F. Gases and Atmospheric Chemistry · Gases and Atmospheric ChemistryStates of Matter and Intermolecular Forces
F3.2 F2.1About 70 minutes
Learning goals- I can describe solids, liquids and gases using the arrangement and motion of their particles.
- I can tell intramolecular bonds apart from intermolecular forces.
- I can identify London dispersion, dipole–dipole and hydrogen bonding forces.
- I can use intermolecular forces to explain and predict boiling points and states at room temperature.
- I can name the changes of state and say whether each absorbs or releases energy.
Words to know
intermolecular force
An attraction between neighbouring molecules. Much weaker than a covalent bond.
intramolecular force
A force inside a molecule, such as a covalent bond, holding its atoms together.
London dispersion force
A weak attraction caused by temporary dipoles from moving electrons. Present in all molecules; stronger with more electrons.
dipole–dipole force
The attraction between the δ+ end of one polar molecule and the δ− end of another.
hydrogen bond
A strong dipole–dipole attraction between an H atom bonded to N, O or F and a lone pair on N, O or F of a neighbouring molecule.
sublimation / deposition
Solid changing directly to gas / gas changing directly to solid.
vaporization
A liquid changing to a gas, by evaporation (at the surface) or boiling (throughout).
Why is air invisible but a puddle isn't?
On a cold January morning in Ontario you might see three states of matter at once: ice on the driveway, liquid water in a puddle by the warm garage, and the water vapour in the air that you cannot see at all. They are all H₂O. The molecules are identical. What changes is how far apart the molecules are, how fast they move and how strongly they pull on each other.
This unit is about gases: how they behave, the laws that predict them, and the air we breathe. To understand gases, we first need to understand why some substances are gases at room temperature and others are not. The answer is intermolecular forces.
Remember from earlier grades
In Grade 8 (Fluids) you used the particle theory: all matter is made of tiny particles that are always moving, with spaces between them, and particles move faster when heated. You saw that gases are much more compressible than liquids. In the unit Matter, Chemical Trends, and Chemical Bonding of this course you learned about electronegativity and polar covalent bonds. Today we put those two ideas together.
The three states, particle by particle
- Solid: particles are packed closely in fixed positions. They vibrate but do not move past each other. A solid has a definite shape and a definite volume.
- Liquid: particles are still close together (so liquids are almost incompressible), but they can slide past each other. A liquid has a definite volume but takes the shape of its container.
- Gas: particles are far apart compared with their own size and move quickly in straight lines until they hit something. A gas has no definite shape or volume. It fills any container and can be squeezed.
When water boils, 18 mL of liquid becomes more than 20 L of steam at 100 °C. The molecules did not get bigger. The space between them grew by more than a thousand times.
Two kinds of forces: inside and between
Intramolecular forces are the bonds inside a molecule (covalent bonds). They are strong. Breaking them is a chemical change.
Intermolecular forces (IMFs) are attractions between neighbouring molecules. They are much weaker. Overcoming them is a physical change: melting, boiling and evaporating. When water boils, the bubbles are H₂O molecules, not hydrogen and oxygen gas. No O–H bonds break.
There are three main intermolecular forces in this course:
- London dispersion forces. Electrons are always moving. For an instant, more of them may be on one side of a molecule, making a tiny temporary dipole. That dipole pulls on the electrons of a neighbour, making another. Every molecule has London forces. They are stronger in molecules with more electrons (larger molecules), because larger electron clouds are easier to distort.
- Dipole–dipole forces. A polar molecule (like HCl) has a permanent partial negative end (δ−) and a partial positive end (δ+). The δ+ end of one molecule attracts the δ− end of the next.
- Hydrogen bonding. A strong kind of dipole–dipole force. It happens when hydrogen is bonded directly to N, O or F (very electronegative atoms). The tiny, very positive H is attracted to a lone pair on an N, O or F atom of a neighbouring molecule.
How intermolecular forces decide the state
At any temperature, molecules have kinetic energy that tends to pull them apart, and intermolecular forces that hold them together.
- If the IMFs are weak compared with the kinetic energy at room temperature, the substance is a gas (N₂, O₂, CH₄, CO₂).
- If they are moderate, it is a liquid (H₂O, ethanol, Br₂).
- If they are strong, it is a solid (I₂, sugar). Ionic compounds like NaCl are solids for a different reason: they are held by very strong ionic bonds throughout the crystal.
So stronger intermolecular forces → higher melting and boiling points. The halogens show London forces in action: all are non-polar, but the number of electrons increases down the group.
| Substance | Molar mass (g/mol) | Strongest IMF | Boiling point (°C) |
|---|
| CH₄ (methane) | 16 | London only | −162 |
| NH₃ (ammonia) | 17 | hydrogen bonding | −33 |
| H₂O (water) | 18 | hydrogen bonding | 100 |
| H₂S (hydrogen sulfide) | 34 | dipole–dipole | −60 |
Changes of state
Every change of state has a name. Energy is absorbed (endothermic) when particles move apart and released (exothermic) when they come together.
- Melting (solid → liquid) and freezing (liquid → solid)
- Vaporization (liquid → gas; evaporation at the surface below the boiling point, boiling throughout the liquid) and condensation (gas → liquid)
- Sublimation (solid → gas directly; dry ice, or snow slowly disappearing on a cold sunny day) and deposition (gas → solid directly; frost forming on a car windshield on a clear winter night)
During a change of state, the temperature of a pure substance stays constant: the added energy goes into overcoming intermolecular forces, not into making particles move faster.
Worked exampleExample 1: Which has the higher boiling point?
Predict which substance in each pair has the higher boiling point, and explain using intermolecular forces: (a) Ne or Ar; (b) ethanol, CH₃CH₂OH, or dimethyl ether, CH₃OCH₃ (both C₂H₆O).
- (a) Neon and argon are single atoms, so the only IMF is London dispersion.
- Argon has 18 electrons; neon has 10. More electrons → stronger London forces.
- (b) The two molecules have the same formula, so the same number of electrons and similar London forces.
- Ethanol has an O–H bond, so it can form hydrogen bonds. In dimethyl ether the O is bonded only to C atoms; there is no H on the O, so it cannot hydrogen bond to itself (it has weaker dipole–dipole forces).
(a) Argon (−186 °C versus −246 °C for neon). (b) Ethanol (78 °C versus −24 °C for dimethyl ether). Same atoms, different arrangement, a 100-degree difference.
Worked exampleExample 2: Name the change and the energy flow
On a February night, frost forms on a car windshield. In the morning sun, some of the frost turns straight into water vapour. Name both changes and say whether each absorbs or releases energy.
- Frost forms from water vapour without becoming liquid first: gas → solid is deposition.
- Molecules come together and IMFs form, so energy is released (exothermic).
- Frost turning directly to vapour: solid → gas is sublimation.
- Molecules move apart and IMFs are overcome, so energy is absorbed (endothermic).
Deposition (releases energy), then sublimation (absorbs energy).
Hands-on activityEvaporation race
You need: a grown-up, a dark plate or baking sheet, a dropper or straw, a timer, water, rubbing alcohol (isopropyl alcohol from a pharmacy), vegetable oil, and a cotton ball.
Safety: rubbing alcohol is flammable and poisonous if swallowed. Work in a room with an open window, far from any flame, stove or heater. A grown-up handles the bottle. Do not taste anything; wash your hands afterwards.
- Predict: which liquid will evaporate fastest? Which slowest? Write your reasons using intermolecular forces.
- Put one drop of each liquid on the plate, well apart. Start the timer.
- Check every 2 minutes for 20 minutes. Record when each drop disappears.
- Dab a little water on the back of one hand and a little rubbing alcohol on the back of the other with the cotton ball. Which feels colder?
What to look for: the rubbing alcohol drop vanishes first (often within a few minutes), water takes much longer, and the oil barely changes. Isopropyl alcohol has only one O–H group for its size, so its hydrogen bonding is weaker than water's network. Oil molecules are very large; their London forces are strong enough that they hardly evaporate at room temperature. The alcohol feels colder because evaporation absorbs energy from your skin, and it evaporates faster.
Just for funTeacher: "What is the chemical formula for water?"
Student: "H, I, J, K, L, M, N, O."
Teacher: "Excuse me?"
Student: "You said it was H to O!"
Your turn
1Describe the arrangement and motion of particles in a gas.
Particles are far apart compared with their size, move quickly in random straight lines, and collide with each other and the container walls.
2What is the difference between an intramolecular force and an intermolecular force?
An intramolecular force (a covalent bond) holds atoms together inside a molecule; an intermolecular force is an attraction between separate molecules. Intramolecular forces are much stronger.
3Which intermolecular force is present in all molecules?
London dispersion forces.
4Which molecules can form hydrogen bonds with each other: HF, CH₄, CH₃OH, H₂S?
HF (H bonded to F) and CH₃OH (H bonded to O). CH₄ has only C–H bonds, and S is not electronegative enough, so H₂S has dipole–dipole forces instead.
5Name the change of state: dry ice becomes a fog-forming gas at a party.
Sublimation (solid CO₂ → gaseous CO₂). The visible fog is tiny droplets of water condensed from the cold air, not the CO₂ itself.
6Why does the temperature of boiling water stay at 100 °C (at sea level) even though the stove keeps adding heat?
The added energy is used to overcome the hydrogen bonds between molecules (to change liquid to gas), not to speed up the molecules, so the temperature stays constant until all the liquid has boiled.
More practice: warm-up
1True or false: when water boils, the bubbles contain hydrogen gas and oxygen gas.
False. The bubbles are water vapour (H₂O molecules). Only intermolecular forces are overcome.
2Name the change: liquid → gas.
Vaporization (evaporation or boiling).
3Name the change: gas → solid.
4Which state has a definite volume but no definite shape?
5Stronger intermolecular forces mean a ___ boiling point.
More practice: on your own
1Explain why chlorine, Cl₂, is a gas at room temperature but iodine, I₂, is a solid.
Both are non-polar, so only London forces act. I₂ has many more electrons (106 versus 34), so its London forces are much stronger, enough to hold the molecules in a solid at room temperature.
2Rank from lowest to highest boiling point: C₃H₈ (propane), CH₄ (methane), C₄H₁₀ (butane). Explain.
CH₄ < C₃H₈ < C₄H₁₀ (−162 °C, −42 °C, about −1 °C). All are non-polar; boiling point rises as molecules get larger and have more electrons, so London forces increase.
3H₂S (34 g/mol) boils at −60 °C, while H₂O (18 g/mol) boils at 100 °C. Explain why the lighter molecule has the higher boiling point.
H₂O forms hydrogen bonds (H bonded to O, with lone pairs on O), which are much stronger than the dipole–dipole and London forces between H₂S molecules. The type of IMF matters more here than size.
4Is the change endothermic or exothermic? (a) a puddle freezes; (b) snow sublimes on a sunny winter day.
(a) Exothermic: molecules come together and IMFs form, releasing energy. (b) Endothermic: molecules separate, absorbing energy (from sunlight).
5CO₂ has two polar C=O bonds, yet it is a non-polar molecule. Explain, and name its strongest IMF.
CO₂ is linear and symmetrical (O=C=O), so the two bond dipoles point in opposite directions and cancel. Its only IMF is London dispersion, which is why it is a gas at room temperature.
Challenge
1Propan-1-ol (CH₃CH₂CH₂OH, 60 g/mol) boils at 97 °C, and butane (C₄H₁₀, 58 g/mol) boils at about −1 °C. Explain the difference in detail.
The molar masses (and numbers of electrons: 34 and 34) are almost the same, so London forces are similar. Propan-1-ol has an O–H group, so its molecules form hydrogen bonds, and its C–O and O–H bonds make it polar. Butane is non-polar with only London forces. The extra hydrogen bonding accounts for the roughly 100-degree difference.
2When 18.0 g of liquid water (about 18 mL) evaporates at 100 °C and normal pressure, the vapour occupies about 30.6 L. By what factor does the volume increase, and what does this tell you about the space between gas particles?
30.6 L = 30 600 mL; 30 600 ÷ 18 ≈ 1700 times larger. The molecules themselves are unchanged, so almost all of the gas's volume is empty space between molecules. That is why gases are so compressible.
3A student says: "Large molecules always have higher boiling points than small molecules." Give a counter-example from this lesson, and rewrite the statement so it is correct.
Water (18 g/mol, 100 °C) boils far higher than H₂S (34 g/mol, −60 °C), or ethanol versus dimethyl ether with the same size. Correct version: "For molecules with the same type of intermolecular forces, larger molecules (more electrons) usually have higher boiling points; hydrogen bonding can outweigh size."
Quiz
1. Which statement best describes particles in a gas?
2. When liquid ethanol boils, what is overcome?
3. Which substance can form hydrogen bonds with itself?
4. Which has the strongest London dispersion forces?
5. Frost forming on a cold window is an example of:
6. Br₂ is a liquid at room temperature while Cl₂ is a gas. The best explanation is that Br₂ has:
Big idea (F3.2): the state of a substance depends on the balance between the kinetic energy of its particles and the intermolecular forces between them. This lesson also builds the vocabulary for the unit (F2.1).
How to help: ask your teen to explain one everyday observation with IMFs: why rubbing alcohol feels cold, why frost forms on the car, why a pot of water takes a long time to boil.
Common difficulties:
- Thinking boiling breaks molecules apart. Ask: "What is in the bubbles?" (Water molecules.)
- Mixing up intra- and intermolecular. "Inter" = between (like an interprovincial highway between provinces).
- Calling any H-containing molecule hydrogen-bonded. The H must be bonded to N, O or F.
Safety: supervise the evaporation activity; rubbing alcohol is flammable and toxic if swallowed, so keep it away from flames and ventilate the room.