The Periodic Table Helps Us Predict Bonding

Atoms rarely behave randomly.

Their chemical behaviour is strongly connected to their electron arrangements, especially the electrons in their outer shells.

The periodic table therefore helps us predict:

• whether an atom tends to gain or lose electrons
• whether atoms are likely to transfer or share electrons
• whether a compound is likely to be ionic or covalent
• the charges of many common ions
• the ratio in which ions combine
• the chemical formula of many simple compounds

The key is to connect:

periodic-table position → valence electrons → electron behaviour → bonding → compound formula

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Why Do Atoms Bond?

Atoms form chemical bonds because interactions between their electrons and nuclei can produce a lower-energy, more stable arrangement.

For many main-group elements, a useful introductory model is that atoms tend toward a complete outer electron shell.

There are two especially important ways this can happen:

Electron transfer → ionic bonding

Electron sharing → covalent bonding

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Start by Identifying Metals and Non-Metals

A quick first step is to locate the elements on the periodic table.

Most metals are located on the left and centre.

Most non-metals are located toward the upper-right.

The staircase-like region between them contains the metalloids, which can show intermediate behaviour.

This gives us a very useful introductory rule.

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A Simple Bonding Prediction

For many simple compounds:

metal + non-metal → usually ionic

non-metal + non-metal → usually covalent

This rule allows us to make a quick prediction before examining the electron arrangements in detail.

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Ionic Bonding

An ionic bond is the strong electrostatic attraction between oppositely charged ions.

Ionic compounds usually form when:

a metal reacts with a non-metal.

The metal loses electrons.

The non-metal gains electrons.

This electron transfer creates:

• positive ions called cations
• negative ions called anions

The oppositely charged ions then attract each other.

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Example: Sodium and Chlorine

Sodium is in Group 1.

Its electron arrangement is:

2,8,1

Chlorine is in Group 17.

Its electron arrangement is:

2,8,7

Sodium loses one electron:

Na → Na⁺ + e⁻

Chlorine gains one:

Cl + e⁻ → Cl⁻

The resulting ions are:

Na⁺

and:

Cl⁻

They attract each other and form:

NaCl

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Why Sodium and Chlorine Form Ions

Before bonding:

Na:

2,8,1

Cl:

2,8,7

After electron transfer:

Na⁺:

2,8

Cl⁻:

2,8,8

Both ions now have complete outer electron shells.

This illustrates how electron arrangement helps explain ionic bonding.

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Predicting Common Ion Charges

For many main-group elements, the periodic table can be used to predict simple ion charges.

Group 1 → +1

Group 2 → +2

Group 13 → +3

Group 15 → −3

Group 16 → −2

Group 17 → −1

Group 18 → generally no common simple ions

These charges are extremely useful when predicting the formulas of ionic compounds.

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Ionic Compounds Must Be Electrically Neutral

An ionic compound contains positive and negative ions.

However, the compound as a whole must have:

total charge = 0

Therefore:

total positive charge = total negative charge

This rule allows us to determine the correct ratio of ions.

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Example: Sodium Chloride

Sodium forms:

Na⁺

Chlorine forms:

Cl⁻

The charges already balance:

(+1) + (−1) = 0

Therefore, the ratio is:

1 Na⁺ : 1 Cl⁻

Formula:

NaCl

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Example: Magnesium Chloride

Magnesium is in Group 2.

Therefore:

Mg²⁺

Chlorine is in Group 17.

Therefore:

Cl⁻

One Mg²⁺ has a charge of:

+2

Two Cl⁻ ions have a total charge of:

−2

Therefore:

MgCl₂

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Why MgCl₂ and Not MgCl?

Suppose we wrote:

MgCl

The charges would be:

Mg²⁺ = +2

Cl⁻ = −1

Total:

+1

The compound would not be electrically neutral.

We need another chloride ion:

(+2) + (−1) + (−1) = 0

Therefore:

MgCl₂

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Example: Magnesium Oxide

Magnesium forms:

Mg²⁺

Oxygen forms:

O²⁻

The charges are equal and opposite:

(+2) + (−2) = 0

Therefore, only one of each ion is required.

Formula:

MgO

Not:

Mg₂O₂

Chemical formulas are written using the simplest whole-number ratio.

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Example: Calcium Fluoride

Calcium belongs to Group 2:

Ca²⁺

Fluorine belongs to Group 17:

F⁻

Two fluoride ions are needed to balance one calcium ion.

(+2) + 2(−1) = 0

Therefore:

CaF₂

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Example: Aluminium Oxide

Aluminium commonly forms:

Al³⁺

Oxygen forms:

O²⁻

We need the smallest combination where the charges cancel.

Two Al³⁺ ions:

2 × (+3) = +6

Three O²⁻ ions:

3 × (−2) = −6

Total:

0

Therefore:

Al₂O₃

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Finding the Lowest Common Charge

For ions with different charge sizes, finding a common total can help.

For:

Al³⁺ and O²⁻

the lowest common multiple of 3 and 2 is:

6

So we need:

2 aluminium ions → +6

3 oxide ions → −6

Formula:

Al₂O₃

This method is reliable and helps explain why the subscripts appear.

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Example: Calcium Nitride

Calcium:

Ca²⁺

Nitrogen:

N³⁻

The lowest common charge is:

6

Three Ca²⁺:

+6

Two N³⁻:

−6

Therefore:

Ca₃N₂

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The Criss-Cross Shortcut

A common shortcut is the criss-cross method.

For:

Mg²⁺ and Cl⁻

the charge numbers become subscripts:

Mg²⁺ and Cl⁻

↓

MgCl₂

For:

Al³⁺ and O²⁻

↓

Al₂O₃

However, the criss-cross method should not replace understanding.

The real rule is:

the total positive and negative charges must balance.

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Always Simplify the Ratio

Consider:

Ca²⁺ and O²⁻

Criss-crossing mechanically might suggest:

Ca₂O₂

But both subscripts can be divided by 2.

The simplest ratio is:

CaO

Therefore, always check whether the formula can be simplified.

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Writing Ionic Formulas: A Reliable Strategy

Use these steps.

Step 1: Identify the metal and non-metal.

Step 2: Determine their ion charges.

Step 3: Find the smallest numbers of each ion that make the total charge zero.

Step 4: Write the cation first.

Step 5: Write the anion second.

Step 6: Use subscripts to show the ratio.

Step 7: Simplify if necessary.

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Transition Metals Require Extra Information

Transition metals can be more complicated because many have multiple oxidation states.

Iron can form:

Fe²⁺

or:

Fe³⁺

Copper can form:

Cu⁺

or:

Cu²⁺

Therefore, we often need additional information before predicting the formula.

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Roman Numerals Tell Us the Charge

Consider:

iron(II) chloride

The Roman numeral II means:

Fe²⁺

Chloride is:

Cl⁻

Therefore:

FeCl₂

Now consider:

iron(III) chloride

Iron is:

Fe³⁺

Therefore:

FeCl₃

The Roman numeral is essential because iron can form more than one ion.

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Covalent Bonding

A covalent bond forms when atoms share pairs of electrons.

Covalent bonding usually occurs between:

non-metal + non-metal

Instead of transferring electrons and forming ions, the atoms share electrons.

This allows each atom to achieve a more stable electron arrangement.

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Example: Hydrogen

Each hydrogen atom has:

1 electron

The first electron shell can hold:

2 electrons

Two hydrogen atoms can share a pair of electrons.

This forms:

H₂

The shared pair is a single covalent bond.

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Example: Chlorine

Each chlorine atom has seven valence electrons.

Each needs one more electron to complete its outer shell.

Two chlorine atoms share one pair of electrons.

This forms:

Cl₂

Each chlorine atom effectively has access to eight outer-shell electrons.

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Example: Oxygen

Oxygen has six valence electrons.

Each oxygen atom needs two additional electrons to complete its outer shell.

Two oxygen atoms can share two pairs of electrons.

This produces a double bond:

O=O

Formula:

O₂

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Example: Nitrogen

Nitrogen has five valence electrons.

Each nitrogen atom needs three additional electrons to complete its outer shell.

Two nitrogen atoms share three pairs of electrons.

This forms a triple bond:

N≡N

Formula:

N₂

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Covalent Bonding in Water

Water contains hydrogen and oxygen.

Both are non-metals.

Therefore, we predict:

covalent bonding

Oxygen needs two additional electrons to complete its outer shell.

Each hydrogen needs one additional electron to complete its first shell.

Oxygen forms two covalent bonds, one with each hydrogen.

Formula:

H₂O

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Predicting Covalent Formulas from Bonding Capacity

For simple molecular compounds, the number of electrons an atom needs for a complete outer shell can help predict how many covalent bonds it commonly forms.

A useful introductory pattern is:

Hydrogen → 1 bond

Group 17 → usually 1 bond

Group 16 → usually 2 bonds

Group 15 → usually 3 bonds

Carbon → usually 4 bonds

This can help predict simple molecular formulas.

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Example: Hydrogen Chloride

Hydrogen typically forms:

1 bond

Chlorine typically forms:

1 bond

Therefore, one hydrogen atom bonds with one chlorine atom.

Formula:

HCl

Bond:

H–Cl

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Example: Hydrogen Sulfide

Sulfur is in Group 16.

It commonly forms two covalent bonds in simple molecules.

Hydrogen forms one bond.

Therefore, sulfur can bond to two hydrogen atoms.

Formula:

H₂S

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Example: Ammonia

Nitrogen has five valence electrons.

It commonly forms three covalent bonds.

Hydrogen forms one bond.

One nitrogen bonds with three hydrogen atoms.

Formula:

NH₃

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Example: Methane

Carbon has four valence electrons.

It commonly forms four covalent bonds.

Each hydrogen forms one bond.

Therefore:

CH₄

Carbon shares one pair of electrons with each of four hydrogen atoms.

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Example: Carbon Dioxide

Carbon commonly forms four bonds.

Oxygen commonly forms two.

One carbon atom can form two double bonds with two oxygen atoms:

O=C=O

Formula:

CO₂

This satisfies the typical bonding requirements of both carbon and oxygen.

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Ionic or Covalent?

A useful first prediction is:

NaCl

Na = metal

Cl = non-metal

Prediction:

ionic

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MgO

Mg = metal

O = non-metal

Prediction:

ionic

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H₂O

H = non-metal

O = non-metal

Prediction:

covalent

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CO₂

C = non-metal

O = non-metal

Prediction:

covalent

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NH₃

N = non-metal

H = non-metal

Prediction:

covalent

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Comparing Ionic and Covalent Bonding

Ionic bonding

Usually metal + non-metal.

Electrons are transferred.

Positive and negative ions form.

The bond is the electrostatic attraction between oppositely charged ions.

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Covalent bonding

Usually non-metal + non-metal.

Electrons are shared.

Shared electron pairs form covalent bonds.

Simple molecular substances contain molecules rather than separate positive and negative ions.

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Electron Arrangements Explain the Difference

Consider sodium:

2,8,1

It is easier for sodium to lose one electron than to gain seven.

Therefore, sodium forms:

Na⁺

Now consider chlorine:

2,8,7

It is easier for chlorine to gain one electron than to lose seven.

Therefore:

Cl⁻

This makes electron transfer between sodium and chlorine favourable in the formation of an ionic solid.

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Why Two Non-Metals Usually Share

Consider two chlorine atoms.

Both strongly attract electrons.

Neither atom readily gives its electron completely to the other.

Instead, they share a pair of electrons.

Therefore:

Cl–Cl

This is a covalent bond.

The difference between electron transfer and electron sharing is central to predicting bonding.

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Electronegativity and Bonding

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

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

When two atoms have similar electronegativities, electron sharing is more likely.

When the difference is large, bonding tends to have greater ionic character.

Therefore, bonding is better understood as a continuum rather than as two completely separate categories.

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Bonding Is a Continuum

In introductory chemistry we often classify compounds as either:

ionic

or:

covalent

This is extremely useful.

However, real chemical bonding can contain different degrees of ionic and covalent character.

For example:

• equal sharing → non-polar covalent
• unequal sharing → polar covalent
• very large electronegativity difference → strong ionic character

The simple metal/non-metal rule is therefore a prediction tool, not an absolute law.

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Predicting an Unfamiliar Compound

Suppose element X is in Group 1 and element Y is in Group 17.

X is expected to form:

X⁺

Y is expected to form:

Y⁻

The charges balance 1:1.

Therefore, the predicted formula is:

XY

The compound is expected to be:

ionic

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Unfamiliar Group 2 + Group 17

Suppose:

X belongs to Group 2.

Y belongs to Group 17.

Predict the ions:

X²⁺

Y⁻

Two Y⁻ ions are required for every X²⁺ ion.

Therefore:

XY₂

The compound should be predominantly:

ionic

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Unfamiliar Group 1 + Group 16

Element X is in Group 1:

X⁺

Element Y is in Group 16:

Y²⁻

Two X⁺ ions are needed:

2(+1) + (−2) = 0

Therefore:

X₂Y

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Unfamiliar Group 2 + Group 16

X is Group 2:

X²⁺

Y is Group 16:

Y²⁻

The charges cancel in a 1:1 ratio.

Therefore:

XY

Not:

X₂Y₂

because formulas use the simplest whole-number ratio.

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Unfamiliar Group 13 + Group 17

X forms:

X³⁺

Y forms:

Y⁻

Three Y⁻ ions are required:

(+3) + 3(−1) = 0

Therefore:

XY₃

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Unfamiliar Group 13 + Group 16

This combination requires more careful charge balancing.

X:

X³⁺

Y:

Y²⁻

Lowest common charge:

6

Two X³⁺:

+6

Three Y²⁻:

−6

Therefore:

X₂Y₃

This follows the same pattern as aluminium oxide:

Al₂O₃

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Unfamiliar Non-Metal + Non-Metal

Suppose X and Y are both non-metals.

A first prediction would be:

covalent bonding

Instead of simply assigning ionic charges, consider:

• number of valence electrons
• typical number of covalent bonds
• electron sharing
• possible molecular structure

For example, a Group 14 non-metal and a Group 17 non-metal often combine through covalent bonding.

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Example: Carbon and Chlorine

Carbon has four valence electrons and commonly forms four covalent bonds.

Chlorine commonly forms one.

One carbon can therefore bond with four chlorine atoms.

Formula:

CCl₄

This is a covalent compound.

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Predicting Properties from Bonding

Predicting bonding can also help us predict some physical properties.

Ionic compounds generally:

• form giant ionic lattices
• have relatively high melting and boiling points
• are often brittle
• conduct electricity when molten or dissolved in water if mobile ions are present
• do not conduct well as solids

These properties result from the strong electrostatic attractions between ions.

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Properties of Simple Molecular Covalent Substances

Simple molecular substances generally:

• contain separate molecules
• have relatively low melting and boiling points compared with giant ionic structures
• do not conduct electricity well
• may exist as gases, liquids, or low-melting solids

The covalent bonds inside molecules are strong, but the attractions between separate molecules are generally much weaker.

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An Important Exception: Giant Covalent Structures

Not every covalent substance consists of small molecules.

Some form enormous networks of covalent bonds.

Examples include:

• diamond
• graphite
• silicon dioxide

These substances can have very high melting points because many strong covalent bonds must be overcome.

Therefore:

covalent does not automatically mean low melting point.

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Using the Periodic Table as a Decision Tool

When predicting a compound, ask:

1. Where are the elements?

Metal or non-metal?

2. How many valence electrons do they have?

Use their groups.

3. Will electrons probably be transferred or shared?

Metal + non-metal → usually transfer.

Non-metal + non-metal → usually share.

4. If ionic, what charges form?

Use group patterns.

5. What ratio makes the compound neutral?

Balance the charges.

6. If covalent, how many bonds are typically needed?

Use valence electrons and stable-shell ideas.

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Worked Example 1: Potassium and Oxygen

Potassium is in Group 1.

Therefore:

K⁺

Oxygen is in Group 16.

Therefore:

O²⁻

Potassium is a metal and oxygen is a non-metal.

Prediction:

ionic

Two K⁺ ions balance one O²⁻:

2(+1) + (−2) = 0

Formula:

K₂O

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Worked Example 2: Magnesium and Nitrogen

Magnesium:

Group 2 → Mg²⁺

Nitrogen:

Group 15 → N³⁻

Metal + non-metal:

ionic

Lowest common charge:

6

Three Mg²⁺:

+6

Two N³⁻:

−6

Formula:

Mg₃N₂

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Worked Example 3: Aluminium and Chlorine

Aluminium:

Al³⁺

Chlorine:

Cl⁻

Metal + non-metal:

ionic

Three chloride ions are needed.

Formula:

AlCl₃

At an introductory level this charge-balancing model correctly predicts the formula, although the bonding in real aluminium chloride can show substantial covalent character depending on its form and conditions.

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Worked Example 4: Carbon and Oxygen

Carbon and oxygen are both non-metals.

Prediction:

covalent

Carbon commonly forms four bonds.

Oxygen commonly forms two.

One carbon can therefore form two double bonds with two oxygen atoms:

O=C=O

Formula:

CO₂

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Worked Example 5: Nitrogen and Hydrogen

Both are non-metals.

Prediction:

covalent

Nitrogen commonly forms three bonds.

Hydrogen forms one.

Therefore, one nitrogen bonds with three hydrogens.

Formula:

NH₃

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Worked Example 6: Unknown Elements

Element X is in Group 2.

Element Y is in Group 17.

X is a metal and forms:

X²⁺

Y is a non-metal and forms:

Y⁻

Prediction:

ionic

Charge balance:

(+2) + 2(−1) = 0

Formula:

XY₂

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Worked Example 7: Identifying Bonding from a Formula

Consider:

CaCl₂

Calcium is a metal.

Chlorine is a non-metal.

Therefore, we predict:

ionic bonding

The formula also makes sense from the ion charges:

Ca²⁺

2 × Cl⁻

Total:

0

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Worked Example 8: Predicting from Position

Suppose an unfamiliar element X is in Group 1 and an unfamiliar element Y is in Group 16.

X forms:

X⁺

Y forms:

Y²⁻

The expected formula is:

X₂Y

Because X is metallic and Y is non-metallic, the compound is predicted to be:

ionic

This prediction can be made without knowing the identities of X and Y.

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Formula vs Charge

Do not include ionic charges in the final formula of an ordinary neutral ionic compound.

Correct ion notation:

Mg²⁺

Cl⁻

Correct compound formula:

MgCl₂

Not:

Mg²⁺Cl₂⁻

The formula shows the ratio of ions, while the charges explain why that ratio is required.

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Subscripts Have Meaning

In:

MgCl₂

the subscript 2 means:

two chloride ions for every magnesium ion.

In:

Al₂O₃

the subscripts mean:

two aluminium ions for every three oxide ions.

Subscripts are therefore not random numbers. They describe the simplest ratio of particles in the compound.

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Common Mistakes

Mistake 1: Assuming metal + metal forms an ionic compound

Simple ionic bonding is usually predicted for:

metal + non-metal

Two metals more commonly form metallic structures or alloys.

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Mistake 2: Assuming non-metal + non-metal is ionic

Two non-metals usually form:

covalent bonds

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Mistake 3: Forgetting to balance charges

Mg²⁺ and Cl⁻ do not produce MgCl.

The correct formula is:

MgCl₂

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Mistake 4: Changing the ion charges to make them balance

The charges stay the same.

Instead, change the number of ions.

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Mistake 5: Writing charges as subscripts

For aluminium oxide:

Al³⁺ and O²⁻

The formula is:

Al₂O₃

not Al₃O₂.

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Mistake 6: Forgetting to simplify

Ca²⁺ and O²⁻ form:

CaO

not Ca₂O₂.

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Mistake 7: Applying simple group charges to transition metals

Transition metals can have multiple oxidation states.

Use information such as:

iron(II) or iron(III)

to determine the correct charge.

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Mistake 8: Thinking covalent bonding means electrons are transferred

Covalent bonding involves:

sharing electron pairs.

Ionic bonding involves:

electron transfer followed by electrostatic attraction between ions.

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Did You Know?

The chemical formula of an ionic compound tells us a ratio, not necessarily that separate molecules exist.

For example:

NaCl

does not mean that a sodium chloride crystal consists of individual NaCl molecules.

Instead, solid sodium chloride contains a giant lattice of Na⁺ and Cl⁻ ions.

The formula NaCl tells us that the ions occur in a:

1 : 1 ratio

This is why the term formula unit is often used for ionic compounds.

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Key Terms

Chemical bond: An attractive interaction that holds atoms or ions together.

Ionic bond: Strong electrostatic attraction between oppositely charged ions.

Covalent bond: A bond formed by a shared pair of electrons.

Cation: A positively charged ion.

Anion: A negatively charged ion.

Valence electron: An electron in the outermost occupied shell.

Ion charge: The electrical charge produced when an atom gains or loses electrons.

Chemical formula: Symbols and subscripts showing the composition of a substance.

Subscript: A small number showing the relative number of atoms or ions in a formula.

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

Formula unit: The simplest whole-number ratio of ions in an ionic compound.

Diatomic: Existing as molecules containing two atoms.

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Key Bonding Patterns

Metal + non-metal

Usually:

ionic

Example:

Na + Cl → NaCl

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Non-metal + non-metal

Usually:

covalent

Example:

H + O → H₂O

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Group 1 + Group 17

X⁺ + Y⁻

Formula:

XY

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Group 2 + Group 17

X²⁺ + Y⁻

Formula:

XY₂

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Group 1 + Group 16

X⁺ + Y²⁻

Formula:

X₂Y

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Group 2 + Group 16

X²⁺ + Y²⁻

Formula:

XY

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Group 13 + Group 17

X³⁺ + Y⁻

Formula:

XY₃

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Group 13 + Group 16

X³⁺ + Y²⁻

Formula:

X₂Y₃

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Key Takeaways

• The periodic table can be used to predict how atoms are likely to bond.
• Metal + non-metal usually suggests ionic bonding.
• Non-metal + non-metal usually suggests covalent bonding.
• Ionic bonding involves electron transfer and electrostatic attraction between oppositely charged ions.
• Covalent bonding involves sharing pairs of electrons.
• Electron arrangements help explain why atoms gain, lose, or share electrons.
• Group 1 metals commonly form +1 ions.
• Group 2 metals commonly 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.
• Ionic compounds must have an overall charge of zero.
• Ion charges can therefore be used to predict the ratio of ions in a compound.
• Na⁺ and Cl⁻ produce NaCl.
• Mg²⁺ and Cl⁻ produce MgCl₂.
• Al³⁺ and O²⁻ produce Al₂O₃.
• Chemical formulas should use the simplest whole-number ratio.
• Transition-metal compounds require additional care because many transition metals have multiple oxidation states.
• For simple covalent compounds, typical bonding capacities can help predict formulas such as H₂O, NH₃, CH₄, and CO₂.
• Bonding is ultimately more complex than a strict ionic/covalent division, but the metal/non-metal rule is a powerful introductory prediction tool.
• For unfamiliar compounds, use:

periodic-table position → metal/non-metal → valence electrons → electron gain, loss, or sharing → bond type → ion charges or bonding capacity → chemical formula.