Difficulty: Introductory | Prerequisites: Basic understanding of atoms and molecules
Source: General Chemistry, Purdue University
Tags: phases of matter, states of matter, kinetic molecular theory, KMT, pressure, viscosity, fluidity, surface tension, solids, liquids, gases, plasma, phase changes, heat of fusion, heat of vaporisation, specific heat, CHEM 101
This topic covers the physical states of matter (solid, liquid, gas, plasma), the properties that distinguish them, and the energy required to move between them. It builds on the particle model of matter you met in earlier chapters and connects forward to thermodynamics and intermolecular forces.
If you are comfortable with the idea that matter is made of particles in motion, you are ready for this material. The maths here is straightforward arithmetic, not algebra.
Matter exists as solid, liquid, gas, or (under extreme conditions) plasma. The kinetic molecular theory explains gas behaviour by treating particles as tiny, perfectly elastic objects in constant random motion. Phase changes require energy input or release, and calculating that energy means adding up "legs" on a heating curve, each with its own formula.
Kinetic Molecular Theory (KMT)
A model stating that all gases consist of tiny particles in constant, random motion. Collisions between particles and with container walls are perfectly elastic (no energy is lost).
Think of it as a room full of bouncy balls flying in every direction, never slowing down.
Pressure
Force per unit of area, arising from gas particles striking the walls of their container. Measured with barometers (atmospheric) or manometers (gas samples).
Barometer
A device for measuring atmospheric pressure. Torricelli's mercury barometer was the first, using a column of mercury in an inverted tube.
Manometer
A device that compares the pressure of a gas sample to atmospheric pressure (open manometer) or to a vacuum (closed manometer).
Fluidity
The ability of a substance to flow. Liquids and gases are both fluids. Fluidity increases with temperature.
In simple terms, the hotter something is, the more easily it flows.
Viscosity
Resistance to flow, the inverse of fluidity. Larger molecules tend to have higher viscosity. Viscosity decreases as temperature increases.
Think of it as "thickness": honey has high viscosity, water has low viscosity.
Surface Tension
The tendency of a liquid surface to contract, caused by molecules at the surface being pulled inward by cohesive forces. This is why water forms droplets. Soap (a surfactant) reduces surface tension.
Cohesion
Attractive forces between molecules of the same substance. Mercury's strong cohesion gives it a convex (upward-pointing) meniscus in glass.
Adhesion
Attractive forces between molecules of different substances. Water's adhesion to glass is stronger than its cohesion, giving it a concave (downward-curving) meniscus.
Capillary Action
The movement of liquid through narrow spaces due to adhesion. This is how water travels up through the xylem in plants (transpiration-cohesion-tension mechanism).
Ionic Crystals
Solids held together by ionic bonds. Strong but brittle (they shatter rather than bend). Do not conduct electricity unless melted, fused, or dissolved in water.
Covalent Network Solids
Solids held together by covalent bonds extending throughout the structure (e.g. diamond, graphite). Very hard, very high melting points.
Amorphous Solids
Solids without a regular crystal structure, formed when molten material cools too quickly for crystals to develop. Examples: glass, rubber, plastic. They have no sharp melting point.
Saturation (Vapour Equilibrium)
The point at which the rate of evaporation equals the rate of condensation. The liquid level stays constant even though molecules are still leaving and returning.
Boiling Point
The temperature at which the vapour pressure of a liquid equals the external atmospheric pressure. Boiling occurs throughout the liquid, not just at the surface (that is evaporation).
Solids have a definite shape and definite volume. Particles vibrate in fixed positions.
Liquids have a definite volume but take the shape of their container. They are fluids.
Gases have neither definite shape nor definite volume. They expand to fill any container. Also fluids.
Plasma forms under extreme conditions (very high temperatures). Not commonly encountered in general chemistry coursework.
All gases are made of particles in constant, random motion.
Collisions between particles, and between particles and container walls, are perfectly elastic (no kinetic energy is lost).
The volume of individual gas particles is negligible compared to the volume of the container.
There are no attractive or repulsive forces between particles.
The average kinetic energy of gas particles is proportional to the absolute temperature (in Kelvin).
Pressure = force / area.
Torricelli invented the mercury barometer to measure atmospheric pressure.
Meniscus in mercury: curves upward (convex), because cohesion between mercury atoms is stronger than adhesion to glass.
Meniscus in water: curves downward (concave), because adhesion to glass is stronger than cohesion between water molecules.
Open manometers compare gas pressure to atmospheric pressure.
Closed manometers compare gas pressure to a vacuum.
Fluidity increases as temperature increases.
Viscosity (resistance to flow) decreases as temperature increases. Larger molecules have higher viscosity.
Surface tension pulls surface molecules inward. Soap reduces surface tension, which is why it helps water spread and clean.
Capillary action (adhesion-driven) moves liquid through narrow spaces, as in plant xylem.
Ionic crystals: Strong bonds, brittle. Conduct electricity only when melted, dissolved, or fused (these are electrolytes in solution).
Metallic solids: Atoms with delocalised (mobile) valence electrons. Good conductors.
Molecular solids: Held together by intermolecular forces (not ionic or metallic bonds). Poor conductors because they lack free ions or electrons.
Covalent network solids: Extended covalent bonding throughout (diamond, graphite). Extremely hard, high melting points.
Amorphous solids: No repeating crystal pattern. Form when molten material cools too quickly. No specific melting point. Examples: glass, rubber, plastic.
For most substances, density increases as temperature decreases (particles pack more closely).
Water is the notable exception: it reaches maximum density at 4 °C. Below that, the hydrogen-bonding structure of ice makes it less dense than liquid water, which is why ice floats.
A heating curve plots temperature (y-axis) against time or energy added (x-axis).
Five segments ("legs") for water going from ice below 0 °C to steam above 100 °C:
Leg 1: Heating ice (solid phase). Use q = m × c_ice × ΔT. Specific heat of ice: 2.03 J/(g·°C).
Leg 2: Melting ice at 0 °C (phase change, temperature constant). Use q = m × heat of fusion. Heat of fusion of water: 334 J/g.
Leg 3: Heating liquid water. Use q = m × c_water × ΔT. Specific heat of water: 4.184 J/(g·°C).
Leg 4: Boiling water at 100 °C (phase change, temperature constant). Use q = m × heat of vaporisation. Heat of vaporisation of water: 2230 J/g.
Leg 5: Heating steam. Use q = m × c_steam × ΔT. Specific heat of steam: 2.06 J/(g·°C).
Total energy = sum of all five legs.
Leg 1 (ice, -50 °C to 0 °C): (50)(50)(2.03) = 5075 J
Leg 2 (melting at 0 °C): (50)(334) = 16 700 J
Leg 3 (water, 0 °C to 100 °C): (50)(100)(4.184) = 20 920 J
Leg 4 (boiling at 100 °C): (50)(2230) = 111 500 J
Leg 5 (steam, 100 °C to 150 °C): (50)(50)(2.06) = 5150 J
Total ≈ 159 345 J ≈ 159 kJ
Heating or cooling a single phase: q = m × c × ΔT (where c is the specific heat of that phase)
Melting or freezing: q = m × heat of fusion (334 J/g for water)
Boiling or condensing: q = m × heat of vaporisation (2230 J/g for water)
Specific heats of water's phases:
Ice: 2.03 J/(g·°C)
Liquid water: 4.184 J/(g·°C)
Steam: 2.06 J/(g·°C)
Total energy for a full phase journey: sum each leg's q value
Ice floating on water is a direct consequence of water's anomalous density behaviour. If ice sank, lakes would freeze from the bottom up, killing aquatic life.
Surface tension is what lets small insects walk on water. Detergents work by lowering surface tension so water can spread into fibres and lift dirt away.
Students often think boiling and evaporation are the same thing. Evaporation happens at the surface at any temperature. Boiling happens throughout the liquid, only at the boiling point.
Students assume temperature keeps rising during a phase change. It does not. During melting or boiling, all added energy goes into breaking intermolecular attractions, not raising temperature.
Students forget that specific heat values differ for each phase of the same substance. Ice, liquid water, and steam each have their own c value.
Students sometimes think amorphous solids are not "real" solids. They are solids (definite shape, definite volume), they just lack a repeating crystal structure.
⚠️ Heating curve calculations are a near-certainty on the exam. You must know which formula to use on each leg and remember that temperature stays flat during phase changes.
⚠️ Know the three specific heat values for water's phases and the two latent heat values (fusion and vaporisation). These are commonly given, but not always.
⚠️ Expect a conceptual question about why ice floats (water's maximum density at 4 °C, hydrogen bonding in ice).
⚠️ Meniscus direction (mercury vs water) and the reason (cohesion vs adhesion) is a favourite multiple-choice item.
⚠️ Types of solids (ionic, metallic, molecular, covalent network, amorphous) and their conductivity properties appear frequently. Know that ionic solids only conduct when melted or dissolved.
True or false: Plasma is the most commonly observed phase of matter in a chemistry lab. (False, plasma requires extreme conditions.)
Fill in the blank: Temperature remains constant during a ______ change. (phase)
True or false: Viscosity increases as temperature increases. (False, viscosity decreases as temperature increases.)
Fill in the blank: Water reaches its maximum density at ______ °C. (4)
True or false: Ionic crystals conduct electricity in their solid state. (False, only when melted, fused, or dissolved.)
Q: How much energy is needed to melt 100 g of ice at 0 °C?
A: q = m × heat of fusion = (100)(334) = 33 400 J = 33.4 kJ.
Q: Why does mercury form a convex meniscus in glass, while water forms a concave one?
A: Mercury's cohesive forces (mercury-to-mercury attraction) are stronger than its adhesion to glass, so the surface curves upward. Water's adhesion to glass is stronger than its internal cohesion, so it climbs the glass walls and curves downward.
Q: What distinguishes an amorphous solid from a crystalline one?
A: Amorphous solids lack a regular, repeating internal structure. They form when molten material cools too rapidly for an ordered lattice to develop, and they soften gradually over a range of temperatures rather than melting at a sharp point.
Q: A student heats 50 g of ice from -20 °C to liquid water at 40 °C. Calculate the total energy required.
A: Three legs. Leg 1 (heating ice): (50)(20)(2.03) = 2030 J. Leg 2 (melting): (50)(334) = 16 700 J. Leg 3 (heating water): (50)(40)(4.184) = 8368 J. Total = 2030 + 16 700 + 8368 = 27 098 J ≈ 27.1 kJ.
Q: Name one property that all fluids share but solids do not.
A: Fluids (liquids and gases) flow and take the shape of their container. Solids have a fixed shape.
The kinetic molecular theory introduced here reappears when you study gas laws (Boyle, Charles, ideal gas law), which are covered in the companion study notes.
Phase change energy calculations connect to thermochemistry (Hess's law, enthalpy of reaction), where you will add energy values across multiple steps in the same way.
Intermolecular forces (hydrogen bonding, dipole-dipole, London dispersion) determine boiling points, viscosity, and surface tension. That topic deepens everything covered here about why different substances behave differently.
Phases of matter, states of matter, solid liquid gas plasma, kinetic molecular theory, KMT, pressure, barometer, manometer, viscosity, fluidity, surface tension, cohesion, adhesion, capillary action, meniscus, ionic crystals, metallic solids, covalent network solids, amorphous solids, phase change, heating curve, specific heat, heat of fusion, heat of vaporisation, latent heat, boiling point, melting point, water density anomaly, CHEM 101, general chemistry, Purdue chemistry