Molecular Structure and Intermolecular Forces
5. Physical Properties of Molecular Substances
Learning outcomes
- I can explain how intermolecular forces affect melting and boiling points.
- I can describe factors influencing evaporation.
- I can compare the physical properties of molecular substances.
- I can relate molecular structure to physical behaviour.
- I can predict properties from intermolecular interactions.
From Molecular Structure to Physical Properties
Molecular substances are made of individual molecules held together by intermolecular forces.
Within each molecule, atoms are joined by strong covalent bonds. Between the molecules, weaker intermolecular attractions act.
These intermolecular forces strongly influence physical properties such as:
- Melting point
- Boiling point
- Physical state
- Evaporation rate
- Volatility
- Viscosity
- Surface tension
- Solubility
A useful relationship is:
molecular structure → intermolecular forces → physical properties
Molecular Substances
A molecular substance consists of separate molecules.
Examples include:
- H₂O
- CO₂
- CH₄
- NH₃
- O₂
- I₂
- Ethanol
- Glucose
The atoms within each molecule are held together by covalent bonds.
The molecules themselves interact through intermolecular forces.
This distinction is essential when explaining physical properties.
Intermolecular Forces
The main intermolecular attractions considered for molecular substances are:
London dispersion forces
Present between all atoms and molecules.
They result from temporary fluctuations in electron distribution.
Dipole-dipole attractions
Occur between polar molecules.
The partially positive region of one molecule attracts the partially negative region of another.
Hydrogen bonding
A particularly strong intermolecular attraction occurring when hydrogen is bonded to:
- N
- O
- F
Water is an important example.
Strength of Intermolecular Forces
For small molecules of broadly similar size, a useful general pattern is:
London dispersion < dipole-dipole < hydrogen bonding
However, this is not an absolute rule.
London dispersion forces can become very strong in large molecules containing many electrons.
Therefore, when comparing molecular substances, we should consider:
- Polarity
- Hydrogen bonding
- Molecular size
- Number of electrons
- Molecular shape
Melting
Melting is the change:
solid → liquid
In a molecular solid, molecules are held in relatively fixed positions by intermolecular attractions.
When the substance is heated:
- Molecules gain kinetic energy.
- Molecular vibrations increase.
- Intermolecular attractions become less able to hold molecules in fixed positions.
- Molecules begin moving past one another.
The substance becomes a liquid.
Melting Does Not Normally Break Covalent Bonds
When a molecular substance melts, the molecules themselves generally remain intact.
For example:
H₂O(s) → H₂O(l)
Ice contains H₂O molecules.
Liquid water also contains H₂O molecules.
The O–H covalent bonds are not normally broken during melting.
Instead, the arrangement and intermolecular interactions between water molecules change.
Melting Point
The melting point is the temperature at which a solid changes into a liquid under a specified pressure.
Stronger intermolecular attractions often require more energy to overcome.
Therefore, stronger intermolecular forces can contribute to:
higher melting points
However, melting points can be difficult to predict from intermolecular force strength alone because the way molecules pack together in a solid crystal is also important.
Molecular symmetry and shape can strongly influence crystal structure.
Boiling
Boiling is the change:
liquid → gas
For molecules to enter the gas phase, they must become sufficiently separated from neighboring molecules.
Energy is required to overcome intermolecular attractions.
Therefore:
stronger intermolecular forces generally → higher boiling point
Boiling Point and Molecular Attractions
Compare two similarly sized substances.
Substance A
Weak intermolecular forces.
Molecules separate relatively easily.
Expected:
lower boiling point
Substance B
Strong intermolecular forces.
More energy is needed to separate the molecules.
Expected:
higher boiling point
This relationship is one of the most useful ways to connect microscopic molecular structure with observable properties.
Boiling Does Not Break Molecules Apart
When water boils:
H₂O(l) → H₂O(g)
The molecules separate from one another.
The O–H covalent bonds remain intact.
Therefore, boiling mainly involves overcoming:
intermolecular forces
not:
covalent bonds
If the covalent bonds were broken, water would undergo a chemical change rather than simply boiling.
Example: Methane and Water
Consider:
CH₄
and:
H₂O
Methane is:
- Nonpolar
- Small
- Mainly affected by relatively weak London dispersion forces
Water is:
- Polar
- Capable of hydrogen bonding
Therefore, water molecules experience much stronger intermolecular attractions.
This helps explain why methane is a gas at room temperature while water is a liquid.
Hydrogen Bonding and Boiling Point
Hydrogen bonding can have a major effect on boiling point.
Water molecules form extensive hydrogen bonds with neighboring molecules.
A significant amount of energy is required to separate water molecules sufficiently for them to enter the gas phase.
This contributes to water having a much higher boiling point than might be expected for such a small molecule.
Molecular Size and Boiling Point
Polarity is not the only factor affecting intermolecular forces.
Molecular size also matters.
Larger molecules generally have:
- More electrons.
- Larger electron clouds.
- More easily distorted electron distributions.
This can produce stronger London dispersion forces.
Therefore, within a similar family of molecules:
larger molecules often have higher boiling points.
Example: The Halogens
Consider:
F₂, Cl₂, Br₂, I₂
These molecules are nonpolar.
Their main intermolecular attractions are London dispersion forces.
Moving down the group:
- Molecular size increases.
- Number of electrons increases.
- Polarizability increases.
- London dispersion forces become stronger.
This contributes to their different physical states around room temperature:
| Substance | Relative molecular size | Typical state near room temperature |
|---|---|---|
| F₂ | Smallest | Gas |
| Cl₂ | Small | Gas |
| Br₂ | Larger | Liquid |
| I₂ | Largest | Solid |
Molecular Shape Matters
The shape of a molecule can influence how closely molecules approach one another.
Long or relatively flat molecules may sometimes have larger areas of contact than compact molecules.
Greater contact can strengthen London dispersion interactions.
Therefore:
molecular shape can influence boiling point and other physical properties.
This is particularly useful when comparing molecules with similar molecular masses.
Evaporation
Evaporation occurs when molecules escape from the surface of a liquid into the gas phase.
Unlike boiling, evaporation can occur:
below the boiling point
and only occurs at the:
surface of the liquid
Why Do Some Molecules Evaporate?
Molecules in a liquid have different kinetic energies.
Some molecules at the surface have enough energy to overcome intermolecular attractions and escape.
These molecules enter the gas phase.
Therefore, evaporation depends partly on:
- Intermolecular force strength.
- Temperature.
- Surface area.
- Air movement.
- Humidity.
Intermolecular Forces and Evaporation
If intermolecular forces are weak:
molecules escape more easily
Therefore:
evaporation tends to be faster.
If intermolecular forces are strong:
more energy is required for molecules to escape
Therefore:
evaporation tends to be slower.
So, in general:
stronger intermolecular forces → slower evaporation
Temperature and Evaporation
Increasing temperature increases the average kinetic energy of molecules.
At a higher temperature, a larger fraction of surface molecules have enough energy to escape.
Therefore:
higher temperature → faster evaporation
This is why wet clothing usually dries more quickly in warm conditions.
Surface Area and Evaporation
Evaporation occurs at the liquid surface.
Increasing the surface area exposes more molecules to the air.
Therefore:
larger surface area → faster evaporation
For example, water spread across a wide tray usually evaporates faster than the same volume of water in a narrow container.
Air Movement and Evaporation
Molecules that evaporate accumulate in the air above a liquid.
Moving air carries some of these vapor molecules away.
This helps maintain conditions that allow continued evaporation.
Therefore:
greater air movement often → faster evaporation
This helps explain why clothes dry faster on a windy day.
Humidity and Evaporation
Humidity describes the amount of water vapor present in the air.
When humidity is high, the air already contains a large amount of water vapor.
Net evaporation from wet surfaces tends to be slower.
When humidity is low, evaporation can occur more readily.
Therefore:
lower humidity generally → faster evaporation of water
Factors Affecting Evaporation
| Factor | Effect on Evaporation |
|---|---|
| Higher temperature | Faster |
| Larger surface area | Faster |
| Greater air movement | Usually faster |
| Lower humidity | Faster for water |
| Weaker intermolecular forces | Faster |
These factors explain many familiar observations.
Volatility
Volatility describes how readily a substance vaporizes.
A highly volatile liquid evaporates readily.
Highly volatile molecular substances generally have:
- Relatively weak intermolecular forces.
- Relatively high vapor pressure.
- Relatively low boiling points.
Therefore:
higher volatility generally corresponds to easier evaporation.
Vapor Pressure
Some molecules continuously escape from a liquid surface.
In a closed container, some gas molecules also return to the liquid.
Eventually, a dynamic equilibrium can develop between:
evaporation ⇌ condensation
The gas above the liquid produces a vapor pressure.
Liquids with weaker intermolecular forces generally have higher vapor pressures at the same temperature.
Connecting Vapor Pressure and Boiling Point
A liquid boils when its vapor pressure becomes equal to the external pressure.
Therefore, substances with weaker intermolecular attractions often:
- Develop higher vapor pressures more easily.
- Reach the boiling condition at lower temperatures.
This gives the general relationship:
weaker intermolecular forces → higher vapor pressure → lower boiling point
and:
stronger intermolecular forces → lower vapor pressure → higher boiling point
Viscosity
Viscosity is a fluid's resistance to flow.
Water has relatively low viscosity compared with substances such as:
- Honey
- Syrup
- Thick oils
Stronger intermolecular attractions can make it more difficult for molecules to move past one another.
Therefore, stronger intermolecular forces can contribute to:
greater viscosity
However, molecular size and shape are also important.
Molecular Shape and Viscosity
Long molecules may become entangled with one another.
This can make them resist flowing past each other.
Therefore, viscosity depends on both:
- Intermolecular attractions.
- Molecular structure.
Large, long molecules can sometimes produce highly viscous liquids even when other factors are similar.
Surface Tension
Surface tension is the tendency of a liquid surface to resist being stretched or disrupted.
Molecules inside a liquid experience attractions from molecules around them.
Surface molecules experience an overall inward attraction.
Strong intermolecular attractions can therefore produce high surface tension.
Water has relatively high surface tension because of hydrogen bonding.
Physical State at Room Temperature
Intermolecular forces help determine whether a molecular substance is:
- Solid.
- Liquid.
- Gas.
If intermolecular attractions are very weak, molecules may separate easily and the substance may be a gas.
Stronger attractions can allow molecules to remain together as liquids or solids.
However, physical state also depends on:
- Temperature.
- Pressure.
- Molecular size.
- Molecular shape.
Comparing Molecular Substances
Consider three hypothetical molecular substances:
| Substance | Main Intermolecular Attraction | Relative Attraction | Expected Boiling Point |
|---|---|---|---|
| A | Weak London forces | Weak | Low |
| B | Dipole-dipole | Moderate | Higher |
| C | Hydrogen bonding | Strong | Often highest |
This prediction assumes the molecules are reasonably similar in size.
If Substance A were dramatically larger than B or C, its London dispersion forces could become very significant.
Predicting Properties from Structure
When given a molecular structure, use this process:
Determine whether the molecule is polar
↓
Identify the intermolecular forces
↓
Consider hydrogen bonding
↓
Consider molecular size and number of electrons
↓
Consider molecular shape
↓
Compare overall intermolecular attraction
↓
Predict physical properties
Worked Example: CH₄ and NH₃
Methane, CH₄
Shape:
tetrahedral
Polarity:
nonpolar
Main intermolecular attraction:
London dispersion
Expected properties:
- Low boiling point.
- High volatility.
- Gas at room temperature.
Ammonia, NH₃
Shape:
trigonal pyramidal
Polarity:
polar
Important intermolecular attraction:
hydrogen bonding
Expected:
- Stronger intermolecular attraction than methane.
- Higher boiling point than methane.
Worked Example: H₂O and H₂S
Water and hydrogen sulfide have somewhat similar molecular structures.
However:
H₂O forms strong hydrogen bonds.
H₂S does not form comparably strong conventional hydrogen bonds.
Therefore, water has much stronger intermolecular attractions.
This contributes to water having a much higher boiling point.
The difference cannot be explained simply by molecular mass.
Worked Example: Br₂ and Cl₂
Both:
Br₂
and:
Cl₂
are nonpolar.
Both rely mainly on London dispersion forces.
However, Br₂:
- Has more electrons.
- Has a larger electron cloud.
- Is more polarizable.
Therefore:
Br₂ has stronger London dispersion forces.
This contributes to Br₂ having a higher boiling point than Cl₂.
Comparing Evaporation
Suppose three similar liquids have the following intermolecular forces:
Liquid A: weak London forces
Liquid B: dipole-dipole attractions
Liquid C: strong hydrogen bonding
At the same temperature, we would generally expect:
A to evaporate fastest
and:
C to evaporate slowest
because the molecules in C require more energy to escape from the liquid.
Comparing Boiling Points
Suppose two molecules have similar molecular masses.
Molecule X is nonpolar.
Molecule Y can form hydrogen bonds.
We would generally predict:
Molecule Y has the higher boiling point.
Reason:
Molecule Y has stronger intermolecular attractions, so more energy is required to separate its molecules.
Real-World Connection: Sweating
Sweating cools the body through evaporation.
The fastest-moving water molecules can escape from the skin.
Energy is required to overcome intermolecular attractions between water molecules.
That energy comes partly from the skin.
As higher-energy molecules leave, the average kinetic energy of the remaining surface decreases.
The skin cools.
Real-World Connection: Drying Clothes
Several factors affect how quickly clothes dry.
Warm weather
Higher temperature increases molecular kinetic energy.
Wind
Moving air carries water vapor away.
Low humidity
Dry air allows greater net evaporation.
Spreading clothes out
Greater surface area exposes more water to the air.
This is why clothes generally dry faster when they are:
warm + spread out + exposed to moving, relatively dry air
Real-World Connection: Perfume
Many perfume ingredients are relatively volatile.
They evaporate from the skin or clothing and enter the air.
Gas molecules then spread through the surroundings.
For a fragrance to reach your nose, at least some of its molecules must enter the gas phase.
The volatility of the substances therefore influences how a fragrance behaves.
Real-World Connection: Refrigeration
Refrigeration systems use substances that repeatedly:
- Evaporate.
- Condense.
During evaporation, the refrigerant absorbs energy.
During condensation, energy is released elsewhere.
The physical properties of the refrigerant must therefore be suitable for these repeated phase changes.
Intermolecular Forces and Solubility
Physical properties are interconnected.
The same molecular features that affect boiling point and evaporation can also affect solubility.
Polar molecules can often interact strongly with polar solvents.
Nonpolar molecules often interact more favorably with nonpolar solvents.
Therefore, understanding molecular structure can help predict several properties at once.
A Useful Property Pattern
For comparable molecular substances:
stronger intermolecular forces
often mean:
↑ boiling point
↓ volatility
↓ evaporation rate
↓ vapor pressure at a given temperature
and sometimes:
↑ viscosity
↑ surface tension
Melting point may also increase, but crystal packing makes melting-point trends less straightforward.
Structure-to-Property Reasoning
A strong chemistry explanation should connect several steps.
Instead of saying:
“Water has a high boiling point because it is polar.”
A stronger explanation is:
Water molecules are polar and can form hydrogen bonds with neighboring water molecules. These relatively strong intermolecular attractions require considerable energy to overcome, contributing to water's relatively high boiling point.
This reasoning pattern can be used for many molecular substances.
Example of a Full Prediction
Suppose a molecule:
- Is small.
- Is nonpolar.
- Cannot form hydrogen bonds.
We predict that its main intermolecular attractions will be relatively weak London dispersion forces.
Therefore, compared with a similarly sized molecule capable of hydrogen bonding, it will probably have:
- Lower boiling point.
- Higher volatility.
- Faster evaporation.
- Higher vapor pressure.
This is how molecular structure can be used to predict macroscopic properties.
Common Mistakes
Saying Boiling Breaks Covalent Bonds
Boiling molecular substances normally overcomes intermolecular forces, not covalent bonds within the molecules.
Assuming Molecular Mass Is the Only Factor
Polarity, hydrogen bonding and molecular shape also matter.
Assuming Hydrogen Bonding Occurs in Every Molecule Containing Hydrogen
For the molecules studied here, hydrogen bonding generally requires H to be directly bonded to:
N, O, or F
Saying Stronger Forces Mean Faster Evaporation
The opposite is generally true.
Stronger intermolecular forces → slower evaporation
Confusing Evaporation and Boiling
Evaporation occurs at the surface and can occur below the boiling point.
Boiling occurs throughout the liquid when the vapor pressure equals the external pressure.
Assuming Melting Point Always Follows Boiling-Point Trends
Crystal packing can strongly affect melting point.
Forgetting London Dispersion Forces
All molecular substances experience London dispersion forces.
Assuming Large Nonpolar Molecules Always Have Weak Attractions
Large electron clouds can produce substantial London dispersion forces.
Check Your Understanding
1. What holds atoms together within a molecular substance?
2. What holds separate molecules together?
3. Name three important intermolecular forces.
4. Explain why stronger intermolecular forces generally produce higher boiling points.
5. What happens to intermolecular attractions when a molecular substance boils?
6. Are covalent bonds normally broken during boiling? Explain.
7. Why does hydrogen bonding increase the boiling point of many substances?
8. Explain why water has a much higher boiling point than methane.
9. Why do London dispersion forces generally become stronger as molecular size increases?
10. Explain why Br₂ has stronger intermolecular forces than Cl₂.
11. Define evaporation.
12. Explain how intermolecular force strength affects evaporation rate.
13. How does increasing temperature affect evaporation?
14. Why does increasing surface area increase evaporation?
15. Explain why wind can make wet clothes dry faster.
16. How does humidity affect the evaporation of water?
17. Define volatility.
18. How are volatility and boiling point generally related?
19. Explain how intermolecular forces affect vapor pressure.
20. Why can molecular shape affect physical properties?
21. Explain how intermolecular forces can influence viscosity.
22. Why does water have relatively high surface tension?
23. A small nonpolar molecule and a similarly sized hydrogen-bonding molecule are compared. Predict which will have the higher boiling point and explain why.
24. A liquid evaporates extremely quickly at room temperature. What might this suggest about its intermolecular forces?
25. Explain the relationship:
molecular structure → intermolecular forces → physical properties
Key Terms
- Molecular substance – substance consisting of separate covalently bonded molecules.
- Intermolecular force – attraction between separate molecules.
- London dispersion force – intermolecular attraction caused by temporary fluctuations in electron distribution.
- Dipole-dipole attraction – attraction between permanent dipoles of polar molecules.
- Hydrogen bonding – relatively strong intermolecular attraction involving H bonded to N, O, or F.
- Melting point – temperature at which a solid becomes a liquid at a specified pressure.
- Boiling point – temperature at which the vapor pressure of a liquid equals the external pressure.
- Evaporation – escape of molecules from the surface of a liquid.
- Volatility – tendency of a substance to vaporize readily.
- Vapor pressure – pressure produced by vapor above a liquid in a closed system.
- Viscosity – resistance of a fluid to flow.
- Surface tension – tendency of a liquid surface to resist disruption.
- Polarizability – ease with which an electron cloud can be distorted.
- Physical property – characteristic that can be observed or measured without changing a substance's chemical identity.
Key Takeaways
- Molecular substances consist of separate molecules.
- Covalent bonds hold atoms together within molecules.
- Intermolecular forces act between molecules.
- Intermolecular forces strongly influence physical properties.
- Stronger intermolecular attractions generally lead to higher boiling points.
- Melting points are influenced by intermolecular forces as well as molecular packing.
- Boiling and melting normally do not break covalent bonds within molecules.
- London dispersion forces occur in all molecular substances.
- Larger, more polarizable molecules generally experience stronger London dispersion forces.
- Polar molecules can experience dipole-dipole attractions.
- Suitable molecules containing H bonded to N, O, or F can form hydrogen bonds.
- Hydrogen bonding can strongly influence physical properties.
- Evaporation occurs at the surface of a liquid and can occur below its boiling point.
- Weak intermolecular forces generally produce faster evaporation and greater volatility.
- Temperature, surface area, air movement and humidity can affect evaporation.
- Stronger intermolecular forces generally produce lower vapor pressure at a given temperature.
- Molecular size and shape can influence intermolecular attractions.
- Intermolecular forces can affect viscosity and surface tension.
- Molecular structure can therefore be used to predict observable physical behaviour.
- Understanding physical properties requires connecting molecular structure to intermolecular forces and then connecting those forces to the behaviour of the substance.