SUPPLEMENTARY TOPIC 1 STATES OF MATTER.

SUPPLEMENTARY TOPIC 1 STATES OF MATTER. States of matter. Matter is generally regarded as existing in three possible phases - solid, liquid or gas. For example, water can be in the solid form (ice) or the familiar liquid form or the gaseous form (water vapour). Each physical state has its own associated properties. The SOLID STATE is characterised by having a fixed shape with rigid boundaries. The LIQUID STATE still has a definite volume and does maintain a boundary but that boundary is not fixed, instead adapting to the container. The GASEOUS STATE has no fixed volume - the component molecules move throughout the container to uniformly fill it. If the gas is not trapped in a container, it spreads throughout the atmosphere. As an example, a smelly chemical demonstration done at the front ultimately makes its presence smelt at the rear of the lecture theatre. These properties which are associated with each state can be explained by the KINETIC MOLECULAR THEORY OF MATTER. Note that when substances undergo changes of state such as those illustrated by water above, the H 2 O molecules themselves are not broken down to component atoms. It is only the weak forces that operate between the molecules which are overcome as the molecules gain more energy, thereby allowing them to separate from each other and escape into the gaseous phase. Thus water is still present as H 2 O molecules in the solid, liquid and gas phases, but in the gaseous phase the molecules are far apart and move about independently of each other instead of having reduced freedom to move as in the liquid state or being anchored in a fixed position as in the solid. The kinetic molecular theory. (a) Gases. In this theory, the molecules of a substance in the gas phase are regarded as being in constant motion, moving at very high speeds and colliding with each other or with the walls of their container many times each second. The speed of the molecules and thus their energy increases with increasing temperature. The molecules are regarded as having no volume of their own and it is assumed there are no attractive forces between them. The molecules are travelling rapidly and have enough energy for them to rebound from collisions without any loss of energy. The collisions with the walls of the container are the cause of gas pressure - the more collisions per second that occur, the greater is the resultant pressure. For example, a H 2 molecule at 0 o C is moving at 5 times the speed of sound and colliding with other molecules at the rate of 5 10 10 collisions per second. The above is only a model. However, this model does approximate well to the behaviour of real gases under normal temperature and pressure conditions. ST1-1

ST1-2 The molecules of all real gases of course do have a finite volume, although it is indeed only a tiny fraction of the container volume at normal pressures but is significant at high pressures. Most of the volume of a container of a gas is actually empty space with the real volume of the gaseous molecules at normal pressures being less than 1%. Thus gases are relatively easy to compress - e.g. using a bicycle pump to blow up a tyre. Real gases also experience attractive forces between their molecules which causes the collisions to be "sticky" to some extent and this becomes significant at low temperatures. However, the kinetic molecular model is adequate to explain and predict the behaviour of gases provided the temperatures are not very low or the pressures not very high. (b) Liquids. In the liquid phase, the component molecules are much closer together than in the gaseous state. This is seen in the fact that 1 ml of liquid water converted to steam at the boiling point provides 1670 ml of the gas. The amount of empty space in a liquid is about 3 % of the total volume and consequently liquids are almost incompressible. In the liquid phase there is still TRANSLATIONAL MOVEMENT of the molecules, but this is very much slower and more restricted than for gases. Because the molecules in the liquid phase are much closer together, the attractions between them are much stronger, giving rise to the phenomenon of surface tension and thus to the defined boundary of the liquid. Surface tension is a consequence of unbalanced attractive forces experienced by the outermost layer of molecules in the liquid. This is well illustrated by the observation that drops of liquid such as water adopt a spherical shape in which the surface area is smallest for any given volume. (c) Solids. The component atoms, molecules or ions in the solid are anchored by even stronger forces than those in the liquid phase. They are held at fixed positions in a crystal lattice and are unable to move within the solid. Instead they can only undergo motions such as vibration which do not require any translational movement. Hence the solid has a rigid boundary and a fixed volume with negligible empty space and cannot be compressed to any significant extent..

ST1-3 Conversion between states - latent heat. Consider water in the liquid state being heated. The temperature rises until the water starts to boil at a temperature of 100 o C at sea level. Measurements would show that 4.18 J of energy would be required for each gram of water heated by 1 o C. Therefore to heat 1 mole (18 g) of water from 25 o C to 100 o C would require 18 4.18 75 J of heat = 5.6 kj. This heat is all used to increase the motion of the water molecules which move faster at higher temperatures. Once the boiling point is attained, if heat continues to be supplied, there is no increase in temperature but instead the liquid converts to the vapour (gaseous) state at 100 o C. The energy being supplied at this point is used by water molecules to escape the attractive forces of other water molecules in the liquid phase and undergo the transition to the gas phase. If heating is carried out for a long enough time, all the liquid sample will be converted to gas and only after this has occurred will the temperature of the steam start to rise again assuming it is trapped. If this experiment were carried out under conditions which allowed the energy supplied to be measured, it would be found that each mole of water required 41 kj for the conversion from liquid to gas at 100 o C. This very large amount of heat energy was used by the water molecules to overcome the attractive forces which exist between them in the liquid state, especially the forces known as hydrogen bonding (see Topic 13). Because no temperature rise is observed, this energy used for the transition from liquid to vapour state is called the "LATENT HEAT OF VAPORISATION". All liquids require the latent heat of vaporisation to be supplied in order to make the transition from the liquid to the vapour, although most have a smaller value than water. The latent heat of vaporisation of water is exceptionally large because of the extensive hydrogen bonding forces which must be broken before individual molecules can escape to the vapour state. Similarly, consider water in the solid phase (ice). Ice melts at 0 o C to form liquid water at the same temperature. Again, heat must be supplied merely to achieve the conversion from solid to liquid phase without any increase in temperature. In this case the amount of heat required is only 6 kj per mole of ice melted, and this is called the "LATENT HEAT OF FUSION". The energy required for ice to melt is much less than that required for water to boil because hydrogen bonds between H 2 O molecules are retained in the melting process but must be broken for the molecules to enter the gaseous state. In both these examples, heat was required for the conversions of liquid to gas or solid to liquid. For the reverse conversions, heat would be released. Thus steam condensing to liquid water at 100 o C releases the same amount of heat as was absorbed in the vaporisation process, viz 41 kj per mole. This is the basis for steam

ST1-4 sterilisation where the large amount of latent heat released is used to kill microorganisms. In order to convert liquid water to ice, heat must be removed, again equal to that amount absorbed when ice melts, viz 6 kj per mole. These processes can be summarised as follows: heat absorbed melting boiling SOLID heat of fusion LIQUID heat of vaporisation GAS freezing condensing heat released In some situations, solids may convert directly to the gaseous state in a process called "SUBLIMATION" in which case the heat involved would be the sum of the heats of vaporisation and fusion. Liquefaction of gases. All gases can be converted to the liquid state. This is achieved by lowering their energy sufficiently for the attractive forces between the molecules to cause them to stick together in large aggregates, forming droplets which ultimately coalesce. The energy of the gaseous molecules is lowered by cooling the gas. High pressure also can be used to force the molecules closer together and by these means allow the attractive forces to suppress the motion of the gaseous molecules. Examples of gases converted to the liquid phase and stored under pressure at room temperature are liquid butane and propane present in LPG cylinders and used as fuel for portable stoves and motor vehicles. Another example of gas liquefaction is seen in the formation of clouds from water in the gaseous form. Clouds contain fine droplets of moisture that have condensed from gaseous water in the atmosphere and which under the appropriate conditions coalesce to form rain drops or even condense to the solid form as hail or snow. The large amount of heat released in the condensation of water vapour to liquid phase has significant effects on the weather. Evaporation compared with boiling. Liquids can convert to the gas phase without being boiled by heating directly. The level of water left in a glass in the open gradually falls and ultimately the water disappears completely. This process is called EVAPORATION and only differs in that the molecules leaving the liquid phase do so from the surface whereas, in boiling, the molecules close to the heat source gain the energy to convert to gas and form bubbles which then rise through the liquid and escape to the surroundings. Both processes require exactly the same amount of heat to occur, but evaporation is much

ST1-5 slower because the heat is gained gradually from the surroundings rather than quickly from an applied heat source. Evaporation is accompanied by cooling of the liquid, and this is the basis for evaporative coolers and also for the body's own cooling mechanism whereby sweat from pores on the surface of the skin evaporates, absorbing 41 kj of heat energy per mole as it does so. Air motion above the surface of a liquid speeds up the process of evaporation by removing gas phase molecules and preventing their return to the liquid. Check your understanding of this section. On a hot day a person may consume an additional litre of water which is subsequently converted to sweat. Calculate the heat removed from the body by the evaporation of this sweat, taking the mass of 1.0 L of water = 1.0 kg. Effect of pressure on boiling point. Because it is ever present, we tend to be oblivious to the high pressure resulting from the atmosphere above us - equivalent to 1.0 kg on each square centimetre of our body. Pressure can be expressed in many units, often as atmospheres but in SI units, pressure is measured in pascals (Pa). The reduction in pressure with increasing height above sea level requires pressurisation in high flying aircraft and presents difficulties for mountain climbers on tall peaks. Athletes competing at venues such as Mexico City require considerable time there to acclimatise to the lower pressure of oxygen available. As the external pressure decreases, molecules can escape from liquid to gas phase more readily, requiring less energy and thus boiling at a lower temperature. At sea level where average atmospheric pressure = 101.3 kpa, pure water boils at 100 o C but at the top of Mt. Kosciusko (2230 m) it boils at 92.6 o C due to the pressure being reduced to 77.3 kpa. Atop Mt.Everest (8840 m) the boiling point would be a mere 69.6 o C at a pressure = 30.7 kpa. The same principle also applies in reverse - if external pressure is increased then the boiling point will increase as the escaping molecules require more energy (ie higher temperature) to leave the surface of the liquid. Use of this is made in pressure cookers and autoclaves where water may be converting to steam at temperatures as high as 121 o C. The combination of the latent heat of vaporisation released when the steam condenses and the high temperature causes cooking to be faster and, in the case of autoclaves, is sufficient to destroy all microorganisms - even heat resistant bacterial spores. Sublimation. SUBLIMATION is the name given to the process whereby a solid is converted directly to the gas phase without first forming the liquid phase. The reverse process, from gas directly to solid, is termed DEPOSITION.

ST1-6 Sublimation occurs when the forces that hold the molecules to each other in the solid state are very weak - thus ionic compounds will not sublime at normal pressures, as the electrostatic forces that bond ions in an ionic crystal lattice are strong. However, many covalently bonded compounds often have only weak forces operating to hold the molecules close to each other in the solid state. Some of these compounds have such weak forces between the molecules that, when heated sufficiently to break down the solid, the molecules can escape immediately into the gaseous phase. Carbon dioxide is a commonly known compound that undergoes sublimation, this property being used in "dry ice" because it converts from the solid state to the gaseous state without passing through the liquid phase. Again, note that when sublimation occurs, the molecules themselves do not break down to their component atoms. Only the weak forces operating between the molecules break, thereby allowing them to separate from each other and escape into the gaseous phase. Unlike water, carbon dioxide does not have hydrogen bonds operating between the CO 2 molecules and so it can undergo the change from solid to gas directly at room conditions without passing through the liquid state. Latent heat and the weather. The heat energy absorbed when water evaporates and released when water vapour condenses is a very significant factor controlling the planet's weather. When water evaporates from the oceans and bodies of fresh water, the latent heat of vaporisation (41 kj for every 18 ml of water evaporating) is absorbed by the evaporating water molecules. This heat energy is obtained mainly as infra-red radiation from the sun. The water vapour so produced rises in the atmosphere and ultimately condenses to form water droplets (rain) or the solid phase (hail, snow). When the water condenses in the upper atmosphere, it necessarily releases heat energy equal to the latent heat of vaporisation which was absorbed when the evaporation occurred. Thus the process of evaporation and condensation of the water transfers vast amounts of energy up into the atmosphere and this energy in turn is the driving force for destructive storms and floods. Gases such as carbon dioxide in the atmosphere assists in increasing the amount of the sun's energy that is available for the evaporation process. Concern about the rise in Earth's temperature underlies attempts to reduce the amount of carbon dioxide released. Molar volumes of liquids and solids. The volume of one mole of any element or compound is called its MOLAR VOLUME. In Topic 9 the interconversion of mass and moles was employed to calculate the masses of reactants and products involved in chemical reactions. The mass and volume of any substance is connected by its DENSITY. The density of any object is defined as the ratio of its mass and volume. Using the symbols m for mass, V for volume and d for density, this relationship can be written as d = m/v. The density

ST1-7 of any pure substance is a characteristic of that substance and tabulations of densities are available for most elements and compounds. The volume of most materials increases as the temperature rises because the motion of the component atoms, molecules or ions increases due to the extra energy supplied as heat. Consequently the density of liquids and solids depends to a small extent on the temperature and this must be specified along with its numerical value. For example. the density of liquid water has the following values at various temperatures: Temperature ( o C) Density (g cm 3 ) 0 0.9998 5 1.0000 10 0.9997 20 0.9982 30 0.9956 100 0.9583 The units used to express density must be mass per unit of volume. Generally in chemistry the units used are grams for mass and cubic centimetres for volume, but millilitres are often used instead of cubic centimetres because they are identical. Molar volumes of gases. While the molar volume of any solid or liquid at a given temperature depends on the specific element or compound, for gases the volume of a mole of any gas at a given temperature and pressure is always the same. This is because the attractive forces between gas phase molecules are so weak that gases expand to completely and uniformly fill their container regardless of which gas is present, so the volume of a mole of gas is determined only by the temperature and pressure. In the next Topic, this property is further examined. Example. Calculate the volume of 1.00 mole of water at 30 o C given the density of water at that temperature is 0.9956 g cm 3. Density = mass volume Molar mass of H 2 O = 18.0 g 0.9956 = 18.0 / molar volume molar volume = 18.0 / 0.9956 = 19.0 cm 3 1 ml = 1 cm 3, molar volume = 19.0 ml

ST1-8 Objectives of this Topic. When you have completed this Topic and its associated tutorial questions, you should have achieved the following goals: 1. Know the three common states of matter and their distinguishing features. 2. Know the principles of the kinetic molecular theory of solids, liquids and gases. 3. Understand why (latent) heat changes accompany phase changes. 4. Understand the processes by which gases liquefy and liquids boil or evaporate and how boiling differs from evaporation. 5. Understand how pressure affects the boiling point of a liquid. 6. Understand why some solids sublime when heated at normal pressures. 7. Recognise the role of latent heat associated with phase changes of water as a factor affecting the weather. 8. Be familiar with the concept of molar volume and understand why the molar volume of a gas is independent of the particular gas examined. SUMMARY. Matter may exist in the solid, liquid or gaseous phases. Each state has its own properties which are a consequence of how far apart the component particles (atoms, molecules or ions) are and how fast they are moving. When molecular substances undergo a phase change, their molecules remain intact and do not decompose to constituent atoms. The kinetic molecular theory for an ideal gas envisages gases as consisting of rapidly moving particles separated by relatively large distances and colliding without loss of energy, due to the absence of attractive forces between them. This model applies to idealised gases where there are no attractive forces operating between the constituent particles and those particles have no volume. Constituent particles of real gases do occupy a volume and do experience attractive forces between themselves. However, while this is only a model, it adequately explains the properties of real gases at normal conditions of temperature and pressure. In the liquid state, the component particles are much closer together with stronger attractions between them. Translational motion is present but restricted in the liquid state. Particles in solids have even less freedom to move with their motion being

ST1-9 restricted to vibrations and rotations with no translational motion. Consequently solids maintain a fixed shape. The conversion between different physical states is accompanied by energy changes - energy absorbed for solid to liquid or liquid to gas phase changes and an identical amount of energy released for the reverse phase changes. These energy changes are called "latent heats" because they are not accompanied by any temperature change, being related solely to the changed distances between the particles as the phase change occurs. Gases are liquefied by lowering the energy of the particles by cooling (to slow the particles) and using high pressure (to force the particles closer together). Ultimately the attractive forces between the atoms or molecules cause them to clump together as a liquid. The reverse process is boiling or evaporation, both of which absorb energy, but evaporation occurs at temperatures below the boiling point and results in cooling of the remaining liquid because the highest energy particles escape from the surface to the gas phase. Boiling occurs at the site where the heat source is applied and the temperature of the boiling liquid remains constant because more heat is supplied to the liquid from the source to replace the energy removed by the escaping gas phase component. The temperature at which boiling occurs is influenced by the external pressure on the surface of the liquid. Higher external pressure raises the boiling point of a liquid and lower pressure decreases it. Some solids can be converted directly to the gas phase at normal pressures in the process called sublimation. This occurs when the forces operating between the molecules are particularly weak. The latent heat associated with the evaporation and condensation of water is a very significant factor associated with the earth's weather. The volume of a given mass of any substance can be calculated from its density. The density of all pure substances is a characteristic and for most elements and compounds, is tabulated at various temperatures. This is necessary because the volume of most materials increases with temperature due to faster motion of the component particles. If the amount of an element or compound is one mole, then the volume is its molar volume. Unlike solids or liquids however, the molar volume of all gases is the same at a specified temperature and pressure because the particles of gas experience such small attractive forces that any gas expands to completely fill its container. Recommended follow up chemcal modules: Section: General Chemistry Module: States of Matter Topic: Vapour pressure of liquids, phase diagrams and their interpretation.

ST1-10 TUTORIAL QUESTIONS - SUPPLEMENTARY TOPIC 1. 1. Name the three common states of matter and give the characteristics of each state. 2. What are the assumptions made as the basis for the kinetic molecular theory of gases? 3. In what ways do real gases deviate from the ideal gases described by the kinetic molecular model? 4. What is the cause of pressure in a gas trapped in a container? How does the temperature of the gas influence the pressure? 5. What is the origin of surface tension? Why are drops of liquid spherical? 6. Why are liquids almost incompressible but gases can be compressed easily? 7. Why do solids maintain a fixed shape whereas liquids and gases do not? 8. Water reaches its boiling point when heated to 100 o C at atmospheric pressure. Further heating causes the water to continue boiling without any temperature increase. What happens to the heat being supplied? 9. Why does water have an exceptionally large heat of vaporisation?

ST1-11 10. To melt 1 mole of ice at 0 o C requires supplying 6 kj of heat. How much heat would have to be removed from 1 mole of water at 0 o C to freeze it? 11. Why are steam burns much more severe than scalding by boiling water? 12. Dry ice (solid carbon dioxide) is often used to keep some types of food frozen. Give a reason why it may be preferred to normal ice? 13. What is the advantage associated with liquefying natural gas? How is this done? 14. Why is water vapour able to condense to droplets after it rises into the upper atmosphere? 15. Compare the processes of evaporation and boiling. Indicate in what ways they are similar and how they differ. 16. How do evaporative coolers work? 17. A can containing a little water is boiled to generate steam. If the cap is then screwed on, the can collapses as it cools. Why does this occur? 18. How could one make a really hot cup of tea on top of Mt.Everest?

ST1-12 19. How could you demonstrate the existence of atmospheric pressure? 20. What property of a compound causes it to sublime when heated? Give an example of a compound that sublimes at atmospheric pressure. 21. Calculate the molar volumes at 25 o C of: (a) iodine (density at 25 o C = 4.9 g cm 3 ) (b) sodium chloride (density at 25 o C = 2.2 g cm 3 ) ANSWERS TO TUTORIAL TOPIC ST1 1. Solid, liquid and gas. In the solid state, the component particles (atoms, molecules or ions) are restricted in their motions to only vibrations due to strong forces between them. Consequently solids retain a fixed shape and are almost incompressible. In the liquid state, the constituent particles have some degree of translational motion because the attractive forces are not as strong as in the solid state. Even so, liquids are also almost incompressible and the forces are strong enough for liquids to maintain defined boundaries although unlike solids, no fixed shape. In the gaseous state the attractive forces between the constituent particles are much weaker that those operating for the same substance in either the solid or liquid states. The constituent particles in the gaseous state are in constant and rapid motion. Consequently gases have no fixed volume or defined boundaries and will expand to completely and uniformly fill their container. 2. The kinetic molecular theory of gases assumes that the component particles are in rapid motion in straight lines and frequently experience perfectly elastic collisions with each other and with the walls of their container with no loss of energy. Pressure from gases is envisaged as a result of those particles colliding

ST1-13 with the walls of the container. Another assumption of the model is that the constituent particles of ideal gases have zero volume. 3. The atoms or molecules of real gases do exhibit some degree of interaction when collisions between them occur, sometimes called sticky collisions. They also do have a finite volume. However, the deviation from ideal gas behaviour caused by these factors is not significant at normal temperature and pressure conditions. 4. Gas atoms or molecules are in rapid motion and collide with the walls of their container, giving rise to the observed pressure attributed to the gas. At higher temperatures the particles have more energy, move faster and impart a greater pressure on the container s walls. 5. In the liquid state, attractive forces operate between adjacent molecules even though the molecules have some degree of translational motion. For molecules that are not near to the surface of the liquid, these forces cancel out in all directions overall so there is no nett force experienced by them. However, those molecules in the outer layer of the liquid experience unbalanced forces due to there being no molecules of the liquid outside that layer. The result is that the outer layer of molecules acts like a skin on the surface of the liquid and this is the origin of surface tension. Liquids typically form spherical drops because this shape contains the largest volume of liquid for the smallest surface area, thereby minimising the overall forces that the outer molecules experience 6. The molecules in a liquid are very close together, typically occupying about 97 % of the total volume. Consequently there is little scope to force them to be closer together by compressing the liquid. In a gas, the component particles are far apart and at room temperature and pressure their actual volume is only about 1% of their container. Consequently there is considerable scope to compress a gas. The more pressure applied to gases, the closer their component particles are forced together and the more difficult it is to further compress the gas into a smaller volume. 7. In a solid, there is almost no void space between the component particles and the forces of interaction between them is so strong that they have no translational motion, so their boundaries do not change. The component particles in liquids and gases are further apart and do have translational motion, much more so in the gas phase than the liquid. In both cases, that translational motion means that they have no fixed shape.

ST1-14 8. After liquid water has been heated to its boiling point, further heating continues to supply energy to the water molecules and they use that energy to escape from the liquid phase to the vapour phase. The energy is used to overcome the attractive forces that exist between the molecules in the liquid, especially the forces called hydrogen bonding. Consequently the additional heat energy supplied does not lead to a rise in temperature because it is being used in the vaporisation process. This heat energy is called the latent heat of vaporisation and for water and is equal to 41 kj per mole of water. 9. Water has a particularly large heat of vaporisation because it has hydrogen bonds attracting adjacent H 2 O molecules to each other as well as the weaker types of intermolecular forces which all molecules in the liquid phase experience. 10. The energy difference between solid water and liquid water at its melting/freezing point is 6 kj. This amount of energy is required to be supplied to the ice to melt it or must be removed from liquid water to freeze it. The actual amount of heat to be transferred is the same, regardless of the direction of the phase change. 11. Burns from steam are more severe than burns from water at its boiling point because steam condensing on the skin releases the latent heat of vaporisation of water, 41 kj per mole. Liquid water at that temperature has none of this latent heat stored so less heat is released to the skin and the burn is less severe (but still painful!). 12. The advantage of using dry ice rather than normal water ice is that the dry ice sublimes at atmospheric pressure and no liquid is formed in the process. This avoids the material being cooled becoming wet as it would when water ice melts to form liquid water. 13. Gases occupy a very large volume for a given quantity of the substance compared with the volume required for the same amount of that substance in the liquid phase. Natural gas, methane, is much more easily transported as a liquid using cooled tankers. It is liquefied by using high pressure to compress the gas, forcing the molecules closer together and also the gas is cooled so that the molecules move more slowly. This combination allows the attractive forces between the CH 4 molecules to cause clumping and finally liquefaction. 14. The temperature of the atmosphere is lower at higher altitudes. The lower temperature allows water molecules in the gaseous phase to lose energy and move more slowly, ultimately condensing as liquid water drops.

ST1-15 15. Evaporation and boiling of a liquid both result in liquid phase molecules escaping to the vapour phase by overcoming the attractive forces that hold them together in the liquid. When a liquid is boiled, heat is supplied from a heat source such as a gas burner and the molecules close to the source of the heat receive the extra energy which allows them to move to the surface and escape as a gas. Evaporation occurs when molecules of the liquid near to its surface are able to absorb sufficient heat from the surroundings to make the transfer to the gas phase. Because the most energetic molecules escape in this way, the temperature of the remaining liquid decreases and the result is evaporative cooling of the liquid, a process used by humans to cool their bodies by sweating. 16. Evaporative coolers make use of the large heat of vaporization (latent heat) of water to cool a stream of air blown over water. The latent heat absorbed by the water molecules as they transfer to the vapour state is enhanced by the greater rate of evaporation resulting from the air travelling across the water s surface. Because the air moving across the water source contains more molecules of water vapour than the surrounding air, the humidity is increased. 17. When the water is boiled, gas phase water molecules displace the molecules in the air within the can. If the can is then sealed and cooled, the water molecules condense back to the liquid phase where they occupy far less volume. This results in the pressure of the atmosphere outside the can no longer being balanced by the pressure inside it, leading to the can being crushed by the weight of air outside. 18. The boiling point of a liquid increases as the external pressure on its surface increases, because the molecules need more energy to escape from the surface. Conversely, if the external pressure is reduced, molecules of a liquid requires less energy to escape to the vapour phase and so will boil at a lower temperature. This is why the boiling point of water on a high mountain is less then the normal 100 o C experienced at sea level. One can elevate the boiling point of water by heating it in a sealed container called a pressure cooker. This is used sometimes in domestic cooking to cook food faster (and destroy it in the process!) and could be used to boil water for the tea at a suitably high temperature. 19. Air pressure can be demonstrated by taking a glass tube that is sealed at one end, filling it with water and inverting the tube into a beaker of water. The bulk of the water remains in the tube, supported by the pressure of the atmosphere on the surface of the water in the beaker. This is the principle underlying the barometer which is used to measure air pressure. Mercury is used in a

ST1-16 barometer because its large density reduces the length of the tube required. At normal atmospheric pressure, the length of the column of mercury required is about 76 cm high. Atmospheric pressure would support a column of water about 16 times this length. 20. Sublimation is the process whereby a solid converts to the gas phase without passing through an intermediate liquid state. For this to happen, the forces of attraction between constituent particles in the solid must be particularly weak or the external pressure very low. Solid carbon dioxide is a common example of a substance that sublimes at atmospheric conditions. 21. (a) density = mass volume volume = mass density mass of 1.0 mole of iodine = 2 126.9 g = 254 g density of iodine at 25 o C = 4.9 g cm 3 volume of 1.0 mole of iodine at this temperature = 254 4.9 = 52 cm 3 (b) mass of 1.0 mole of sodium chloride = 23.0 + 35.5 = 58.5 g density if sodium chloride at 25 o C = 2.2 g cm 3 volume of 1.0 mole of sodium chloride = 58.5 2.2 = 27 cm 3