Chapter   13 (Chapter 10 in your book) “States of Matter”
Section 13.1                               The Nature of Gases   OBJECTIVES:         Describe the assumptions of the “kin...
Measuring PressureThe first device for measuring atmosphericpressure was developed by Evangelista Torricelliduring the 17t...
 What happens when a substance is heated? _______________________________________________________________________________...
Name _______________________________________ Date ______________________                               13-1 Section Review...
Section 13.2                        The Nature of LiquidsOBJECTIVES:      Identify factors that determine physical proper...
 Particles left behind have _______________________ average kinetic energies;  thus the temperature _____________________...
 Normal boiling point- defined as the bp of a liquid ____________________________________________________________________...
Name _______________________________________ Date ______________________                                13-2 Section Revie...
Section 13.3                          The Nature of SolidsOBJECTIVES:      Evaluate how the way particles are organized e...
 When a liquid is cooled, a temperature is eventually reached at which the liquid                 begins to freeze. It ch...
Time in minutesName _______________________________________ Date ______________________                                13-...
Section 13.4                            Changes of StateOBJECTIVES:      Identify the conditions necessary for sublimatio...
a)P                       Critical(k                      PointresuesPr     Temperature (oC)                            14
Name _______________________________________ Date ______________________                               13-4 Section Review...
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Chemistry - Chp 13 - States of Matter - Notes

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Chemistry - Chp 13 - States of Matter - Notes

  1. 1. Chapter 13 (Chapter 10 in your book) “States of Matter”
  2. 2. Section 13.1 The Nature of Gases OBJECTIVES:  Describe the assumptions of the “kinetic theory” as it applies to gases.  Interpret gas pressure in terms of kinetic theory.  Interpret gas pressure in terms of kinetic theory.  Define the relationship between Kelvin temperature and average kinetic energy.  Kinetic refers to ____________________  The energy an object has because of it’s motion is called ____________________  The kinetic theory states that the tiny particles in all forms of matter are in __________________________________________  Three basic assumptions of the kinetic theory as it applies to gases: #1. Gas is ______________________________________- usually molecules or atoms  Small, hard spheres  Insignificant volume; relatively far apart from each other  No attraction or repulsion between particles #2. Particles in a gas move rapidly in ________________________________ motion  Move in straight paths, changing direction only when colliding with one another or other objects  Average speed of O2 in air at 20 oC is an amazing 1700 km/h! #3. Collisions are ______________________________________- meaning kineticenergy is transferred without loss from one particle to another- the total kinetic energyremains constant  Gas Pressure – defined as ____________________________________________ ____________________________________________________________________  Due to: a) ___________________________________________ b) ___________________________________________  No particles present? Then there cannot be any collisions, and thus no pressure – called a ________________________________  Atmospheric pressure results from the collisions of air molecules with objects  Decreases as you climb a mountain because the air layer thins out as elevation increases  ________________________________ is the measuring device for atmospheric pressure, which is dependent upon weather & altitude 2
  3. 3. Measuring PressureThe first device for measuring atmosphericpressure was developed by Evangelista Torricelliduring the 17th century.The device was called a “barometer” • Baro = _________________ • Meter = ________________  The SI unit of pressure is the _______________________________  At sea level, atmospheric pressure is about 101.3 kilopascals (kPa)  Older units of pressure include millimeters of mercury (mm Hg), and atmospheres (atm) – both of which came from using a mercury barometer  Mercury Barometer – Fig. 10.2, page 269 – a straight glass tube filled with Hg, and closed at one end; placed in a dish of Hg, with the open end below the surface  At sea level, the mercury would rise to 760 mm high at 25 oC- called one ______________________________________ (atm)  Equal pressures:1 atm = 760 mm Hg = 101.3 kPaExample ProblemA gas is at a pressure of 1.50 atm. Convert this pressure to a) kilopascals b) millimetersof mercury  Most modern barometers do not contain mercury- too dangerous  These are called aneroid barometers, and contain a sensitive metal diaphragm that responds to the number of collisions of air molecules  For gases, it is important to relate measured values to standards  Standard values are defined as a temperature of 0 oC and a pressure of 101.3 kPa, or 1 atm  This is called Standard Temperature and Pressure, or ____________ 3
  4. 4.  What happens when a substance is heated? ___________________________________________________________________________________________________  Some of the energy is stored within the particles- this is potential energy, and does not raise the temperature  Remaining energy speeds up the particles (increases average kinetic energy)- thus ________________________________________________ The particles in any collection have a wide range of kinetic energies, from very low to very high- but most are somewhere in the middle, thus the term ________________________________________________________is used  The higher the temperature, the wider the range of kinetic energies An increase in the average kinetic energy of particles causes the temperature to rise.  As it cools, the particles tend to move more slowly, and the average K.E. declines.  Is there a point where they slow down enough to stop moving? The particles would have no kinetic energy at that point, because they would have no motion  ________________________________ ( ____K, or __________ oC) is the temperature at which the motion of particles theoretically ceases  This has never been reached, but about 0.5 x 10-9 K has been achieved The Kelvin temperature scale reflects a __________________________________ between temperature and average kinetic energy  Particles of He gas at 200 K have twice the average kinetic energy as particles of He gas at 100 K Solids and liquids differ in their response to temperature  However, at any given temperature the particles of all substances, regardless of their physical state, have the same average kinetic energy What happens to the temperature of a substance when the average kinetic energy of its particles decreases? 4
  5. 5. Name _______________________________________ Date ______________________ 13-1 Section Review 1. According to the assumptions of kinetic theory, how do the particles in a gas move? 2. Use kinetic theory to explain what causes gas pressure. 3. Express the pressure 545mm Hg in kilopascals. 4. How do you raise the average kinetic energy of the water molecules in a glass of water? 5. A cylinder of oxygen gas is cooled from 300 K to 150 K. By what factor does the average kinetic energy of the oxygen molecules in the cylinder decrease? 5
  6. 6. Section 13.2 The Nature of LiquidsOBJECTIVES:  Identify factors that determine physical properties of a liquid.  Define “evaporation” in terms of kinetic energy.  Describe the equilibrium between a liquid and its vapor.  Identify the conditions at which boiling occurs. Liquid particles are also in motion.  Liquid particles are free to slide past one another  Gases and liquids can both FLOW  However, liquid particles ______________________________________ to each other, whereas gases are not Particles of a liquid spin and vibrate while they move, thus contributing to their average kinetic energy  But, most of the particles ________________________ have enough energy to escape into the gaseous state; they would __________________ ___________________________ their intermolecular attractions with other particles The intermolecular attractions also reduce the amount of space between particles of a liquid  Thus, liquids are more __________________ than gases  Increasing pressure on liquid has hardly any effect on it’s volume Increasing the pressure also has little effect on the volume of a solid  For that reason, liquids and solids are known as the ___________________________________________________________ _ Water in an open vessel or puddle eventually goes into the air The conversion of a liquid to a gas or vapor is called _______________________  When this occurs at the surface of a liquid that is not boiling, the process is called ________________________________  Some of the particles break away and enter the gas or vapor state; but only those with the minimum kinetic energy A liquid will also evaporate faster when heated  Because the added heat increases the average kinetic energy needed to overcome the attractive forces  But, evaporation is a __________________________________________ Cooling occurs because those with the highest energy escape first 6
  7. 7.  Particles left behind have _______________________ average kinetic energies; thus the temperature _________________________  Similar to removing the fastest runner from a race- the remaining runners have a lower average speed Evaporation helps to keep our skin cooler on a hot day, unless it is very humid on that day. Why? Evaporation of a liquid in a closed container is somewhat different  When some particles do vaporize, these collide with the walls of the container producing ___________________________________________ Eventually, some of the particles will return to the liquid, or _________________ After a while, the number of particles evaporating will equal the number condensing- the space above the liquid is now saturated with vapor  A dynamic equilibrium exists  Rate of evaporation = rate of condensation Note that there will still be particles that evaporate and condense  But, there will be no NET change An __________________________________________________ of a contained liquid increases the vapor pressure- the particles have an increased kinetic energy, thus more minimum energy to escape The vapor pressure of a liquid can be determined by a device called a “manometer”- Figure 10.2, p.277 The vapor pressure of the liquid will push the mercury into the U-tube A barometer is a type of manometer We now know the rate of evaporation from an open container increases as heat is added  The heating allows larger numbers of particles at the liquid’s surface to overcome the attractive forces  Heating allows the average kinetic energy of all particles to increase The boiling point (bp) is the temperature at which the ________________________________________________________________________________________________________________________________________________________________  Bubbles form throughout the liquid, rise to the surface, and escape into the air Since the boiling point is where the vapor pressure equals external pressure, the bp changes if the external pressure changes 7
  8. 8.  Normal boiling point- defined as the bp of a liquid _____________________________________________________________________________________________ Normal bp of water = ______________  However, in Denver = ______________, since Denver is 1600 m above sea level and average atmospheric pressure is about 85.3 kPa (Recipe adjustments?)  In pressure cookers, which reduce cooking time, water boils above 100 oC due to the increased pressure Autoclaves, devices often used in the past to sterilize medical instruments, operated much in a similar way – higher pressure, thus higher boiling point Boiling is a __________________________________ much the same as evaporation  Those particles with highest KE escape first Turning down the source of external heat drops the liquid’s temperature below the boiling point Supplying more heat allows particles to acquire enough KE to escape- the _________________________________________________________, the liquid only boils at a faster rate The heat of vaporization of a liquid ______________________________________________________________________________________________________________________________________________________________________________  For water at its normal boiling point of 100 degrees Celcius, the heat of vaporization is __________________________ The heat of condensation is _________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________  Condensation is exothermic  For water, the heat of condensation is is 2259 Joules per gram – the same as the heat of vaporization 8
  9. 9. Name _______________________________________ Date ______________________ 13-2 Section Review 1. Describe the nature of liquids. Refer to the role of attractive forces in your answer. 2. Use kinetic theory to explain the differences between the particles in a gas and those in a liquid 3. Use kinetic theory to explain the differences between evaporation and boiling of a liquid 4. Explain why the boiling point of a liquid varies with atmospheric pressure 5. Why does evaporation lower the temperature of a liquid? 6. What is the difference between heat of vaporization and the heat of condensation? 9
  10. 10. Section 13.3 The Nature of SolidsOBJECTIVES:  Evaluate how the way particles are organized explains the properties of solids.  Identify the factors that determine the shape of a crystal.  Explain how allotropes of an element are different. Particles in a liquid are relatively free to move  Solid particles are not solid particles tend to ___________________________________________, rather than sliding from place to place Most solids have particles packed against one another in a highly organized pattern  Tend to be dense and incompressible  Do not flow, nor take the shape of their container Are still able to move, unless they would reach absolute zero When a solid is heated, the particles vibrate more rapidly as the kinetic energy increases  The organization of particles within the solid breaks down, and eventually the solid melts The melting point (mp) is ____________________________________________ At the melting point, the disruptive vibrations are strong enough to overcome the interactions holding them in a fixed position  Melting point can be reversed by cooling the liquid so it freezes  Solid ↔ liquid Generally, most ionic solids have ______________________________________, due to the relatively strong forces holding them together  Sodium chloride (an ionic compound) has a melting point = 801 oC Molecular compounds have relatively low melting points Hydrogen chloride (a molecular compound) has a mp = -112 oC Not all solids melt- wood and cane sugar tend to decompose when heated Most solid substances are crystalline in structure 10
  11. 11.  When a liquid is cooled, a temperature is eventually reached at which the liquid begins to freeze. It changes into a solid  This temperature, which remains constant until all the liquid has solidified at 1 atmosphere pressure, is called the ____________________________________ of the liquid  While the liquid is cooling, the average kinetic energy of its particles decreases until it is low enough for the attractive forces to be able to hold the particles in the fixed positions characteristic of the solid phase Heat of solidification  The amount of heat needed to change ___________________________________ ____________________________________________________________________ ____________________________________________________________________ __  Heat of solidification = 334 Joules Heat of fusion  The amount of heat needed to change a unit mass of a substance from solid to liquid at STP  Heat of fusion of ice = 334 Joules Heating and cooling curvesuslsiCe 200in 180e 160 140urratpe 120m 100Te 80 60 40 20 0 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 11
  12. 12. Time in minutesName _______________________________________ Date ______________________ 13-3 Section Review 1. Explain the nature of solids and tell why they differ from liquids. Refer to the organization of particles in your answer. 2. Would you expect MgCl2 to have a high or low melting point? Explain your answer. 3. What is the difference between the heat of fusion and the heat of solidification? 4. Describe what happens when a solid is heated to its melting point. 12
  13. 13. Section 13.4 Changes of StateOBJECTIVES:  Identify the conditions necessary for sublimation.  Describe how equilibrium conditions are represented in a phase diagram.  Describe how equilibrium conditions are represented in a phase diagram. Sublimation- ________________________________________________________________________________________________________________________________________________________________________________________________  Examples: iodine, dry ice (-78 oC); mothballs; solid air fresheners Sublimation is useful in situations such as freeze-drying foods- such as by freezing the freshly brewed coffee, and then removing the water vapor by a vacuum pump Also useful in separating substances - organic chemists use it separate mixtures and purify materials The relationship among the solid, liquid, and vapor states (or phases) of a substance in a sealed container are best represented in a single graph called a phase diagram Phase diagram- _______________________________________________________________________________________________________________________________________________________________________________________________ The diagram below shows the phase diagram for water  Each region represents a pure phase  Line between regions is where the two phases exist in equilibrium  _________________________________ is where all 3 curves meet, the conditions where all 3 phases exist in equilibrium 13
  14. 14. a)P Critical(k PointresuesPr Temperature (oC) 14
  15. 15. Name _______________________________________ Date ______________________ 13-4 Section Review 1. What general information can you get from a phase diagram for water at various temperatures and pressures? 2. Describe the process of sublimation. What is a practical use of this process? 3. Explain triple point. 15

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