1. Chapter 13 States of Matter

2. The Nature of Gases Kinetic – means motion
Kinetic Energy – the energy an object has because of its motion.
Kinetic Theory – all matter consists of tiny particles that are in constant motion.

3. Kinetic Theory as Applied to Gases
Fundamental assumptions about gases:
The particles in a gas are considered to be small, hard spheres with an insignificant volume.
Between particles in a gas there is empty space.
No attractive or repulsive forces exist between the particles.

4. Kinetic Theory as Applied to Gases
Fundamental assumptions about gases:
The motion of the particles in a gas is rapid, constant, and random.
Gases fill their container regardless of the shape and volume of the container.
Particles travel in straight-line paths until they collide with another particle or another object such as the wall of their container.

5. Kinetic Theory as Applied to Gases
The average speed of oxygen molecules in air at 20º is 1700km/h
At this speed, the odor from a hot pizza in Washington, D.C., should reach Mexico city in about 115 minutes. Does this happen?
No, the odor molecules are constantly striking molecules in air and rebounding in other directions.

6. Kinetic Theory as Applied to Gases
Fundamental assumptions about gases:
All collisions between particles in a gas are perfectly elastic.
during a perfectly elastic collision, kinetic energy is transferred from one particle to another and
the total kinetic energy remains constant.

7. Gas Pressure
Gas pressure is the result of simultaneous collisions of billions of rapidly moving particles in a gas with an object.
Ex – a helium-filled balloon maintains its shape because of the pressure of the gas within it.
Vacuum – an empty space with no particles and no pressure. (no particles, no collisions)

8. Atmospheric Pressure
Atmospheric pressure results from the collisions of atoms and molecules in air with objects.
Atmospheric pressure decreases as you climb a mountain because the density of Earth’s atmosphere decreases as elevation increases.
less particles, less pressure

9. Atmospheric Pressure
Barometer a device that is used to measure atmospheric pressure.
Atmospheric pressure depends on weather and on altitude.
At sea level and with fair weather, the atmospheric pressure is sufficient to support a mercury column about 760 mm Hg high
On Mount Everest the air exerts only enough pressure to support 253mm Hg

10. SI unit of pressure is the pascal (Pa)
Gas Pressure
SI unit of pressure is the pascal (Pa)
Represents a very small amount of pressure, for example normal atmospheric pressure is about 100,000 Pa or 100 kPa
Two other units of pressure are commonly used
Millimeters of mercury (mm Hg)
Standard atmosphere (atm)

11. Gas Pressure
1 atm = 760 mm Hg = kPa

12. Kinetic Energy and Temperature
As a substance is heated, its particles absorb energy, some of which is stored within the particles.
The stored portion of the energy (potential energy) does not raise the temperature of the substance.

13. Average Kinetic Energy
Particles in any collection of atoms or molecules at a given temperature have a wide range of kinetic energies.
Most of the particles have kinetic energies somewhere in the middle
Average kinetic energy is used when discussing the kinetic energy of a collection of particles in a substance.

14. Average Kinetic Energy & Temperature
At any given temperature the particles of all substances, regardless of physical state, have the same average kinetic energy.
Ions in table salt (s), molecules in water (l) and atoms in helium (g) all have the same average kinetic energy at room temperature even though the three substances are in different physical states.

15. Average Kinetic Energy & Temperature
An increase in the average kinetic energy of the particles causes the temperature of a substance to rise.
As a substance cools, the particles tend to move more slowly and their average kinetic energy declines.
Absolute zero (0K or ºC or -459ºF) is the temperature at which the motion of particles theoretically ceases.

16. Average Kinetic Energy & Kelvin Temperature
The Kelvin temperature of a substance is directly proportional to the average kinetic energy of the particles of the substance.
Particles in helium gas at 200K have twice the average kinetic energy as the particles in helium gas at 100K

17. Air rises near low-pressure areas.
As air rises, it cools and often condenses into clouds and precipitation.
Usually, when forecasters say a low-pressure area is moving toward your region, cloudy weather and precipitation often result as the low-pressure area approaches.

18. End of section 13.1

19. The Nature of Liquids
Kinetic Theory says both the particles in gases and liquids have kinetic energy allowing them to flow past one another.
Substances that flow are referred to as liquids
Ability of gases and liquids to flow allows them to conform to the shape of their containers.

20. Key difference between gases and liquids
The Nature of Liquids
Key difference between gases and liquids
kinetic theory says there are no attractions between the particles in a gas
particles in a liquid are attracted to each other
intermolecular attractions keep the particles in a liquid close together

21. The Nature of Liquids
The interplay between the disruptive motions of particles in a liquid and the attractions among the particles determines the physical properties of liquids.

22. Properties of Liquids
Intermolecular attractions reduce the amount of space between particles in a liquid.
liquids are much more dense than gases
Increasing the pressure on a liquid has hardly any effect on its volume.
Known as a condensed state of matter

23. Evaporation Vaporization – conversion of a liquid to a gas or vapor
Evaporation – when conversion from a liquid to a gas or vapor occurs at the surface of a liquid that is not boiling.
Most molecules in a liquid don’t have enough KE to overcome the attractive forces and escape into the gaseous state.

24. Evaporation
During evaporation, only those molecules with a certain minimum KE can escape from the surface of the liquid.
Liquid evaporates faster when heated because heating increases the average KE
Added energy of heating enables more particles to overcome the attractive forces keeping them in the liquid state.
Particles with the highest KE tend to escape first.

25. Evaporation
Particles left in the liquid have a lower average KE than the particles that escaped
As evaporation takes place, temperature decreases
Added energy of heating enables more particles to overcome the attractive forces keeping them in the liquid state.

26. Vapor Pressure
Evaporation in a closed container differs from evaporation in an open container.
No particles can escape into the outside air from the closed container
When a partially filled container of liquid is sealed, some particles at the surface vaporize.
Particles collide with the walls of the sealed container and produce a vapor pressure.

27. Vapor Pressure
Vapor Pressure – is a measure of the force exerted by a gas above a liquid.
Over time, the number of particles entering the vapor increases and some of the particles condense and return to the liquid state.

28. Vapor Pressure
Eventually, the number of particles condensing will equal the number of particles vaporizing.
Vapor pressure will remain constant
In a system at constant vapor pressure, a dynamic equilibrium exists between the vapor and the liquid.
The system is in equilibrium because the rate of evaporation of liquid equals the rate of condensation of vapor.

29. Vapor Pressure & Temperature
Increase in temperature of a contained liquid increases the vapor pressure.
Particles in the warmed liquid have increased KE.
More particles will have the minimum KE necessary to escape the surface of the liquid.
Vapor pressure of substances indicates how easily it evaporates and also how volatile it is

30. Vapor Pressure & Measurements
Vapor pressure of a liquid can be determined with a manometer.

31. Manometer
One end - U shaped glass tube with mercury is attached to a container
Other end of tube is open to the atmosphere.
When a liquid is added to the container, the pressure in the container increased due to vapor pressure of the liquid.

32. Boiling Point
Rate of evaporation of a liquid from an open container increases as the liquid is heated.
Heating allows a greater number of particles to overcome the attractive forces that keep them in the liquid state.
The remaining particles in the liquid move faster as they absorb the added energy from heating.
Average KE of the particles in the liquid increases and the temperature increases.

33. Boiling Point
When a liquid is heated to a temperature at which particles throughout the liquid have enough kinetic energy to vaporize, the liquid begins to boil
Bubbles of vapor form, rise to the surface, and escape into the air.
Boiling Point – the temperature at which the vapor pressure of the liquid is just equal to the external pressure on the liquid

34. Boiling Point & Pressure Changes
Liquids don’t always boil at the same temperature
atmospheric pressure is lower at higher altitudes, boiling points decrease at higher altitudes.

35. Boiling point is a cooling process similar to evaporation.
During boiling, particles with highest KE escape first.
Temperature of the boiling liquid never rises above its boiling point
Vapor produced is at the same temperature as that of the boiling liquid.

36. Boiling Point
Although the vapor has the same average KE as the liquid, its potential energy is much higher.
Burn from a steam is more severe than one from boiling water at the same temperature
Normal Boiling Point – boiling point of a liquid at a pressure of 101.3kPa.

37. End of Section 13.2

38. Nature of Solids The general properties of solids reflect their
orderly arrangement of their particles
fixed locations of their particles.
Atoms, ions, or molecules are packed tightly together
Dense, not easy to compress
Do not flow

39. Nature of Solids
When you heat a solid, particles vibrate more rapidly as their KE increases.
Organization of particles within breaks down
Eventually it melts
Melting Point (mp) – temperature at which a solid changes into a liquid.
At mp temperature, disruptive vibrations of particles is strong enough to overcome the attractions that hold them in fixed positions.

40. Melting Point
Melting point and freezing points of a substance are at the same temperature.
At that temperature, the liquid and solid phases are in equilibrium.
melting
Liquid
Solid
freezing

41. Crystal Structure and Unit Cells
Most solid substances are crystalline.
In a crystal, particles are arranges in an orderly, repeating, three-dimensional pattern called crystal lattice.
Shape of a crystal reflects the arrangement of the particles within the solid
Sodium chloride
Crystal lattice

42. Ions usually formed from a metal and a nonmetal
Crystal Structure and Unit Cells
Type of bonding that exists between particles in crystals determines their melting points.
In general, ionic solids have high melting points because relatively strong forces hold them together.
Calcium Fluoride
ionic solid
Ions usually formed from a metal and a nonmetal

43. Crystal Structure and Unit Cells
Molecular Solids have relatively low melting points
Molecular Solid
Ice
molecules held together by relatively weak intermolecular forces
Nonmetallic elements

44. Crystal Systems A crystal has sides, or faces.
The angles at which the faces of a crystal intersect are always the same for a given substance and are characteristic of that substance.
Crystals are classified into seven groups or crystal systems.
The crystal systems differ in terms of the angles between the faces and the number of edges of equal length on each face.

46. Crystal Systems
Shape of the crystal depends on the arrangement of the particles within it.
Unit Cell – the smallest group of particles within a crystal that retains the geometric shape of the crystal
A crystal lattice is a repeating array of unit cells. Ex: wallpaper

47. Allotropes
Some solid substances can exist in more than one form. Ex: carbon
One crystalline form of carbon – graphite
Another crystalline form of carbon - diamond
The physical properties of diamond and graphite are quite different.
Diamond – high density and hard
Graphite – low density, soft and slippery

48. Allotropes
Allotropes – two or more different molecular forms of the same element in the same physical state.
Diamond and graphite are allotropes of carbon
Even though allotropes are composed of atoms of the same element, they have different properties because their structures are different.
Only a few elements have allotropes
phosphorus  sulfur  oxygen

49. Non-Crystalline Solids
Not all solids are crystalline in form, some are amorphous.
Amorphous Solid – lacks an ordered internal structure.
Rubber  plastic  asphalt
Atoms of amorphous solids are arranged randomly.

50. Glass is another amorphous solid
Glass is a substance that has cooled to a rigid state without crystallizing.
Sometimes called supercooled liquids.
Internal structure of glass is intermediate between a crystalline solid and a free-flowing liquid.
Glass does not melt at a definite temperature

51. Glass gradually softens when heated
When crystalline solids shatter, fragments tend to have same surface angles as the original solid.
When amorphous solids shatter, fragments have irregular angles and jagged edges.

52. End of Section 13.3

53. Sublimation
Sublimation – the change of a substance from a solid to a vapor without passing through the liquid state.
Sublimation can occur because solids, like liquids, have vapor pressure.
Sublimation occurs in solids with vapor pressures that exceed atmospheric pressure at or near room temperature.

54. Sublimation
At normal pressures, most chemical compounds and elements possess three different states at different temperatures.
Some elements or substances at some pressures may pass directly from solid to the gaseous state.
Can occur if the atmospheric pressure exerted on the substance is too low to stop the molecules from escaping from the solid state.

55. Sublimation Applications
Solid carbon dioxide (dry ice) sublimes at atmospheric pressure.
Used as a coolant. It does not produce a liquid as ordinary ice does when it melts.

56. Sublimation Applications
Snow and other water ices also sublime, although more slowly, at below-freezing temperatures.
Used in freeze drying (ex: coffee), allows wet clothes to be hung outdoors in freezing weather and retrieved later in a dry state

57. Sublimation Applications
Solid air fresheners contain a variety of substances that sublime at room temperature.

58. Sublimation Applications
Chemists use sublimation to separate mixtures and to purify compounds
Solid is placed in a vessel which is heated under vacuum.
Under reduced pressure, solid volatilizes & condenses as a purified compound.
Leaving the non-volatile residue (impurities) behind.

59. Phase Diagram
Relationships among the solid, liquid, and vapor phases of a substance in a sealed container can be represented in a single graph.
Phase diagram – gives the conditions of temperature and pressure at which a substance exists as solid, liquid and gas.

60. Phase Diagram
Triple point – point in the phase diagram where all three lines separating the phases meet.
Describes the only set of conditions at which all three phases can exist in equilibrium with one another.
The conditions of pressure and temperature at which two phases exist in equilibrium are indicated by a line separating the phases.

61. Phase Diagram Questions
At the triple point of water, what are the values of temperature and pressure?
0.016ºC and 0.61 kPa
What states of matter are present at the triple point of water?
Solid, liquid and vapor
Assuming standard pressure, at what temperature is there an equilibrium between water vapor and liquid water?
100ºC
Between liquid water and ice?
0ºC

62. Phase Diagram
A decrease in pressure lowers the boiling point and raises the melting point
An increase in pressure will raise the boiling point and lower the melting point.
Rain is the condensation product of water vapor in the atmosphere.

63. Questions What properties must a solid have to undergo sublimation?
Sublimation occurs in solids that have vapor pressures that exceed atmospheric pressure at or near room temperature.
What do the curved lines on a phase diagram represent?
The lines show the conditions of temperature and pressure at which two phases exist in equilibrium

64. Questions Describe on practical use of sublimation.
Freeze-dried coffee, dry ice as a coolant, air fresheners, separating mixtures and purifying substances.
What does the triple point on a phase diagram describe?
The only set of conditions at which three phases can exist in equilibrium.
Using phase diagram for water, estimate the boiling point of water at a pressure of 50kPa.
60ºC

65. Questions
At the following temperature-pressure values, which phase of water is most stable?
90ºC, 40 kPa
vapor
20ºC, 10 kPa
liquid
70ºC, kPa
120ºC, 40 kPa
50ºC, kPa

66. End of Chapter 13