1. What Is an Element?

Introduction

Everything around us is made of matter, from the air we breathe to the water we drink and the devices we use every day. Scientists have discovered that all matter is built from a relatively small number of basic substances called elements. These elements are the building blocks from which all compounds and mixtures are formed.

Understanding elements is one of the foundations of chemistry. By learning what an element is and how it differs from compounds and mixtures, we can better understand how substances are formed, how chemical reactions occur, and how matter is organised.

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What Is an Element?

An element is a pure substance made of only one type of atom.

Every atom in an element has the same number of protons.

Examples include:

• Hydrogen
• Oxygen
• Carbon
• Iron
• Gold
• Copper

Although atoms of the same element may have different numbers of neutrons (forming isotopes), they are still the same element because they all contain the same number of protons.

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Figure 1. An element is made entirely of one type of atom.

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Elements Are Pure Substances

A pure substance contains only one type of particle.

For an element:

• Every atom is the same type.
• The composition is always the same.
• The properties are consistent.

For example:

A piece of pure copper contains only copper atoms.

A sample of pure helium contains only helium atoms.

Pure substances have fixed properties such as:

• Melting point
• Boiling point
• Density

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Why Can't Elements Be Broken Down Chemically?

Elements are the simplest chemical substances.

They cannot be broken down into simpler substances by chemical reactions.

For example:

Water can be broken down into:

• Hydrogen
• Oxygen

However:

Hydrogen cannot be chemically broken down into anything simpler because it is already an element.

Only nuclear reactions can change one element into another by altering the number of protons in the nucleus.

Ordinary chemical reactions do not change one element into another.

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Figure 2. Compounds can be broken down chemically into elements, but elements cannot be broken down further by chemical means.

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Elements, Compounds, and Mixtures

Although these terms are related, they describe different kinds of matter.

Element	Compound	Mixture
One type of atom	Two or more elements chemically bonded.	Two or more substances physically mixed
Pure substance	Pure substance	Not a pure substance
Cannot be chemically broken down.	Can be chemically broken down	Can be separated by physical methods
Example: Gold	Example: Water	Example: Air

Understanding these differences is essential in chemistry.

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Figure 3. Elements, compounds, and mixtures differ in the types of particles they contain.

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Element Names and Symbols

Each element has:

• A unique name.
• A unique chemical symbol.

Symbols contain one or two letters.

The first letter is always capitalised.

Examples:

Element	Symbol
Hydrogen	H
Oxygen	O
Carbon	C
Nitrogen	N
Iron	Fe
Copper	Cu
Gold	Au
Silver	Ag

Some symbols come from Latin names.

For example:

• Fe comes from ferrum (iron).
• Au comes from aurum (gold).
• Ag comes from argentum (silver).

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Atomic Structure and Elements

Every element is identified by its atomic number.

The atomic number equals the number of protons in the nucleus.

For example:

Element	Protons (Atomic Number)
Hydrogen.	1
Carbon	6
Oxygen	8
Sodium	11
Iron	26

No two different elements have the same number of protons.

This is what makes each element unique.

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Figure 4. Elements are identified by the number of protons in their atoms.

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Classifying Common Substances

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

Elements

Examples:

• Gold (Au)
• Copper (Cu)
• Oxygen (O₂)
• Nitrogen (N₂)
• Helium (He)

Although oxygen and nitrogen usually exist as molecules containing two atoms, each molecule contains only one type of atom, so they are still elements.

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Compounds

Examples:

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

These contain atoms of different elements chemically bonded together.

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Mixtures

Examples:

• Air
• Seawater
• Soil
• Salad
• Brass (an alloy)

These contain different substances that are physically mixed rather than chemically bonded.

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Why Elements Are Important

Elements are the building blocks of all matter.

They:

• Form compounds.
• Build living organisms.
• Make up rocks and minerals.
• Produce metals used in technology.
• Participate in every chemical reaction.

There are currently 118 known elements, and every substance around us is made from one or more of them.

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Figure 5. A wide variety of elements combine to form the materials found throughout our world.

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Worked Example

Question

Classify each substance as an element, compound, or mixture.

Substance	Classification
Oxygen gas (O₂)	Element
Water (H₂O)	Compound
Air	Mixture
Gold (Au)	Element
Salt water	Mixture
Carbon dioxide (CO₂).	Compound

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Real-World Connection

Many everyday products depend on the unique properties of different elements. Copper is used for electrical wiring because it conducts electricity well, aluminium is used in aircraft because it is strong and lightweight, and silicon is used to manufacture computer chips. Understanding the properties of elements allows scientists and engineers to design new technologies and materials.

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Did You Know?

Although scientists have discovered 118 elements, only about 90 occur naturally in significant amounts on Earth. The remaining elements have been created artificially in laboratories, and many exist for only fractions of a second before they decay into other elements.

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Key Terms

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

Atomic number – The number of protons in the nucleus of an atom.

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

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

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

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

Proton – A positively charged particle found in the nucleus of an atom.

Pure substance – A material containing only one type of particle with a fixed composition.

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Key Takeaways

• An element is a pure substance made of only one type of atom.
• Every atom of an element has the same number of protons, which is its atomic number.
• Elements cannot be broken down into simpler substances by ordinary chemical reactions.
• Each element has a unique name and chemical symbol.
• Elements differ from compounds and mixtures in the types of particles they contain.
• All matter is ultimately built from one or more of the 118 known chemical elements.

 

 

 

2. Chemical Symbols

Introduction

Imagine trying to write the full names of every element during a chemistry experiment. Instead of writing hydrogen, oxygen, carbon, and sodium repeatedly, chemists use short abbreviations called chemical symbols. These symbols make chemical formulas and equations much simpler to write and understand.

Chemical symbols are used by scientists all over the world, regardless of the language they speak. Whether a scientist is working in Canada, Brazil, Japan, or South Africa, the symbol O always represents oxygen and Fe always represents iron. This international system allows scientists to communicate clearly and accurately.

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What Is a Chemical Symbol?

A chemical symbol is a one- or two-letter abbreviation used to represent an element.

Every element has its own unique symbol.

Examples:

Element	Symbol
Hydrogen.	H
Oxygen	O
Carbon	C
Nitrogen	N
Helium	He
Sodium	Na
Iron	Fe

A chemical symbol represents:

• One atom of that element.
• The element itself.

For example:

• H represents hydrogen.
• O represents oxygen.
• Cu represents copper.

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Figure 1. Every element has a unique chemical symbol that is recognised worldwide.

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Why Do Chemists Use Symbols?

Chemical symbols make chemistry:

• Faster to write.
• Easier to read.
• More accurate.
• Internationally understood.

For example:

Instead of writing:

Hydrogen reacts with oxygen to produce water.

Chemists write:

2H₂ + O₂ → 2H₂O

Symbols allow complex chemical reactions to be written using only a few characters.

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Rules for Writing Chemical Symbols

Chemical symbols follow strict rules.

One-Letter Symbols

If the symbol contains one letter:

• Always use a capital letter.

Examples:

• H
• O
• C
• N
• F

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Two-Letter Symbols

If the symbol contains two letters:

• The first letter is capitalised.
• The second letter is lowercase.

Examples:

• He
• Li
• Na
• Fe
• Cu
• Ag

Incorrect examples:

• HE
• he
• FE
• cu

Correct capitalisation is essential because different capitalisation can change the meaning.

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Figure 2. Chemical symbols follow strict rules for capital and lowercase letters.

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Where Do Chemical Symbols Come From?

Many symbols are based on the English names of elements.

Examples:

Element	Symbol
Carbon	C
Oxygen	O
Nitrogen	N
Hydrogen.	H

However, some symbols come from Latin names.

Examples:

Element	Latin Name	Symbol
Sodium	Natrium	Na
Potassium.	Kalium	K
Iron	Ferrum	Fe
Copper	Cuprum	Cu
Silver	Argentum	Ag
Gold	Aurum	Au
Mercury	Hydrargyrum.	Hg
Lead	Plumbum	Pb

Learning these symbols takes practice, but they become familiar with regular use.

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Symbols and Atomic Numbers

Every element has:

• A name.
• A symbol.
• An atomic number.

The atomic number equals the number of protons in the atom.

Examples:

Element	Symbol.	Atomic Number
Hydrogen	H	1
Helium	He	2
Carbon	C	6
Oxygen	O	8
Sodium	Na	11
Iron	Fe	26

The symbol and atomic number together uniquely identify an element.

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Figure 3. Each element has a unique symbol and atomic number on the periodic table.

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Using Symbols in Chemical Formulas

Chemical symbols combine to form chemical formulas.

Examples:

Formula.	Meaning
H₂O	Water
CO₂	Carbon dioxide
NaCl	Sodium chloride
NH₃	Ammonia
CH₄	Methane

The subscript (small number) shows how many atoms of each element are present.

Example:

H₂O contains:

• 2 hydrogen atoms
• 1 oxygen atom

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Using Symbols in Chemical Equations

Chemical equations use symbols to describe chemical reactions.

Example:

Hydrogen + Oxygen → Water

Written using symbols:

2H₂ + O₂ → 2H₂O

Another example:

Magnesium + Oxygen → Magnesium oxide

2Mg + O₂ → 2MgO

Chemical symbols make reactions much easier to represent.

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Figure 4. Chemical symbols are combined to write formulas and chemical equations.

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Common Chemical Symbols to Know

Students should become familiar with the symbols of common elements.

Element	Symbol
Hydrogen	H
Helium	He
Carbon	C
Nitrogen	N
Oxygen	O
Sodium	Na
Magnesium	Mg
Aluminium	Al
Silicon	Si
Phosphorus	P
Sulfur	S
Chlorine	Cl
Potassium	K
Calcium	Ca
Iron	Fe
Copper	Cu
Zinc	Zn
Silver	Ag
Gold	Au
Mercury	Hg

Learning these symbols makes reading formulas and equations much easier.

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Figure 5. Familiarity with common chemical symbols is essential for studying chemistry.

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Worked Example

Question

Complete the table.

Element	Symbol.	Atomic Number
Oxygen	O	8
Sodium	Na	11
Iron	Fe	26
Carbon	C	6
Magnesium	Mg	12

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Real-World Connection

Chemical symbols appear everywhere in modern life. Nutrition labels list minerals such as Ca (calcium) and Fe (iron), batteries contain Li (lithium), electrical wiring is made from Cu (copper), and jewellery often contains Ag (silver) or Au (gold). Understanding these symbols helps people interpret scientific information, product labels, and medical reports.

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Did You Know?

The chemical symbol for tungsten is W, even though the element's English name begins with T. This is because the symbol comes from its historical name Wolfram, which is still used in several languages. This is one reason why some chemical symbols may seem unusual at first.

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Key Terms

Atomic number – The number of protons in the nucleus of an atom.

Chemical equation – A symbolic representation of a chemical reaction using chemical formulas.

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

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

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

Periodic table – A chart that organises all known chemical elements according to their atomic number and properties.

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

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Key Takeaways

• Every element has a unique chemical symbol used by scientists around the world.
• Chemical symbols make formulas and equations easier to write and understand.
• One-letter symbols use a capital letter, while two-letter symbols use a capital first letter and a lowercase second letter.
• Some symbols come from English names, while others come from Latin names such as Fe (ferrum) for iron and Na (natrium) for sodium.
• Chemical symbols are combined to write chemical formulas and chemical equations.
• Learning common chemical symbols is an essential skill for studying chemistry.

 

 

 

3. Metals and Non-Metals

Introduction

The periodic table contains over one hundred elements, but these elements are not all alike. Most can be grouped into two broad categories: metals and non-metals. Although both are made of atoms, they have very different physical and chemical properties.

Understanding the differences between metals and non-metals helps explain why different materials are used for different purposes. Copper is used for electrical wiring because it conducts electricity well, while oxygen is essential for respiration because it exists as a non-metal gas. By studying their properties, scientists can predict how elements will behave and choose suitable materials for specific applications.

What Are Metals and Non-Metals?

Elements can be divided into three broad groups:

• Metals
• Non-metals
• Metalloids (elements with properties of both)

Most elements on the periodic table are metals.

Only a relatively small number are non-metals.

 

 

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Figure 1. Most elements are metals, while non-metals are mainly found on the upper right side of the periodic table.

Location on the Periodic Table

The periodic table helps identify whether an element is a metal or non-metal.

Metals

Metals are found:

• On the left side.
• In the centre of the periodic table.

Examples include:

• Sodium (Na)
• Magnesium (Mg)
• Aluminium (Al)
• Iron (Fe)
• Copper (Cu)
• Gold (Au)

Non-Metals

Non-metals are found:

• On the upper right side of the periodic table.

Examples include:

• Hydrogen (H)
• Carbon (C)
• Nitrogen (N)
• Oxygen (O)
• Sulfur (S)
• Chlorine (Cl)

Metalloids

Between the metals and non-metals is a zigzag "staircase."

Elements along this line are called metalloids.

Examples include:

• Silicon (Si)
• Boron (B)
• Germanium (Ge)

Metalloids have properties of both metals and non-metals.

 

 

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Figure 2. The zigzag line separates most metals from non-metals, with metalloids lying along the boundary.

Physical Properties of Metals

Most metals share several common properties.

Good Conductors

Metals conduct:

• Electricity
• Heat

This is why copper and aluminium are widely used in electrical wiring.

Shiny (Lustrous)

Freshly cut metals usually have a shiny surface called metallic lustre.

Malleable

Malleable metals can be hammered into thin sheets without breaking.

Example:

Aluminium foil.

Ductile

Ductile metals can be drawn into wires.

Example:

Copper electrical wire.

High Melting Points

Many metals melt only at high temperatures.

Examples:

• Iron
• Tungsten

However, there are exceptions, such as mercury, which is liquid at room temperature.

Strong and Dense

Many metals are:

• Strong
• Hard
• Dense

This makes them useful in construction and engineering.

 

 

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Figure 3. Metals are generally good conductors, shiny, malleable, and ductile.

Physical Properties of Non-Metals

Non-metals have very different properties.

Most non-metals are:

• Poor conductors of heat.
• Poor conductors of electricity.
• Dull in appearance.
• Brittle if solid.
• Lower density than metals.

Many non-metals are gases at room temperature.

Examples:

• Oxygen
• Nitrogen
• Chlorine

One non-metal, bromine, is a liquid at room temperature.

Comparing Metals and Non-Metals

These are general trends, although there are exceptions.

 

 

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Figure 4. Metals and non-metals differ in several important physical properties.

Chemical Properties

Metals and non-metals also behave differently in chemical reactions.

Metals

Metals often:

• Lose electrons to form positive ions.
• React with oxygen to form metal oxides.
• React with acids to produce hydrogen gas.
• React with water (some metals only).

Non-Metals

Non-metals often:

• Gain electrons to form negative ions.
• Form acidic oxides with oxygen.
• React by sharing electrons in covalent bonds.

These differences help determine the types of compounds they form.

Properties and Everyday Uses

The properties of elements determine how they are used.

Metals

Non-Metals

Choosing materials based on their properties is an important part of engineering and technology.

 

 

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Figure 5. The unique properties of metals and non-metals make them suitable for different everyday uses.

Classifying Unfamiliar Elements

Scientists can often identify whether an unknown element is a metal or a non-metal by examining its properties.

For example:

Scientists use evidence rather than memorisation to classify elements.

Worked Example

Question

An unknown element has the following properties:

• Shiny
• Conducts electricity
• Can be hammered into thin sheets
• High melting point

Would it most likely be a metal or a non-metal?

Solution

Answer: Metal

Explanation:

These are all typical properties of metals.

Real-World Connection

Modern technology depends on both metals and non-metals. Smartphones contain metals such as copper, gold, and lithium, which conduct electricity and store energy, while non-metals such as silicon (a metalloid) are used to make computer chips. Engineers choose materials carefully based on their physical and chemical properties to produce safe, reliable, and efficient devices.

Did You Know?

Mercury (Hg) is the only metal that is liquid at room temperature, while bromine (Br) is the only non-metal that is liquid at room temperature. Most other metals are solids, and many non-metals are gases under the same conditions.

Key Terms

Brittle – Easily broken or shattered when struck.

Conductor – A material that allows heat or electricity to pass through easily.

Ductile – Able to be drawn into wires.

Lustre – The shiny appearance of a material.

Malleable – Able to be hammered or rolled into thin sheets without breaking.

Metalloid – An element with properties intermediate between those of metals and non-metals.

Metal – An element that is usually shiny, conducts heat and electricity, and is malleable and ductile.

Non-metal – An element that is generally a poor conductor of heat and electricity and is often dull and brittle if solid.

Periodic table – A chart that organises elements according to their atomic number and properties.

Key Takeaways

• Most elements are classified as metals or non-metals, with a small group of metalloids between them.
• Metals are generally shiny, good conductors, malleable, ductile, and have high melting points.
• Non-metals are generally poor conductors, often dull, brittle if solid, and many exist as gases at room temperature.
• The location of an element on the periodic table provides clues about whether it is a metal or a non-metal.
• The physical and chemical properties of elements determine their uses in everyday life.
• Scientists classify unfamiliar elements by examining their measurable properties rather than relying only on their names or positions on the periodic table.

4. Introduction to the Periodic Table

Introduction

The periodic table is one of the most important tools in chemistry. It organises all known chemical elements into a logical pattern that allows scientists to understand their properties, compare them, and predict how they will behave in chemical reactions. Although it may appear to be simply a chart of symbols and numbers, the periodic table contains a huge amount of scientific information.

The modern periodic table includes 118 known elements, ranging from hydrogen, the lightest element, to oganesson, one of the heaviest. By learning how to read the periodic table, students can quickly identify an element's name, symbol, atomic number, and many of its physical and chemical properties.

What Is the Periodic Table?

The periodic table is a chart that organises all known chemical elements according to their atomic number.

Each box on the table represents one element.

The periodic table allows scientists to:

• Organise elements.
• Compare properties.
• Predict chemical behaviour.
• Identify patterns among elements.

Because elements with similar properties are grouped together, the table is a powerful tool for studying chemistry.

 

 

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Figure 1. The periodic table organises all known elements according to their atomic number and chemical properties.

How Is the Periodic Table Organised?

Elements are arranged in order of increasing atomic number.

The atomic number is the number of protons in the nucleus of an atom.

For example:

Each element has one more proton than the element before it.

No two elements have the same atomic number.

Information Found in an Element Box

Each element has its own box on the periodic table.

A typical element box contains:

• Atomic number
• Chemical symbol
• Element name
• Relative atomic mass (atomic weight)

For example:

Different versions of the periodic table may include additional information such as electron configuration or oxidation states.

 

 

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Figure 2. Each element box contains important information used by chemists.

Periods

The horizontal rows of the periodic table are called periods.

The modern periodic table has 7 periods.

Elements within the same period:

• Have the same number of occupied electron shells.
• Show gradual changes in their properties as you move from left to right.

For example:

Period 2 contains:

• Lithium
• Beryllium
• Boron
• Carbon
• Nitrogen
• Oxygen
• Fluorine
• Neon

As you move across a period, the properties of the elements change gradually.

Groups

The vertical columns are called groups.

The periodic table has 18 groups.

Elements within the same group often have:

• Similar chemical properties.
• Similar numbers of outer (valence) electrons.
• Similar reactions.

Examples include:

• Group 1 – Alkali metals
• Group 2 – Alkaline earth metals
• Group 17 – Halogens
• Group 18 – Noble gases

Members of the same group often react in similar ways because they have similar electron arrangements.

 

 

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Figure 3. Periods are horizontal rows, while groups are vertical columns on the periodic table.

Locating Elements

Elements can be located using:

• Their name
• Their symbol
• Their atomic number

For example:

Scientists often identify elements by their symbols because they are faster to read and write.

Predicting Properties

The periodic table allows scientists to predict many properties of an element.

For example:

An element found:

• On the left side is probably a metal.
• On the upper right side is probably a non-metal.
• In Group 18 is likely to be very unreactive.
• In Group 1 is likely to be highly reactive.

Elements in the same group often have similar:

• Chemical reactions.
• Physical properties.
• Electron arrangements.

This is one of the reasons the periodic table is such a powerful scientific tool.

 

 

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Figure 4. The position of an element on the periodic table helps predict its properties.

Why Is It Called the "Periodic" Table?

The word periodic means repeating at regular intervals.

As elements are arranged by increasing atomic number, many of their physical and chemical properties repeat in a regular pattern.

Examples include:

• Reactivity
• Melting point
• Atomic size
• Electrical conductivity

This repeating pattern is called the Periodic Law.

The periodic table is organised to display these repeating trends clearly.

Why the Periodic Table Is Important

Scientists use the periodic table to:

• Identify elements.
• Predict chemical behaviour.
• Understand chemical reactions.
• Compare physical properties.
• Discover new elements.

The periodic table is often called the "map of chemistry" because it organises all known elements into one logical system.

 

 

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Figure 5. Chemists use the periodic table to identify elements and predict their behaviour in scientific investigations.

Worked Example

Question

Use the periodic table to answer the following.

Real-World Connection

The periodic table is used by scientists, engineers, doctors, pharmacists, and environmental scientists every day. It helps doctors understand the role of minerals such as calcium and iron in the human body, enables engineers to choose suitable metals for buildings and aircraft, and allows chemists to develop new medicines, batteries, and advanced materials.

Did You Know?

When the Russian chemist Dmitri Mendeleev created the first widely accepted periodic table in 1869, he deliberately left empty spaces for elements that had not yet been discovered. He even predicted many of their properties with remarkable accuracy. As those elements were later discovered, they closely matched his predictions, providing strong evidence that the periodic table was organised correctly.

Key Terms

Atomic number – The number of protons in the nucleus of an atom.

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

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

Group – A vertical column of elements in the periodic table whose members have similar chemical properties.

Periodic table – A chart that organises all known elements according to increasing atomic number and recurring chemical properties.

Period – A horizontal row of elements in the periodic table.

Periodic Law – The principle that the physical and chemical properties of elements repeat in a regular pattern when elements are arranged by increasing atomic number.

Relative atomic mass – The average mass of an atom of an element compared with one-twelfth of the mass of a carbon-12 atom.

Key Takeaways

• The periodic table is an organised arrangement of all known chemical elements.
• Elements are arranged by increasing atomic number, which is the number of protons in each atom.
• Periods are horizontal rows, while groups are vertical columns.
• Each element box contains important information such as the element's name, symbol, atomic number, and relative atomic mass.
• The position of an element on the periodic table helps predict whether it is a metal or non-metal and provides clues about its chemical properties.
• The periodic table is an essential tool that allows scientists to organise elements and predict their behaviour.

5. Atomic and Ionic Diagrams

Introduction

Atoms are far too small to see, even with the most powerful light microscopes. To understand their structure and how they form chemical bonds, scientists use a variety of atomic diagrams. These simplified models help us visualise the arrangement of electrons, predict how atoms behave, and explain how compounds are formed.

Three of the most common diagrams used in chemistry are Bohr diagrams, Lewis electron-dot diagrams, and dot-and-cross diagrams. Each diagram highlights different aspects of atomic structure and chemical bonding. Learning how to draw and interpret these diagrams provides a foundation for understanding ionic and covalent bonding.

Bohr Diagrams

A Bohr diagram is a simple model that shows:

• The nucleus
• Protons
• Neutrons
• Electrons arranged in shells (energy levels)

Electrons occupy shells around the nucleus.

For the first 20 elements, the maximum number of electrons in the first few shells is commonly represented as:

• First shell: 2 electrons
• Second shell: 8 electrons
• Third shell: 8 electrons (for introductory chemistry)
• Fourth shell: Remaining electrons

Bohr diagrams are useful for showing how electrons are arranged around the nucleus.

 

 

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Figure 1. Bohr diagrams show electrons arranged in shells around the nucleus.

Drawing a Bohr Diagram

To draw a Bohr diagram:

Step 1

Find the atomic number.

This tells you the number of:

• Protons
• Electrons (for a neutral atom)

Step 2

Draw the nucleus.

Label:

• Number of protons
• Number of neutrons (if known)

Step 3

Place electrons into shells.

Example:

Carbon (Atomic Number = 6)

• First shell = 2 electrons
• Second shell = 4 electrons

Electron arrangement:

2,4

More Examples

Electron Shells

Electrons occupy different energy levels, often called electron shells.

Important ideas:

• Electrons fill the inner shells first.
• Outer-shell electrons are called valence electrons.
• Valence electrons determine how atoms react chemically.

For example:

Atoms with full outer shells are generally very stable.

 

 

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Figure 2. The number of valence electrons determines how atoms react and form chemical bonds.

Forming Ions

Atoms become more stable by gaining or losing electrons.

When this happens, they form ions.

Positive Ions (Cations)

Metals usually:

• Lose electrons.
• Form positive ions.

Example:

Sodium

Na → Na⁺ + e⁻

Sodium loses one electron and becomes Na⁺.

Negative Ions (Anions)

Non-metals usually:

• Gain electrons.
• Form negative ions.

Example:

Cl + e⁻ → Cl⁻

Chlorine gains one electron and becomes Cl⁻.

The number of protons does not change—only the number of electrons changes.

 

 

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Figure 3. Metals lose electrons to form positive ions, while non-metals gain electrons to form negative ions.

Lewis Electron-Dot Diagrams

A Lewis electron-dot diagram shows:

• The element symbol.
• The valence electrons only.

The nucleus and inner electrons are not shown.

Examples:

Hydrogen

H•

Carbon

  •
• C •
  •

Oxygen

 ••
• O •
 ••

Chlorine

 ••
••Cl••
  •

Lewis diagrams help predict how atoms will bond.

 

 

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Figure 4. Lewis diagrams show only the valence electrons that participate in chemical bonding.

Lewis Diagrams for Ions

When atoms become ions, the Lewis diagram changes.

Sodium Ion

Na⁺

• No valence dots shown.
• Written inside brackets.

[Na]⁺

Chloride Ion

Chlorine gains one electron.

  ••
••Cl••
  ••

Written as:

[Cl]⁻

with eight electrons around the symbol.

This represents a complete outer shell.

Dot-and-Cross Diagrams

A dot-and-cross diagram shows:

• Which electrons belong to each atom.
• How electrons are transferred or shared.

Different symbols are used:

• Dots (•)
• Crosses (×)

This makes it easy to identify the origin of each electron.

Ionic Bond Example

Sodium chloride (NaCl)

Sodium transfers one electron to chlorine.

The diagram shows:

• Sodium losing one electron.
• Chlorine gaining that electron.
• Both ions achieving full outer shells.

Covalent Bond Example

Hydrogen molecule (H₂)

Each hydrogen contributes one electron.

The shared pair forms a covalent bond.

Dot-and-cross diagrams clearly show this sharing.

 

 

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Figure 5. Dot-and-cross diagrams show how electrons are transferred in ionic bonds and shared in covalent bonds.

Comparing Atomic Diagrams

Each diagram highlights different information about atoms and chemical bonding.

Why These Diagrams Are Important

Atomic diagrams help scientists:

• Understand electron arrangement.
• Predict chemical reactions.
• Explain ionic and covalent bonding.
• Understand the periodic table.
• Design new materials and compounds.

Although these diagrams are simplified models, they remain valuable tools for learning chemistry.

Worked Example

Question

Draw the Bohr electron arrangement and Lewis diagram for magnesium (Atomic Number = 12).

Solution

Bohr Diagram

Electron arrangement:

2,8,2

Magnesium has:

• 12 protons
• 12 electrons
• 2 valence electrons

Lewis Diagram

• Mg •

Two dots represent magnesium's two valence electrons.

Real-World Connection

Chemists use electron diagrams to understand how atoms bond together to form the substances we use every day. Whether designing stronger alloys, developing new medicines, producing batteries, or creating advanced electronic materials, understanding electron arrangements is essential. Engineers also rely on these models when developing semiconductors, solar cells, and rechargeable batteries.

Did You Know?

The Bohr model, proposed by Niels Bohr in 1913, was a major breakthrough in understanding atomic structure. Although modern quantum mechanics provides a more accurate description of electrons, Bohr diagrams are still widely used in schools because they clearly illustrate electron shells and help explain chemical bonding.

Key Terms

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

Bohr diagram – A model showing electrons arranged in shells around the nucleus.

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

Dot-and-cross diagram – A diagram showing how electrons are transferred or shared during bonding using dots and crosses to distinguish electrons from different atoms.

Electron shell – An energy level around the nucleus where electrons are found.

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

Lewis electron-dot diagram – A diagram showing an element's symbol surrounded by its valence electrons.

Valence electrons – The electrons in the outermost shell of an atom that are involved in chemical bonding.

Key Takeaways

• Bohr diagrams show the nucleus and the arrangement of electrons in shells.
• Electrons fill the inner shells first, and the outermost valence electrons determine how atoms react.
• Atoms become ions by gaining or losing electrons to achieve more stable electron arrangements.
• Lewis electron-dot diagrams show only the valence electrons involved in bonding.
• Dot-and-cross diagrams illustrate how electrons are transferred in ionic bonds or shared in covalent bonds.
• These diagrams are important tools for understanding atomic structure and predicting how elements form compounds.