1. Periods and Groups

Periods and Groups

The periodic table is arranged in a way that helps scientists understand and predict the properties of elements. An element’s position is not random. The row and column in which an element appears provide important information about its atomic structure and chemical behaviour.

Periods

The horizontal rows of the periodic table are called periods.

There are 7 periods on the modern periodic table.

• Period 1 contains hydrogen (H) and helium (He).
• Period 2 begins with lithium (Li) and ends with neon (Ne).
• Period 3 begins with sodium (Na) and ends with argon (Ar).
• Periods generally become longer as you move down the periodic table.

The period number tells us how many occupied electron shells an atom has.

For example:

Therefore, sodium is in Period 3 because its electrons occupy three electron shells.

Periods and Atomic Structure

As you move from left to right across a period, the number of protons and electrons increases.

For example, consider part of Period 3:

Na → Mg → Al → Si → P → S → Cl → Ar

Each element has one more proton than the element before it. For a neutral atom, it also has one more electron.

However, all of these elements have the same number of occupied electron shells: three.

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Groups

The vertical columns of the periodic table are called groups.

There are 18 groups in the modern periodic table.

Elements in the same group often have similar chemical properties because they have similar arrangements of valence electrons.

Valence electrons are the electrons in the outermost occupied electron shell. These electrons are especially important because they are involved in chemical bonding and reactions.

For many main-group elements, the group can therefore help us predict the number of valence electrons.

For example:

*Helium has 2 electrons in its only shell; the other Group 18 elements typically have 8 valence electrons.

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Why Do Elements in the Same Group Behave Similarly?

Chemical reactions usually involve an atom's outer electrons.

Elements in the same group have similar outer-electron arrangements, so they tend to react in similar ways.

For example, lithium, sodium and potassium are all Group 1 metals.

They each have one valence electron:

• lithium: 2, 1
• sodium: 2, 8, 1
• potassium: 2, 8, 8, 1

Because each atom has one outer electron, these elements tend to lose one electron during chemical reactions.

This is why Group 1 elements share many properties and form ions with a +1 charge.

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Position and Atomic Structure

An element's position on the periodic table gives us clues about its atomic structure.

For many of the elements students encounter first:

Period → number of occupied electron shells

Group → information about the valence electrons

Consider magnesium (Mg).

Magnesium has the electron arrangement:

2, 8, 2

This tells us:

• it has 3 occupied shells, so it is in Period 3
• it has 2 valence electrons, so it is in Group 2

Now consider chlorine (Cl):

2, 8, 7

This tells us:

• it has 3 occupied shells, so it is in Period 3
• it has 7 valence electrons, so it is in Group 17

This relationship allows us to connect the periodic table directly to the structure of atoms.

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Using Groups to Predict Properties

Groups can also help us make predictions about unfamiliar elements.

Group 1 – Alkali Metals

Examples include lithium, sodium and potassium.

They:

• are metals
• have 1 valence electron
• usually lose one electron
• form +1 ions
• are chemically reactive

Group 17 – Halogens

Examples include fluorine, chlorine and bromine.

They:

• are non-metals
• have 7 valence electrons
• usually gain one electron
• form −1 ions
• are reactive

Group 18 – Noble Gases

Examples include helium, neon and argon.

They:

• have full outer electron shells
• are very stable
• are generally very unreactive

Because elements within a group have related electron arrangements, knowing the group allows scientists to make useful predictions about an element's behaviour.

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Locating an Element

We can describe an element's position using both its period and group.

For example:

Calcium (Ca) is in Period 4, Group 2.

This immediately tells us that calcium:

• has four occupied electron shells
• has two valence electrons
• is a metal
• is likely to lose two electrons during reactions
• commonly forms a Ca²⁺ ion

Similarly, bromine (Br) is in Period 4, Group 17.

We can predict that bromine:

• has four occupied electron shells
• has seven valence electrons
• is a non-metal
• tends to gain one electron
• commonly forms a Br⁻ ion

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A Useful Way to Remember

Think of the periodic table like a large apartment building:

• Periods are the floors running horizontally.
• Groups are vertical stacks of apartments.
• Elements living in the same vertical group belong to the same "family" and tend to behave similarly.

So:

Across = Period

Down = Group

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Did You Know?

The periodic table is powerful because it can be used to predict properties of elements that have not yet been studied in detail.

In the 1800s, Dmitri Mendeleev deliberately left gaps in his periodic table for elements that had not yet been discovered. He predicted several of their properties based on the groups in which the missing elements should appear. Later discoveries closely matched many of his predictions.

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Key Terms

Period – A horizontal row on the periodic table.

Group – A vertical column on the periodic table.

Valence electron – An electron in the outermost occupied electron shell.

Electron shell – A region around the nucleus occupied by electrons.

Chemical property – A characteristic describing how a substance behaves during chemical reactions.

Periodic table – An arrangement of the elements according to atomic number and repeating patterns in their properties.

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Key Takeaways

• The horizontal rows of the periodic table are called periods.
• The vertical columns are called groups.
• The period number indicates the number of occupied electron shells.
• Elements in the same group have similar outer-electron arrangements.
• Similar valence-electron arrangements cause elements in the same group to have similar chemical properties.
• An element's position can therefore provide information about its atomic structure.
• Periods and groups can be used to locate elements and predict basic chemical properties.

2. Atomic Number and Position

Atomic Number and Position

Every element has its own unique atomic number. The atomic number is one of the most important pieces of information about an element because it tells us the number of protons in its nucleus and determines where the element appears on the periodic table.

What Is Atomic Number?

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

It is usually represented by the symbol Z.

For example:

Therefore:

Atomic number = Number of protons

So an oxygen atom has atomic number 8 because every oxygen atom contains 8 protons.

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Atomic Number Identifies the Element

The number of protons determines the identity of an element.

Every atom with 6 protons is carbon.

Every atom with 8 protons is oxygen.

Every atom with 17 protons is chlorine.

If the number of protons changes, the atom becomes a different element.

For example:

11 protons → Sodium (Na)

12 protons → Magnesium (Mg)

13 protons → Aluminium (Al)

This means that no two different elements can have the same atomic number.

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Atomic Number on the Periodic Table

Each element's atomic number is shown in its box on the periodic table.

A typical element box might look like this:

11
Na
Sodium
22.99

Here:

• 11 is the atomic number.
• Na is the chemical symbol.
• Sodium is the element name.
• 22.99 is the relative atomic mass.

Be careful not to confuse atomic number with relative atomic mass.

Suggested placement: Place an enlarged element square here so students can practise identifying the atomic number, symbol, element name and relative atomic mass.

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How Is the Periodic Table Arranged?

The modern periodic table is arranged in order of increasing atomic number.

Starting from hydrogen:

H → He → Li → Be → B → C → N → O → F → Ne

Their atomic numbers increase:

1 → 2 → 3 → 4 → 5 → 6 → 7 → 8 → 9 → 10

The pattern continues throughout the periodic table.

This means that as you move through the periodic table from one element to the next, the atomic number normally increases by one.

For example:

Atomic number therefore gives every element a specific position in the periodic table.

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Why Is the Periodic Table Arranged This Way?

Arranging elements by atomic number creates repeating patterns in their electron arrangements and chemical properties.

When one period is completed, a new electron shell begins to fill and a new period starts.

This produces the repeating pattern that gives the periodic table its name.

Elements with similar outer-electron arrangements appear in the same groups, which is why elements in the same group often have similar chemical properties.

Atomic number therefore does more than simply number the elements. It provides the foundation for the entire organisation of the modern periodic table.

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Atomic Number and Protons

The relationship between atomic number and protons is always:

Atomic number = Number of protons

Consider chlorine.

Chlorine has atomic number 17.

Therefore:

Atomic number = 17

Number of protons = 17

This is true for every chlorine atom. An atom cannot be chlorine unless it contains 17 protons.

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Atomic Number and Electrons

Atoms also contain negatively charged electrons.

A proton has a charge of +1, while an electron has a charge of −1.

In a neutral atom, the positive and negative charges balance.

Therefore:

Number of electrons = Number of protons

For a neutral atom:

Atomic number = Protons = Electrons

For example, magnesium has atomic number 12.

A neutral magnesium atom therefore contains:

• 12 protons
• 12 electrons

The positive charge from the 12 protons is balanced by the negative charge from the 12 electrons.

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Using Atomic Number to Predict Atomic Structure

If we know the atomic number, we can immediately determine the number of protons and, for a neutral atom, the number of electrons.

Example 1: Carbon

Carbon has atomic number 6.

Therefore:

• Protons = 6
• Electrons = 6

Example 2: Aluminium

Aluminium has atomic number 13.

Therefore:

• Protons = 13
• Electrons = 13

Example 3: Calcium

Calcium has atomic number 20.

Therefore:

• Protons = 20
• Electrons = 20

This gives us a useful rule:

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From Atomic Number to Position

Because the periodic table follows atomic number, we can also use atomic number to locate an element.

Suppose an unknown element has atomic number 16.

We can find atomic number 16 on the periodic table and identify the element as sulfur (S).

Its position tells us that sulfur is:

• in Period 3
• in Group 16
• a non-metal

Similarly, atomic number 19 identifies potassium (K), which is found in Period 4 and Group 1.

The atomic number therefore connects an element's identity, atomic structure and position on the periodic table.

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What About Neutrons?

Atomic number tells us the number of protons, but it does not directly tell us the number of neutrons.

For example, carbon atoms always have 6 protons, but different carbon atoms can contain different numbers of neutrons.

These different forms of the same element are called isotopes.

The number of protons remains the same because changing the number of protons would change the identity of the element.

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Did You Know?

The earliest periodic tables were not arranged using atomic number because scientists did not yet understand the structure of the atom.

In 1913, physicist Henry Moseley showed that elements could be organised according to their nuclear charge, corresponding to their atomic number. This helped establish the organisation used by the modern periodic table.

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Key Terms

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

Proton – A positively charged particle found in the nucleus.

Electron – A negatively charged particle found outside the nucleus.

Neutral atom – An atom containing equal numbers of protons and electrons.

Nucleus – The central region of an atom containing protons and neutrons.

Element – A substance made of atoms that all contain the same number of protons.

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Key Takeaways

• The atomic number is equal to the number of protons in an atom.
• Every element has a unique atomic number.
• The number of protons determines the identity of an element.
• The modern periodic table is arranged in order of increasing atomic number.
• Atomic number determines where an element appears on the periodic table.
• In a neutral atom, the number of electrons equals the number of protons.
• Therefore, for a neutral atom: Atomic number = Number of protons = Number of electrons
• Atomic number can be used to identify an element and predict basic information about its atomic structure.

3. Metals, Non-Metals, and Metalloids

Metals, Non-Metals, and Metalloids

The elements on the periodic table can be divided into three broad categories: metals, non-metals, and metalloids.

An element's location on the periodic table provides useful clues about its properties. These properties also help determine how different elements are used in everyday life, technology, construction, medicine, and industry.

Where Are Metals, Non-Metals, and Metalloids?

The three categories occupy different regions of the periodic table.

• Metals are found mainly on the left and centre.
• Non-metals are found mainly on the right side.
• Metalloids are found along the zigzag or staircase line separating metals and non-metals.

There is one important exception: hydrogen is located on the left side of the periodic table but is a non-metal.

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Metals

Most elements on the periodic table are metals.

Common examples include:

• iron (Fe)
• copper (Cu)
• aluminium (Al)
• sodium (Na)
• calcium (Ca)
• gold (Au)
• silver (Ag)

Although metals do not all behave exactly the same way, they share several characteristic physical properties.

Physical Properties of Metals

Metals are generally:

• shiny when freshly cut or polished
• good conductors of electricity
• good conductors of thermal energy
• malleable – they can be hammered or pressed into different shapes
• ductile – they can be drawn into wires
• strong and dense
• solid at room temperature

There are exceptions. For example, mercury (Hg) is a metal but is liquid at room temperature.

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Chemical Properties of Metals

During chemical reactions, metal atoms often lose electrons.

When an atom loses electrons, it forms a positive ion, called a cation.

For example, sodium can lose one electron:

Na → Na⁺ + electron

Magnesium can lose two electrons:

Mg → Mg²⁺ + 2 electrons

Many metals also react with other substances.

For example:

Metal + oxygen → metal oxide

Some metals can also react with acids:

Metal + acid → salt + hydrogen gas

However, metals vary greatly in their reactivity. Potassium and sodium are very reactive, while metals such as gold are much less reactive.

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Non-Metals

Non-metals occupy a much smaller region of the periodic table.

Examples include:

• hydrogen (H)
• carbon (C)
• nitrogen (N)
• oxygen (O)
• phosphorus (P)
• sulfur (S)
• chlorine (Cl)

Non-metals can exist as solids, liquids, or gases at room temperature.

For example:

• carbon is a solid
• bromine is a liquid
• oxygen is a gas

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Physical Properties of Non-Metals

Solid non-metals are generally:

• dull rather than shiny
• poor conductors of electricity
• poor conductors of thermal energy
• brittle rather than malleable
• not ductile
• often less dense than metals

Again, there are exceptions. For example, graphite, a form of carbon, can conduct electricity even though carbon is a non-metal.

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Chemical Properties of Non-Metals

Non-metal atoms often gain electrons or share electrons when they form chemical bonds.

For example, chlorine can gain one electron:

Cl + electron → Cl⁻

The chlorine atom becomes a negatively charged chloride ion.

Non-metals can also share electrons with other non-metals to form covalent bonds.

For example, hydrogen and oxygen share electrons when they form water molecules.

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Comparing Metals and Non-Metals

There are several important differences between the typical properties of metals and non-metals.

These are general patterns, not rules without exceptions.

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Metalloids

Between the metals and non-metals is a small group of elements called metalloids.

Metalloids have properties that are intermediate between metals and non-metals.

Commonly recognised metalloids include:

• boron (B)
• silicon (Si)
• germanium (Ge)
• arsenic (As)
• antimony (Sb)
• tellurium (Te)

The exact classification of a few elements near the staircase can vary between sources.

Properties of Metalloids

Metalloids may:

• have a metallic or shiny appearance
• be brittle like non-metals
• conduct electricity better than most non-metals
• conduct electricity less effectively than typical metals
• behave chemically like metals in some situations and non-metals in others

One particularly important property is their ability to act as semiconductors.

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What Is a Semiconductor?

A semiconductor is a material whose electrical conductivity lies between that of a conductor and an insulator.

This property can be carefully controlled, making semiconductors extremely useful in electronics.

Silicon is one of the most important examples.

Silicon is used to manufacture:

• computer chips
• smartphones
• electronic sensors
• calculators
• solar cells
• many other electronic devices

Suggested placement: Silicon wafers or microchips make a strong visual connection between metalloids and modern technology.

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Properties Determine Uses

Scientists and engineers select materials because of their properties.

The properties of an element determine whether it is suitable for a particular job.

Copper – Electrical Wiring

Copper is an excellent electrical conductor and can be drawn into thin wires.

Therefore, copper is widely used in:

• electrical cables
• motors
• electronics

Aluminium – Aircraft and Transport

Aluminium has a relatively low density and is resistant to corrosion.

It is therefore useful for:

• aircraft
• vehicles
• drink cans
• window frames

Iron – Construction

Iron is used to produce steel, which can be very strong.

Steel is widely used in:

• buildings
• bridges
• vehicles
• machinery

Silicon – Electronics

Silicon is a semiconductor, allowing its electrical behaviour to be carefully controlled.

It is therefore essential for many electronic components.

Oxygen – Medicine and Respiration

Oxygen is a reactive non-metal gas required for aerobic respiration.

It is used in:

• medical oxygen systems
• breathing equipment
• industrial processes

This gives us an important relationship:

Structure → Properties → Uses

The atomic structure of an element influences its properties, and those properties influence how the element can be used.

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Classifying an Unfamiliar Element

Even if we do not know the name of an element, we may be able to classify it from information about its properties.

Example 1

An unknown element:

• is shiny
• conducts electricity
• can be hammered into sheets
• forms positive ions

These are typical properties of a metal.

Example 2

Another element:

• is dull
• is brittle when solid
• does not conduct electricity well
• tends to gain electrons

These properties suggest a non-metal.

Example 3

A third element:

• is shiny
• is brittle
• conducts electricity moderately well
• is used in electronic components

This combination suggests a metalloid.

When classifying an unfamiliar element, look for several pieces of evidence rather than relying on only one property.

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Did You Know?

More than three-quarters of the known elements are metals.

However, some of the elements most important to living organisms are non-metals. Carbon, hydrogen, oxygen, nitrogen, phosphorus, and sulfur are all non-metals and are essential components of living things.

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Key Terms

Metal – An element that generally conducts heat and electricity well and tends to lose electrons during reactions.

Non-metal – An element that generally conducts heat and electricity poorly and tends to gain or share electrons.

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

Malleable – Able to be hammered or pressed into different shapes without breaking.

Ductile – Able to be drawn into a wire.

Brittle – Likely to break or shatter rather than bend.

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

Semiconductor – A material with electrical conductivity between that of a conductor and an insulator.

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Key Takeaways

• Metals are located mainly on the left and centre of the periodic table.
• Non-metals are located mainly on the right side, with hydrogen as an important exception.
• Metalloids occur along the staircase between metals and non-metals.
• Metals are generally shiny, malleable, ductile, and good conductors.
• Non-metals are generally poor conductors and are often brittle when solid.
• Metals often lose electrons, while non-metals often gain or share electrons.
• Metalloids have a mixture of metallic and non-metallic properties.
• Some metalloids, especially silicon, are important semiconductors.
• The properties of elements determine their practical uses.
• The properties and periodic-table position of an unfamiliar element can provide evidence for classifying it as a metal, non-metal, or metalloid.

4. Electron Arrangements and Group Number

Electron Arrangements and Group Number

The periodic table provides much more information than simply the names and symbols of elements. An element's position can help us predict how its electrons are arranged, how many valence electrons it has, and how it is likely to behave during chemical reactions.

For the first 20 elements, these relationships are especially useful because their electron arrangements can be represented using simple Bohr diagrams.

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Electron Shells

Electrons occupy regions around the nucleus called electron shells or energy levels.

For the first 20 elements, we can use a simplified model:

• first shell holds up to 2 electrons
• second shell holds up to 8 electrons
• third shell holds up to 8 electrons in this introductory model
• electrons then begin entering the fourth shell

For example, a sodium atom has 11 electrons.

Its electron arrangement is:

2, 8, 1

This means:

• 2 electrons in the first shell
• 8 electrons in the second shell
• 1 electron in the third shell

Suggested placement: A Bohr model of sodium works well here because students can clearly see the relationship between the notation 2,8,1 and the three occupied shells.

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Atomic Number and Electron Arrangement

The atomic number tells us the number of protons in an atom.

In a neutral atom:

Number of electrons = Number of protons = Atomic number

Therefore, we can use the atomic number to determine how many electrons need to be placed into the shells.

Example: Oxygen

Oxygen has atomic number 8.

Therefore, a neutral oxygen atom has:

• 8 protons
• 8 electrons

Fill the shells:

• first shell → 2 electrons
• second shell → remaining 6 electrons

Electron arrangement:

2, 6

Example: Calcium

Calcium has atomic number 20, so it has 20 electrons.

Its electron arrangement is:

2, 8, 8, 2

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Period Number and Electron Shells

The period number tells us the number of occupied electron shells.

For example:

Notice the pattern.

Sodium has:

2,8,1 → 3 occupied shells → Period 3

Potassium has:

2,8,8,1 → 4 occupied shells → Period 4

So we can remember:

Period number → Number of occupied electron shells

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Group Number and Valence Electrons

The electrons in the outermost occupied shell are called valence electrons.

These electrons are particularly important because they are involved in chemical bonding and reactions.

For the main-group elements, the position of an element provides information about its valence electrons.

For Groups 13–18, a useful introductory rule is:

Valence electrons = Group number − 10

For example:

Group 16 → 16 − 10 = 6 valence electrons

Group 17 → 17 − 10 = 7 valence electrons

This simple relationship applies to the main-group elements, not the transition metals in Groups 3–12.

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Why Do Elements in the Same Group Behave Similarly?

Elements in the same group have the same number of valence electrons.

Consider Group 1:

All three atoms have one valence electron.

Because chemical reactions mainly involve valence electrons, lithium, sodium, and potassium tend to behave in similar ways.

The same pattern occurs in Group 17:

• fluorine → 2,7
• chlorine → 2,8,7

Both have 7 valence electrons, so they have similar chemical properties.

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Stable Electron Arrangements

Atoms often become more stable by achieving a full outer electron shell.

For many elements, this means having 8 electrons in the outer shell.

This is sometimes called the octet rule.

The noble gases in Group 18 already have full outer shells.

For example:

• helium → 2
• neon → 2,8
• argon → 2,8,8

Because their outer shells are already full, noble gases are generally very unreactive.

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Electron Arrangements and Ion Formation

Atoms can become more stable by losing or gaining electrons.

When this happens, the atom becomes an electrically charged particle called an ion.

Metals Usually Lose Electrons

Sodium has the electron arrangement:

2,8,1

It has one valence electron.

Sodium can lose this electron:

2,8,1 → 2,8

It now has 10 electrons but still has 11 protons.

The result is a +1 ion:

Na⁺

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Non-Metals Usually Gain Electrons

Chlorine has the electron arrangement:

2,8,7

It needs one additional electron to complete its outer shell.

It can gain one electron:

2,8,7 → 2,8,8

The resulting ion has a −1 charge:

Cl⁻

Suggested placement: A sodium-to-chlorine electron-transfer diagram is useful here because students can visually see why Na forms Na⁺ and Cl forms Cl⁻.

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Predicting Ion Charges from Groups

For many main-group elements, the group can help us predict the charge of the ion they commonly form.

Group 14 is less straightforward because losing or gaining four electrons is generally energetically difficult, so these elements commonly form covalent bonds instead of simple ions.

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Using Position to Predict Electron Arrangement

Suppose an unfamiliar element is located in:

Period 3, Group 2

We can predict:

• Period 3 → 3 occupied electron shells
• Group 2 → 2 valence electrons

Its electron arrangement must therefore end with 2 electrons in its third shell.

For the Period 3 element in Group 2:

2,8,2

This element is magnesium (Mg).

We can also predict that magnesium is likely to lose its two valence electrons and form:

Mg²⁺

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Drawing a Bohr Diagram

A Bohr diagram is a simplified model showing electrons arranged in shells around the nucleus.

To draw one using the periodic table:

Step 1 – Find the atomic number

For example, aluminium has atomic number 13.

Therefore, a neutral aluminium atom has 13 electrons.

Step 2 – Fill the first shell

Place 2 electrons in the first shell.

Remaining electrons:

13 − 2 = 11

Step 3 – Fill the second shell

Place 8 electrons in the second shell.

Remaining electrons:

11 − 8 = 3

Step 4 – Place the remaining electrons

Place the final 3 electrons in the third shell.

Electron arrangement:

2,8,3

Aluminium is therefore in:

• Period 3 because it has 3 occupied shells
• Group 13 because it has 3 valence electrons

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Interpreting a Bohr Diagram

We can also work backwards.

Imagine a Bohr diagram shows:

2,8,6

Add the electrons:

2 + 8 + 6 = 16 electrons

For a neutral atom:

16 electrons → 16 protons → atomic number 16

Atomic number 16 is sulfur (S).

We can also determine:

• 3 occupied shells → Period 3
• 6 valence electrons → Group 16
• needs 2 electrons for a full outer shell
• likely ion → S²⁻

A simple electron diagram can therefore provide a surprising amount of information about an element.

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Putting It All Together

Consider chlorine:

Atomic number: 17

From this one element we can connect several ideas:

Atomic number 17

↓

17 protons and 17 electrons

↓

Electron arrangement: 2,8,7

↓

3 occupied shells → Period 3

↓

7 valence electrons → Group 17

↓

Needs 1 electron for a full outer shell

↓

Usually forms Cl⁻

↓

Chemically reactive non-metal

This demonstrates why the periodic table is such a powerful tool for predicting atomic structure and chemical behaviour.

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Did You Know?

The Bohr model is a simplified model of electron arrangement.

Electrons do not actually travel around the nucleus in fixed circular paths like planets orbiting the Sun. Modern atomic theory describes electrons using orbitals and probability distributions.

However, Bohr diagrams remain extremely useful for learning about electron shells, valence electrons, bonding, and periodic-table patterns.

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Key Terms

Electron arrangement – The distribution of electrons among an atom's electron shells.

Electron shell – An energy level occupied by electrons around the nucleus.

Valence electron – An electron in the outermost occupied shell.

Bohr diagram – A simplified representation showing electrons arranged in shells around a nucleus.

Ion – An atom or group of atoms with an electrical charge caused by gaining or losing electrons.

Cation – A positively charged ion.

Anion – A negatively charged ion.

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

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Key Takeaways

• The atomic number tells us how many electrons a neutral atom contains.
• Electrons occupy different electron shells.
• For the first 20 elements, simple arrangements such as 2,8,1 can be used.
• The period number tells us the number of occupied electron shells.
• The group number provides information about the number of valence electrons for main-group elements.
• Elements in the same group have similar chemical properties because they have similar valence-electron arrangements.
• Metals commonly lose electrons and form positive ions.
• Non-metals commonly gain electrons and form negative ions.
• Electron arrangements can help predict ion formation and chemical behaviour.
• Periodic-table information can be used to draw and interpret simple Bohr diagrams.

5. Predicting Element Properties

Predicting Element Properties

One of the greatest strengths of the periodic table is that it allows us to make predictions. Even if we have never studied a particular element before, its position on the periodic table can provide clues about its atomic structure, physical properties, chemical behaviour, and the ions it may form.

This works because the properties of elements follow repeating patterns called periodic trends.

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Position Provides Information

When we locate an element on the periodic table, we can immediately ask:

• What group is it in?
• What period is it in?
• Is it a metal, non-metal, or metalloid?
• How many valence electrons is it likely to have?
• Is it likely to gain or lose electrons?
• What other elements are in the same group?

These clues allow us to make predictions even about unfamiliar elements.

Suggested placement: A periodic table showing groups, periods, and the metal/non-metal staircase provides a useful reference for the predictions throughout this section.

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Predicting Metal or Non-Metal Properties

An element's location gives us an immediate clue about whether it is a metal or non-metal.

Metals

Most elements on the left and centre of the periodic table are metals.

We can generally predict that a metal will:

• conduct electricity
• conduct thermal energy
• have a shiny appearance
• be malleable
• be ductile
• be solid at room temperature
• tend to lose electrons
• form positive ions

Non-Metals

Most elements on the right side are non-metals.

We can generally predict that a non-metal will:

• be a poor electrical conductor
• be a poor thermal conductor
• be dull if solid
• be brittle if solid
• exist as a solid, liquid, or gas
• tend to gain or share electrons
• often form negative ions

These are general patterns, so there are some exceptions.

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Predicting Ion Formation

Atoms often gain or lose electrons to achieve a more stable outer electron arrangement.

The element's group can help us predict what it is likely to do.

This pattern is most useful for the main-group elements.

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Predicting Positive Ions

Metals generally form positive ions, or cations, because they lose electrons.

For example, magnesium is in Group 2.

Its electron arrangement is:

2,8,2

It can lose its two outer electrons:

Mg → Mg²⁺ + 2 electrons

We can therefore predict that other Group 2 metals, such as calcium, will also commonly form 2+ ions.

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Predicting Negative Ions

Non-metals commonly form negative ions, or anions, by gaining electrons.

Chlorine is in Group 17 and has 7 valence electrons.

It can gain one electron:

Cl + electron → Cl⁻

We can therefore predict that other Group 17 elements, such as fluorine and bromine, will also commonly form 1− ions.

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Why Do Elements in the Same Group Behave Similarly?

Elements in the same group have similar numbers of valence electrons.

Chemical reactions mainly involve these outer electrons.

Consider Group 1:

• lithium → 2,1
• sodium → 2,8,1
• potassium → 2,8,8,1

Each element has one valence electron.

All three therefore tend to lose one electron and form +1 ions.

This gives them similar chemical behaviour.

Same group → Similar valence-electron arrangement → Similar chemical properties

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Predicting Properties of Important Groups

Some groups have especially recognisable patterns.

Group 1 – Alkali Metals

Group 1 elements include lithium, sodium, and potassium.

We can predict that they:

• are metals
• have 1 valence electron
• form +1 ions
• conduct electricity
• are reactive
• react with water
• become generally more reactive down the group

For example:

Li → Na → K

Reactivity generally increases.

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Group 2 – Alkaline Earth Metals

Group 2 includes magnesium and calcium.

We can predict that they:

• are metals
• have 2 valence electrons
• form +2 ions
• conduct electricity
• are reactive, although generally less reactive than nearby Group 1 metals

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Group 17 – Halogens

Group 17 includes fluorine, chlorine, bromine, and iodine.

We can predict that they:

• are non-metals
• have 7 valence electrons
• commonly form −1 ions
• are reactive
• become generally less reactive down the group

For example:

F → Cl → Br → I

Reactivity generally decreases.

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Group 18 – Noble Gases

Group 18 includes helium, neon, and argon.

They have full outer electron shells.

We can therefore predict that they:

• are non-metals
• are gases at room temperature
• are very unreactive
• rarely form ions
• rarely form compounds under ordinary conditions

Their chemical stability is directly related to their electron arrangements.

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Physical Trends Across the Periodic Table

The periodic table also shows broad patterns in physical properties.

Metallic Character

Metallic character generally:

increases ← toward the left

and

increases ↓ down the periodic table

Elements toward the lower-left region are therefore generally strongly metallic.

Elements toward the upper-right are generally strongly non-metallic.

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Atomic Size

Atomic size also follows a useful general trend.

Across a period:

Atomic radius generally decreases →

Down a group:

Atomic radius generally increases ↓

Why?

Moving down a group adds additional electron shells, making atoms larger.

Across a period, the number of protons increases while electrons are added to the same main shell. The stronger attraction between the nucleus and electrons generally pulls the electrons closer.

So:

Atomic size increases ↓ and ←

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Chemical Trends

The periodic table can also help predict how strongly atoms attract or lose electrons.

Broadly:

• metals toward the left tend to lose electrons
• non-metals toward the right tend to gain or share electrons
• noble gases already have stable outer shells and are generally unreactive

This helps explain why compounds frequently form between metals and non-metals.

For example:

sodium + chlorine → sodium chloride

Sodium tends to lose an electron, while chlorine tends to gain one.

Their positions on the periodic table allow us to predict this behaviour before seeing the reaction.

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Making Predictions About an Unfamiliar Element

Imagine you are given an unfamiliar element called Element X.

You are told:

Element X is in Group 2 and Period 4.

Even without knowing its name, we can make several predictions.

Step 1 – Use the period

Period 4 means the atom has:

4 occupied electron shells

Step 2 – Use the group

Group 2 means:

2 valence electrons

Step 3 – Predict the type of element

Group 2 elements are metals.

Therefore, Element X is likely to:

• conduct electricity
• conduct thermal energy
• have a metallic appearance

Step 4 – Predict ion formation

It has two valence electrons, so it will probably lose them.

Therefore:

X → X²⁺ + 2 electrons

We can predict that Element X commonly forms a +2 ion.

In fact, the element in Period 4, Group 2 is calcium (Ca).

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Using Evidence to Make Predictions

Scientific predictions should be based on evidence, not guesses.

Suppose an unfamiliar element is located directly below chlorine in Group 17.

A strong prediction would be:

The element is likely to form a −1 ion because it is in Group 17 and therefore has seven valence electrons.

This is stronger than simply saying:

"It probably forms a negative ion."

A good scientific prediction contains:

Prediction + Evidence + Scientific reasoning

For example:

Prediction: Element X will probably form a +1 ion.

Evidence: Element X is located in Group 1.

Reasoning: Group 1 elements have one valence electron and usually lose this electron to achieve a stable outer electron arrangement.

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Predicting an Unknown Element

Consider this information:

Element Y

• Period 3
• Group 17

What can we predict?

From Period 3:

• it has 3 occupied electron shells

From Group 17:

• it has 7 valence electrons
• it is a non-metal
• it is likely to gain 1 electron
• it commonly forms a −1 ion
• it should have properties similar to other halogens

The actual element is chlorine, but we were able to predict many of its properties before identifying it.

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Important: Trends Are Patterns, Not Perfect Rules

Periodic trends are extremely useful, but they should not be treated as rules that apply perfectly in every situation.

For example:

• mercury is a metal but is liquid at room temperature
• graphite is a non-metal but conducts electricity
• transition metals can form ions with different charges
• some elements do not form the simple ions predicted by basic group patterns

Periodic trends provide evidence for predictions, but experimental evidence is still important.

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Did You Know?

The periodic table has been used to predict the properties of elements before those elements were discovered.

In the 1800s, Dmitri Mendeleev left gaps in his periodic table where he believed undiscovered elements should exist. He predicted their properties by comparing them with nearby elements.

When some of these elements were later discovered, their properties were remarkably close to his predictions.

This demonstrated one of the most powerful features of the periodic table: position can predict properties.

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Key Terms

Periodic trend – A repeating pattern in the properties of elements across the periodic table.

Cation – A positively charged ion formed when electrons are lost.

Anion – A negatively charged ion formed when electrons are gained.

Valence electron – An electron in the outermost occupied electron shell.

Metallic character – The degree to which an element displays typical metallic properties.

Atomic radius – A measure of the size of an atom.

Reactivity – How readily a substance undergoes a chemical reaction.

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Key Takeaways

• An element's position on the periodic table can be used to predict many of its properties.
• Elements on the left are generally metals, while those on the right are generally non-metals.
• Metals generally lose electrons and form positive ions.
• Non-metals often gain electrons and form negative ions.
• Elements in the same group have similar chemical properties because they have similar valence-electron arrangements.
• Group 1 elements commonly form +1 ions.
• Group 2 elements commonly form +2 ions.
• Group 17 elements commonly form −1 ions.
• Group 18 elements are generally very unreactive.
• Atomic size generally increases down a group and decreases across a period.
• Good predictions should use an element's group, period, electron arrangement, and location as evidence.
• Periodic trends are useful patterns, but they have exceptions.