The properties of GASES, LIQUIDS and SOLIDS - Application of the particle model to the three states of matter

Using the particle models to describe and explain the properties of gases, liquids and solids and state changes between them

Sub–index for Parts 0 to 3 (this page):

0 Introduction - What are the three states of matter?

1.1 Three states of matter - what can we expect from particle models and are there limitations?

1.1a Properties of gases - particle model and properties explained and diffusion experiments

What is diffusion? Examples of demonstrating diffusion in gases

1.1b Properties of liquids - particle model and properties explained and diffusion experiment

1.1c Properties of solids - particle model and properties explained

2. State changes - a summary diagram

2a Evaporation and boiling - explained using particle model

2b Condensation - explained using particle model

2c Distillation - explained using particle model

2d Melting - explained using particle model

2e Freezing-solidifying - explained using particle model

2f Cooling and heating curves - state changes and relative energy changes

2g Sublimation - explained using particle model

2h Comparison of latent heat changes in physical changes of state for different substances

3a-d. (a) Dissolving, (b) Solutions, (c) Miscible liquids & immiscible liquids, (d) Separating funnel

Appendix 1. Particle pictures of elements, compounds & mixtures

GCSE multiple choice QUIZ on states of matter – gases, liquids & solids

See also P-V-T pressure-volume-temperature gas law calculations and lots more

KEYWORD index for Part 1 (this page): Boiling * Boiling point * Brownian motion * Changes of state * Condensing * Cooling curve * Diffusion * Dissolving * Evaporation * Energy changes & change of state * Freezing * Freezing point  * Gas particle picture * Heating curve * Liquid particle picture * Melting * Melting point * miscible/immiscible liquids * Particle pictures of elements, compounds & mixtures * Properties of gases * Properties of liquids * Properties of solids * solutions * sublimation * Solid particle picture

Sub–index for Part 4 (on two separate more advanced pages): Some of these sections are for Advanced A level students only: Introduction to the kinetic particle theory of an ideal gas * Kelvin temperature scale * Kelvin temperature scale and Boyle's Law * Charles's–Gay Lussac's Law and the combined gas law equation * A Level only 4. The ideal gas equation PV=nRT * Dalton's Law of partial pressures * Graham's Law of diffusion * The deviations of a gases from ideal behaviour and their causes * The Van der Waals equation of state * Compressibility factors * The Critical Point – The Critical Temperature and Critical Pressure *

Part 0 Introduction

You should know that the three states of matter are solid, liquid and gas. Melting and freezing take place at the melting point, boiling and condensing take place at the boiling point. The three states of matter can be represented by a simple model in which the particles are represented by small solid spheres. Particle theory can help to explain melting, boiling, freezing and condensing.

The amount of energy needed to change state from solid to liquid and from liquid to gas depends on the strength of the forces between the particles of the substance and the nature of the particles involved depends on the type of bonding and the structure of the substance. The stronger the forces between the particles the higher the melting point and boiling point of the substance. For details see  structure and bonding notes.

The strength of the forces between particles depends on the material (structure and type of bonding), the temperature (affects the energy of the particles) and pressure (how close the particles are compressed together e.g. in a gas).

The physical state a material adopts depends on its structure, temperature and pressure.

State symbols used in equations: (g) gas       (l) liquid      (aq) aqueous solution      (s) solid

aqueous solution means something dissolved in water,

a good example of how to use the state symbols correctly is calcium carbonate dissolving in hydrochloric acid:

CaCO₃(s)  +  2HCl(aq)  ====>  CaCl₂(aq)  +  H₂O(l)  +  CO₂(g)

Most diagrams of particles on this page are 2D representations of their structure and state

EXAMPLES OF THE THREE PHYSICAL STATES OF MATTER

GASES

e.g. the air mixture around us (including the oxygen needed for combustion) and the high pressure steam in the boiler and cylinders of the steam locomotive. All of the gases in air are 'invisible', being colourless and transparent. Note that the 'steam' you see outside of a kettle or steam locomotive is actually fine liquid droplets of water, formed from the expelled steam gas condensing when it meets the cold air – the 'state change' of gas to liquid (same effect in mist and fog formation).

LIQUIDS

e.g. water is the most common example, but so are, milk, hot butter, petrol, oil, mercury or alcohol in a thermometer.

SOLIDS

e.g. stone, all metals at room temperature (except mercury), rubber of walking boots and the majority of physical objects around you. In fact most objects are useless unless they have a solid structure!

On this page

The basic physical properties of gases, liquids and solids are described in terms of structure, particle movement (kinetic particle theory), effects of temperature and pressure changes, and particle models used to explain these properties and characteristics. Hopefully, theory and fact will match up to give students a clear understanding of the material world around them in terms of gases, liquids and solids – referred to as the three physical states of matter.

The changes of state known as melting, fusing, boiling, evaporating, condensing, liquefying, freezing, solidifying, crystallising are described and explained with particle model pictures to help understanding. There is also a mention of miscible and immiscible liquids and explaining the terms volatile and volatility when applied to a liquid.

• WHAT IS THE GASEOUS STATE OF MATTER?
• WHAT ARE THE PROPERTIES OF A GAS?
• HOW DO GASEOUS PARTICLES BEHAVE?
• How does the kinetic particle theory of gases explain the properties of gases?
• A gas has no fixed shape or volume, but always spreads out to fill any container - the gas molecules will diffuse into any space available.
• There are almost no forces of attraction between the particles so they are completely free of each other.
• The particles are widely spaced and scattered and always moving rapidly at random throughout the container so there is no order in the system.
• The particles move linearly and rapidly in all directions, and frequently collide with each other and the side of the container.
• The collision of gas particles with the surface of a container causes gas pressure, on bouncing off a surface they exert a force in doing so.
• With increase in temperature, the particles move faster as they gain kinetic energy, the rate of collisions between the particles themselves and the container surface increases and this increases gas pressure eg in a steam locomotive or the volume of the container if it can expand eg like a balloon.

• Gases have a very low density (‘light’) because the particles are so spaced out in the container (density = mass / volume).
  • Density order: solid > liquid >>> gases
• Gases flow freely because there are no effective forces of attraction between the gaseous particles – molecules.
  • Ease of flow order: gases > liquids >>> solids (no real flow in solid unless you finely powder it!)
  • Because of this gases and liquids are described as fluids.
• Gases have no surface, and no fixed shape or volume, and because of lack of particle attraction, they always spread out and fill any container (so gas volume = container volume).
• Gases are readily compressed because of the ‘empty’ space between the particles.
  • Ease of compression order: gases >>> liquids > solids (almost impossible to compress a solid)
• Gas pressure
  • When a gas is confined in a container the particles will cause and exert a gas pressure which is measured in atmospheres (atm) or Pascals (1.0 Pa = 1.0 N/m²), pressure is force/area i.e. the effect of all the collisions on the surface of the container.
    • All particles have mass and their movement gives them kinetic energy and momentum.
    • The gas pressure is caused by the force created by millions of impacts of the tiny individual gas particles on the sides of a container.
    • For example – if the number of gaseous particles in a container is doubled, the gas pressure is doubled because doubling the number of molecules doubles the number of impacts on the side of the container so the total impact force per unit area is also doubled.
      • This doubling of the particle impacts doubling the pressure is pictured in the two diagrams below.

• If the ‘container’ volume can change, gases readily expand* on heating because of the lack of particle attraction, and readily contract on cooling.
  • On heating, gas particles gain kinetic energy, move faster and hit the sides of the container more frequently, and significantly, they hit with a greater force.
  • Depending on the container situation, either or both of the pressure or volume will increase (reverse on cooling).
  • Note: * It is the gas volume that expands NOT the molecules, they stay the same size!
  • If there is no volume restriction the expansion on heating is much greater for gases than liquids or solids because there is no significant attraction between gaseous particles. The increased average kinetic energy will make the gas pressure rise and so the gas will try to expand in volume if allowed to e.g. balloons in a warm room are significantly bigger than the same balloon in a cold room!
  • For gas volume–temperature calculations see Part 2 Charles's/Gay–Lussac's Law
• DIFFUSION in Gases:
  • The natural rapid and random movement of the particles in all directions means that gases readily ‘spread’ or diffuse.
    • The net movement of a particular gas will be in the direction from lower concentration to a higher concentration, down the so–called diffusion gradient.
    • Diffusion continues until the concentrations are uniform throughout the container of gases, but ALL the particles keep moving with their ever present kinetic energy!
• Diffusion is faster in gases than liquids where there is more space for them to move (experiment illustrated below) and diffusion is negligible in solids due to the close packing of the particles.
  • Diffusion is responsible for the spread of odours even without any air disturbance e.g. use of perfume, opening a jar of coffee or the smell of petrol around a garage.
  • The rate of diffusion increases with increase in temperature as the particles gain kinetic energy and move faster.
  • Other evidence for random particle movement including diffusion:
    • When smoke particles are viewed under a microscope they appear to 'dance around' when illuminated with a light beam at 90^o to the viewing direction. This is because the smoke particles show up by reflected light and 'dance' due to the millions of random hits from the fast moving air molecules. This is called 'Brownian motion' (see below in liquids). At any given instant of time, the particle hits will not be evenly distributed over the surface, so the smoke particle get a greater bashing in a random direction and then another, so they appear to dance and zig-zag around at random.
    • 
    • A two gaseous molecule diffusion experiment is illustrated above and explained below!
    • A long glass tube (2–4 cm diameter) is filled at one end with a plug of cotton wool soaked in conc. hydrochloric acid sealed in with a rubber bung (for health and safety!) and the tube is kept perfectly still, clamped in a horizontal position. A similar plug of conc. ammonia solution is placed at the other end. The soaked cotton wool plugs will give off fumes of HCl and NH₃ respectively, and if the tube is left undisturbed and horizontal, despite the lack of tube movement, e.g. NO shaking to mix and the absence of convection, a white cloud forms about ¹/₃rd along from the conc. hydrochloric acid tube end.
      • Explanation: What happens is the colourless gases, ammonia and hydrogen chloride, diffuse down the tube and react to form fine white crystals of the salt ammonium chloride.
      • ammonia + hydrogen chloride ===> ammonium chloride
        • NH_3(g) + HCl_(g) ===> NH₄Cl_(s)
      • Note the rule: The smaller the molecular mass, the greater the average speed of the molecules (but all gases have the same average kinetic energy at the same temperature).
        • Therefore the smaller the molecular mass, the faster the gas diffuses.
        • e.g. M_r(NH₃) = 14 + 1x3 = 17, moves faster than M_r(HCl) = 1 + 35.5 = 36.5
        • AND that's why they meet nearer the HCl end of the tube!
        • So the experiment is not only evidence for particle movement, it is also evidence that molecules of different molecular masses move/diffuse at different speeds.
        • See other page for a mathematical treatment of Graham's Law of Diffusion

A demonstration of diffusion A coloured gas, heavier than air (greater density), is put into the bottom gas jar and a second gas jar of lower density colourless air is placed over it separated with a glass cover. Diffusion experiments should be enclosed at constant temperature to minimise disturbance by convection. If the glass cover is removed then (i) the colourless air gases diffuses down into the coloured brown gas and (ii) bromine diffuses up into the air. The random particle movement leading to mixing cannot be due to convection because the more dense gas starts at the bottom! No 'shaking' or other means of mixing is required. The random movement of both lots of particles is enough to ensure that both gases eventually become completely mixed by diffusion (spread into each other). This is clear evidence for diffusion due to the random continuous movement of all the gas particles and, initially, the net movement of one type of particle from a higher to a lower concentration ('down a diffusion gradient'). When fully mixed, no further colour change distribution is observed BUT the random particle movement continues! See also other evidence in the liquid section after the particle model for diffusion diagram below. A particle model of diffusion in gases:  Imagine the diffusion gradient from left to right for the green particles added to the blue particles on the left. So, for the green particles, net migration is from left to right and will continue, in a sealed container, until all the particles are evenly distributed in the gas container (as pictured). Diffusion is faster in gases compared to liquids/solutions because there is more space between the particles for other particles to move into at random. ==> ==>

 1.1b. The particle model of a LIQUID

• WHAT IS THE LIQUID STATE OF MATTER?
• WHAT ARE THE PROPERTIES OF A LIQUID?
• HOW DO LIQUID PARTICLES BEHAVE?
• How does the kinetic particle theory of liquids explain the properties of liquids?
• A liquid has a fixed volume at a given temperature but its shape is that of the container which holds the liquid.
• There are much greater forces of attraction between the particles in a liquid compared to gases, but not quite as much as in solids and the particles are sufficiently free to move past each other.
  If there were no intermolecular forces, liquids could not exist!
• The particles are quite close together but still arranged at random throughout the container due to their constant random movement, there is a little close range order as you can get clumps of particles clinging together temporarily (as in the diagram above).
• As well as moving rapidly in all directions, they collide more frequently with each other than in gases due to shorter distances between particles – much greater density - particles closer together.
• With increase in temperature, the particles move faster as they gain kinetic energy, so increased collision rates, increased collision energy, increased rates of particle diffusion, expansion leading to decrease in density.

• Liquids have a much greater density than gases (‘heavier’) because the particles are much closer together because of the attractive forces.
• Most liquids are just a little less dense than when they are solid
  • Water is a curious exception to this general rule, which is why ice floats on water.
• Liquids usually flow freely despite the forces of attraction between the particles but liquids are not as ‘fluid’ as gases.
  • Note 'sticky' or viscous liquids have much stronger attractive forces between the molecules BUT not strong enough to form a solid.
• Liquids have a surface, and a fixed volume (at a particular temperature) because of the increased particle attraction, but the shape is not fixed and is merely that of the container itself.
  • Liquids seem to have a very weak 'skin' surface effect which is caused by the bulk molecules attracting the surface molecules disproportionately.
• Liquids are not readily compressed because there is so little ‘empty’ space between the particles, so increase in pressure has only a tiny effect on the volume of a solid, and you need a huge increase in pressure to see any real contraction in the volume of a liquid.
• Liquids will expand on heating but nothing like as much as gases, but more than solids, because of the greater particle attraction restricting the expansion (will contract on cooling).
  • The expansion of a liquid is due to the higher average kinetic energy of the particles and the more energetic collisions cause the expansion. BUT, they are still held together by the intermolecular forces, which restricts the expansion - this is not part of the kinetic particle theory!
  • Note: When heated, the liquid particles gain kinetic energy and hit the sides of the container more frequently, and more significantly, they hit with a greater force, so in a sealed container the pressure produced in a liquid can be considerable!
• DIFFUSION: The natural rapid and random movement of the particles means that liquids ‘spread’ - diffuse. Diffusion is much slower in liquids compared to gases because there is less space for the particles to move in and more ‘blocking’ collisions happen.
  • Just dropping lumps/granules/powder of a soluble solid (preferably coloured!) will resulting in a dissolving followed by an observable diffusion effect.
  • Again, the net flow of dissolved particles will be from a higher concentration to a lower concentration until the concentration is uniform throughout the container.
• Diffusion in liquids – evidence for random particle movement in liquids:
  • If coloured crystals of e.g. the highly coloured salt crystals of potassium manganate(VII) are dropped into a beaker of water and covered at room temperature.
    • Despite the lack of mixing due to shaking or convection currents from a heat source etc. the bright purple colour of the dissolving salt slowly spreads throughout all of the liquid but it is much slower than the gas experiment described above because of the much greater density of particles slowing the spreading due to close proximity collisions.
    • The same thing happens with dropping copper sulphate crystals (blue, so observable) or coffee granules into water and just leaving the mixture to stand.
    • Experiment to show the slower diffusion in liquids eg water
    • 
    • You start with a beaker of still pure colourless water and drop a few crystals of ANY highly coloured soluble crystals into it and put on a lid cover to prevent any air disturbance.
    • The beaker is left to stand, preferably at a constant temperature to prevent mixing due to convention. Immediately the crystals are added they will begin to dissolve and due to natural random particle motion the coloured molecules will begin to spread from an area of high concentration to one of low concentration and in all directions. You could take a series of photographs to record the spreading. The spreading is self-evident and direct experimental evidence for the natural constant random movement of particles (molecules or ions).
    • After many hours all of the crystals will have dissolved AND due to the random movement of ALL the particles, everything dissolved becomes evenly distributed giving an evenly coloured solution. Note that although the colour doesn't seem to spread anymore, ALL the particles are still moving with a random motion, nothing stops!

• A particle model of diffusion in liquids: Imagine the diffusion gradient from left to right for the green particles added to the blue particles on the left. So, for the green particles, net migration is from left to right and will continue, in a sealed container, until all the particles are evenly distributed (as pictured). Diffusion is slower in liquids because there is less space between the particles for other particles to move into and random collisions will occur more frequently slowing down the particle spreading effect down a diffusion gradient.

  ==> ==>

  • When pollen grains suspended in water are viewed under a microscope they appear to 'dance around' when illuminated with a light beam at 90^o to the viewing direction.
    • This is because the pollen grains show up by reflected light and 'dance' due to the millions of random hits from the fast moving water molecules.
    • This phenomenon is called 'Brownian motion' after a botanist called Brown first described the effect (see gases above).
    • At any given instant of time, the particle hits will not be even all round the surface of the pollen grains, so they get a greater number of hits in a random direction and then another, hence the pollen grains zig-zag around in all directions at random.
• Heat conduction in liquids
  • Most liquids are poor conductors of thermal energy, energy which is due to the kinetic energy of the moving particles.
  • Heat energy is transferred by 'hotter' higher kinetic energy liquid particles colliding with 'cooler' lower kinetic energy particles so raising their kinetic energy and spreading the heat energy.
  • However, the density of liquids is much greater than gases (particles much closer together), so the density or rate of 'collision transfer' is much higher, so liquids are better heat conductors than gases.
  • Liquid metals are very good heat conductors because of the freely moving electrons that can carry the kinetic energy rapidly through the liquid. For more details see 'metal structure'.
• Electrical conduction in liquids
  • Electrical conduction requires the presence of free IONS or free ELECTRONS i.e. particles that can carry an electrical charge.
  • Most liquids are poor conductors of electricity (good insulators), but there are important exceptions.
  • For example, if a liquid contains ions e.g. salt solutions, then electrical conduction can take place
  • Liquid metals are very good electrical conductors because of the freely moving electrons that can carry the electrical current rapidly through the liquid metal.
  • For more details see 'electrolysis' and 'metal structure'.

• WHAT IS THE SOLID STATE OF MATTER?
• WHAT ARE THE PROPERTIES OF A SOLID?
• HOW DO SOLID PARTICLES BEHAVE?
• How does the kinetic particle theory of solids explain the properties of solids?
• A solid has a fixed volume and shape at a particular temperature unless physically subjected to some force.
• The greatest forces of attraction are between the particles in a solid and they pack together as tightly as possible in a neat and ordered arrangement called a lattice.
• The particles are too strongly held together to allow movement from place to place but the particles vibrate about their position in the structure.
• With increase in temperature, the particles vibrate faster and more strongly as they gain kinetic energy, so the vibration increases causing expansion.
• More on the kinetic particle theory of an ideal gas

• Solids have the greatest density (‘heaviest’) because the particles are closest together.
• Solids cannot flow freely like gases or liquids because the particles are strongly held in fixed positions.
• Solids have a fixed surface and volume (at a particular temperature) because of the strong particle attraction.
• Solids are extremely difficult to compress because there is no real ‘empty’ space between the particles, so increase in pressure has virtually no effect on the volume of a solid.
• Solids will expand a little on heating but nothing like as much as liquids because of the greater particle attraction restricting the expansion and causing the contraction occurs on cooling.
  • The expansion is caused by the increased kinetic energy of particle vibration, forcing them further apart causing an increase in volume and corresponding decrease in density.
  • Although the expansion of a solid is due to the higher average kinetic energy of the particles and the more energetic vibrations, they are still held together by the intermolecular bonding forces (or much stronger strong ionic or covalent bonds), which restricts the expansion - this is not part of the kinetic particle theory!
• Diffusion is almost impossible in solids because the particles are too closely packed and strongly held together in a lattice. The immobile particles cannot move around because there is no random movement into ‘empty space’ for them to move through.
Its quite a different situation in gases and liquids where diffusion readily takes place because of the freedom of the particles to move around at random and 'bash' each other aside!
• Heat conduction in solids
  • Apart from metals, most solids are poor conductors of heat energy, energy which is due to the kinetic energy of the vibrating particles in the crystal structure – remember, unlike gases and liquids, the particles can't move around, they just vibrate about a fixed point.
  • Heat energy is transferred by 'hotter' higher kinetic energy vibrating particles colliding against 'cooler' lower kinetic energy vibrating particles so raising their kinetic energy and spreading the heat energy through the solid structure.
  • The density of solids and order of particles is are greater than liquids (particles closest together), so the density or rate of 'collision transfer' vibration is much higher, so solids are better heat conductors than liquids (and much greater than gases).
  • However, although most non-metal solids are poor heat conductors, metals are exceptionally good heat conductors because of the freely moving electrons that can carry the kinetic energy rapidly through the crystal structure.
  • For more details see 'metal structure'.
• Electrical conduction in solids
  • Electrical conduction requires the presence of free IONS or free ELECTRONS i.e. particles that can carry an electrical charge within a solid structure. Which of course is impossible in most solids (except metals) because ALL particles can't move around, so even solid ionic compounds cannot conduct electricity.
  • Most non-metal solids are poor conductors of electricity (good insulators), but there are important exceptions.
  • All metals are relatively good electrical conductors because of the freely moving electrons that can carry the electrical current rapidly through the liquid metal. For more details see 'metal structure'.
  • Graphite and graphene, forms (allotropes) of the non–metallic element carbon, are electrical conductors due to free moving electrons in the solid structure, a rare exception of conducting solids apart from metals.

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2. Changes of State for gas <=> liquid <=> solid

 You need to be able to predict the state of a substance at different temperatures given appropriate data.

Below its melting point a substance is a solid.

Between its melting point and boiling point, the substance is a liquid.

Above its boiling point, a substance is a gas/vapour.

		FREEZING MELTING
SUBLIMING -the opposite is deposition or 'reverse sublimation'				BOILING or EVAPORATING
SUMMARY of the CHANGES of STATE between a gas, liquid and solid All mass conserved in these PHYSICAL CHANGES				CONDENSING These are NOT chemical changes !

A change of state means an interconversion between two states of matter, namely gas <=> liquid <=> solid

A 'triangular' summary of important state changes is illustrated above.

e.g. solid ==> liquid is melting or fusing

liquid ==> gas/vapour (vapor) is boiling, evaporation or vapourisation (vaporisation)

and the reverse processes

gas/vapour (vapor) ==> liquid is condensation, liquefaction/liquefying

liquid ==> solid is freezing, solidifying or crystallising

and there is also

solid ==> gas is sublimation

We can use the state particle models and diagrams to explain changes of state and the energy changes involved.

These are NOT chemical changes BUT PHYSICAL CHANGES, e.g. the water molecules H₂O are just the same in ice, liquid water, steam or water vapour. What is different, is how they are arranged, and how strongly they are held together by intermolecular forces in the solid, liquid and gaseous states.

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2a. Evaporation and Boiling (liquid to gas)

• Evaporation is when particles of a liquid escape to form a gas or vapour i.e. water evaporating into the air.
• Because of random collisions, the particles in a liquid have a variety of speeds and kinetic energies. On heating, particles gain kinetic energy and move faster and are more able to overcome the intermolecular forces between the molecules i.e. some particles will have enough kinetic energy to overcome the attractive forces holding the particles together in the bulk liquid.
  • Even without further heating, evaporation occurs all the time from volatile liquids, but it is still the higher kinetic energy particles that can overcome the attractive forces between the molecules in the bulk of the liquid and escape from the surface into the surrounding air.
• In evaporation and boiling (both are vaporisation) it is the highest kinetic energy molecules that can ‘escape’ from the attractive forces of the other liquid particles.
  • The particles lose any order and become completely free to form a gas or vapour.
  • Also, because the highest kinetic energy particles have escaped, the liquid is cooler, because the lower kinetic energy particles are left.
  • This is equivalent to energy being used to evaporate a liquid (see below).
  • 
  • The graph above shows how the distribution of kinetic energy and speed of particles changes with changes in temperature - with increase in temperature, the average speed and kinetic energy of the particles increases.

  • Note that the random movement and collisions of the particles creates a wide range of speeds/kinetic energies.

  • When the temperature is increased, more particles have a greater kinetic energy and greater speed, but only the highest speed/kinetic energy particles can escape from the surface (only the very right-hand section of the graph curves)
  • Below is a particle model of evaporation.
  • 
• Energy is needed to overcome the attractive forces between particles in the liquid and is taken in from the surroundings.
  • In boiling, heat energy must be continually supplied e.g. from an electrical heating element or Bunsen burner etc.
  • In the case of evaporation, the heat is taken from the liquid, so an evaporating liquid cools - the lower speed/kinetic energy particles are left behind.
• This means heat is taken in, so evaporation and boiling are endothermic processes (ΔH +ve).
• If the temperature is high enough boiling takes place and bubbles of gas form in the bulk liquid – something you don't see in evaporation, because that can only occur on the surface of a liquid.
• Boiling is rapid vapourisation anywhere in the bulk liquid and at a fixed temperature called the boiling point and requires continuous addition of heat.
  • Boiling point depends on the ambient pressure, the lower the gas pressure above the liquid, the lower the boiling point of the liquid.
  • This is why tea brewed on the top of high mountain isn't quite as good as at sea level, the water boils at a lower temperature and doesn't extract substances from the tea leaves as efficiently!
  • In the past, measuring the boiling point of water was used to estimate the height of land above sea level!
• The rate of boiling is limited by the rate of heat transfer into the liquid.
• Evaporation takes place more slowly than boiling at any temperature between the melting point and boiling point, and only from the surface, and results in the liquid becoming cooler due to loss of higher kinetic energy particles.
• Factors affecting the rate of evaporation of a liquid.
  • The higher the temperature of the liquid, the faster it evaporates, because more particles have sufficient kinetic energy to overcome the intermolecular forces of the bulk liquid and can escape from the liquid surface.
  • The larger the surface area of given volume of liquid, the faster it evaporates, because there is a greater probability of particles escaping.
  • The greater the airflow over a liquid the faster it evaporates because its stops a build–up of vapour particles which may hit the surface and condense! The airflow lowers the concentration of evaporated particles by sweeping them away and so more readily replaced by freshly evaporated particles.
  • Please note that the best conditions for drying washing are a warm sunny day, a good breeze, and spreading the clothes out as much as possible to increase their surface area (I get told off about this one!).
• More details on the energy changes for these physical changes of state for a range of substances are dealt with in a section of the Energetics Notes.

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2b. Condensing (gas to liquid) – the process of condensation

Explained using the kinetic particle theory of gases and liquids

• On cooling, gas particles lose kinetic energy, they slow down and eventually become attracted together via intermolecular forces to form a liquid i.e. they haven't enough kinetic energy to remain free in the gaseous state.
• There is an increase in order as the particles are much closer together and can form clumps of molecules.
• The process requires heat to be lost to the surroundings i.e. heat given out, so condensation is exothermic (ΔH –ve).
  • This is why steam has such a scalding effect, its not just hot, but you get extra heat transfer to your skin due to the exothermic condensation on your surface!
  • In your home you see condensation on cold windows and steam is invisible, and what you refer to as steam coming out of a kettle is actually a cloud of water droplets from the condensation of steam vapour in the cooler air.
• Factors affecting the rate of condensation of a gas–vapour
  • The lower the temperature of the gas the faster it condenses because the particles on average have less kinetic energy to overcome the attractive intermolecular forces i.e. they gas particles are more likely to aggregate into drops of liquid.
  • The colder the surface the gas condenses on, the faster the heat transfer to reduce the kinetic energy of the gas particles, so the faster the gas/vapour can condense.
  • The higher the concentration of vapour in air, the faster condensation can take place. The particles are closer together and more chance of combining to form liquid droplets.

• Simple and fractional distillation involve the processes of boiling and condensation and are described on the Elements, Compounds and Mixtures

• Part 2 page, where other methods of separation are also described.

• The process of distillation involves boiling (liquid ==> gas/vapor) and the reverse process of condensation (gas/vapour ==> liquid)

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2d. Melting (solid to liquid)

Explained using the kinetic particle theory of liquids and solids

• When a solid is heated the particles vibrate more strongly as they gain kinetic energy and the particle attractive forces are weakened.
• Eventually, at the melting point, the attractive forces are too weak to hold the particles in the structure together in an ordered way and so the solid melts.
  • Note that the intermolecular forces are still there to hold the bulk liquid together – but the effect is not strong enough to form an ordered crystal lattice of a solid.
• The particles become free to move around and lose their ordered arrangement.
• Energy is needed to overcome the attractive forces and give the particles increased kinetic energy of vibration.
• So heat is taken in from the surroundings and melting is an endothermic process (ΔH +ve).
• Energy changes for these physical changes of state for a range of substances are dealt with in a section of the Energetics Notes.

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2e. Freezing (liquid to solid)

Explained using the kinetic particle theory of liquids and solids

• On cooling, liquid particles lose kinetic energy and so can become more strongly attracted to each other.
• When the temperature is low enough, the kinetic energy of the particles is insufficient to prevent the particle attractive forces causing a solid to form.
• Eventually at the freezing point the forces of attraction are sufficient to remove any remaining freedom of movement (in terms of one place to another) and the particles come together to form the ordered solid arrangement (though the particles still have vibrational kinetic energy.
• Since heat must be removed to the surroundings, so strange as it may seem, freezing is an exothermic process (ΔH –ve).

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2f. Cooling and Heating Curves and the comparative energy changes for changes of state: gas <=> liquid <=> solid Below the melting/freezing point, the substance is a liquid. Between the melting/freezing point and the boiling point, the substance is a liquid. Above the boiling point, the substance is a gas/vapour.

2f(i) Cooling curve: What happens to the temperature of a substance if it is cooled from the gaseous state to the solid state? Note the temperature stays constant during the state changes of condensing at temperature Tc, and freezing/solidifying at temperature Tf. This is because all the heat energy removed on cooling at these temperatures (the latent heats or enthalpies of state change), allows the strengthening of the inter–particle forces (intermolecular bonding)  without temperature fall. The heat loss is compensated by the exothermic increased intermolecular force attraction. In between the 'horizontal' state change sections of the graph, you can see the energy 'removal' reduces the kinetic energy of the particles, lowering the temperature of the substance. See section 2. for detailed description of the state changes.   A cooling curve summarises the changes: gas ==> liquid ==> solid For each change of state, energy must be removed, known as the 'latent heat' Actual energy values for these physical changes of state for a range of substances are dealt with in more detail in the Energetics Notes.

2f(ii) Heating curve:  What happens to the temperature of a substance if it is heated from the solid state to the gaseous state? Note the temperature stays constant during the state changes of melting at temperature Tm and boiling at temperature Tb. This is because all the energy absorbed in heating at these temperatures (the latent heats or enthalpies of state change), goes into weakening the inter–particle forces (intermolecular bonding) without temperature rise The heat gain equals the endothermic/heat absorbed energy required to reduce the intermolecular forces. In between the 'horizontal' state change sections of the graph, you can see the energy input increases the kinetic energy of the particles and raising the temperature of the substance.  See section 2. for detailed description of the state changes.   A heating curve summarises the changes: solid ==> liquid ==> gas For each change of state, energy must be added, known as the 'latent heat' Actual energy values for these physical changes of state for a range of substances are dealt with in more detail in the Energetics Notes.   A comparison of cooling and heating graph curves. SPECIFIC LATENT HEATS - refer to diagram below The latent heat for the state changes solid <=> liquid is called the specific latent heat of fusion (for melting or freezing). The latent heat for the state changes liquid <=> gas is called the specific latent heat of vaporisation (for condensing, evaporation or boiling) For more on latent heat see my physics notes on specific latent heat

• Sublimation:

  • This is when a solid, on heating, directly changes into a gas without melting, AND the gas on cooling re–forms a solid directly without condensing to a liquid. Sublimation usually just involves a physical change BUT its not always that simple (see ammonium chloride!).

  • The opposite of sublimation is sometimes referred to as deposition or 'reverse sublimation'.

• Theory in terms of particles:

  • When the solid is heated the particles vibrate with increasing force from the added thermal energy.

    • If the particles have enough kinetic energy of vibration to partially overcome the particle–particle attractive forces you would expect the solid to melt.

    • HOWEVER, if the particles at this point have enough energy at this point that would have led to boiling, the liquid will NOT form and the solid turns directly into a gas.

      • Overall endothermic change, energy absorbed and 'taken in' to the system.

  • On cooling, the particles move slower and have less kinetic energy.

    • Eventually, when the particle kinetic energy is low enough, it will allow the particle–particle attractive forces to produce a liquid.

    • BUT the energy may be low enough to permit direct formation of the solid, i.e. the particles do NOT have enough kinetic energy to maintain a liquid state!

      • Overall exothermic change, energy released and 'given out' to the surroundings.

• Examples:

  1. Even at room temperature bottles of solid iodine show crystals forming at the top of the bottle above the solid. The warmer the laboratory, the more crystals form when it cools down at night!

     • I_{2 (s)} I_{2 (g)}   (physical change only)

     • If you gently heat iodine in a test tube you see the iodine readily sublime and recrystallise on the cooler surface near the top of the test tube.

  2. The formation of a particular form of frost involves the direct freezing of water vapour (gas).  Frost can also evaporate directly back to water vapour (gas) and this happens in the 'dry' and extremely cold winters of the Gobi Desert on a sunny day.

     • H₂O _(s) H₂O _(g)   (physical change only)

     • See pictures of 'hoar frost' below.

  3. Solid carbon dioxide (dry ice) is formed on cooling the gas down to less than –78^oC. On warming it changes directly to a very cold gas!, condensing any water vapour in the air to a 'mist', hence its use in stage effects.
     • CO_{2 (s)} CO_{2 (g)}   (physical change only)
  4. On heating strongly in a test tube, white solid ammonium chloride, decomposes into a mixture of two colourless gases ammonia and hydrogen chloride. On cooling the reaction is reversed and solid ammonium chloride reforms at the cooler top surface of the test tube.

     • Ammonium chloride + heat energy ammonia + hydrogen chloride

     • NH₄Cl_(s) NH_3(g) + HCl_(g)     

     • This involves both chemical and physical changes and is so is more complicated than examples 1. to 3. In fact the ionic ammonium chloride crystals change into covalent ammonia and hydrogen chloride gases which are naturally far more volatile (covalent substances generally have much lower melting and boiling points than ionic substances).

  The liquid particle picture does not figure here, but the other models fully apply apart from state changes involving liquid formation. GAS particle model and SOLID particle model links.

  PLEASE NOTE, At a higher level of study, you need to study the g–l–s phase diagram for water and the vapour pressure curve of ice at particular temperatures. For example, if the ambient vapour pressure is less than the equilibrium vapour pressure at the temperature of the ice, sublimation can readily take place. The snow and ice in the colder regions of the Gobi Desert do not melt in the Sun, they just slowly 'sublimely' disappear!

The formation of hoar frost - the reverse of sublimation

Frost is a thin layer of ice on a solid surface.

Hoar frost forms directly from water vapour in air above 0^oC, coming in contact with a solid surface whose temperature is below freezing (<0^oC).

The water vapor changes directly from gas (vapour) to solid (ice) as it comes into contact with the solid surface.

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3. Dissolving solids, solutions and miscible/immiscible liquids

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• 3a. WHAT HAPPENS TO PARTICLES WHEN A SOLID DISSOLVES IN A LIQUID SOLVENT?

• What do the words SOLVENT, SOLUTE and SOLUTION mean?

• When a solid (the solute) dissolves in a liquid (the solvent) the resulting mixture is called a solution.

  • In general: solute + solvent ==> solution

  • So, the solute is what dissolves in a solvent, a solvent is a liquid that dissolves things and the solution is the result of dissolving something in a solvent.

  • The solid loses all its regular structure and the individual solid particles (molecules or ions) are now completely free from each other and randomly mix with the original liquid particles, and all particles can move around at random.

  • This describes salt dissolving in water, sugar dissolving in tea or wax dissolving in a hydrocarbon solvent like white spirit.

  • It does not usually involve a chemical reaction, so it is generally an example of a physical change.

  • Whatever the changes in volume of the solid + liquid, compared to the final solution, the Law of Conservation of Mass still applies.

  • This means: mass of solid solute + mass of liquid solvent = mass of solution after mixing and dissolving.

  • You cannot create mass or lose mass, but just change the mass of substances into another form.

  • If the solvent is evaporated, then the solid is reformed e.g. if a salt solution is left out for a long time or gently heated to speed things up, eventually salt crystals form, the process is called crystallisation.

• 3b. WHAT HAPPENS TO PARTICLES WHEN TWO LIQUIDS COMPLETELY MIX WITH EACH OTHER?

• WHAT DOES THE WORD MISCIBLE MEAN?

• Using the particle model to explain miscible liquids.

• If two liquids completely mix in terms of their particles, they are called miscible liquids because they fully dissolve in each other. This is shown in the diagram below where the particles completely mix and move at random. The process can be reversed by fractional distillation.

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• 3c. WHAT HAPPENS TO PARTICLES WHEN TWO LIQUIDS DO NOT MIX WITH EACH OTHER?

• WHAT DOES THE WORD IMMISCIBLE MEAN?

• WHY DO THE LIQUIDS NOT MIX?

• Using the particle model to explain immiscible liquids.

• If the two liquids do NOT mix, they form two separate layers and are known as immiscible liquids, illustrated in the diagram below where the lower purple liquid will be more dense than the upper layer of the green liquid.

  • You can separate these two liquids using a separating funnel.

  • The reason for this is that the interaction between the molecules of one of the liquids alone is stronger than the interaction between the two different molecules of the different liquids.

  • For example, the force of attraction between water molecules is much greater than either oil–oil molecules or oil–water molecules, so two separate layers form because the water molecules, in terms of energy change, are favoured by 'sticking together'.

3d. How a separating funnel is used 1. The mixture is put in the separating funnel with the stopper on and the tap closed and the layers left to settle out. 2. The stopper is removed, and the tap is opened so that you can carefully run the lower grey layer off first into a beaker. 3. The tap is then closed again, leaving behind the upper yellow layer liquid, so separating the two immiscible liquids.

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Appendix 1 some SIMPLE particle pictures of ELEMENTS, COMPOUNDS and MIXTURES

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OTHER ASSOCIATED PAGES

GCSE/IGCSE multiple choice QUIZ on states of matter – gases, liquids & solids

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QCA 7G "Particle model of solids, liquids and gases" Multiple Choice Questions for Science  revision on gases, liquids and solids – particle models, properties, explaining the differences between them.

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