States of MatterSolids, Liquids, Gases, Plasma, and

Phase Changes

Is there any movement within the structure of a solid?

3

A Model for Solids

• In most solids, the atoms, ions, or molecules are packed tightly together.

• Solids are dense and not easy to compress.• Because the particles in solids tend to vibrate

about fixed points, solids do not flow.

The general properties of solids reflect the orderly arrangement of their particles and the fixed locations of their particles.

Crystal Structure and Unit Cells

The shape of a crystal reflects the arrangement of the particles within the solid.

• In sodium chloride, sodium ions and chloride ions are closely packed in a regular array.

• The ions vibrate about fixed points in the crystal.

5

A Model for SolidsWhen you heat a solid, its particles vibrate more rapidly as their kinetic energy increases.

• The melting point (mp) is the temperature at which a solid changes into a liquid.

6

A Model for SolidsWhen you heat a solid, its particles vibrate more rapidly as their kinetic energy increases.

• The melting point (mp) is the temperature at which a solid changes into a liquid.

– At this temperature, the disruptive vibrations of the particles are strong enough to overcome the attractions that hold them in fixed positions.

7

A Model for Solids• In general, ionic solids have high melting points

because relatively strong forces hold them together.

• Freezing Point: The temperature at which a liquid turns into a solid. (Reverse of melting point)

– Sodium chloride, an ionic compound, has a rather high melting point of 801°C.

8

A Model for Solids• In general, ionic solids have high melting points

because relatively strong forces hold them together.

– Sodium chloride, an ionic compound, has a rather high melting point of 801°C.

• Molecular solids have relatively low melting points.– Hydrogen chloride, a molecular compound,

melts at –112°C.

9

A Model for Solids• Different forces hold substances together.

– Ionic compounds are held together by the attraction of anions to cations.

– Covalent compounds are held together by several different intermolecular forces, the forces that hold covalent compounds together

• Hydrogen Bonds (positive and negative ends attract)• Dipole – Dipole forces (positive and negative ends

attract)• Dispersion forces (temporary positive and negative

ends form that attract)

10

Crystal Structure and Unit Cells

Some substances can exist in more than one form.• Diamond is one crystalline form of carbon.• A different form of carbon is graphite.• In 1985, a third crystalline form of carbon was

discovered. This form is called buckminsterfullerene.

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Crystal Structure and Unit Cells

In diamond, each carbon atom in the interior of the diamond is strongly bonded to four others. The array is rigid and compact.

In graphite, the carbon atoms are linked in widely spaced layers of hexagonal arrays.

In buckminster-fullerene, 60 carbon atoms form a hollow sphere. The carbons are arranged in penta-gons and hexagons.

12

Crystal Structure and Unit Cells

The physical properties of diamond, graphite, and fullerenes are quite different.• Diamond has a high density and is very hard.• Graphite has a relatively low density and is soft

and slippery.• The hollow cages in fullerenes give them

strength and rigidity.

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Crystal Structure and Unit Cells• Allotropes

Diamond, graphite, and fullerenes are crystalline allotropes of carbon.• Allotropes are two or more different molecular

forms of the same element in the same physical state.

14

Crystal Structure and Unit Cells• Allotropes

Only a few elements have allotropes.

• In addition to carbon, these include phosphorus, sulfur, oxygen (O2 and O3), boron, and antimony.

15

Crystal Structure and Unit Cells• Non-Crystalline Solids

Not all solids are crystalline in form; some solids are amorphous.• An amorphous solid lacks an ordered internal

structure.• Rubber, plastic, and asphalt are amorphous

solids.• Their atoms are randomly arranged.

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Crystal Structure and Unit Cells• Non-Crystalline SolidsOther examples of amorphous solids are glasses.

• A glass is a transparent fusion product of inorganic substances that have cooled to a rigid state without crystallizing.

• Glasses are sometimes called supercooled liquids.

• The irregular internal structures are intermediate between those of a crystalline solid and those of a free-flowing liquid.

17

•What is the difference between an amorphous solid and a crystalline solid?

18

•What is the difference between an amorphous solid and a crystalline solid?

Particles in a crystalline solid are arranged in an orderly, repeating pattern or lattice. Particles in an amorphous solid are arranged randomly.

19

Kinetic Theory

• The word kinetic refers to motion.

• The energy an object has because of its motion is called kinetic energy.

• According to the kinetic theory, all matter consists of tiny particles that are in constant motion.

• The particles in a gas are usually molecules or atoms.

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Kinetic Energy and Temperature•As a substance is heated, its particles absorb energy, some of which is stored within the particles.• This stored portion of the energy, or potential

energy, does not raise the temperature of the substance.

• The remaining absorbed energy does speed up the particles, that is, increases their kinetic energy.

– This increase in kinetic energy results in an increase in temperature.

21

Kinetic Energy and Temperature• Average Kinetic EnergyThe particles in any collection of atoms or molecules at a given temperature have a wide range of kinetic energies.• Most have kinetic energies somewhere in the

middle of this range.• We use average kinetic energy when discussing

the kinetic energy of a collection of particles in a substance.

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Kinetic Energy and Temperature• Average Kinetic EnergyAt any given temperature, the particles of all substances, regardless of physical state, have the same average kinetic energy.

23

•The figure below shows the distribution of kinetic energies of water molecules at two different temperatures.

Interpret Graphs

• The green curve shows the distribution of kinetic energy in cold water.

• The purple curve shows the distribution of kinetic energy in hot water.

24

Kinetic Energy and Temperature• Average Kinetic EnergyThe average kinetic energy of the particles in a substance is directly related to the substance’s 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 decreases.

25

Kinetic Energy and Temperature• Average Kinetic EnergyAbsolute zero (0 K, or –273.15oC) is the temperature at which the motion of particles theoretically ceases.• No temperature can be lower than absolute zero.• Absolute zero has never been produced in the

laboratory.

– A near-zero temperature of about 0.000 000 000 1 K, which is 0.1 nanokelvin, has been achieved.

26

Kinetic Energy and Temperature• Average Kinetic EnergyThe coldest temperatures recorded outside the laboratory are from space.• Astronomers used a

radio telescope to measure the temperature of the boomerang nebula.

• At about 1 K, it is the coldest known region of space.

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Kinetic Energy and Temperature

•Average Kinetic Energy and Kelvin Temperature

The Kelvin temperature of a substance is directly proportional to the average kinetic energy of the particles of the substance.

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What is the result of increasing the temperature of a gas sample? A. A decrease in the average kinetic energy of

the sample

B. No effect on the sample

C. An increase in the average kinetic energy of the sample

D. The particles slow down.

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What is the result of increasing the temperature of a gas sample? A. A decrease in the average kinetic energy of

the sample

B. No effect on the sample

C. An increase in the average kinetic energy of the sample

D. The particles slow down.

30

Kinetic Theory and a Model for Gases

•The kinetic theory as it applies to gases includes the following fundamental assumptions about gases.

The particles in a gas are considered to be small, hard spheres with an insignificant volume.

– Within a gas, the particles are relatively far apart compared with the distance between particles in a liquid or solid.

– Between the particles, there is empty space.

– No attractive or repulsive forces exist between the particles.

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Kinetic Theory and a Model for Gases

•The kinetic theory as it applies to gases includes the following fundamental assumptions about gases.

Bromine molecule

The motion of particles in a gas is rapid, constant, and random.

– Gases fill their containers regardless of the shape and volume of the containers.

– An uncontained gas can spread out into space without limit.

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Kinetic Theory and a Model for Gases

•The kinetic theory as it applies to gases includes the following fundamental assumptions about gases.

All collisions between particles in a gas are perfectly elastic.

– During an elastic collision, kinetic energy is transferred without loss from one particle to another.

– The total kinetic energy remains constant.

33

Gas Pressure

Gas pressure results from the force exerted by a gas per unit surface area of an object.

• Moving bodies exert a force when they collide with other bodies.

34

Gas Pressure

• If no particles are present, no collisions can occur. Consequently, there is no pressure.

• An empty space with no particles and no pressure is called a vacuum.

Gas pressure is the result of billions of rapidly moving particles in a gas simultaneously colliding with an object.

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Gas PressureAir exerts pressure on Earth because gravity holds the particles in air within Earth’s atmosphere.• The collisions of atoms and molecules in air

with objects results in atmospheric pressure.*Observe the effects of atmospheric pressure in the following demo

• Atmospheric pressure decreases as you climb a mountain because the density of Earth’s atmosphere decreases as the elevation increases.

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Gas PressureA barometer is a device that is used to measure atmospheric pressure.• At sea level, air exerts

enough pressure to support a 760-mm column of mercury.

• On top of Mount Everest, at 9000 m, the air exerts only enough pressure to support a 253-mm column of mercury.

Vacuum

Atmospheric pressure

760 mm Hg (barometric pressure)

253 mm Hg

Sea level On top of Mount Everest

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•When weather forecasters state that a low-pressure system is moving into your region, it usually means that a storm is coming. What do you think happens to the column of mercury in a barometer as a storm approaches? Why?

CHEMISTRY & YOU

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•When weather forecasters state that a low-pressure system is moving into your region, it usually means that a storm is coming. What do you think happens to the column of mercury in a barometer as a storm approaches? Why?

CHEMISTRY & YOU

When a storm approaches, the column of mercury goes down, indicating a decrease in atmospheric pressure.

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Gas PressureThe SI unit of pressure is the pascal (Pa).• Normal atmospheric pressure is about 100,000 Pa,

that is, 100 kilopascals (kPa).• Two older units of pressure are commonly used.

– millimeters of mercury (mm Hg)

– atmospheres (atm)

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Gas PressureOne standard atmosphere (atm) is the pressure required to support 760 mm of mercury in a mercury barometer at 25°C.• The numerical relationship among the three units

is

1 atm = 760 mm Hg = 101.3 kPa.• Standard temperature and pressure (STP) are

defined as a temperature of 0°C and a pressure of 101.3 kPa, or 1 atm.

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A Model for Liquids•Substances that can flow are referred to as fluids.• Both liquids and gases can flow.

– The ability of gases and liquids to flow allows them to conform to the shape of their containers.

42

A Model for Liquids

•Gases and liquids have a key difference between them.• According to kinetic theory, there are no

attractions between the particles in a gas.• The particles in a liquid are attracted to each

other.

– These intermolecular attractions keep the particles in a liquid close together, which is why liquids have a definite volume.

43

A Model for Liquids• Different forces hold substances together.

– Ionic compounds are held together by the attraction of anions to cations.

– Covalent compounds are held together by several different intermolecular forces, the forces that hold covalent compounds together

• Hydrogen Bonds (positive and negative ends attract)• Dipole – Dipole forces (positive and negative ends

attract)• London Dispersion forces (temporary positive and

negative ends form that attract)

44

A Model for Liquids

Class Activity: Try fitting as many drops of water as possible on a penny without any dripping off the penny, and then do the same with drops of acetone.

Which could you drop more drops of? Why?

45

A Model for Liquids

Class Activity: Try fitting as many drops of water as possible on a penny without any dripping off the penny, and then do the same with drops of acetone.

Which could you drop more drops of? Why?Both have surface tension that allows for the addition

of more drops, but water has stronger intermolecular forces holding it together so you can add more water.

46

A Model for Liquids

• Surface tension: A property of liquids where liquids can resist force on the surface.

• Results from the net force of the particles in a liquid pulling down on the surface molecules.

47

A Model for Liquids

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

48

A Model for Liquids

•Liquids are much more dense than gases.• Increasing the pressure on a liquid has hardly

any effect on its volume.• The same is true for solids.

– Liquids and solids are known as condensed states of matter.

49

•Describe one way in which liquids and gases are similar, and one way in which they are different.

50

Evaporation

•The conversion of a liquid to a gas or vapor is called vaporization.• When this conversion occurs at the surface of

a liquid that is not boiling, the process is called evaporation.

51

Evaporation•The process of evaporation has a different outcome in an open system, such as a lake or an open container, than in a closed system, such as a sealed container.

• In an open system, molecules that evaporate can escape from the system.

• In a closed system, the molecules collect as a vapor above the liquid. Some condense back into a liquid.

52

Evaporation

During evaporation, only those molecules with the greatest kinetic energy can escape from the surface of the liquid. • This is why evaporation has a cooling

effect

53

Evaporation

• Heating the liquid increases the average kinetic energy of its particles.

• The added energy enables more particles to overcome the attractive forces keeping them in the liquid state. The particles with the greatest kinetic energy evaporate.

A liquid evaporates faster when heated.

54

•Explain why heating a liquid causes evaporation to occur faster.

Why does sweating cool you down?

56

Vapor PressureThe evaporation of a liquid in a closed system differs from evaporation in an open system.

• When a partially filled container of liquid is sealed, some of the particles at the surface of the liquid vaporize.

• These particles collide with the walls of the sealed container, producing pressure.

– A measure of the force exerted by a gas above a liquid is called vapor pressure.

57

Vapor Pressure

Over time, the number of particles entering the vapor increases and some of the particles condense and return to the liquid state.

Liquid Vapor (gas)evaporation

condensation

Condensation: When a gas returns to the liquid state. (reverse of evaporation)

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Vapor Pressure

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.

59

Vapor Pressure

• Vapor Pressure and Temperature ChangeAn increase in the temperature of a contained liquid increases the vapor pressure.

• This happens because the particles in the warmed liquid have increased kinetic energy.

– More of the particles will reach the minimum kinetic energy necessary to escape the surface of the liquid.

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Vapor Pressure

• Observe the following demonstration and explain your observations using what you know about vapor pressure.

61

Vapor Pressure

• Observe the following demonstration and explain your observations using what you know about vapor pressure.– Different substances have different vapor

pressures. The weaker the intermolecular forces in the substance, the more particles in the gas phase and the higher the vapor pressure.

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• The vapor pressure data indicates how volatile a given liquid is, or how easily it evaporates.

Vapor Pressure (in kPa) of Three Substances at Different Temperatures

Substance 0°C 20°C 40°C 60°C 80°C 100°C

Water 0.61 2.33 7.37 19.92 47.34 101.33Ethanol 1.63 5.85 18.04 47.02 108.34 225.75Diethyl ether 24.70 58.96 122.80 230.65 399.11 647.87

Interpret Data

• Of the three liquids shown, diethyl ether is the most volatile and water is the least volatile.

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Vapor Pressure• Vapor Pressure MeasurementsThe vapor pressure of a liquid can be determined with a device called a manometer.

12.2 mm Hg or 1.63 kPa

43.9 mm Hg or 5.85 kPa

Air at standard temperature and pressure

Ethanol at 0°C Ethanol at room temperature (20°C)

AirMercury

EthanolMercury

EthanolMercury

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Vapor Pressure• Vapor Pressure MeasurementsThe vapor pressure is equal to the difference in height of the mercury in the two arms of the U-tube.

12.2 mm Hg or 1.63 kPa

43.9 mm Hg or 5.85 kPa

Air at standard temperature and pressure

Ethanol at 0°C Ethanol at room temperature (20°C)

AirMercury

EthanolMercury

EthanolMercury

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•In a sealed gas-liquid system at a constant temperature, eventually

A. there will be no more evaporation.

B. the rate of condensation decreases to zero.

C. the rate of condensation exceeds the rate of evaporation.

D. the rate of evaporation equals the rate of condensation.

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•In a sealed gas-liquid system at a constant temperature, eventually

A. there will be no more evaporation.

B. the rate of condensation decreases to zero.

C. the rate of condensation exceeds the rate of evaporation.

D. the rate of evaporation equals the rate of condensation.

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Boiling Point

• Boiling Point

What happens to the temperature of a substance as it boils? To answer this each group will record the temperature of water as they heat it every 30 seconds. When you finish graph the results.

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Boiling Point

• The vapor produced is at the same temperature as that of the boiling liquid.

Boiling is a cooling process, similar to evaporation.

– Although the vapor has the same average kinetic energy as the liquid, its potential (or stored) energy is much higher.

– Thus, a burn from steam is more severe than one from an equal mass of boiling water, even though they are both at the same temperature.

69

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 throughout the liquid, rise

to the surface, and escape into the air.• The boiling point (bp) is the temperature at which

the vapor pressure of the liquid is just equal to the external pressure on the liquid.

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Boiling Point• Boiling Point and Pressure Changes

Because a liquid boils when its vapor pressure is equal to the external pressure, liquids don’t always boil at the same temperature.

• Because atmospheric pressure is lower at higher altitudes, boiling points decrease at higher altitudes.

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Boiling Point• Boiling Point and Pressure Changes

Atmospheric pressure at the surface of water at 70°C is greater than its vapor pressure. Bubbles of vapor cannot form in the water, and it does not boil.

At the boiling point, the vapor pressure is equal to the atmospheric pressure. Bubbles of vapor form in the water, and it boils.

At higher altitudes, the atmospheric pressure is lower than it is at sea level. Thus, the water boils at a lower temperature.

101.3 kPa 101.3 kPa 34 kPa 70°C 70°C

100°C

Sea Level Sea Level Atop Mount Everest

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•You can use this graph to show how the boiling point of a liquid is related to vapor pressure.

Interpret Graphs

• At a lower external pressure, the boiling point decreases.

• At a higher external pressure, the boiling point increases.

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• The normal boiling point is defined as the boiling point of a liquid at a pressure of 101.3 kPa.

Normal Boiling Points of Several Substances

Substance Boiling Point (°C)Carbon disulfide (CS2) 46.0Chloroform (CHCl3) 61.7Methanol (CH4O) 64.7Carbon tetrachloride (CCl4) 76.8Ethanol (C2H6O) 78.5Water (H2O) 100.0

Interpret Data

Normal Boiling Point

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Is the boiling point of water at the top of Mount McKinley (the highest point in North America) higher or lower than it is in Death Valley (the lowest point in North America)?

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Is the boiling point of water at the top of Mount McKinley (the highest point in North America) higher or lower than it is in Death Valley (the lowest point in North America)?

The boiling point of water decreases as altitude increases. Therefore, the boiling point of water is lower atop Mount McKinley than it is in Death Valley.

What happens if you raise the temperature or matter to super-

high levels…between

1000°C and 1,000,000,000°C ?

Will everything just be a gas?

STATES OF MATTERPLASMA

A plasma is an ionized gas.

A plasma is a very good conductor of electricity and is affected by magnetic fields.

Plasmas, like gases have an indefinite shape and an indefinite volume.

• Plasma is a common state of matter

STATES OF MATTER

SOLID LIQUID GAS PLASMA

Tightly packed, in a regular pattern

Vibrate, but do not move from place to

place

Close together with no regular

arrangement.Vibrate, move

about, and slide past each other

Well separated with no regular

arrangement.Vibrate and move

freely at high speeds

Has no definite volume or shape

and is composed of electrical charged

particles

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Sublimation

•The change of a substance from a solid to a vapor without passing through the liquid state is called sublimation.

• Sublimation can occur because solids, like liquids, have a vapor pressure.

81

Sublimation

•Sublimation occurs in solids with vapor pressures that exceed atmospheric pressure at or near room temperature.

Solid Vaporsublimation

deposition

Deposition: A gas directly becomes a solid. (Reverse of sublimation)

82

Sublimation•Iodine is an example of a substance that undergoes sublimation.• The outer test tube contains solid

iodine that is being gently heated.• The inner test tube contains liquid

water and ice.

• The iodine crystals at the bottom of the outer test tube change directly to iodine vapor.

• When the vapor reaches the cool surface of the inner test tube, it goes directly from the gaseous to the solid state.

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Sublimation•Sublimation has many useful applications.

• Solid carbon dioxide (dry ice) is often used as a coolant for goods such as ice cream, which must remain frozen during shipment.– Because it sublimes, it does

not produce a liquid as ordinary ice does when it melts.

– As it changes state, dry ice absorbs heat, keeping materials nearby cool and dry.

84

•Which of the following describes a solid undergoing sublimation? A. The vapor pressure exceeds the atmospheric

pressure.

B. The vapor pressure equals the atmospheric pressure.

C. The vapor pressure is less than the atmospheric pressure.

D. The vapor pressure is less than half the atmospheric pressure.

85

•Which of the following describes a solid undergoing sublimation? A. The vapor pressure exceeds the atmospheric

pressure.

B. The vapor pressure equals the atmospheric pressure.

C. The vapor pressure is less than the atmospheric pressure.

D. The vapor pressure is less than half the atmospheric pressure.

86

Phase Diagrams•The relationships among the solid, liquid, and vapor states (or phases) of a substance in a sealed container can be represented in a single graph.• The graph is called a phase diagram.

87

Phase Diagrams•The relationships among the solid, liquid, and vapor states (or phases) of a substance in a sealed container can be represented in a single graph.• The graph is called a phase diagram.

– A phase diagram gives the conditions of temperature and pressure at which a substance exists as a solid, liquid, or gas (vapor).

88

•The phase diagram of water shows the relationship among pressure, temperature, and the physical states of water.

Interpret Graphs

• In each of the curving regions of the phase diagram, water is in a single phase.

89

•The phase diagram of water shows the relationship among pressure, temperature, and the physical states of water.

Interpret Graphs

• The curving line that separates water’s vapor phase from its liquid phase describes the equilibrium conditions for liquid and vapor.

90

•The phase diagram of water shows the relationship among pressure, temperature, and the physical states of water.

Interpret Graphs

• The other two lines describe the conditions for equilibrium between liquid water and ice and between water vapor and ice.

91

•The phase diagram of water shows the relationship among pressure, temperature, and the physical states of water.

Interpret Graphs

• The point on the diagram at which all three lines meet is called the triple point.

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Phase Diagrams•The triple point describes the only set of conditions at which all three phases can exist in equilibrium with one another.

• For water, the triple point is a temperature of 0.016°C and a pressure of 0.61 kPa.

• This flask is at the triple point. Freezing, melting, boiling, and condensation are all occurring at the same time in the flask.

93

Phase DiagramsBy referring to the phase diagram of water, you can determine what happens if you melt ice or boil water at pressures less than 101.3 kPa.• 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.

94

Phase DiagramsBelow the triple point, the vapor and liquid cannot exist in equilibrium.• Increasing the pressure

won’t change the vapor to a liquid.

• The solid and vapor are in equilibrium at temperatures below 0.016°C.

• With an increase in pressure, the vapor begins to behave more like a solid.

95

Phase DiagramsFor years, the accepted hypothesis for how ice-skaters move along the ice was the following.

• The blades of the skates exert pressure, which lowers the melting point of the ice.

• The ice melts, and a film of water forms under the blades of the skates.

• This film acts as a lubricant, enabling the skaters to glide gracefully over the ice.

96

Phase DiagramsThis hypothesis fails to explain why skiers glide nicely on another solid form of water—snow.

• Wide skis exert much less pressure per unit area of snow than narrow skate blades exert on ice.

• Recent research shows that the surface of ice has a slippery, water-like surface layer that exists well below ice’s melting point.

• The liquid-like surface layer provides lubrication needed for smooth skating and skiing.

97

•Describe the meaning of a line in a phase diagram.

98

•Describe the meaning of a line in a phase diagram.

Along a line in a phase diagram, two phases are in equilibrium with each other.