become more stable when their outer electron shell is full.

Most atoms achieve this by obtaining a noble gas electron arrangement.

Metals

Metals usually have:

• 1, 2, or 3 valence electrons.

They tend to:

• Lose electrons.
• Form positive ions (cations).

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Non-Metals

Non-metals usually have:

• 5, 6, or 7 valence electrons.

They tend to:

• Gain electrons.
• Form negative ions (anions).

Both atoms become more stable after the electron transfer.

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Formation of Ions

When electrons are transferred, ions are formed.

Sodium

Electron arrangement:

2,8,1

Sodium loses one electron.

Na → Na⁺ + e⁻

New arrangement:

2,8

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Chlorine

Electron arrangement:

2,8,7

Chlorine gains one electron.

Cl + e⁻ → Cl⁻

New arrangement:

2,8,8

The sodium ion and chloride ion are now attracted to each other.

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Figure 2. Electron transfer gives both ions stable outer electron shells.

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Predicting Common Ions

Many elements form predictable ions based on their group in the periodic table.

Examples:

These charges help determine the formulas of ionic compounds.

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Using Electron Arrangements

Electron arrangements explain why atoms form ions.

Example:

Magnesium

Electron arrangement:

2,8,2

Loses two electrons.

Becomes:

Mg²⁺

Electron arrangement:

2,8

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Oxygen

Electron arrangement:

2,6

Gains two electrons.

Becomes:

O²⁻

Electron arrangement:

2,8

Because both ions now have full outer shells, they are more stable.

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Figure 3. Electron arrangements help explain why atoms lose or gain electrons to form ions.

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Dot-and-Cross Diagrams

A dot-and-cross diagram shows:

• Which electrons belong to each atom.
• Which electrons are transferred.
• The completed outer electron shells.

Different symbols are used:

• Dots (•)
• Crosses (×)

This makes it easy to identify the transferred electrons.

Example:

Sodium chloride (NaCl)

• Sodium donates one electron.
• Chlorine accepts one electron.
• Both ions achieve full outer shells.

Dot-and-cross diagrams are also useful for compounds such as:

• Magnesium oxide (MgO)
• Calcium chloride (CaCl₂)

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Figure 4. Dot-and-cross diagrams show electron transfer during ionic bond formation.

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The Structure of Ionic Compounds

Ionic compounds do not exist as separate molecules.

Instead, the ions are arranged in a giant ionic lattice.

An ionic lattice is:

• A repeating three-dimensional arrangement of positive and negative ions.
• Held together by strong electrostatic forces acting in all directions.

Every positive ion is surrounded by negative ions, and every negative ion is surrounded by positive ions.

This arrangement makes ionic compounds very stable.

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Figure 5. Ionic compounds form giant crystal lattices made of alternating positive and negative ions.

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Properties of Ionic Compounds

The structure of an ionic lattice explains the properties of ionic compounds.

High Melting and Boiling Points

Strong electrostatic attractions require a large amount of energy to overcome.

As a result, ionic compounds usually have high melting and boiling points.

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Hard but Brittle

Ionic compounds are:

• Hard
• Brittle

If layers of ions are forced to slide, ions with the same charge are brought next to each other.

They repel strongly, causing the crystal to crack.

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Conduct Electricity When Molten or Dissolved

Solid ionic compounds do not conduct electricity because the ions are fixed in place.

When melted or dissolved in water:

• The ions become free to move.
• They can carry electric current.

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Crystalline Solids

Most ionic compounds form:

• Regular crystals.
• Ordered lattice structures.

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Examples of Ionic Compounds

These compounds all consist of positive and negative ions arranged in giant ionic lattices.

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Why Ionic Bonding Is Important

Ionic compounds are found throughout everyday life.

They are used in:

• Food (table salt)
• Fertilisers
• Medicines
• Batteries
• Ceramics
• Building materials

Many naturally occurring minerals are also ionic compounds.

Understanding ionic bonding helps explain the properties and uses of these materials.

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Worked Example

Question

Explain how magnesium oxide (MgO) forms.

Solution

1. Magnesium has the electron arrangement 2,8,2 and loses two electrons to form Mg²⁺.
2. Oxygen has the electron arrangement 2,6 and gains two electrons to form O²⁻.
3. Both ions now have full outer electron shells.
4. The oppositely charged ions attract each other, forming a strong ionic bond.
5. Magnesium oxide forms a giant ionic lattice.

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Real-World Connection

Table salt (NaCl) is one of the most familiar ionic compounds. It is essential for many biological processes, including nerve function and muscle contraction. Ionic compounds are also used in rechargeable batteries, where the movement of ions allows electrical energy to be stored and released. In medicine, many intravenous fluids contain dissolved ionic compounds that help maintain the body's electrolyte balance.

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Did You Know?

A single grain of table salt contains billions upon billions of sodium and chloride ions arranged in a perfectly repeating crystal lattice. There are no individual "NaCl molecules" in solid table salt—only one enormous three-dimensional network of alternating positive and negative ions extending throughout the crystal.

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Key Terms

Anion – A negatively charged ion formed when an atom gains electrons.

Cation – A positively charged ion formed when an atom loses electrons.

Dot-and-cross diagram – A diagram that uses dots and crosses to show the transfer of electrons during bond formation.

Electron transfer – The movement of one or more electrons from one atom to another.

Electrostatic attraction – The force of attraction between oppositely charged particles.

Ion – A charged particle formed when an atom gains or loses electrons.

Ionic bond – The electrostatic attraction between oppositely charged ions.

Ionic compound – A compound made of positive and negative ions held together by ionic bonds.

Ionic lattice – A giant three-dimensional arrangement of alternating positive and negative ions.

Valence electron – An electron in the outermost shell of an atom that participates in chemical bonding.

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Key Takeaways

• Ionic bonding occurs when electrons are transferred from a metal atom to a non-metal atom.
• Metals form positive ions (cations) by losing electrons, while non-metals form negative ions (anions) by gaining electrons.
• Atoms transfer electrons to achieve full outer electron shells, making them more stable.
• Dot-and-cross diagrams show how electrons are transferred during ionic bond formation.
• Ionic compounds form giant ionic lattices rather than individual molecules.
• The strong electrostatic attractions within the lattice explain why ionic compounds have high melting points, are hard but brittle, and conduct electricity when molten or dissolved in water.

Introduction

Not all atoms form compounds by transferring electrons. Many substances, including water, oxygen, carbon dioxide, methane, and the molecules found in living organisms, are formed when atoms share electrons. This type of bonding is called covalent bonding.

Covalent bonding occurs mainly between non-metal atoms. By sharing electrons, each atom can achieve a full outer electron shell, becoming more stable. Covalent bonds are responsible for forming millions of different molecules, from the oxygen we breathe to the DNA inside our cells.

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What Is Covalent Bonding?

A covalent bond is a shared pair of electrons between two atoms.

Instead of transferring electrons, the atoms share them.

Covalent bonding usually occurs between:

• Non-metal and non-metal atoms.

Examples include:

• Hydrogen (H₂)
• Oxygen (O₂)
• Water (H₂O)
• Carbon dioxide (CO₂)
• Methane (CH₄)

By sharing electrons, both atoms achieve more stable outer electron shells.

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Figure 1. Covalent bonds form when non-metal atoms share pairs of electrons.

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Why Do Atoms Share Electrons?

Most atoms become more stable when their outer electron shell is full.

Rather than gaining or losing electrons, non-metals often share electrons.

For example:

Hydrogen has:

• 1 electron.

By sharing one electron with another hydrogen atom, both atoms achieve a full first shell containing 2 electrons.

Oxygen shares electrons so that each oxygen atom effectively has 8 electrons in its outer shell.

This stable arrangement is often called the octet rule (or the duet rule for hydrogen).

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Lewis Structures

A Lewis structure shows:

• The chemical symbols.
• The shared pairs of electrons (bonds).
• The unshared (lone) pairs of electrons.

Each shared pair represents one covalent bond.

Examples:

Hydrogen (H₂)

H—H

One shared pair of electrons.

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Chlorine (Cl₂)

Cl—Cl

Each chlorine atom shares one electron.

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Water (H₂O)

Oxygen forms two covalent bonds.

H—O—H

Oxygen also has two lone pairs of electrons.

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Figure 2. Lewis structures show shared electron pairs (bonds) and lone pairs around atoms.

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Dot-and-Cross Diagrams

A dot-and-cross diagram shows:

• Which electrons come from each atom.
• Which electrons are shared.

Dots (•) represent electrons from one atom.

Crosses (×) represent electrons from the other atom.

The shared pair contains:

• One dot.
• One cross.

Dot-and-cross diagrams help illustrate exactly how covalent bonds form.

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Figure 3. Dot-and-cross diagrams show how atoms share electrons to form covalent bonds.

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How Many Bonds Do Common Atoms Form?

Different non-metals usually form a predictable number of covalent bonds.

For example:

Carbon has four valence electrons.

It usually shares four pairs of electrons, allowing it to form four covalent bonds.

This ability allows carbon to form millions of different compounds.

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Single, Double, and Triple Bonds

Atoms may share:

• One pair of electrons.
• Two pairs of electrons.
• Three pairs of electrons.

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Single Bond

One shared pair.

Example:

H—H

or

Cl—Cl

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Double Bond

Two shared pairs.

Example:

O=O

Each oxygen atom shares two pairs of electrons.

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Triple Bond

Three shared pairs.

Example:

N≡N

Each nitrogen atom shares three pairs of electrons.

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Figure 4. Covalent bonds may involve one, two, or three shared pairs of electrons.

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Properties of Covalent Substances

The type of bonding helps explain the properties of covalent substances.

Low Melting and Boiling Points

Many covalent substances consist of small molecules.

The forces between these molecules are relatively weak.

As a result, many have low melting and boiling points.

Examples:

• Oxygen
• Carbon dioxide
• Water (compared with ionic compounds)

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Poor Electrical Conductors

Most covalent substances do not conduct electricity because they contain:

• No free ions.
• No free-moving electrons.

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Can Exist as Solids, Liquids, or Gases

Examples:

Solid:

• Iodine

Liquid:

• Water

Gas:

• Oxygen
• Carbon dioxide

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Some Exceptions

Some giant covalent structures behave differently.

Examples:

• Diamond
• Graphite
• Silicon dioxide

These have:

• Very high melting points.

Graphite also conducts electricity because it contains mobile electrons.

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Figure 5. Most covalent substances are made of small molecules, but giant covalent structures such as diamond and graphite have very different properties.

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Ionic and Covalent Bonding Compared

Both types of bonding help atoms achieve stable electron arrangements, but they do so in different ways.

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Why Covalent Bonding Is Important

Covalent compounds are found everywhere.

Examples include:

• Water
• Oxygen
• Carbon dioxide
• Sugars
• Proteins
• DNA
• Plastics
• Medicines

Almost every molecule found in living organisms is held together by covalent bonds.

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Worked Example

Question

Explain how a water molecule (H₂O) forms.

Solution

1. Oxygen has 6 valence electrons and needs 2 more to complete its outer shell.
2. Each hydrogen atom has 1 electron and needs 1 more to fill its first shell.
3. Oxygen shares one pair of electrons with each hydrogen atom.
4. Two single covalent bonds are formed.
5. Oxygen now has 8 electrons in its outer shell, and each hydrogen has 2 electrons in its first shell.

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Real-World Connection

The molecules that make up living organisms are held together by covalent bonds. Water, DNA, proteins, carbohydrates, and lipids all depend on covalent bonding for their structure and function. Chemists also design new covalent compounds to produce medicines, plastics, synthetic fibres, fuels, and advanced materials used in electronics and aerospace engineering.

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Did You Know?

The human body contains trillions of water molecules, and each water molecule is held together by two covalent bonds between oxygen and hydrogen atoms. Without these covalent bonds, water—the substance essential for all known life—could not exist.

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Key Terms

Covalent bond – A chemical bond formed when two atoms share one or more pairs of electrons.

Double bond – A covalent bond in which two pairs of electrons are shared.

Dot-and-cross diagram – A diagram showing which electrons each atom contributes to a covalent bond.

Lewis structure – A diagram showing atoms, shared electron pairs (bonds), and lone pairs of electrons.

Lone pair – A pair of valence electrons that is not involved in bonding.

Molecule – A group of atoms held together by covalent bonds.

Octet rule – The tendency of atoms to achieve eight electrons in their outer shell through bonding.

Single bond – A covalent bond involving one shared pair of electrons.

Triple bond – A covalent bond involving three shared pairs of electrons.

Valence electron – An electron in the outermost shell of an atom that participates in chemical bonding.

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Key Takeaways

• Covalent bonding occurs when non-metal atoms share electrons to achieve stable outer electron shells.
• Lewis structures and dot-and-cross diagrams show how electrons are shared in molecules.
• Different atoms typically form a predictable number of covalent bonds based on their valence electrons.
• Covalent bonds can be single, double, or triple, depending on the number of shared electron pairs.
• Most covalent substances have low melting and boiling points and do not conduct electricity, although giant covalent structures such as diamond and graphite are important exceptions.
• Covalent bonding is responsible for forming the molecules that make up living organisms and countless everyday materials.