What Is Metallic Bonding?

Metals have a distinctive type of bonding called metallic bonding.

In a solid metal, the atoms are arranged closely together in a regular structure. Their outer electrons become delocalized, meaning they are not associated with just one particular atom.

This produces a structure consisting of:

• positive metal ions arranged in a lattice
• delocalized electrons moving throughout the structure

Metallic bonding is the strong electrostatic attraction between positive metal ions and delocalized electrons.

This simple model helps explain many characteristic properties of metals.

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The Metallic Lattice

A piece of metal contains an enormous number of metal particles packed closely together.

The positive metal ions occupy positions in a regular three-dimensional arrangement called a lattice.

The delocalized electrons move throughout this lattice.

The attraction between the positive ions and negative electrons holds the metal together.

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From Metal Atoms to Metallic Bonding

Consider a simplified metal atom with outer-shell electrons.

In metallic bonding, these outer electrons become delocalized.

They are no longer considered to belong to one particular atom.

The remaining metal particles can therefore be represented as positive ions surrounded by mobile electrons.

The structure is sometimes described as:

positive ions in a sea of delocalized electrons

This is useful as a simplified model, although the real electronic structure of metals is more sophisticated.

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What Does Delocalized Mean?

Localized means restricted to a particular location.

Delocalized means spread throughout a larger structure rather than belonging to one particular atom or bond.

In metals, the outer electrons are delocalized throughout the metallic lattice.

They can move through the structure.

This mobility is extremely important because it helps explain:

• electrical conductivity
• thermal conductivity

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Electrostatic Attraction

Positive and negative charges attract each other.

The positive metal ions attract the negatively charged delocalized electrons.

This attraction occurs throughout the structure rather than between just one pair of atoms.

Metallic bonding is therefore often described as non-directional compared with many covalent bonds.

This helps explain why layers of metal ions can change position without the entire bonding structure immediately breaking apart.

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Metallic Bonding Is Strong

In many metals, there is strong attraction between:

• the positive ions
• the delocalized electrons

A large amount of energy may therefore be required to separate the particles.

This contributes to properties such as:

• strength
• relatively high melting points for many metals
• relatively high boiling points for many metals

However, metallic bond strength varies between different metals.

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What Affects Metallic Bond Strength?

The strength of metallic bonding depends on several factors, including:

• charge of the metal ions
• number of delocalized electrons
• size of the metal ions
• distance between positive ions and electrons

A stronger electrostatic attraction generally produces stronger metallic bonding.

This can help explain differences in physical properties among metals.

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Example: Sodium and Magnesium

Sodium and magnesium provide a useful simplified comparison.

Sodium contributes approximately one outer electron per atom to the delocalized electron system.

Magnesium contributes approximately two.

Magnesium forms ions with a greater positive charge and contributes more delocalized electrons.

The electrostatic attraction in magnesium is therefore stronger than in sodium in this simplified model.

This contributes to magnesium having stronger metallic bonding and a higher melting point than sodium.

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Electrical Conductivity

One of the most important properties of metals is that they conduct electricity.

Why?

Because the delocalized electrons can move through the metallic structure.

When a potential difference is applied across a metal, the mobile electrons acquire a net drift through the structure.

This movement of charge produces an electric current.

Therefore:

mobile delocalized electrons → electrical conductivity

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Metals Conduct as Solids

This is an important difference between metals and ionic compounds.

Metals conduct electricity while solid because their delocalized electrons can move.

In a solid ionic compound, the ions are fixed in the lattice and cannot move freely.

Therefore:

solid metal → conducts

solid ionic compound → generally does not conduct

When an ionic compound is molten or dissolved in water, its ions can move and it can conduct.

The mobile charge carriers are different:

Metal → electrons

Molten/aqueous ionic compound → ions

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Example: Copper Wiring

Copper is widely used for electrical wiring.

Its usefulness is connected to several properties:

• high electrical conductivity
• ductility
• relatively good strength
• practical durability

Metallic bonding helps explain both its conductivity and its ability to be drawn into wires.

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Thermal Conductivity

Metals are also generally good conductors of thermal energy.

When one part of a metal is heated, energy can be transferred through the structure.

Delocalized electrons can carry energy through the metal, and vibrations of the lattice also contribute to thermal conduction.

This is why metals are commonly used in:

• cooking pans
• heat exchangers
• radiators
• cooling systems

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Why Are Metals Malleable?

Malleability is the ability of a material to be hammered, pressed, or rolled into different shapes without breaking.

Many metals are highly malleable.

Examples include:

• aluminium
• copper
• gold

The metallic bonding model helps explain this property.

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Layers Can Slide

When force is applied to a metal, layers of positive ions can move relative to one another.

In many metals, the structure can rearrange while the delocalized electrons continue to attract the positive ions.

The bonding is therefore maintained as the layers shift.

This allows the metal to change shape rather than immediately shattering.

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Why Doesn't the Metal Simply Break?

Metallic bonding is not restricted to fixed, directional bonds between particular pairs of atoms.

The delocalized electrons continue interacting with surrounding positive ions even after the ions change position.

This allows many metals to undergo substantial deformation while maintaining cohesion.

This is very different from the behaviour of many brittle ionic crystals.

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Metals Compared with Ionic Solids

An ionic crystal contains alternating positive and negative ions.

If layers shift far enough, ions with the same charge may become positioned near one another.

Like charges repel.

The crystal can fracture.

A metal behaves differently because the positive ions remain surrounded by delocalized electrons as layers shift.

Therefore:

many ionic solids → brittle

many metals → malleable

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What Is Ductility?

Ductility is the ability of a material to be drawn into a wire.

Copper is a familiar example.

The metal can be stretched and reshaped without immediately breaking because its structure can rearrange while metallic bonding continues to hold the material together.

Therefore:

mobile layers + continuing metallic attraction → ductility

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Malleability vs Ductility

These terms are related but not identical.

Malleability

Ability to be hammered or rolled into sheets.

Example:

aluminium foil

Ductility

Ability to be drawn into wires.

Example:

copper electrical wire

Both properties are connected to the ability of the metallic structure to deform without immediately losing cohesion.

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Why Are Many Metals Strong?

The strong electrostatic attraction between positive ions and delocalized electrons holds the metallic lattice together.

This can make many metals strong materials.

However, not all metals are equally strong.

For example:

• sodium is soft
• aluminium is stronger
• iron is much stronger in many practical situations

Metal strength depends on metallic bonding, crystal structure, defects, processing, temperature, and whether the material is a pure metal or an alloy.

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Melting Metals

To melt a metal, enough energy must be supplied to disrupt the ordered solid structure and overcome enough of the attractions holding it together so that the particles can move past one another as a liquid.

Metals with stronger metallic bonding often have higher melting points.

However, melting points vary considerably among metals.

For example:

• mercury is liquid at room temperature
• sodium has a relatively low melting point
• iron has a much higher melting point
• tungsten has an extremely high melting point

So "metals have high melting points" is a useful general trend, not an absolute rule.

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Metallic Bonding and Metal Properties

The metallic bonding model connects microscopic structure with observable properties.

Structure:

Positive metal ions + delocalized electrons

This helps explain:

Electrical conductivity

Electrons can move and carry charge.

Thermal conductivity

Mobile electrons and lattice vibrations transfer energy.

Strength

Electrostatic attraction holds the lattice together.

Malleability

Layers can shift while metallic attraction remains.

Ductility

The structure can deform while maintaining cohesion.

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Using Structure to Explain Properties

In chemistry, it is not enough to simply state:

"Metals conduct electricity."

A stronger explanation connects the property to the structure.

For example:

Copper conducts electricity because it contains delocalized electrons that are mobile throughout the metallic lattice and can carry electric charge.

This follows the reasoning:

structure → particle behaviour → observable property

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Example 1: Explain Electrical Conductivity

Question:

Why does aluminium conduct electricity?

Weak answer:

"Because aluminium is a metal."

Better answer:

"Aluminium has metallic bonding with delocalized electrons that can move through the lattice."

Strong answer:

"Aluminium contains positive metal ions surrounded by mobile delocalized electrons. When a potential difference is applied, these electrons can move through the structure and carry electric charge."

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Example 2: Explain Malleability

Question:

Why can gold be hammered into thin sheets?

Answer:

Gold has metallic bonding. Its positive metal ions are surrounded by delocalized electrons. Layers of ions can shift relative to one another while the attraction between the positive ions and delocalized electrons continues to hold the structure together.

Therefore, gold can change shape without immediately fracturing.

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Example 3: Explain Ductility

Question:

Why can copper be drawn into wires?

Answer:

Copper's metallic structure can deform as layers of ions change position. The delocalized electrons continue attracting the positive ions, so the metallic bonding remains effective as the metal is reshaped.

Therefore, copper is ductile.

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Example 4: Explain Strength

Question:

Why can many metals resist large forces?

Answer:

There is strong electrostatic attraction between positive metal ions and the delocalized electrons throughout the metallic lattice.

A significant amount of energy may therefore be required to separate the particles.

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Metallic Bonding and Aluminium

Aluminium demonstrates several useful metallic properties.

Aluminium is:

• electrically conductive
• thermally conductive
• malleable
• ductile
• relatively strong for its density

This contributes to its use in:

• aircraft
• electrical cables
• drink cans
• foil
• construction materials

Its low density is also important and is not explained by metallic bonding alone.

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Metallic Bonding and Copper

Copper combines:

• excellent electrical conductivity
• excellent thermal conductivity
• high ductility
• good malleability

These properties make it useful in:

• electrical wires
• motors
• generators
• plumbing
• electronics
• heat exchangers

The metallic bonding model explains several of these properties using the same underlying structure.

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Metallic Bonding and Iron

Iron has strong metallic bonding and is widely used as the basis of steels.

Pure iron and different steels do not have identical properties, but iron-based materials are important because they can combine:

• strength
• toughness
• workable mechanical properties
• useful thermal and electrical behaviour

Alloying can significantly modify these properties.

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Pure Metals and Alloys

A pure metal contains atoms of one metallic element.

An alloy contains a metal mixed with one or more other elements.

Examples include:

• steel
• stainless steel
• brass
• bronze

Alloys still involve metallic bonding, but their structures are less uniform than those of pure metals.

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Why Can Alloys Be Harder?

In a pure metal, similarly sized atoms can form relatively regular layers.

These layers may slide past one another comparatively easily.

In an alloy, atoms of different sizes can disrupt the regular arrangement.

This can make it more difficult for layers and defects in the crystal to move.

Therefore, alloys can be:

• harder
• stronger
• less easily deformed

than the pure metal.

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Example: Steel

Steel is primarily iron containing carbon and often other elements.

Carbon atoms affect the iron structure and can hinder the movement of defects through the lattice.

This can make steel harder and stronger than pure iron, depending on its composition and processing.

This is why alloys are so important in engineering.

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Metallic Bonding vs Ionic Bonding

Metallic bonding

• positive metal ions
• delocalized electrons
• electrons can move
• conducts as a solid
• often malleable
• often ductile

Ionic bonding

• positive and negative ions
• electrons are not freely mobile through the solid lattice
• solid usually does not conduct
• often brittle
• conducts when molten or aqueous because ions can move

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Metallic Bonding vs Covalent Bonding

In a simple metallic model, electrons are delocalized throughout the metal.

In covalent bonding, electrons are shared in bonds between particular atoms or across a network structure.

These different arrangements help produce very different physical properties.

For example:

Copper → metallic → excellent conductor

Diamond → giant covalent → extremely hard but poor electrical conductor

Sodium chloride → ionic → brittle and non-conducting as a solid

Structure strongly influences properties.

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Why Metals Are Shiny

Many metals have a characteristic metallic lustre.

Their electronic structure allows them to interact strongly with incoming electromagnetic radiation, including visible light, and much of that light can be reflected.

This gives polished metals their shiny appearance.

Metallic lustre is another property related to the electronic structure of metals.

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Why Metals Ring When Struck

Many metals are sonorous, meaning they produce a ringing sound when struck.

The combination of:

• strong bonding
• elasticity
• ability to transmit vibrations

allows vibrations to travel through the metal.

This property has historically contributed to the use of metals in:

• bells
• cymbals
• musical instruments

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Metallic Bonding in Engineering

Metallic bonding gives metals combinations of properties that are extremely valuable in engineering.

Engineers can select metals based on required properties.

Electrical cable:

needs conductivity + ductility

Bridge:

needs strength + toughness

Car body:

needs strength + formability

Cooking pan:

needs thermal conductivity

Aircraft:

needs low density + strength

No single metal is ideal for every application.

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Using the Metallic Bonding Model

When explaining a metal property, use this sequence:

1. Describe the structure.

Positive metal ions arranged in a lattice.

2. Describe the electrons.

Outer electrons are delocalized.

3. Describe what can move.

Electrons can move throughout the structure; ions can shift positions during deformation.

4. Identify the electrostatic attraction.

Positive ions are attracted to the delocalized electrons.

5. Connect this to the property.

For example:

mobile electrons → conductivity

strong attraction → strength

layers shifting while bonding remains → malleability and ductility

This produces a complete scientific explanation.

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Worked Analysis 1: Electrical Cable

A company needs a material for an electrical cable.

Why might copper be suitable?

Copper contains delocalized electrons that can move through the metallic lattice.

Therefore, it conducts electricity well.

Copper is also ductile because its metallic structure can deform while metallic attraction remains.

Therefore, it can be drawn into long, thin wires.

This connects two useful properties to metallic bonding.

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Worked Analysis 2: Aluminium Foil

Why can aluminium be rolled into very thin sheets?

Aluminium has metallic bonding.

Its positive ions are surrounded by delocalized electrons.

When pressure is applied, the metallic structure can deform and layers can shift while the delocalized electrons continue to maintain attraction between the positive ions.

Therefore, aluminium is malleable.

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Worked Analysis 3: Why Use an Alloy?

Suppose a pure metal is too soft for an engineering structure.

Adding atoms of different sizes can disrupt the regular lattice.

This can make movement of defects and deformation more difficult.

The resulting alloy may therefore be harder and stronger.

The engineer has modified the microscopic structure to change the macroscopic property.

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Common Mistakes

Mistake 1: Saying metals contain free protons

They do not.

Metallic bonding involves positive metal ions and delocalized electrons.

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Mistake 2: Saying the positive ions move through a wire to carry electricity

In solid metals, electrical conduction is primarily due to mobile electrons.

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Mistake 3: Saying electrons leave the metal during metallic bonding

The electrons are delocalized within the metallic structure. They are not simply lost from the entire piece of metal.

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Mistake 4: Saying metallic bonding is attraction between positive ions

Positive ions repel one another.

The bonding comes from electrostatic attraction between the positive ions and delocalized electrons.

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Mistake 5: Saying metals conduct because their atoms move

Electrical conduction is primarily due to the movement of delocalized electrons.

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Mistake 6: Saying metals are malleable because the bonds break easily

A better explanation is that the structure can deform while metallic attraction continues to hold it together.

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Mistake 7: Assuming all metals are very hard

Some metals, such as sodium, are quite soft.

Metallic bonding strength and structure vary between metals.

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Mistake 8: Assuming all metals have high melting points

Many do, but there are important exceptions.

Mercury is liquid at room temperature, and several metals have relatively low melting points.

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Did You Know?

Gold is extraordinarily malleable.

It can be hammered into extremely thin sheets called gold leaf.

Copper, meanwhile, can be drawn into long thin wires.

These two familiar applications demonstrate two different consequences of metallic bonding:

gold leaf → malleability

copper wire → ductility

Both properties arise because metallic structures can undergo considerable deformation while remaining bonded.

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Key Terms

• Metallic bonding: Strong electrostatic attraction between positive metal ions and delocalized electrons.
• Delocalized electron: Electron that is not associated with one particular atom or bond and can move throughout the metallic structure.
• Metallic lattice: Regular arrangement of positive metal ions surrounded by delocalized electrons.
• Electrostatic attraction: Attraction between opposite electric charges.
• Electrical conductivity: Ability of a material to allow electric charge to move through it.
• Thermal conductivity: Ability of a material to transfer thermal energy.
• Malleability: Ability to be hammered or rolled into sheets.
• Ductility: Ability to be drawn into wires.
• Alloy: Mixture containing a metal and one or more other elements.
• Lustre: Shiny appearance characteristic of many metals.

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Structure–Property Connections

Positive metal ions + delocalized electrons

↓

Strong electrostatic attraction

↓

Strong metallic structure

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Mobile delocalized electrons

↓

Charge can move

↓

Electrical conductivity

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Mobile electrons + lattice vibrations

↓

Energy transfer

↓

Thermal conductivity

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Layers can shift while metallic attraction remains

↓

Structure can deform without immediately fracturing

↓

Malleability and ductility

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Different-sized atoms disrupt regular layers

↓

Movement through the structure becomes more difficult

↓

Many alloys are harder than pure metals

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Key Takeaways

• Metals are held together by metallic bonding.
• A metallic structure consists of positive metal ions and delocalized electrons.
• Metallic bonding is the electrostatic attraction between these positive ions and delocalized electrons.
• Delocalized electrons are not associated with one particular atom.
• These electrons can move throughout the metallic structure.
• Mobile electrons allow metals to conduct electricity.
• Delocalized electrons also contribute to thermal conductivity.
• Strong electrostatic attraction contributes to the strength of many metals.
• Metallic bond strength varies among metals.
• Many metals have relatively high melting and boiling points because significant energy is required to disrupt their structures, although there are important exceptions.
• Metals are malleable because their structures can deform while metallic attraction continues to hold the ions and electrons together.
• Metals are ductile for the same general reason.
• Malleability allows metals to be formed into sheets.
• Ductility allows metals to be drawn into wires.
• Metallic bonding explains why copper is useful for electrical wiring.
• Metallic bonding helps explain why aluminium can be rolled into foil.
• Alloys can be harder than pure metals because differently sized atoms disrupt the regular structure and hinder deformation.
• The metallic bonding model allows microscopic structure to explain macroscopic properties.
• A strong explanation should follow:

structure → delocalized electrons → electrostatic attraction/mobility → particle behaviour → observable property.