Doc Brown's Chemistry  States of Matter gas-liquid-solid revision notes

Part 1 The particle model and properties of the gases, liquids and solids, state changes and solutions

Revision KS4 Science IGCSE/O level/GCSE Chemistry Information Study Notes for revising for AQA GCSE Science, Edexcel 360Science/IGCSE Chemistry & OCR 21stC Science, OCR Gateway Science  (revise courses equal to US grades 9-10)

GCSE/IGCSE m/c QUIZ on states of matter

Part II Advanced work on gas law calculations, kinetic model theory of an IDEAL GAS and non-ideal gases

Sub-index for Part I sections (this page): 1.1 Three states of matter: 1.1a gases, 1.1b liquids, 1.1c solids * 2. State changes: 2a evaporation and boiling, 2b condensation, 2c distillation, 2d melting, 2e freezing, 2f cooling and heating curves and relative energy changes, 2g sublimation * 3. Dissolving, solutions. miscible/immiscible liquids

GCSE-AS(basic) KEYWORD index for Part I (this page): Boiling * Boiling point * Brownian motion * Changes of state * Condensing * Cooling curve * Diffusion * Dissolving * Evaporation * Freezing * Freezing point  * Gas particle picture * Heating curve * Liquid particle picture * Melting * Melting point * miscible/immiscible liquids * Properties of gases * Properties of liquids * Properties of solids * solutions * sublimation * Solid particle picture

Sub-index for Part II (on separate Advanced page): Section 4 Ideal gas behaviour and the gas laws: Introduction-the kinetic particle theory of an ideal gas * 4a. Boyle's Law * 4b. Charles's-Gay Lussac's Law and the combined gas law equation * 4c. The ideal gas equation PV=nRT * 4d. Dalton's Law of partial pressures * 4e. Graham's Law of diffusion * Section 5. Non-ideal real gas behaviour and Van der Waal's Equation: 5a. The deviations of a gases from ideal behaviour and their causes * 5b. The Van der Waals equation of state * 5c Compressibility factors * 5d The Critical Point - The Critical Temperature and Critical Pressure *

• A gas has no fixed shape or volume, but always spreads out to fill any container.
• 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 at random throughout the container so there is no order in the system.
• The particles move rapidly in all directions, frequently colliding with each other and the side of the container.
• With increase in temperature, the particles move faster as they gain kinetic energy.

• 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 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 (Pa = N/m²) - pressure is force/area on which force is exerted.
    • 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 and the net movement will be in the direction of lower concentration from a higher concentration, down the so-called diffusion gradient.
• 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 hits will not be even, so the smoke particle get a greater bashing in a random direction.
    • 
    • 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!). 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 molecule movement, it is also evidence that molecules of different molecular masses move/diffuse at different speeds.

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. 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 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. 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 below.

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A note on 'forces'

• Forces between particles are mentioned on this page and some ideas will seem more abstract than others - but think about it ...

  • A gas spreads everywhere in a given space, so there can't be much attraction between the molecules/particles.

  • Something must hold liquid molecules together or how can a liquid form from a gas?

  • In fact between liquid molecules there are actually weak electrical forces of attraction called intermolecular forces, but they can't be strong enough to create a rigid solid structure.

  • However, in solids, these forces must be stronger to create the rigid structure.

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 1.1b. The particle model of a Liquid

• 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.
• Particles quite close together but still arranged at random throughout the container, there is a little close range order as you can get clumps of particles clinging together temporarily.
• Particles moving rapidly in all directions but more frequently collisions with each other than in gases due to shorter distances between particles.
• With increase in temperature, the particles move faster as they gain kinetic energy, so increased collision rates, increased collision energy and increased rate of diffusion.

Using the particle model to explain the properties of a Liquid

• Liquids have a much greater density than gases (‘heavier’) because the particles are much closer together because of the attractive forces.
• 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 of the lack of ‘empty’ space between the particles.
• Liquids will expand on heating but nothing like as much as gases because of the greater particle attraction restricting the expansion (will contract on cooling).
  • 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 can be considerable!
• The natural rapid and random movement of the particles means that liquids ‘spread’ or 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.
• 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, convection 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.
  • The same thing happens with dropping copper sulphate crystals or coffee granules into water and just leaving the mixture to stand.
  • When pollen grains 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 is called 'Brownian motion' after a botanist called Brown first described the effect (see gases above). At any given instant of time, the hits will not be even all round the pollen grain, so they get a greater number of hits in a random direction.

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 1.1c. The particle model of a Solid

• 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.
• 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.

Using the particle model to explain the properties of a Solid

• 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.
• Solids will expand a little on heating but nothing like as much as liquids because of the greater particle attraction restricting the expansion and contraction occurs on cooling.
  • The expansion is caused by the increased energy of particle vibration, forcing them further apart causing an increase in volume and corresponding decrease in density.
• Diffusion is almost impossible in solids because the particles are too closely packed and strongly held together with no ‘empty space’ for particles to move through.

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

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)

• On heating particles gain kinetic energy and move faster.
• In evaporation* and boiling the highest kinetic energy molecules 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.
• Energy is needed to overcome the attractive forces in the liquid and is taken in from the surroundings.
• This means heat is taken in, so evaporation and boiling are endothermic processes (ΔH +ve).
• If the temperature is high enough boiling takes place.
• Boiling is rapid evaporation anywhere in the bulk liquid and at a fixed temperature called the boiling point and requires continuous addition of heat.
• 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.
• Energy changes for these physical changes of state for a range of substances are dealt with in a section of the Energetics Notes.

2b. Condensing (gas to liquid)

• On cooling, gas particles lose kinetic energy and eventually become attracted together to form a liquid.
• 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!

2c. Distillation

• 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.

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2d. Melting (solid to liquid)

• 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.
• 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.

2e. Freezing (liquid to solid)

• On cooling, liquid particles lose kinetic energy and so can become more strongly attracted to each other.
• Eventually at the freezing point the forces of attraction are sufficient to remove any remaining freedom and the particles come together to form the ordered solid arrangement.
• 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 of state changes gas <=> liquid <=> solid

2f(i) Cooling curve: 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 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. A cooling curve summarises the changes: gas ==> liquid ==> solid Energy changes for these physical changes of state for a range of substances are dealt with in a section of the Energetics Notes.

2f(ii) Heating curve: 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 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. A heating curve summarises the changes: solid ==> liquid ==> gas 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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2g. Sublimation

• 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. They usually involve just a physical change BUT its not always that simple!

• 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)

  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)

  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, the 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 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 (e.g. UK A2 advanced level), 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 Gobi Desert do not melt in the Sun, they just slowly 'sublimely' disappear!

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3. Dissolving solids, solutions and miscible/immiscible liquids

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• 3a. When a solid (the solute) dissolves in a liquid (the solvent) the resulting mixture is called a solution.

  • In general: solute + solvent ==> solution

  • 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 applies.

  • This means: mass of solid solute + mass of liquid solvent = mass of solution

  • You cannot make mass or lose mass, but just change it 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. 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. 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. This leaves behind the upper yellow layer liquid, so separating the two immiscible liquids.

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