1. Predicting Chemical Behaviour

The Periodic Table as a Predictive Tool

One of the most powerful features of the periodic table is that it allows us to predict chemical behaviour.

An element's position provides information about:

• its number of valence electrons
• whether it is a metal or non-metal
• whether it tends to gain or lose electrons
• the charge of the ion it is likely to form
• how strongly it holds its outer electrons
• how it may react with other elements

We therefore do not need to memorize the behaviour of every element individually.

Instead, we can use periodic patterns.

---

Atoms and Electrical Charge

An ordinary atom is electrically neutral because it contains equal numbers of:

protons and electrons

Protons have a charge of:

+1

Electrons have a charge of:

−1

Neutrons have:

no charge

If the number of electrons changes, the particle is no longer electrically neutral.

It becomes an ion.

---

What Is an Ion?

An ion is an atom or group of atoms with an overall electrical charge.

Ions form when electrons are:

• lost
• gained

The number of protons in the nucleus does not change during ordinary ion formation.

If electrons are lost:

positive ion forms

If electrons are gained:

negative ion forms

---

Positive Ions: Cations

A positively charged ion is called a cation.

Cations form when atoms lose electrons.

For example:

A sodium atom contains 11 protons and 11 electrons.

If it loses one electron, it has:

11 protons

10 electrons

The overall charge becomes:

+1

Therefore:

Na → Na⁺ + e⁻

---

Negative Ions: Anions

A negatively charged ion is called an anion.

Anions form when atoms gain electrons.

For example, chlorine has 17 protons and 17 electrons.

If chlorine gains one electron, it has:

17 protons

18 electrons

The overall charge becomes:

−1

Therefore:

Cl + e⁻ → Cl⁻

---

Why Do Atoms Form Ions?

Ion formation is strongly connected to electron arrangement.

Many main-group atoms become more stable when they achieve a complete outer electron shell.

For many elements, this means reaching an electron arrangement similar to a nearby noble gas.

Atoms may therefore gain or lose electrons if doing so produces a more energetically favourable arrangement.

---

Noble Gas Stability

Noble gases have complete outer electron shells.

Examples:

Helium:

2

Neon:

2,8

Argon:

2,8,8

These electron arrangements are particularly stable.

This helps explain why other main-group atoms often form ions that have the same number of electrons as a nearby noble gas.

---

Metals Tend to Lose Electrons

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

Main-group metals typically have relatively few valence electrons.

For example:

Lithium:

2,1

Sodium:

2,8,1

Magnesium:

2,8,2

Aluminium:

2,8,3

For these atoms, losing a small number of electrons can reveal a complete shell underneath.

Therefore, metals generally tend to:

lose electrons → form positive ions

---

Why Metals Form Positive Ions

Consider magnesium.

Its electron arrangement is:

2,8,2

Magnesium can lose its two outer electrons:

Mg → Mg²⁺ + 2e⁻

The resulting Mg²⁺ ion has:

2,8

This is the same electron arrangement as neon.

Because magnesium has lost two negatively charged electrons, its overall charge becomes:

+2

---

Non-Metals Tend to Gain Electrons

Many reactive non-metals are found on the right side of the periodic table.

They often have nearly complete outer shells.

For example:

Nitrogen:

2,5

Oxygen:

2,6

Fluorine:

2,7

Chlorine:

2,8,7

These atoms can reach a complete outer shell by gaining a relatively small number of electrons.

Therefore, many non-metals tend to:

gain electrons → form negative ions

---

Group Number and Ion Charge

For the main-group elements, the group position can often be used to predict the most common simple ion charge.

A useful pattern is:

Group 1 → +1

Group 2 → +2

Group 13 → +3

Group 15 → −3

Group 16 → −2

Group 17 → −1

Group 18 → generally no common ion

These patterns work particularly well for many common introductory chemistry examples.

---

Group 1: +1 Ions

Group 1 elements have one valence electron.

They usually lose that electron.

General pattern:

M → M⁺ + e⁻

Examples:

Li → Li⁺ + e⁻

Na → Na⁺ + e⁻

K → K⁺ + e⁻

Therefore:

Group 1 → +1

---

Why Group 1 Does Not Usually Form +2 Ions

Consider sodium:

2,8,1

Removing one electron produces:

2,8

This is a stable noble-gas arrangement.

Removing a second electron would require taking an electron from the now-complete inner shell.

That requires much more energy.

Therefore, sodium overwhelmingly forms:

Na⁺

rather than:

Na²⁺

---

Group 2: +2 Ions

Group 2 elements have two valence electrons.

They tend to lose both.

Examples include:

• magnesium
• calcium
• strontium
• barium

For magnesium:

Mg → Mg²⁺ + 2e⁻

Therefore:

Group 2 → +2

---

Calcium Example

Calcium has the electron arrangement:

2,8,8,2

It loses two electrons:

Ca → Ca²⁺ + 2e⁻

The resulting ion has:

2,8,8

This is the electron arrangement of argon.

Therefore, calcium commonly forms:

Ca²⁺

---

Group 13: +3 Ions

Group 13 elements have three valence electrons.

For introductory ionic chemistry, aluminium is the most important example.

Aluminium has:

2,8,3

It can lose three electrons:

Al → Al³⁺ + 3e⁻

The resulting electron arrangement is:

2,8

Therefore:

Al³⁺

has a stable noble-gas-like arrangement.

A useful introductory pattern is:

Group 13 metals → +3

although heavier Group 13 elements can show more complicated chemistry.

---

Group 15: −3 Ions

Group 15 elements have five valence electrons.

For some common non-metals, gaining three electrons produces a complete outer shell.

For example, nitrogen:

2,5

can gain three electrons:

N + 3e⁻ → N³⁻

The nitride ion has:

2,8

Therefore:

Group 15 non-metals can form −3 ions.

---

Group 16: −2 Ions

Group 16 elements have six valence electrons.

They need two additional electrons to complete the outer shell.

Oxygen:

2,6

gains two electrons:

O + 2e⁻ → O²⁻

The resulting ion is called:

oxide, O²⁻

Sulfur can similarly form:

S²⁻

Therefore:

Group 16 → −2

for many common simple ions.

---

Group 17: −1 Ions

Halogens have seven valence electrons.

They need only one additional electron.

For chlorine:

Cl + e⁻ → Cl⁻

For fluorine:

F + e⁻ → F⁻

For bromine:

Br + e⁻ → Br⁻

Therefore:

Group 17 → −1

These negative ions are called halide ions.

---

Group 18: No Common Ions

Noble gases already have complete outer electron shells.

Therefore, they generally have little tendency to gain or lose electrons.

Examples:

Neon:

2,8

Argon:

2,8,8

As a result:

Group 18 elements generally do not form common simple ions.

This is closely connected to their low chemical reactivity.

---

A Simple Ion-Charge Pattern

For many main-group elements, you can think about how many electrons must be lost or gained to reach a complete outer shell.

1 outer electron

Lose 1 → +1

2 outer electrons

Lose 2 → +2

3 outer electrons

Lose 3 → +3

5 outer electrons

Gain 3 → −3

6 outer electrons

Gain 2 → −2

7 outer electrons

Gain 1 → −1

Full outer shell

Usually no ion.

---

Why Not Group 14?

Group 14 elements have four valence electrons.

Examples include:

• carbon
• silicon
• germanium
• tin
• lead

Losing four or gaining four electrons is often energetically difficult.

Therefore, Group 14 elements frequently form covalent bonds instead of simple ±4 monatomic ions.

This is why there is no simple universal:

Group 14 = ±4 ion

rule for introductory chemistry.

---

Metals vs Non-Metals

The broad periodic pattern can be summarized as:

Metals

Usually lose electrons.

Form:

cations

Non-metals

Often gain electrons.

Form:

anions

Noble gases

Usually neither gain nor lose electrons.

Remain:

neutral atoms

This pattern is closely related to the position of elements on the periodic table.

---

The Metal–Non-Metal Boundary

Metals occupy most of the left and centre of the periodic table.

Non-metals are concentrated toward the upper-right region.

Elements near the boundary can show intermediate or more complicated behaviour.

These elements are often called metalloids.

Examples include:

• boron
• silicon
• germanium

Their chemistry is often less easily described using simple ion-charge rules.

---

Transition Metals Are More Complicated

Transition metals form positive ions, but their charges cannot usually be predicted using a simple group-number rule.

For example:

Iron can form:

Fe²⁺

and:

Fe³⁺

Copper can form:

Cu⁺

and:

Cu²⁺

This happens because transition-metal chemistry involves outer s electrons and nearby d electrons.

Therefore, many transition metals have variable oxidation states.

---

Common Transition-Metal Ions

Some common examples include:

Fe²⁺ – iron(II)

Fe³⁺ – iron(III)

Cu⁺ – copper(I)

Cu²⁺ – copper(II)

Co²⁺ – cobalt(II)

Ni²⁺ – nickel(II)

Cr³⁺ – chromium(III)

For transition metals, the ion charge is often given in the compound's name using a Roman numeral.

For example:

iron(III) chloride

means that iron has oxidation state:

+3

---

Electron Arrangements and Stability

Consider these examples:

Na: 2,8,1 → Na⁺: 2,8

Mg: 2,8,2 → Mg²⁺: 2,8

Al: 2,8,3 → Al³⁺: 2,8

All three produce ions with the electron arrangement of:

neon

Now consider:

N: 2,5 → N³⁻: 2,8

O: 2,6 → O²⁻: 2,8

F: 2,7 → F⁻: 2,8

Again, all three reach the same stable noble-gas arrangement.

---

Isoelectronic Ions

Particles with the same number of electrons are described as isoelectronic.

For example:

N³⁻

O²⁻

F⁻

Ne

Na⁺

Mg²⁺

Al³⁺

all contain 10 electrons.

They are different elements because they contain different numbers of protons, but their electron counts are the same.

---

Ion Formation and Ionic Bonding

Ion formation becomes especially important when metals react with non-metals.

For example:

Sodium loses one electron:

Na → Na⁺ + e⁻

Chlorine gains one:

Cl + e⁻ → Cl⁻

The resulting ions are oppositely charged.

They attract each other through a strong electrostatic attraction.

This forms an ionic bond.

---

Predicting Ionic Formulas

Once we can predict ion charges, we can also predict the formulas of ionic compounds.

The total positive charge must balance the total negative charge.

For example:

Na⁺ and Cl⁻ combine in a:

1 : 1 ratio

Therefore:

NaCl

---

Magnesium Chloride

Magnesium forms:

Mg²⁺

Chlorine forms:

Cl⁻

One Mg²⁺ ion requires two Cl⁻ ions:

(+2) + 2(−1) = 0

Therefore:

MgCl₂

---

Aluminium Oxide

Aluminium forms:

Al³⁺

Oxygen forms:

O²⁻

We need the charges to balance.

Two Al³⁺ ions give:

+6

Three O²⁻ ions give:

−6

Therefore:

Al₂O₃

This shows why knowing common ion charges is so useful.

---

Calcium Nitride

Calcium is in Group 2.

Therefore:

Ca²⁺

Nitrogen is in Group 15.

Therefore:

N³⁻

The lowest common total charge is 6.

Three Ca²⁺ ions:

+6

Two N³⁻ ions:

−6

Therefore:

Ca₃N₂

---

Periodic Trends and Ion Formation

Ion formation is not determined only by counting outer electrons.

It is also connected to periodic trends such as:

• atomic radius
• shielding
• ionization energy
• electronegativity

These trends help explain how easily atoms gain or lose electrons.

---

Ionization Energy and Metals

Ionization energy is the energy required to remove an electron from an isolated gaseous atom.

Metals that form positive ions must lose electrons.

Therefore, relatively low ionization energies generally make cation formation easier.

Group 1 metals have relatively low first ionization energies.

This helps explain why they readily form:

+1 ions

---

Electronegativity and Non-Metals

Electronegativity describes an atom's tendency to attract bonding electrons.

Electronegativity generally increases toward the upper-right region of the periodic table.

Reactive non-metals such as fluorine and chlorine strongly attract electrons.

This is consistent with their tendency to gain electrons or attract shared electrons strongly in chemical bonds.

---

Atomic Radius and Electron Loss

Atomic radius generally increases down a group.

As atoms become larger:

• the outer electron is farther from the nucleus
• shielding generally increases
• the outer electron experiences weaker effective nuclear attraction

This can make electron removal easier.

This is especially important for explaining Group 1.

---

Group 1 Down the Group

Consider:

Li → Na → K → Rb → Cs

All have one valence electron.

Moving down the group:

• atomic radius increases
• shielding increases
• first ionization energy generally decreases

Therefore, the outer electron becomes easier to remove.

All still form:

+1 ions

but the ease of ion formation helps explain why their reactivity increases down the group.

---

Group 17 Down the Group

Consider:

F → Cl → Br → I

All have seven valence electrons.

They all tend to form:

−1 ions

However, moving down the group:

• atomic radius increases
• shielding increases
• attraction for an incoming electron becomes less effective

Therefore, halogen reactivity generally:

decreases down the group.

The typical ion charge remains −1, but the tendency to react changes.

---

Charge vs Reactivity

This distinction is important.

Ion charge tells us how many electrons are typically gained or lost.

Reactivity tells us how readily the chemical process occurs.

For example:

Lithium and potassium both form:

+1 ions

but potassium is more reactive.

Chlorine and iodine both form:

−1 ions

but chlorine is more reactive.

Therefore:

same ion charge does not mean same reactivity.

---

Predicting an Unfamiliar Group 1 Element

Suppose element X is located below potassium in Group 1.

We can predict:

• X is a metal
• X has one valence electron
• X loses one electron
• X forms X⁺
• X should be highly reactive
• X should be more reactive than potassium

Why?

Its outer electron experiences greater shielding and is farther from the nucleus.

Therefore, it should be relatively easy to remove.

---

Predicting an Unfamiliar Group 2 Element

Element Y is below calcium in Group 2.

We can predict:

• Y is metallic
• Y has two valence electrons
• Y tends to lose two electrons
• Y forms Y²⁺

The exact reactivity may differ from calcium, but the common ion charge can be predicted from its group membership.

---

Predicting an Unfamiliar Group 16 Element

Element Z is a non-metal in Group 16.

It has:

six valence electrons

It needs:

two more electrons

to complete its outer shell.

Therefore:

Z + 2e⁻ → Z²⁻

We predict:

Z²⁻

---

Predicting an Unfamiliar Halogen

Suppose element Q lies below bromine in Group 17.

We can predict:

• Q has seven valence electrons
• Q is a non-metal
• Q tends to gain one electron
• Q forms Q⁻
• Q should be less reactive than bromine

This prediction combines:

group membership + electron arrangement + periodic trend

---

Predicting an Unknown Element from Its Ion

Suppose an unknown main-group element forms:

X²⁺

What might we predict?

If X follows the common main-group pattern, it may belong to:

Group 2

because Group 2 metals lose two valence electrons.

Now suppose:

Y²⁻

Y may belong to:

Group 16

because Group 16 non-metals commonly gain two electrons.

---

Worked Example 1

Predict the ion formed by potassium.

Potassium is in:

Group 1

Group 1 elements have one valence electron.

Potassium loses one electron.

Therefore:

K → K⁺ + e⁻

Answer:

K⁺

---

Worked Example 2

Predict the ion formed by calcium.

Calcium is in:

Group 2

It has two valence electrons.

It loses both.

Therefore:

Ca → Ca²⁺ + 2e⁻

Answer:

Ca²⁺

---

Worked Example 3

Predict the ion formed by oxygen.

Oxygen is in:

Group 16

It has six valence electrons.

It needs two more electrons to complete its outer shell.

Therefore:

O + 2e⁻ → O²⁻

Answer:

O²⁻

---

Worked Example 4

Predict the ion formed by bromine.

Bromine belongs to:

Group 17

It has seven valence electrons.

It gains one electron.

Therefore:

Br + e⁻ → Br⁻

Answer:

Br⁻

---

Worked Example 5

Predict the ion formed by aluminium.

Aluminium is in:

Group 13

It has three valence electrons.

It tends to lose those three electrons.

Therefore:

Al → Al³⁺ + 3e⁻

Answer:

Al³⁺

---

Worked Example 6

An unknown element has the electron arrangement:

2,8,6

Predict its ion.

The atom has six valence electrons.

It needs two additional electrons to reach:

2,8,8

Therefore:

X + 2e⁻ → X²⁻

The predicted ion is:

X²⁻

---

Worked Example 7

An unknown element has:

2,8,2

Predict its behaviour.

It has two valence electrons.

It is likely to lose both:

X → X²⁺ + 2e⁻

Its ion has:

2,8

Therefore, we would predict that X behaves like a Group 2 metal.

---

Worked Example 8

An element lies in Group 18.

Predict its ion.

Group 18 elements already have complete outer electron shells.

Therefore, they generally do not need to gain or lose electrons.

Answer:

No common simple ion is expected.

---

Solving Unfamiliar Ion Questions

When given an unfamiliar element, use this process.

Step 1: Find the group.

Identify its position on the periodic table.

Step 2: Determine the number of valence electrons.

For the main groups, group position provides this information.

Step 3: Decide whether gaining or losing electrons is more reasonable.

Metals usually lose.

Non-metals often gain.

Step 4: Determine how many electrons are transferred.

Aim for a stable outer electron arrangement.

Step 5: Determine the ion charge.

Lose electrons → positive.

Gain electrons → negative.

Step 6: Check the periodic trend.

Consider ionization energy, atomic radius, shielding, and electronegativity when explaining how readily the process occurs.

---

A Quick Prediction Pattern

For common main-group ions:

Group 1 → lose 1 → +1

Group 2 → lose 2 → +2

Group 13 → lose 3 → +3

Group 14 → usually shares electrons rather than forming simple ±4 ions

Group 15 → gain 3 → −3

Group 16 → gain 2 → −2

Group 17 → gain 1 → −1

Group 18 → usually no ion

This pattern is extremely useful for predicting unfamiliar main-group ions.

---

Common Mistakes

Mistake 1: Thinking positive ions gain electrons

Positive ions form when electrons are:

lost

---

Mistake 2: Thinking negative ions lose electrons

Negative ions form when electrons are:

gained

---

Mistake 3: Changing the number of protons

Ordinary chemical ion formation changes the number of:

electrons

not protons.

Changing the proton number would change the identity of the element.

---

Mistake 4: Using the group number directly as the charge

Group 16 does not form +16 ions.

Group number helps determine the number of valence electrons, which is then used to predict electron gain or loss.

---

Mistake 5: Predicting Group 14 as automatically +4 or −4

Group 14 frequently forms covalent bonds rather than simple monatomic ±4 ions.

---

Mistake 6: Applying simple group rules to transition metals

Transition metals often have:

variable oxidation states

Their charges may need to be given or determined from the compound.

---

Mistake 7: Confusing ion charge with reactivity

Potassium and lithium both form +1 ions, but potassium is more reactive.

Ion charge and reactivity are related to electron structure but are not the same property.

---

Did You Know?

The ion-charge pattern creates a useful symmetry across the main-group periodic table.

Moving from the left toward Group 18:

+1 → +2 → +3 → ... → −3 → −2 → −1 → 0

This pattern is not an arbitrary set of numbers.

It reflects how many electrons atoms typically need to lose or gain to reach more stable electron arrangements.

---

Key Terms

Ion: A particle with an overall electrical charge.

Cation: A positively charged ion.

Anion: A negatively charged ion.

Valence electron: An electron in the outermost occupied shell.

Ion formation: The process of gaining or losing electrons to produce an ion.

Noble-gas configuration: An electron arrangement corresponding to a noble gas, often associated with a complete outer shell.

Ionization energy: The energy required to remove an electron from an isolated gaseous atom.

Electronegativity: An atom's tendency to attract bonding electrons.

Atomic radius: A measure of atomic size.

Shielding: The reduction in effective nuclear attraction caused by inner electrons.

Isoelectronic: Having the same number of electrons.

Oxidation state: A value used to describe the formal distribution of electrons in a compound or ion.

---

Key Ion Patterns

Group 1

1 valence electron

Loses 1 electron

+1

Examples: Li⁺, Na⁺, K⁺

---

Group 2

2 valence electrons

Loses 2 electrons

+2

Examples: Mg²⁺, Ca²⁺

---

Group 13

3 valence electrons

Often loses 3 electrons

+3

Important example: Al³⁺

---

Group 15

5 valence electrons

Often gains 3 electrons

−3

Example: N³⁻

---

Group 16

6 valence electrons

Gains 2 electrons

−2

Examples: O²⁻, S²⁻

---

Group 17

7 valence electrons

Gains 1 electron

−1

Examples: F⁻, Cl⁻, Br⁻, I⁻

---

Group 18

Complete outer shell

Generally neither gains nor loses electrons

No common simple ion

---

Key Takeaways

• The periodic table can be used to predict the ions formed by many elements.
• Ion formation involves changing the number of electrons, not protons.
• Losing electrons produces a positive ion or cation.
• Gaining electrons produces a negative ion or anion.
• Metals generally form positive ions because they tend to lose valence electrons.
• Non-metals often form negative ions because they tend to gain electrons.
• Many main-group atoms form ions with stable, noble-gas-like electron arrangements.
• Group 1 elements normally form +1 ions.
• Group 2 elements normally form +2 ions.
• Aluminium commonly forms Al³⁺.
• Group 15 non-metals can form −3 ions.
• Group 16 non-metals commonly form −2 ions.
• Group 17 halogens commonly form −1 ions.
• Group 18 noble gases generally do not form common simple ions.
• Group 14 does not follow a simple ±4 monatomic-ion rule because its elements frequently form covalent bonds.
• Transition metals often have multiple oxidation states, so simple group-number rules do not usually predict their charges.
• Periodic trends such as atomic radius, shielding, ionization energy, and electronegativity help explain how readily electrons are gained or lost.
• Ion charge and reactivity are not the same thing: elements can form the same charge while reacting at different rates.
• Once ion charges are known, they can also be used to predict the formulas of ionic compounds.
• For unfamiliar elements, use the reasoning chain:

group position → valence electrons → metal/non-metal behaviour → electron gain or loss → stable electron arrangement → ion charge → periodic trend → predicted chemical behaviour.