2 What You Need to MasterHow to use the kinetic-molecular theory to explain the physical properties of gasses, liquids, and solidsCompare types of intermolecular forcesExplain how kinetic energy and intermolecular forces combine to determine the state of a substanceDescribe the role of energy in phase changes
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 waterThey could now measure properties of gasses1860 Ludwig Boltzmann and James Maxwell proposed models to explain the properties of gassesKinetic-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 attractionParticle motion:particles are in constant, random motion.Particles move in a straight line until they collide with another particle or the container wallParticle collisions are elastic (no energy is lost, but may be transferred)
7 Kinetic-Molecular Theory Assumptions:Particle Energy: 2 factors, mass & velocityKE = ½ mv2All particles do NOT have the same energyTemperature is a measure of the average kinetic energy of the particles
8 How K-M Theory Explains Gas Behavior Low DensityDensity is mass/volumeCl2 density at 20°C is 2.95 x 10-3 g/mLAu density is 19.3 g/mLBecause 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 ExpansionBecause there is a lot of empty space between molecules/atoms, you can compress or squeeze together the atoms/moleculesRemoves some of the empty spaceAllowing the atoms to return to normal is expansion
10 How K-M Theory Explains Gas Behavior Diffusion and EffusionDiffusion – describes the movement of one material through anotherEffusion – gas escaping through a tiny openingLike puncturing a balloon or tireGraham’s Law of DiffusionRate of effusion ͒1/Ã molar mass͒means ‘is proportional to’Larger particles diffuse slower than smaller particles
11 How K-M Theory Explains Gas Behavior Rate Relationships:RateaRatebÃMolar massaMolar 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 pressureBut there are about 1022 particles in one literPressure can be substantialEarth’s air pressure varies – gets lower with altitude
13 Measuring Air Pressure Evangelista Torricelli (Italian – before 1650) first demonstrated air pressureDesigned 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/m2PSI (pounds per square inch) is still usedBarometers and manometers still report millimeters of mercury (mm Hg)At sea level, air pressure is 760 mm Hg at 0°CAir pressure is frequently reported as 1 atmosphere (atm) = 760 mm Hg or 760 torr or kilopascals
16 Units of PressureThe units 1 atm, 760 mm Hg, 760 torr, and kilopascals are considered to be defined unitsThey 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 mixturePartial 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 + …. PNDalton’s Law can be used for:Determine amount of gas produced by a reactionPartial pressures of gasses at same temperature are related to concentrations.Look up values in reference tablesAt 20° C, partial pressure of water is 2.3 kPa
20 “Inter” vs “Intra” “Intra” means within. “Inter” means across or between.SportsIt’s really intramural sportsWithin the same schoolIt’s interscholastic sportsAcross or between schools“Intramolecular forces” are ionic, covalent, metallic bondsIntramolecular forces are from molecule to molecule
21 Intramolecular Forces Intermolecular forces are very weak compared to intramolecular forcesDispersion (or London) forcesDipole-Dipole forcesHydrogen 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 sizeThe cause of Bromine being liquid and Iodine being solid
23 Dispersion (or London) Forces Discoverer was Fritz London
24 Dipole-Dipole ForcesDipole-dipole forces exist between neutral, polar moleculesThe 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 weakSubstances 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.
26 Hydrogen BondingThe positively charged hydrogen end of one molecule is attracted to the negatively charged end of another moleculeNegative end must be an extremely electronegative element (fluorine, oxygen, or nitrogen - FON)Hydrogen bonds are the strongest weak bondthe 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 temperaturesHydrogen 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.
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 containerK-M Theory:Molecules still in motion, but closer togetherNo fixed positionsForces of attraction between particles limit their range of motion
30 Density and Compression Liquids are much denser than gassesAt 1 atm and 25C, water is 1250 times more denser than water vaporThey are the same temperature so they have the same average kinetic energySo density has to do with the intermolecular forcesLiquids CAN be compressed, but:Change in volume is much smallerNot much room for atoms to be pushed together
31 Fluidity Fluidity is the ability to flow Both gasses and liquids are fluidsLiquids can flow through liquids so they can diffuse, but liquids diffuse more slowly than gasses at the same temperatureLiquids 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 viscousSize and shape of molecules affect viscosityMolecules with long chains will have higher viscosity than shorter chainsMotor oil is very long chains of carbon and hydrogen
33 Viscosity Viscosity decreases with Temperature Increases with lower temperatureExample: Different motor oilsAdded energy gives molecules more kinetic energy to overcome intermolecular forces
34 Surface TensionIntermolecular forces do NOT have an equal effect on all particlesParticles in the middle can be attracted to particles all around themParticles at surface have no attraction at top to balance forces from belowSurface has smallest possible area for and acts like a stretched drum headTakes energy to break attractions from interiorResult: SURFACE TENSIONSurface Tension is Energy required to overcome attractions from interior of liquid
35 Surface TensionStronger the intermolecular forces, higher the surface tensionWater has a high surface tension:Hydrogen bonding provided high surface tensionDrops are shaped like spheres because surface area of sphere is lowerSome spiders can walk on water because of surface tensionFalling into water at high speed feels like hitting concreteCompounds to lower surface tension are called “surfactants”Used in cleaning supplies and soaps
36 Capillary ActionWater in a narrow container (like graduated cylinder) has a surface that is not straightConcave ‘meniscus’Two forces at workCohesion – force of attraction between two identical moleculesAdhesion – force of attraction between different moleculesAdhesion of water to glass is higher than water to waterSo water looks like it is ‘climbing’ the walls of the cylinder
37 Capillary ActionIf cylinder is very small, a thin film of water will be drawn upwardNarrow tubes are called capillary tubesUpward movement is called capillary actionCapillary action is what makes paper towels absorb water upwardWater forms hydrogen bonds with cellulose in paperSame action helps keep disposable diapers dry.
39 SolidsAccording 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 SolidsParticles in solids are more densely packed than liquidMOST Solids are more dense than liquidsWhen 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 SolidsMost solids form a solid packed in a crystalline solid formPredictable pattern formIndividual pieces are called crystalsSmallest arrangement of a piece is a unit cellShape of crystal is determined by shape of the unit cellThere are 7 categories of crystalsCrystal shapes differ based on faces, angles, and length of edges
43 Solids Molecular solids – held together by intermolecular forces only Usually not solids at room temperature or have low melting points like sugarCovalent network solids – like carbon and silicon form multiple covalent bondsVery hard, very high melting points (diamond, quartz, glass)Ionic solids – like salt with alternating atoms of different ionic chargeVery hard, brittle, high melting points, poor conductivity
44 SolidsMetallic Solids – all metallic elements have atoms surrounded by mobile electronsStrength varies, high melting points, good conductorsAmorphous Solids – particles are not arranged in regular, repeating pattern“amorphous”, Greek, means without shapeForms when liquids cool too quickly so crystals cannot formExamples: glass, rubber, plastic
46 Phase ChangesMost substances can exist in 3 states (Solid, Liquid, Gas)Depending on temperature and pressureEach state is referred to as a “Phase” when they coexist as physically different parts of a mixtureIce water is a heterogeneous mixture of 2 phasesWhen energy is added or removed, one phase can change into another
47 Phase Changes Requiring Energy MeltingVaporizationSublimation
48 MeltingAmount of energy needed to melt a substance depends on forces keeping particles together.Water takes a relatively high amount of energy because of hydrogen bondingAdding energy allows molecules to move faster, breaking the hydrogen bondsTemperature at which the forces holding crystal lattice together are broken and it becomes liquid is melting point
49 VaporizationIn 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 vaporVaporization is the process of changing a liquid to a gas
50 Vaporization (Cont)Vaporization that occurs at the surface of a liquid is called evaporationEvaporation is gradualParticles at surface are attracted to fewer other molecules than those in centerEven at cold temperatures, some molecules have enough energy to break the attractions and become gas.Evaporation is how your body cools itselfWater in sweat absorbs heat from your bodyWater 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 pressureTemperature at which the vapor pressure of a liquid equals the external atmospheric pressure is boiling pointWhat happens to the boiling point if the pressure is reduced? Increased?Where/when do we experience this?
52 SublimationSublimation is when a substance goes from solid directly to gas without becoming a liquidSolid iodineFrozen carbon dioxide (dry ice)Moth balls (p-dichlorobenzene)Solid air freshenersIce 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 CondensationDepositionFreezing
55 Condensation Condensation – process of changing from gas to liquid When molecule loses is energy,velocity is reducedIntermolecular forces take over againWhen hydrogen bonds form again, energy is released (heat comes out)Different causes for condensation, all involve loss of energy:Contact with cold item
56 DepositionDeposition – changing from gas directly to solidSnowflakes
57 Freezing Freezing – process of changing from liquid to solid Remove energy from liquidMolecules don’t move past each other any longerMolecules stay in fixed, set positionBecome solid
58 Phase DiagramsPhase diagram shows phase of matter at different temperatures and pressuresEach substance uniqueX-axis usually temperatureY-axis usually pressureThere is usually a “Triple Point” where all three phases can coexist“Critical Point” – temperature and pressure at which above substance cannot exist as liquid