chapter 10 states of matter 10.1 the kinetic-molecular theory of matter

Chapter 10Chapter 10States of MatterStates of Matter

10.1 The Kinetic-Molecular 10.1 The Kinetic-Molecular Theory of MatterTheory of Matter

Kinetic-Molecular Theory of GasesKinetic-Molecular Theory of Gases

• Particles of matter are ALWAYS in motion; constant, rapid motion. (kinetic energy!)

• Particles are very small & relatively far apart.

• Collisions of particles with container walls cause pressure exerted by gas.

• Volume of individual particles is zero.

• Particles exert no attractive or repulsive forces on each other.

• Gas particles undergo elastic collisions: Collisions in which no energy is lost

• Average kinetic energy is directly proportional to Kelvin temperature of a gas.

Kinetic-Molecular Theory of GasesKinetic-Molecular Theory of Gases

Air Hockey Table

Ideal GasIdeal Gas

• An An imaginaryimaginary gas that perfectly fits all the gas that perfectly fits all the assumptions of the kinetic-molecular theory assumptions of the kinetic-molecular theory

• A gas with its particles in constant random A gas with its particles in constant random motion without attraction for each other is motion without attraction for each other is called an called an Ideal GasIdeal Gas. These particles undergo . These particles undergo elastic collisions. elastic collisions.

• Nearly all real gases behave as ideal gases Nearly all real gases behave as ideal gases EXCEPT at very low temperatures or high EXCEPT at very low temperatures or high pressures.pressures.

Real GasesReal Gases• A gas that does not behave completely according

to the assumptions of the kinetic-molecular theory.

• Real gases occupy space and exert attractive forces on one another

Likely to behave nearly Likely to behave nearly ideally:ideally:

Likely not to behave ideally:Likely not to behave ideally:

Gases @ high temp. & low Gases @ high temp. & low pressurepressure

Gases @ low temp. & high Gases @ low temp. & high pressurepressure

Small non-polar gas moleculesSmall non-polar gas molecules Large polar gas moleculesLarge polar gas molecules

Kinetic-Molecular Theory of the Kinetic-Molecular Theory of the Nature of GasesNature of Gases

• Expansion Expansion Gases do not have a definite shape or volume Gases take the shape of their containers Gases evenly distribute themselves within a container

• Fluidity Fluidity Gas particles easily flow past one another

• Low Density Low Density A substance in the gaseous state has 1/1000 the density of the same substance in the liquid or solid state

• Compressibility Compressibility Gases can be compressed, decreasing the distance between particles, and decreasing the volume occupied by the gas


• Diffusion Diffusion

Spontaneous mixing of particles of two substances caused by their random motion – Rate of diffusion is dependent

upon: • speed of particles speed of particles • diameter of particles diameter of particles • attractive forces between the attractive forces between the

particles particles

• EffusionEffusion

Process by which particles under pressure pass through a tiny opening – Rate of effusion is

dependent upon: • speed of particles speed of particles

(small molecules have (small molecules have greater speed than greater speed than large molecules at the large molecules at the same temperature, so same temperature, so the effuse more rapidly) the effuse more rapidly)



10.2 Liquids10.2 Liquids

Some Properties of a LiquidSome Properties of a LiquidSurface TensionSurface Tension: :

The resistance to an increase in its surface area (polar molecules, liquid metals).

A force that tends to pull A force that tends to pull adjacent parts of a adjacent parts of a liquid's surface together, liquid's surface together, thereby decreasing thereby decreasing surface area to the surface area to the smallest possible size.smallest possible size.

Some Properties of a LiquidSome Properties of a Liquid

Capillary Action:Capillary Action: Spontaneous rising of a liquid in a narrow tube.

Some Properties of a LiquidSome Properties of a Liquid

ViscosityViscosity: : Resistance to flow (molecules with large intermolecular forces).

Some Properties of LiquidsSome Properties of Liquids

Volatility

• Liquids that have weak forces of attraction and evaporate easily

Nonvolatile Liquids

• Liquids that have strong forces of attraction and do not evaporate easily

Properties of FluidsProperties of FluidsRelative High Density Relative High Density • 10% less dense than solids (average) • Water is an exception • 1000x more dense than gases

Relative Incompressibility Relative Incompressibility • The volume of liquids doesn't change

appreciably when pressure is applied

Ability to Diffuse Ability to Diffuse • Liquids diffuse and mix with other liquids • Rate of diffusion increases with temperature


10.3 Solids10.3 Solids

Types of SolidsTypes of Solids

Crystalline SolidsCrystalline Solids: : highly regular arrangement of their components

[table salt (NaCl), pyrite (FeS2)].

Types of SolidsTypes of SolidsAmorphous solids aka

supercooled liquids: considerable disorder in their structures (glass).

• Greek for "without shape"

• Formation of amorphous solids:

• Rapid cooling of molten materials can prevent the formation of crystals

* They do not have definite * They do not have definite melting pointsmelting points

Representation of Components in Representation of Components in a Crystalline Solida Crystalline Solid

Lattice: A 3-dimensional system of points designating the centers of components (atoms, ions, or molecules) that make up the substance.

Types of Crystalline SolidsTypes of Crystalline Solids

Ionic Solid: contains ions at the points of the lattice that describe the structure of the solid (NaCl).

Unit Cell Unit Cell • The smallest portion of a crystal

lattice that shows the three-dimensional pattern of the entire lattice

Types of Crystalline Types of Crystalline SolidsSolids

Molecular SolidMolecular Solid: discrete : discrete covalently bonded covalently bonded molecules at each of its molecules at each of its lattice points (sucrose, lattice points (sucrose, ice).ice).

Packing in MetalsPacking in Metals

ModelModel: Packing uniform, hard spheres to best use : Packing uniform, hard spheres to best use available space. This is called available space. This is called closest packingclosest packing. . Each atom has 12 nearest neighbors.Each atom has 12 nearest neighbors.

Closest Packing Closest Packing HolesHoles

Metal AlloysMetal Alloys

Substitutional AlloySubstitutional Alloy: : some metal atoms some metal atoms replacedreplaced by others of by others of similar size.similar size.

• brass = Cu/Znbrass = Cu/Zn

Metal AlloysMetal Alloys(continued)(continued)

Interstitial Alloy:Interstitial Alloy: Interstices (holes) Interstices (holes) in in closest packed metal closest packed metal structure are occupied structure are occupied by by smallsmall atoms. atoms.

steel = iron + carbonsteel = iron + carbon

Network SolidsNetwork Solids

•Composed of strong directional Composed of strong directional covalent covalent bondsbonds that are best viewed as a “giant that are best viewed as a “giant molecule”.molecule”.

- brittle (non-flexible)brittle (non-flexible)- do not conduct heat or electricitydo not conduct heat or electricity- carbon, silicon-basedcarbon, silicon-based

•graphite, diamond, ceramics, glassgraphite, diamond, ceramics, glass

Sulfur – SSulfur – S88

Phosphorus – PPhosphorus – P44

DiamondDiamond

GraphiteGraphite

ZirconiaZirconia


10.4 Changes of State10.4 Changes of State

EquilibriumEquilibrium

• Dynamic condition in which two opposing changes occur at equal rates in a closed system

• A closed system at constant temperature will reach an equilibrium position at which the rates of evaporation and condensation will be the same

Equilibrium Vapor PressureEquilibrium Vapor Pressure• The pressure of the The pressure of the

vapor present vapor present at at equilibriumequilibrium..

• Determined principally Determined principally by the size of the by the size of the intermolecular forces in intermolecular forces in the liquid.the liquid.

• Increases significantly Increases significantly with temperature.with temperature.

• Volatile liquidsVolatile liquids have have high vapor pressures. high vapor pressures.

Increasing the temperature will move more particles into the vapor phase to compensate for the new

energy

BoilingBoiling

Boiling PointBoiling Point • The temperature at which the equilibrium vapor The temperature at which the equilibrium vapor

pressure of the liquid equals the atmospheric pressure of the liquid equals the atmospheric pressure pressure

Water boils at 100 °C at 1 atm pressure Water boils at 100 °C at 1 atm pressure

Water boils above 100 °C at higher pressures Water boils above 100 °C at higher pressures

Water boils below 100 °C at lower pressuresWater boils below 100 °C at lower pressures

The conversion of a liquid to a vapor within the liquid as well as at its surface. It

occurs when the equilibrium vapor pressure

of the liquid equals the atmospheric pressure

LeChatelier’s LeChatelier’s PrinciplePrinciple

When a system at equilibrium is placed

under stress, the system will undergo a change in such a way as to relieve

that stress.

When you take something away from a system When you take something away from a system at equilibrium, the system at equilibrium, the system shiftsshifts in such a way in such a way as to as to replace what you’ve taken away.replace what you’ve taken away.

Translation:Translation:

When you add something to a system at When you add something to a system at equilibrium, the system equilibrium, the system shiftsshifts in such a way in such a way as to as to use up what you’ve added.use up what you’ve added.

LeChatelier’s Example #1LeChatelier’s Example #1

A closed container of ice and water at A closed container of ice and water at equilibrium. The temperature is raised.equilibrium. The temperature is raised.

Ice + Energy Ice + Energy Water Water

The equilibrium of the system shifts to The equilibrium of the system shifts to the _______ to use up the added energy.the _______ to use up the added energy.rightright


A closed container of NA closed container of N22OO44 and NO and NO22 at at

equilibrium. NOequilibrium. NO22 is added to the container. is added to the container.

NN22OO44 + Energy + Energy 2 NO 2 NO22

The equilibrium of the system shifts to The equilibrium of the system shifts to the _______ to use up the added NOthe _______ to use up the added NO22..leftleft


A closed container of water and its vapor at A closed container of water and its vapor at equilibrium. Vapor is removed from the system.equilibrium. Vapor is removed from the system.

water + Energy water + Energy vapor vapor

The equilibrium of the system shifts to The equilibrium of the system shifts to the _______ to produce more vapor.the _______ to produce more vapor.rightright

Water phase changesWater phase changesTemperature remains __________during a phase change.

constant

Phase DiagramPhase DiagramRepresents phases as a function of temperature Represents phases as a function of temperature and pressure.and pressure.Critical temperatureCritical temperature: temperature above which : temperature above which the vapor can not be liquefied.the vapor can not be liquefied.Critical pressureCritical pressure: pressure required to liquefy : pressure required to liquefy ATAT the critical temperature. the critical temperature.Critical pointCritical point: critical temperature and pressure : critical temperature and pressure (for water, (for water, TTcc = 374°C and 218 atm). = 374°C and 218 atm).

WaterWaterWaterWater

Phase changes by NamePhase changes by Name

Carbon dioxideCarbon dioxideCarbonCarbondioxidedioxide

CarbonCarbonCarbonCarbon

SulfurSulfur


10.5 Water10.5 Water

Water’s Water’s PropertiesProperties

Sea IceSea Ice

• Ice forms on top of the ocean in a thin layer & acts to insulate the warmer waters below from the colder air temperatures.

• This occurs in the polar regions, the Artic & Antarctic

• Since ice is less dense than liquid water, it will float on top, instead of sinking which would kill all life below the surface. •Sea ice is not the same as an iceberg. Icebergs are pieces of glaciers which are formed by snowfall on land. • Sea ice is not salty, as the hydrogen bonds that hold ice together will not form properly if salt remains in the structure.

Sea Water DensitySea Water Density• There are 2 main factors that affect the density of There are 2 main factors that affect the density of

sea water:sea water:Temperature & SalinityTemperature & Salinity

• As the temperature decreases, density increases.As the temperature decreases, density increases.• As salinity increases, density increases.As salinity increases, density increases.• Since the least dense layer of a liquid will float Since the least dense layer of a liquid will float

above a more dense layer, the warmest & lowest above a more dense layer, the warmest & lowest salinity layer will be on top. salinity layer will be on top.

• However, temperature has a greater effect on However, temperature has a greater effect on density than salinity. density than salinity.

• That means that a higher salinity layer can float on That means that a higher salinity layer can float on top of a lower salinity layer if it is considerably top of a lower salinity layer if it is considerably warmer in temperature. warmer in temperature.

Molar Heat of FusionMolar Heat of Fusion

• The amount of heat energy required to melt one mole of solid at its melting point.

• 6.009 kJ per 1 mole

Practice!

How much energy is absorbed when 16.3 g of ice melts?

16.3 g x 1 mole x 6.009 kJ = 5.44 kJ

18 g 1 mol

Molar Heat of Vaporization Molar Heat of Vaporization • The amount of heat energy required to vaporize

one mole of a liquid at its boiling point • Strong attractive forces between particles result in

high molar heat of vaporization • 40.79 kJ per 1 mole

Practice!

Find the mass of liquid water required to absorb 5.23 x 104 kJ of energy upon boiling.

5.23 x 104 kJ x 1 mole x 18 g = 2.31 x 104 g 40.79kJ 1 mol

Specific HeatSpecific Heat

• The amount of energy required to change the temperature of 1 gram of a substance by 1 degree Celsius is known as specific heat capacity. (Ch.16 p.533)

Q = s * m * Q = s * m * ΔΔTT

Where Q = energy (heat) required

s = specific heat capacity

m = mass of sample, grams

ΔT = change in temperature, Celsius

Specific Heat Capacities of Some Specific Heat Capacities of Some Common SubstancesCommon SubstancesSubstanceSubstance Specific Heat Capacity (J/gSpecific Heat Capacity (J/g°C)°C)

Water, LiquidWater, Liquid 4.1844.184

Water, IceWater, Ice 2.032.03

Water, SteamWater, Steam 2.02.0

Aluminum, sAluminum, s 0.890.89

Iron, sIron, s 0.450.45

Mercury, lMercury, l 0.140.14

Carbon, sCarbon, s 0.710.71

Silver, sSilver, s 0.240.24

Gold, sGold, s 0.130.13

Practice!Practice!A piece of copper alloy with a mass of 85.0 g is heated from 30.°C to 45°C. In the process, it absorbs 523 J of energy as heat.

a.What is the specific heat for this copper alloy?

b.How much energy will the same sample lose if it is cooled from 45°C to 25°C?

Answers…

a.0.41 J/g∙K

b.7.0 x 102 J

Specific GravitySpecific Gravity

• A function of density.A function of density.• The ratio of density of a material to the density of The ratio of density of a material to the density of

water at a specified temperature. water at a specified temperature. • The density of water is usually 1 g/cmThe density of water is usually 1 g/cm33

• When dividing density by density, the units cancel When dividing density by density, the units cancel and therefore, and therefore, specific gravity has no units!specific gravity has no units!

• If the density of the material is in the units g/cmIf the density of the material is in the units g/cm33 (or g/mL) when dividing by the density of water, (or g/mL) when dividing by the density of water, the numeric value of the material’s density is the the numeric value of the material’s density is the same, but the units cancel out. same, but the units cancel out.

Intermolecular ForcesIntermolecular Forces((Review from Chapter 6)Review from Chapter 6)

Dipole-dipole attractionDipole-dipole attractionHydrogen bondsHydrogen bondsDispersion forcesDispersion forces

Forces of attraction between different Forces of attraction between different molecules rather than bonding forces molecules rather than bonding forces within the same molecule.within the same molecule.

Dipole-Dipole AttractionDipole-Dipole Attraction

dipole-dipole attractiondipole-dipole attraction: molecules with : molecules with dipoles orient themselves so that “+” and “dipoles orient themselves so that “+” and “” ” ends of the dipoles are close to each other.ends of the dipoles are close to each other.

Dipole Dipole ForcesForces

Hydrogen BondingHydrogen Bonding

hydrogen bondshydrogen bonds: dipole-dipole attraction in : dipole-dipole attraction in which hydrogen is bound to a highly which hydrogen is bound to a highly electronegative atom. (electronegative atom. (F, O, NF, O, N))

Hydrogen Bonding in WaterHydrogen Bonding in Water

Hydrogen Bonding in DNAHydrogen Bonding in DNA

London Dispersion ForcesLondon Dispersion Forces

Dispersion forcesDispersion forces: relatively weak: relatively weak forces forces caused by instantaneous dipole, in which caused by instantaneous dipole, in which electron distribution becomes asymmetrical.electron distribution becomes asymmetrical.

The ONLY forces of attractionThe ONLY forces of attraction that exist among that exist among noble gas atoms and nonpolar molecules. (Ar, noble gas atoms and nonpolar molecules. (Ar, CC88HH1818))

London Dispersion ForcesLondon Dispersion Forces

Boiling point as a measure of intermolecular Boiling point as a measure of intermolecular attractive forcesattractive forces

chapter 10 states of matter 10.1 the kinetic-molecular theory of matter

Documents

collisions of particles

gas kineticmolecular

volume of individual

particles effusion process

volume gases

ideal gases

real gasesa gas

containers gases