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States of Matter Chapter 13.

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Presentation on theme: "States of Matter Chapter 13."— Presentation transcript:

1 States of Matter Chapter 13

2 What You Need to Master How to use the kinetic-molecular theory to explain the physical properties of gasses, liquids, and solids Compare types of intermolecular forces Explain how kinetic energy and intermolecular forces combine to determine the state of a substance Describe the role of energy in phase changes

3 Gases Chapter 13.1

4 Kinetic-Molecular Theory
Jan Baptista Van Helmont (Flemish) used the word Chaos (without order) to describe products of reactions that had no fixed shape or volume. By 18th century, scientists knew how to collect gasses by displacing water They could now measure properties of gasses 1860 Ludwig Boltzmann and James Maxwell proposed models to explain the properties of gasses Kinetic-molecular theory

5 Kinetic-Molecular Theory
Kinetic – Greek word meaning “to move” any object in motion has “kinetic energy” Kinetic-Molecular Theory describes the behavior of gasses in terms of particles in motion. Model makes several assumptions about size, motion, and energy of particles

6 Kinetic-Molecular Theory
Assumptions: Particle Size : particles are so small and so far apart that there no significant forces of attraction Particle motion: particles are in constant, random motion. Particles move in a straight line until they collide with another particle or the container wall Particle collisions are elastic (no energy is lost, but may be transferred)

7 Kinetic-Molecular Theory
Assumptions: Particle Energy: 2 factors, mass & velocity KE = ½ mv2 All particles do NOT have the same energy Temperature is a measure of the average kinetic energy of the particles

8 How K-M Theory Explains Gas Behavior
Low Density Density is mass/volume Cl2 density at 20°C is 2.95 x 10-3 g/mL Au density is 19.3 g/mL Because Chlorine is a gas, K-M theory says there must be a lot of empty space between molecules. Thus, fewer molecules in the same volume

9 How K-M Theory Explains Gas Behavior
Compression and Expansion Because there is a lot of empty space between molecules/atoms, you can compress or squeeze together the atoms/molecules Removes some of the empty space Allowing the atoms to return to normal is expansion

10 How K-M Theory Explains Gas Behavior
Diffusion and Effusion Diffusion – describes the movement of one material through another Effusion – gas escaping through a tiny opening Like puncturing a balloon or tire Graham’s Law of Diffusion Rate of effusion ͒1/Ã molar mass ͒means ‘is proportional to’ Larger particles diffuse slower than smaller particles

11 How K-M Theory Explains Gas Behavior
Rate Relationships: Ratea Rateb à Molar massa Molar massb =

12 Gas Pressure Pressure = force/area
An elephant has less pressure on its foot than a human. Why? Gasses exert pressure when they collide with walls of the container. Small mass means small pressure But there are about 1022 particles in one liter Pressure can be substantial Earth’s air pressure varies – gets lower with altitude

13 Measuring Air Pressure
Evangelista Torricelli (Italian – before 1650) first demonstrated air pressure Designed equipment to measure air pressure using mercury in a tube. Now called a barometer

14 Measuring Air Pressure
A ‘manometer’ measures air pressure in a u-shaped tube, in a closed container.

15 Units of Pressure SI unit is ‘Pascal’
1 Pascal = 1 newton/m2 PSI (pounds per square inch) is still used Barometers and manometers still report millimeters of mercury (mm Hg) At sea level, air pressure is 760 mm Hg at 0°C Air pressure is frequently reported as 1 atmosphere (atm) = 760 mm Hg or 760 torr or kilopascals

16 Units of Pressure The units 1 atm, 760 mm Hg, 760 torr, and kilopascals are considered to be defined units They have as many significant digits as necessary

17 Dalton’s Law of Partial Pressure
Dalton studied mixtures of gasses: Found that each gas exerts pressure independently. Law of Partial Pressures: The total pressure is equal to the sum of pressures by each gas in the mixture Partial pressure depends on the moles of gas, size of container and temperature. Does NOT depend on the identity of the gas

18 Dalton’s Law of Partial Pressure
At a given temperature and pressure, PTOTAL = PA + PB + PC + PD + …. PN Dalton’s Law can be used for: Determine amount of gas produced by a reaction Partial pressures of gasses at same temperature are related to concentrations. Look up values in reference tables At 20° C, partial pressure of water is 2.3 kPa

19 Intermolecular Forces
Chapter 13-2

20 “Inter” vs “Intra” “Intra” means within.
“Inter” means across or between. Sports It’s really intramural sports Within the same school It’s interscholastic sports Across or between schools “Intramolecular forces” are ionic, covalent, metallic bonds Intramolecular forces are from molecule to molecule

21 Intramolecular Forces
Intermolecular forces are very weak compared to intramolecular forces Dispersion (or London) forces Dipole-Dipole forces Hydrogen bonding

22 Dispersion (or London) Forces
Dispersion is an intermolecular attraction force that exists between all molecules. These forces are the result of the movement of electrons which cause slight polar moments. Dispersion forces are generally very weak but as the molecular weight increases so does their strength. Most important between nonpolar molecules of the same size The cause of Bromine being liquid and Iodine being solid

23 Dispersion (or London) Forces
Discoverer was Fritz London

24 Dipole-Dipole Forces Dipole-dipole forces exist between neutral, polar molecules The positive end of one molecule is attracted to the negative end of another molecule. The greater the polarity (difference in electronegativity of the atoms in the molecule), the stronger the dipole-dipole attraction. Dipole-dipole attractions are very weak Substances held together by these forces have low melting and boiling point temperatures.  Generally, substances held together by dipole-dipole attractions are gases at room temperature.

25 Dipole-Dipole Forces

26 Hydrogen Bonding The positively charged hydrogen end of one molecule is attracted to the negatively charged end of another molecule Negative end must be an extremely electronegative element (fluorine, oxygen, or nitrogen - FON) Hydrogen bonds are the strongest weak bond the H atom essentially gives its single electron to form a bond and is therefore left unshielded.   The relatively strength of hydrogen bonds results in higher melting and boiling point temperatures Hydrogen bonding can explain why water is less dense in the solid phase than it is in the liquid phase (contrary to most other substances).   The hydrogen bonds between water molecules in ice to form a crystal structure, keeping them further apart than they are in the liquid phase.

27 Hydrogen Bonding

28 Liquids and Solids Chapter 13.3

29 Liquids From Chapter 3 K-M Theory: Liquids cannot be compressed
Liquids take the shape of their container, but cannot expand to fill the container K-M Theory: Molecules still in motion, but closer together No fixed positions Forces of attraction between particles limit their range of motion

30 Density and Compression
Liquids are much denser than gasses At 1 atm and 25C, water is 1250 times more denser than water vapor They are the same temperature so they have the same average kinetic energy So density has to do with the intermolecular forces Liquids CAN be compressed, but: Change in volume is much smaller Not much room for atoms to be pushed together

31 Fluidity Fluidity is the ability to flow
Both gasses and liquids are fluids Liquids can flow through liquids so they can diffuse, but liquids diffuse more slowly than gasses at the same temperature Liquids are less fluid than gasses

32 Viscosity Viscosity is a measure of the resistance of a fluid to flow.
Highly viscous flows very slowly. Molecules are close enough together for intermolecular forces to slow them down as they move past each other. More attraction, more viscous Size and shape of molecules affect viscosity Molecules with long chains will have higher viscosity than shorter chains Motor oil is very long chains of carbon and hydrogen

33 Viscosity Viscosity decreases with Temperature
Increases with lower temperature Example: Different motor oils Added energy gives molecules more kinetic energy to overcome intermolecular forces

34 Surface Tension Intermolecular forces do NOT have an equal effect on all particles Particles in the middle can be attracted to particles all around them Particles at surface have no attraction at top to balance forces from below Surface has smallest possible area for and acts like a stretched drum head Takes energy to break attractions from interior Result: SURFACE TENSION Surface Tension is Energy required to overcome attractions from interior of liquid

35 Surface Tension Stronger the intermolecular forces, higher the surface tension Water has a high surface tension: Hydrogen bonding provided high surface tension Drops are shaped like spheres because surface area of sphere is lower Some spiders can walk on water because of surface tension Falling into water at high speed feels like hitting concrete Compounds to lower surface tension are called “surfactants” Used in cleaning supplies and soaps

36 Capillary Action Water in a narrow container (like graduated cylinder) has a surface that is not straight Concave ‘meniscus’ Two forces at work Cohesion – force of attraction between two identical molecules Adhesion – force of attraction between different molecules Adhesion of water to glass is higher than water to water So water looks like it is ‘climbing’ the walls of the cylinder

37 Capillary Action If cylinder is very small, a thin film of water will be drawn upward Narrow tubes are called capillary tubes Upward movement is called capillary action Capillary action is what makes paper towels absorb water upward Water forms hydrogen bonds with cellulose in paper Same action helps keep disposable diapers dry.


39 Solids According to K-M theory, at the same temperature, 1 mole of solid has same energy as 1 mole of liquid. So why do solids have definite shape and volume? Must be strong attraction forces acting on the particles: Forces will limit the motion of the particles to vibrations around fixed locations. Solids are not classified as fluids

40 Density of Solids Particles in solids are more densely packed than liquid MOST Solids are more dense than liquids When solid and liquid coexist, solid almost always sinks in the liquid. Water is an exception because water forms a rigid 3-dimensional structure keeping molecules further apart.

41 Crystalline Solids Most solids form a solid packed in a crystalline solid form Predictable pattern form Individual pieces are called crystals Smallest arrangement of a piece is a unit cell Shape of crystal is determined by shape of the unit cell There are 7 categories of crystals Crystal shapes differ based on faces, angles, and length of edges

42 Crystalline Shape Categories

43 Solids Molecular solids – held together by intermolecular forces only
Usually not solids at room temperature or have low melting points like sugar Covalent network solids – like carbon and silicon form multiple covalent bonds Very hard, very high melting points (diamond, quartz, glass) Ionic solids – like salt with alternating atoms of different ionic charge Very hard, brittle, high melting points, poor conductivity

44 Solids Metallic Solids – all metallic elements have atoms surrounded by mobile electrons Strength varies, high melting points, good conductors Amorphous Solids – particles are not arranged in regular, repeating pattern “amorphous”, Greek, means without shape Forms when liquids cool too quickly so crystals cannot form Examples: glass, rubber, plastic

45 Phase Changes Chapter 13.4

46 Phase Changes Most substances can exist in 3 states (Solid, Liquid, Gas) Depending on temperature and pressure Each state is referred to as a “Phase” when they coexist as physically different parts of a mixture Ice water is a heterogeneous mixture of 2 phases When energy is added or removed, one phase can change into another

47 Phase Changes Requiring Energy
Melting Vaporization Sublimation

48 Melting Amount of energy needed to melt a substance depends on forces keeping particles together. Water takes a relatively high amount of energy because of hydrogen bonding Adding energy allows molecules to move faster, breaking the hydrogen bonds Temperature at which the forces holding crystal lattice together are broken and it becomes liquid is melting point

49 Vaporization In liquid water, some particles will have more kinetic energy than others. At a certain point, the particles will have enough energy to overcome the forces of attraction and will escape the liquid as a gas. A substance that is liquid at room temperature and becomes gas is called a vapor Vaporization is the process of changing a liquid to a gas

50 Vaporization (Cont) Vaporization that occurs at the surface of a liquid is called evaporation Evaporation is gradual Particles at surface are attracted to fewer other molecules than those in center Even at cold temperatures, some molecules have enough energy to break the attractions and become gas. Evaporation is how your body cools itself Water in sweat absorbs heat from your body Water evaporates leaving less heat in your body and a lower ‘average kinetic energy’ (lower temperature)

51 Vaporization (Cont.) The pressure of a vapor over a liquid is called vapor pressure Temperature at which the vapor pressure of a liquid equals the external atmospheric pressure is boiling point What happens to the boiling point if the pressure is reduced? Increased? Where/when do we experience this?

52 Sublimation Sublimation is when a substance goes from solid directly to gas without becoming a liquid Solid iodine Frozen carbon dioxide (dry ice) Moth balls (p-dichlorobenzene) Solid air fresheners Ice cubes left in a freezer for a long time

53 Sublimation At low pressure, ice (water) sublimes faster
This is freeze drying used for food preservation for campers, space, etc. Freeze drying has no bacteria

54 Phase Changes the Release Energy
Condensation Deposition Freezing

55 Condensation Condensation – process of changing from gas to liquid
When molecule loses is energy, velocity is reduced Intermolecular forces take over again When hydrogen bonds form again, energy is released (heat comes out) Different causes for condensation, all involve loss of energy: Contact with cold item

56 Deposition Deposition – changing from gas directly to solid Snowflakes

57 Freezing Freezing – process of changing from liquid to solid
Remove energy from liquid Molecules don’t move past each other any longer Molecules stay in fixed, set position Become solid

58 Phase Diagrams Phase diagram shows phase of matter at different temperatures and pressures Each substance unique X-axis usually temperature Y-axis usually pressure There is usually a “Triple Point” where all three phases can coexist “Critical Point” – temperature and pressure at which above substance cannot exist as liquid

59 Phase Diagram - Water

60 Phase Diagram CO2

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