Properties of Metals

Metallic bonding helps explain why many metals have:

• High electrical conductivity.
• High thermal conductivity.
• Malleability.
• Ductility.
• Strength.
• Relatively high melting points.

Different metals vary considerably, but these are common metallic properties.

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Why Metals Conduct Electricity

The delocalized electrons can move throughout the metallic structure.

When a potential difference is applied, the electrons develop an overall drift.

Therefore:

mobile electrons → movement of charge → electrical current

Metals can conduct while solid because the electrons are already mobile.

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Why Metals Are Malleable

Layers of positive metal ions can move relative to one another.

The delocalized electrons remain throughout the structure.

Therefore, the electrostatic attraction between the ions and electrons can continue even after the layers move.

The metal can change shape without immediately shattering.

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Comparing What Happens to Electrons

The three bonding types can be distinguished by what happens to valence electrons.

Ionic

electrons transferred

Covalent

electrons shared between atoms

Metallic

electrons delocalized throughout the structure

A useful summary is:

Ionic → transfer

Covalent → share

Metallic → delocalize

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Comparing the Particles

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Comparing Typical Element Combinations

A useful first clue is the types of elements present.

Ionic

Usually:

metal + nonmetal

Example:

MgO

Covalent

Usually:

nonmetal + nonmetal

Example:

CO₂

Metallic

Metal atoms only, including mixtures of metals in alloys.

Example:

Cu

These rules are useful for introductory classification, although real chemical bonding can be more complex than three completely separate categories.

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Identifying Bonding from Formulas

Consider:

NaCl

Na is a metal.

Cl is a nonmetal.

Therefore:

ionic

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Consider:

CO₂

C is a nonmetal.

O is a nonmetal.

Therefore:

covalent

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Consider:

Mg

Magnesium is a metal element.

Therefore:

metallic

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Consider:

MgCl₂

Mg is a metal.

Cl is a nonmetal.

Therefore:

ionic

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Consider:

CH₄

Carbon and hydrogen are nonmetals.

Therefore:

covalent

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A Quick Classification Method

When given a simple chemical formula:

Step 1: Identify the elements.

Step 2: Determine whether each element is a metal or nonmetal.

Then use:

metal + nonmetal → usually ionic

nonmetal + nonmetal → usually covalent

metal only → metallic

This method works well for many introductory examples.

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Important Exception: Polyatomic Ions

Some ionic compounds contain groups of covalently bonded atoms.

For example:

NaNO₃

The compound contains:

Na⁺ and NO₃⁻

The attraction between these ions is ionic.

However, the atoms inside the nitrate ion are covalently bonded.

Therefore, a substance can contain more than one type of bonding.

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Comparing Melting Points

Ionic substances

Often high because strong electrostatic attractions must be overcome.

Simple molecular substances

Often relatively low because melting mainly involves overcoming intermolecular attractions.

Giant covalent substances

Often very high because many strong covalent bonds are involved in the structure.

Metals

Often moderate to high, although values vary widely, because metallic attractions hold the lattice together.

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Comparing Electrical Conductivity

Graphite is a major exception because it contains mobile delocalized electrons.

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Comparing Mechanical Properties

Ionic

Often:

hard and brittle

Simple molecular

Often:

soft or easily melted, depending on the substance.

Giant covalent

Can be:

extremely hard

as in diamond.

Metallic

Often:

strong, malleable and ductile

These properties arise from differences in bonding and structure.

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Bonding and Solubility

Solubility is more complicated than simply identifying the bond type.

Many ionic compounds dissolve in water, but not all do.

Many molecular substances have different solubilities depending on their polarity.

A useful principle for molecular substances is:

like dissolves like

Polar substances tend to dissolve more readily in polar solvents.

Nonpolar substances tend to dissolve more readily in nonpolar solvents.

Therefore, bonding and molecular structure both influence solubility.

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Worked Example: NaCl

Formula:

NaCl

Elements:

sodium = metal

chlorine = nonmetal

Classification:

ionic

Expected properties:

• Crystalline solid.
• Relatively high melting point.
• Brittle.
• Does not conduct as a solid.
• Conducts when molten.

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Worked Example: CO₂

Formula:

CO₂

Elements:

carbon = nonmetal

oxygen = nonmetal

Classification:

covalent molecular

Expected properties:

• Exists as separate molecules.
• Relatively weak attractions between molecules.
• Low melting and boiling temperatures compared with ionic lattices.
• Poor electrical conductivity.

At room conditions, carbon dioxide is a gas.

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Worked Example: Copper

Symbol:

Cu

Copper is a metal.

Classification:

metallic

Expected properties include:

• High electrical conductivity.
• High thermal conductivity.
• Malleability.
• Ductility.

This makes copper especially useful for electrical wiring.

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Worked Example: Diamond

Diamond contains only carbon.

Carbon is a nonmetal.

Bonding:

covalent

However, diamond does not consist of separate C molecules.

Instead, it forms a:

giant covalent structure

Therefore, diamond has very different properties from simple molecular substances.

Its strong three-dimensional network makes it extremely hard.

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Worked Example: Magnesium Oxide

Formula:

MgO

Magnesium:

metal

Oxygen:

nonmetal

Therefore:

ionic bonding

Magnesium transfers electrons and forms Mg²⁺ ions.

Oxygen gains electrons and forms O²⁻ ions.

The oppositely charged ions attract in a giant ionic lattice.

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Worked Example: Aluminum

Aluminum contains only metal atoms.

Therefore:

metallic bonding

Its structure contains positive metal ions surrounded by delocalized electrons.

This helps explain why aluminum:

• Conducts electricity.
• Conducts thermal energy.
• Is malleable.
• Is ductile.

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Worked Example: Which Conducts as a Solid?

Consider:

NaCl, Cu and CO₂

NaCl

Ions cannot move in the solid lattice.

Does not conduct.

Cu

Contains mobile delocalized electrons.

Conducts.

CO₂

Does not contain mobile charged particles.

Does not conduct.

Therefore, copper is the conductor when these substances are in their ordinary solid/molecular forms.

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Worked Example: Predicting Brittleness

Compare sodium chloride and copper.

Sodium chloride

If layers shift, like-charged ions may become aligned.

Repulsion causes the lattice to fracture.

Therefore:

brittle

Copper

Layers can move while attraction between metal ions and delocalized electrons remains.

Therefore:

malleable and ductile

The difference in mechanical properties can be explained by bonding.

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Similarities Among the Bond Types

Although ionic, covalent and metallic bonding differ, they also have similarities.

All involve:

• Electrons.
• Electrostatic forces.
• Interactions involving atoms or ions.
• Arrangements that can lower the energy of the system compared with separated particles.

All three can form stable substances.

The major differences involve how the electrons and charged particles are arranged.

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Bonding Is a Model

The categories ionic, covalent and metallic are extremely useful models.

However, real bonding is not always perfectly divided into three completely separate types.

For example, many bonds have both:

ionic character + covalent character

Bonding is better understood as a continuum in more advanced chemistry.

For introductory chemistry, however, classifying substances as predominantly ionic, covalent or metallic is extremely useful.

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From Bonding to Properties

One of the most important skills in chemistry is explaining properties from structure.

Instead of memorizing:

“metals conduct electricity”

explain:

metals contain mobile delocalized electrons → electrons carry charge → metals conduct electricity

Instead of memorizing:

“ionic compounds are brittle”

explain:

layers shift → like charges align → electrostatic repulsion occurs → lattice fractures

This connects microscopic structure to observable behavior.

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Structure → Property → Application

Bonding can ultimately explain why materials are useful.

Copper

metallic bonding

↓

mobile electrons + movable layers

↓

conductive + ductile

↓

electrical wire

Sodium chloride

ionic lattice

↓

strong attractions between ions

↓

high melting point

Diamond

giant covalent network

↓

strong bonds throughout structure

↓

extreme hardness

↓

cutting and abrasive applications

This type of reasoning is central to materials chemistry.

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Common Mistakes

Saying Ionic Bonds Involve Shared Electrons

Ionic bonding involves attraction between oppositely charged ions formed after electron transfer.

Saying Covalent Bonds Involve Electron Transfer

Covalent bonds involve shared electron pairs.

Saying Metals Have No Bonds

Metals are held together by strong metallic bonding.

Saying All Covalent Substances Have Low Melting Points

Giant covalent structures can have extremely high melting temperatures.

Saying Ionic Solids Conduct Electricity

Their ions are charged but cannot move while held in the solid lattice.

Saying Ionic Compounds Are Made of Molecules

They generally form giant ionic lattices rather than individual molecules.

Saying All Covalent Substances Are Electrical Insulators

Graphite is an important exception.

Classifying a Substance from Its Properties Alone

Use its composition and structure whenever possible.

Forgetting That Some Substances Contain More Than One Bond Type

Compounds containing polyatomic ions can contain both ionic and covalent bonding.

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Check Your Understanding

1. Name the three major bonding types discussed here.

2. What happens to electrons during ionic bonding?

3. What happens to electrons during covalent bonding?

4. What happens to electrons during metallic bonding?

5. What combination of elements commonly produces ionic bonding?

6. What combination commonly produces covalent bonding?

7. Describe the structure of a metal.

8. Classify NaCl as ionic, covalent or metallic.

9. Classify CO₂.

10. Classify Cu.

11. Classify MgO.

12. Classify H₂O.

13. Why do ionic compounds often have high melting points?

14. Why does solid NaCl not conduct electricity?

15. Why does molten NaCl conduct electricity?

16. Why does copper conduct electricity while solid?

17. Why are many ionic solids brittle?

18. Why are many metals malleable?

19. Why do many simple molecular substances have relatively low boiling points?

20. Why does diamond have very different properties from many other covalent substances?

21. Why can graphite conduct electricity?

22. Explain why NaNO₃ contains both ionic and covalent bonding.

23. A substance is a solid conductor that can be drawn into wire. Which bonding type is most likely?

24. A substance is brittle, has a high melting point and conducts when molten but not when solid. Which bonding type is most likely?

25. Explain how bonding can be used to predict material properties.

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Key Terms

• Ionic bond – electrostatic attraction between oppositely charged ions.
• Covalent bond – attraction involving a shared pair of electrons between atoms.
• Metallic bond – electrostatic attraction between positive metal ions and delocalized electrons.
• Ion – charged particle formed when an atom or group gains or loses electrons.
• Cation – positively charged ion.
• Anion – negatively charged ion.
• Valence electron – electron in the outer energy level of an atom.
• Delocalized electron – electron able to move throughout a larger structure.
• Ionic lattice – repeating three-dimensional arrangement of oppositely charged ions.
• Molecule – discrete group of atoms joined by covalent bonds.
• Giant covalent structure – large network of atoms connected by covalent bonds.
• Intermolecular force – attraction acting between molecules.
• Electrical conductivity – ability of a substance to allow electrical charge to move.
• Brittle – likely to fracture with relatively little plastic deformation.
• Malleable – able to be hammered, pressed or rolled into shape.
• Ductile – able to be drawn into wire.

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Key Takeaways

• Ionic, covalent and metallic bonding all involve electrostatic interactions involving charged particles.
• Ionic bonding usually occurs between metals and nonmetals.
• Ionic bonding involves electron transfer followed by attraction between oppositely charged ions.
• Covalent bonding usually occurs between nonmetals.
• Covalent bonds contain shared pairs of electrons.
• Metallic bonding involves positive metal ions and delocalized electrons.
• Ionic compounds usually form giant ionic lattices.
• Covalent substances can form individual molecules or giant covalent structures.
• Metals form giant metallic structures.
• Ionic substances often have high melting points because of strong electrostatic attractions.
• Ionic solids do not conduct because their ions cannot move.
• Molten ionic substances conduct because their ions become mobile.
• Metals conduct because they contain mobile delocalized electrons.
• Metals are often malleable and ductile because layers can move while metallic attraction remains.
• Simple molecular substances often have relatively low melting and boiling points because intermolecular attractions are weaker than the covalent bonds within molecules.
• Giant covalent substances can have very different properties from molecular substances.
• Graphite is an important covalent substance that conducts electricity.
• Some compounds contain more than one type of bonding.
• Chemical formulas and structures can often be used to identify the dominant bonding type.
• Bonding models allow us to explain and predict material properties.
• A useful summary is:

ionic → electrons transferred

covalent → electrons shared

metallic → electrons delocalized