Structure Determines Properties
2. Bonding and Physical Properties
Learning outcomes
- I can explain how bonding affects melting point and boiling point.
- I can compare conductivity among different substances.
- I can predict physical properties from bonding type.
- I can explain why substances with different structures behave differently.
- I can use bonding to explain observed material properties.
Why Bonding Affects Physical Properties
The physical properties of a substance depend strongly on:
- The particles it contains.
- How those particles are arranged.
- The forces or bonds between the particles.
- Whether charged particles are able to move.
This means that properties such as melting point, boiling point, electrical conductivity, hardness, brittleness, malleability, and physical state can often be predicted by examining bonding and structure.
A useful reasoning pattern is:
bonding → structure → forces between particles → physical properties
Melting and Boiling
Melting and boiling require particles to move farther apart or move more freely.
Melting
solid → liquid
Boiling
liquid → gas
Energy must be supplied to overcome some of the attractive forces holding particles together.
Therefore, a useful general principle is:
stronger attractions that must be overcome → more energy required → higher melting or boiling temperature
However, we must identify which attractions are actually being overcome.
That depends on the structure of the substance.
Ionic Substances
Ionic substances consist of:
positive ions + negative ions
arranged in a giant ionic lattice.
Strong electrostatic attractions act between oppositely charged ions throughout the structure.
Examples include:
- NaCl
- MgO
- CaCl₂
- KBr
Melting Ionic Substances
When an ionic solid melts, ions must become able to move relative to one another.
This requires enough energy to overcome a significant amount of the strong electrostatic attraction within the lattice.
Therefore, ionic compounds generally have:
relatively high melting points
and:
relatively high boiling points
Comparing Ionic Bond Strength
Not all ionic compounds have the same melting point.
The strength of electrostatic attraction depends partly on:
- The charges of the ions.
- The sizes of the ions.
- The distances between them.
Greater ionic charges generally produce stronger attractions.
For example, ions carrying charges of 2+ and 2− can experience stronger electrostatic attractions than similarly sized ions carrying 1+ and 1− charges.
Therefore, ionic structure allows us to explain general trends, although actual melting points also depend on the details of the crystal structure.
Electrical Conductivity of Ionic Substances
Ionic substances contain charged particles.
However, having charged particles is not enough to conduct electricity.
The charged particles must also be:
mobile
Solid ionic compound
The ions are held in fixed lattice positions.
Therefore:
no mobile charge carriers → does not conduct
Molten ionic compound
The ions are free to move.
Therefore:
mobile ions → conducts
Ionic compound dissolved in water
If the compound dissolves and produces mobile ions:
mobile dissolved ions → conducts
This gives an important principle:
electrical conductivity requires mobile charged particles.
Simple Molecular Covalent Substances
Many covalent substances consist of individual molecules.
Examples include:
- H₂O
- CO₂
- CH₄
- NH₃
- O₂
Atoms within each molecule are joined by strong covalent bonds.
However, separate molecules are held near one another by weaker:
intermolecular forces
Melting Molecular Substances
When a molecular solid melts, the molecules do not normally break apart.
Instead, enough intermolecular attraction is overcome to allow the molecules to move more freely.
For example:
H₂O(s) → H₂O(l)
The H₂O molecules remain H₂O molecules.
The O–H covalent bonds remain intact.
It is mainly the interactions between molecules that change.
Boiling Molecular Substances
The same principle applies when a molecular liquid boils.
For example:
H₂O(l) → H₂O(g)
Water molecules move farther apart.
The covalent bonds within the molecules remain intact.
Therefore, the boiling point depends strongly on the:
intermolecular forces between molecules
Why Many Molecular Substances Have Low Melting and Boiling Points
Intermolecular forces are generally much weaker than ionic or covalent bonds.
Therefore, many simple molecular substances require relatively little energy to separate their molecules.
This often results in:
- Low melting points.
- Low boiling points.
- Gases or liquids at room temperature.
However, intermolecular forces vary greatly in strength.
Intermolecular Forces Matter
Different molecular substances can have very different boiling points because their intermolecular attractions differ.
Important intermolecular forces include:
- London dispersion forces.
- Dipole-dipole attractions.
- Hydrogen bonding.
For comparable molecules:
stronger intermolecular attraction → higher boiling point
Example: Water and Methane
Water, H₂O, and methane, CH₄, are both small molecular substances.
However:
Methane
- Nonpolar.
- Mainly London dispersion forces.
Water
- Polar.
- Forms hydrogen bonds.
Water therefore has much stronger intermolecular attractions.
This contributes to water having a much higher boiling point than methane.
Molecular Size Matters Too
London dispersion forces generally become stronger as molecular size and polarizability increase.
Consider:
F₂ → Cl₂ → Br₂ → I₂
All are nonpolar molecules.
However, their electron clouds become larger and more polarizable down the group.
London dispersion forces become stronger.
Their physical states near room temperature change:
- F₂ – gas
- Cl₂ – gas
- Br₂ – liquid
- I₂ – solid
This shows that molecular substances do not all have identical physical properties.
Conductivity of Molecular Substances
Most simple molecular substances are poor electrical conductors.
Their electrons are generally localized in covalent bonds, and they do not contain freely moving ions or electrons.
Therefore:
no mobile charged particles → poor electrical conductivity
There are exceptions when substances react or ionize in solution, so the actual conductivity of a solution depends on the particles present.
Giant Covalent Structures
Not every covalent substance consists of separate molecules.
Some form enormous networks of covalently bonded atoms.
These are called:
giant covalent structures
Examples include:
- Diamond.
- Graphite.
- Silicon dioxide.
Their properties differ greatly from those of simple molecular substances.
Diamond
Diamond consists of carbon atoms joined in a giant three-dimensional covalent network.
Each carbon atom is covalently bonded to four other carbon atoms.
Strong covalent bonds extend throughout the structure.
A very large amount of energy is required to disrupt this network.
Therefore, diamond has:
- Extremely high hardness.
- Very high thermal stability.
- Very high sublimation/melting conditions.
- Poor electrical conductivity under ordinary conditions.
Why Diamond Does Not Conduct Electricity
The valence electrons in diamond are involved in covalent bonds.
There are no freely moving charged particles available to carry an electrical current.
Therefore:
diamond does not normally conduct electricity
despite being made entirely of carbon.
Graphite
Graphite is also made entirely of carbon.
However, its structure is different.
Each carbon atom forms covalent bonds with three other carbon atoms.
This produces layers.
Some electrons are delocalized and can move along these layers.
Therefore, graphite:
conducts electricity
This demonstrates an important principle:
bond type alone is not always enough — structure also matters.
Comparing Diamond and Graphite
| Property | Diamond | Graphite |
|---|---|---|
| Element | Carbon | Carbon |
| Main bonding | Covalent | Covalent |
| Structure | 3D network | Layers |
| Delocalized electrons | No mobile electron system like graphite | Yes |
| Electrical conductivity | Poor | Conducts along layers |
| Mechanical behavior | Extremely hard | Layers slide relatively easily |
Both contain carbon.
Both contain strong covalent bonds.
Their different structures produce very different physical properties.
Metallic Substances
Metals have a giant metallic structure consisting of:
positive metal ions + delocalized electrons
The electrostatic attraction between these particles produces metallic bonding.
Melting Metals
To melt a metal, enough energy must be supplied for the ordered metallic structure to lose its fixed arrangement.
Metallic attractions must be sufficiently disrupted for the ions to move relative to one another.
Many metals therefore have:
relatively high melting points
However, metallic bond strength varies significantly.
Therefore, metal melting points vary widely.
Mercury, for example, is liquid at room temperature.
Why Metals Conduct Electricity
The delocalized electrons in metals are mobile.
When a potential difference is applied, these electrons develop an overall drift through the metallic structure.
Therefore:
mobile electrons → electrical conductivity
Unlike ionic compounds, metals can conduct while solid.
Why Metals Conduct Thermal Energy
Delocalized electrons also help transfer thermal energy rapidly through the metallic structure.
Vibrations of the lattice contribute as well.
Therefore, metals are generally:
good thermal conductors
This is why metals are commonly used in:
- Cooking equipment.
- Heat exchangers.
- Radiators.
- Electronic heat sinks.
Why Metals Are Malleable
Metallic bonding also explains why many metals can be shaped.
Layers of positive metal ions can move relative to one another.
The delocalized electrons continue attracting the positive ions.
Therefore:
layers move → metallic attraction remains → metal changes shape
This produces malleability.
Why Metals Are Ductile
The same basic structure allows metals to be drawn into wires.
As the structure changes shape, metallic attraction continues throughout the material.
Therefore, many metals are:
ductile
This is one reason copper is useful for electrical wiring.
Comparing the Main Structures
| Structure | Particles | Main Attraction | Typical Melting Point | Electrical Conductivity |
|---|---|---|---|---|
| Ionic | Positive and negative ions | Electrostatic attraction | High | Solid: no; molten: yes |
| Simple molecular | Molecules | Intermolecular forces between molecules | Usually low | Usually no |
| Giant covalent | Atoms | Covalent bonds | Very high | Usually no, but exceptions |
| Metallic | Positive ions + delocalized electrons | Metallic attraction | Often moderate to high | Yes |
These are general patterns, not absolute rules.
Predicting Properties from Bonding
Suppose you are given an unknown substance.
It has:
- High melting point.
- Brittle crystals.
- No conductivity when solid.
- Conductivity when molten.
This strongly suggests:
Why?
high melting point → strong lattice attractions
brittle → characteristic ionic lattice behavior
molten conductivity → mobile ions
Prediction Example: A Metal
An unknown solid:
- Conducts electricity.
- Conducts thermal energy.
- Can be hammered into sheets.
- Can be drawn into wires.
These properties strongly suggest:
metallic bonding
The mobile electrons explain conductivity.
The non-directional metallic attraction helps explain malleability and ductility.
Prediction Example: Molecular Substance
An unknown substance:
- Melts at a low temperature.
- Boils at a relatively low temperature.
- Does not conduct electricity.
- Exists as separate molecules.
These observations suggest:
simple molecular covalent structure
Relatively weak intermolecular forces explain the low melting and boiling temperatures.
Prediction Example: Giant Covalent Substance
An unknown material:
- Has extremely high thermal stability.
- Is extremely hard.
- Does not conduct electricity.
A possible explanation is:
giant covalent structure
Strong covalent bonds extend throughout the material.
Diamond is a classic example.
Using Experimental Evidence
Chemists can sometimes use physical properties to help identify the likely structure of an unknown substance.
Useful tests include:
- Melting point.
- Electrical conductivity as a solid.
- Electrical conductivity when molten.
- Solubility.
- Hardness.
- Brittleness.
- Malleability.
No single property should always be used by itself.
A combination of evidence gives a stronger conclusion.
Worked Example: Substance A
Observations:
- Solid at room temperature.
- Melting point is high.
- Brittle.
- Does not conduct as a solid.
- Conducts when molten.
Reasoning
Conductivity when molten suggests mobile charged particles.
Lack of solid conductivity suggests those particles are fixed in the solid.
Brittleness is consistent with an ionic lattice.
Conclusion
Substance A is likely:
ionic
Worked Example: Substance B
Observations:
- Solid.
- Shiny.
- Conducts electricity.
- Can be bent and shaped.
- Conducts thermal energy.
Reasoning
Electrical conductivity suggests mobile charge carriers.
Malleability suggests layers can move while bonding remains.
Conclusion
Substance B is likely:
metallic
Worked Example: Substance C
Observations:
- Gas at room temperature.
- Very low boiling point.
- Does not conduct electricity.
Reasoning
A very low boiling point suggests relatively weak attractions between particles.
Lack of conductivity indicates no mobile charged particles.
Conclusion
Substance C is likely:
simple molecular
Worked Example: Substance D
Observations:
- Extremely hard.
- Very high thermal stability.
- Does not conduct electricity.
- Consists only of carbon.
A likely structure is:
diamond
Its giant covalent network explains its hardness and thermal stability.
The lack of mobile charged particles explains its poor electrical conductivity.
Bonding and Physical State
Bonding can even help explain whether a substance is likely to be a solid, liquid, or gas at room temperature.
Strong extended bonding
Ionic, metallic and giant covalent structures are usually solids at room temperature.
Molecular substances
These can be:
- Solids.
- Liquids.
- Gases.
Their state depends strongly on the strength of intermolecular forces and molecular size.
For example:
CH₄ → gas
H₂O → liquid
I₂ → solid
All are molecular substances.
Structure Matters as Much as Bond Type
It is not enough to say:
“The substance is covalent.”
You should ask:
What kind of covalent structure?
A covalent substance could be:
simple molecular
or:
giant covalent
These structures have dramatically different physical properties.
Similarly, two substances with the same element can behave differently because of different structures.
Diamond and graphite are the classic example.
A Better Way to Explain Properties
Weak answer:
“NaCl has a high melting point because it is ionic.”
Better answer:
“NaCl forms a giant ionic lattice containing oppositely charged ions. Strong electrostatic attractions act throughout the lattice, so considerable energy is required to allow the ions to move freely. Therefore, NaCl has a high melting point.”
The second answer connects:
structure → attraction → energy → property
Another Strong Explanation
Weak answer:
“Copper conducts because it is a metal.”
Better answer:
“Copper has a metallic structure containing delocalized electrons. These electrons can move through the structure and carry electrical charge, so copper conducts electricity.”
Again:
structure → mobile particles → observed property
Comparing Bonding and Properties
Ionic
Structure: giant ionic lattice
Strong attraction: between positive and negative ions
Typical result:
- High melting point.
- High boiling point.
- Brittle.
- Conducts when molten or dissolved.
Simple Molecular
Structure: individual molecules
Attractions between molecules: intermolecular forces
Typical result:
- Lower melting and boiling points.
- Usually poor electrical conductivity.
Giant Covalent
Structure: enormous network of covalent bonds
Typical result:
- Very high thermal stability.
- Often very hard.
- Usually poor electrical conductivity.
Metallic
Structure: positive ions and delocalized electrons
Typical result:
- Good electrical conductivity.
- Good thermal conductivity.
- Malleability.
- Ductility.
- Often relatively high melting point.
Real-World Connection: Electrical Cables
Copper is used inside many electrical cables because:
metallic bonding → delocalized electrons → high conductivity
The outside is usually a polymer because:
covalent molecular structure → no freely moving charged particles → electrical insulation
Different bonding structures allow two materials to perform completely different jobs in the same product.
Real-World Connection: Cookware
Metals are useful in cookware because they transfer thermal energy efficiently.
However, the handle may be made from a polymer or another poor thermal conductor.
Therefore:
metal body → thermal conductor
polymer handle → thermal insulator
Material properties determine how each part of the product is designed.
Real-World Connection: Cutting Tools
Diamond is extremely hard because of its giant covalent structure.
Strong covalent bonds extend throughout the three-dimensional network.
Diamond or diamond-like materials can therefore be useful in:
- Cutting.
- Grinding.
- Drilling.
- Abrasive surfaces.
The application follows directly from the structure and bonding.
Common Mistakes
Saying Strong Bonds Always Mean High Boiling Point
For molecular substances, boiling normally overcomes intermolecular forces, not the covalent bonds inside molecules.
Saying All Covalent Substances Have Low Melting Points
Giant covalent structures can have extremely high thermal stability.
Saying Ionic Solids Conduct Because They Contain Ions
The ions must be able to move.
Solid ionic compounds generally do not conduct.
Saying Metals Conduct Because the Positive Ions Move
The mobile charge carriers in solid metals are delocalized electrons.
Saying All Molecular Substances Are Gases
Molecular substances can be solids, liquids, or gases.
Assuming All Metals Have High Melting Points
Metallic bond strengths vary considerably.
Forgetting Graphite
Graphite is covalent but conducts electricity because it contains delocalized electrons.
Looking Only at Bond Type
The structure of the substance is also essential.
Check Your Understanding
1. Why does bonding affect physical properties?
2. What must generally happen to particles when a substance melts?
3. Why do ionic substances often have high melting points?
4. Why does solid NaCl not conduct electricity?
5. Why does molten NaCl conduct?
6. What forces are mainly overcome when a simple molecular substance boils?
7. Are covalent bonds normally broken when water boils?
8. Why do many simple molecular substances have low boiling points?
9. Why does water have a higher boiling point than methane?
10. Why does molecular size affect London dispersion forces?
11. What is a giant covalent structure?
12. Why is diamond extremely hard?
13. Why does diamond not normally conduct electricity?
14. Why can graphite conduct electricity?
15. Explain why graphite and diamond have different properties even though both contain only carbon.
16. Why do metals conduct electricity?
17. Why are metals generally good thermal conductors?
18. Explain why metals are malleable.
19. An unknown substance conducts when molten but not when solid. What bonding type might it have?
20. An unknown solid conducts electricity and can be drawn into wire. What bonding type is likely?
21. A substance has a very low boiling point and does not conduct electricity. What structure might it have?
22. Why should several physical properties be considered when identifying an unknown substance?
23. Explain the relationship:
bonding → structure → physical properties
Key Terms
- Physical property – characteristic that can be observed or measured without changing chemical identity.
- Melting point – temperature at which a solid becomes a liquid at a specified pressure.
- Boiling point – temperature at which a liquid's vapor pressure equals the external pressure.
- Ionic lattice – giant repeating structure of oppositely charged ions.
- Simple molecular substance – substance consisting of individual covalently bonded molecules.
- Giant covalent structure – large network of atoms connected by covalent bonds.
- Metallic structure – positive metal ions surrounded by delocalized electrons.
- Intermolecular force – attraction between separate molecules.
- Delocalized electron – electron able to move through a larger structure.
- Electrical conductivity – ability to allow electrical charge to move.
- Thermal conductivity – ability to transfer thermal energy.
- Charge carrier – mobile charged particle responsible for electrical conduction.
- Brittle – tending to fracture with relatively little plastic deformation.
- Malleable – able to be hammered, rolled, or pressed into shape.
- Ductile – able to be drawn into wire.
Key Takeaways
- Physical properties arise from both bonding and structure.
- Strong attractions generally require more energy to overcome.
- Ionic substances contain strong electrostatic attractions throughout a giant lattice.
- Ionic substances therefore commonly have high melting and boiling points.
- Ionic solids do not conduct because their ions cannot move.
- Molten ionic compounds conduct because their ions are mobile.
- Simple molecular substances contain strong covalent bonds within molecules but weaker intermolecular forces between molecules.
- Melting and boiling molecular substances generally involve overcoming intermolecular attractions rather than covalent bonds.
- Stronger intermolecular forces generally produce higher boiling points.
- Giant covalent structures contain strong covalent bonds throughout large networks.
- Diamond's structure explains its extreme hardness and poor electrical conductivity.
- Graphite conducts because it contains mobile delocalized electrons.
- Metals contain positive ions and delocalized electrons.
- Mobile electrons allow metals to conduct electricity.
- Metallic structure also helps explain thermal conductivity, malleability, and ductility.
- Different structures can produce very different properties even when substances contain the same element.
- Observed physical properties can provide evidence about bonding and structure.
- Predictions are strongest when several properties are considered together.
- A strong scientific explanation should connect:
bonding → structure → forces/mobile particles → observed physical property