Ch. 15.1 Kinetic TheoryCh. 15.1 Kinetic Theory

1.All matter is made of atoms and molecules that act like tiny particles.

Kinetic TheoryKinetic Theory

2.These tiny particles are always in motion. The higher the temperature, the faster the particles move.

Kinetic TheoryKinetic Theory

3.The more massive the particles, the slower the particles will diffuse and flow.

Gas PressureGas Pressure– The result of

simultaneous collisions of billions of gas particles with an object.

– The more collisions, the greater the pressure

Near Death

VacuumVacuum•A controlled condition where no gas particles are present.

•So no gas pressure can exist.

Atmospheric PressureAtmospheric Pressure•Results from collisions of air molecules with objects.

• Decreases as you climb a mountain because the air thins out at higher elevations

•Measured by a barometer

Measuring PressureMeasuring Pressure

•STP (Standard Temp. Pressure)•Standard Temperature at sea level is 00C or 273 K

•Standard Pressure is 101.3 kPa, 760 torr, 760 mm Hg, or 1 atm

Pressure ConversionsPressure Conversions• How many kPa’s are in 1.50 atm?1 atm = 101.3 kPa1.50 atm x 101.3kPa =

1 atm• How many kPa’s are in 690 mm Hg?690 mm Hg x 101.3 kPa =

760 mm Hg

152 kPa

92 kPa

Energy and TemperatureEnergy and Temperature• Kinetic Energy measures the

average speed of particles.• The higher the temperature the

greater the particle speed.• Temperature measures kinetic

Energy• SI base unit is Kelvin (K)

Converting from CelsiusConverting from Celsius• 0 degree Celsius is equal to 273 K• K = 0C + 273• 0C = K – 273• Convert 191 K to Celsius• 0C = 191 – 273 = -82 0C

Absolute ZeroAbsolute Zero• Theoretical temperature at which all

motion stops.• Scientists have experimented a tenth

of degree to 0 K, but have never gotten 0 K.

Ch18.1Ch18.1 Ideal/Real GasesIdeal/Real Gases

•no definite shape. (both)•no definite volume. (both)•Particles move rapidly in constant random motion (both)

•Low density (both)•All collisions perfectly elastic, no attractive forces (Ideal)

Ideal/Real GasesIdeal/Real Gases•Real gas particles will stick together and attract

•A real gas will behave like an ideal gas at low pressures or high temperatures

•Hydrogen and helium always behave ideally due to there small masses.

Gas RelationshipsGas Relationships• Relationships between pressure,

volume, temperature and number of moles (amount)

• While examining relationships, two measurements will always be constant (unchanged)

Pressure Pressure vsvs VolumeVolume

Real Gas

1. Pressure 1. Pressure vsvs VolumeVolume• Boyle’s Law• For a given mass of gas at a constant

temperature, the volume of the gas varies inversely with pressure.

• Pressure increases, volume decreases• As volume (space) decreases, the

particles become closer and collide (pressure) more often.

• P1 V1 =P2 V2

Practice ProblemPractice Problem• If you had a gas that exerted 202 kPa of

pressure and took up a space of 3000.0 mL. If you decide to expand the tank to 7.00 L, what would be the new pressure? (Assume constant temperature)

• P1 V1 =P2 V2

• Check units• 202 kPa x 3.00 liters = P2 x 7.00 liters• 606 = P2 x 7.00 liters• P2 = 86.6 kPa

Ideal gas

2. Temperature 2. Temperature vsvs VolumeVolume• Charles Law• For a given mass of gas at a constant

pressure, the volume of the gas varies directly with its Kelvin temperature.

• Temperature increases, volume increases• As temperature (speed of particles)

increases, the particles move farther apart increasing volume (space) while maintaining a constant pressure.

• V1 /T1 =V2 /T2 or V1 T2 =V2 T1

Practice ProblemPractice Problem• If you took a balloon outside that was at

20.00C at 2.0 liters and heated up to 29.00C, what volume would the balloon occupy now? (Assume constant pressure)

• V1 T2 =V2 T1

• Check units(Remember KELVIN)• 2.0 L x 302 K = V2 x 293 K• 604 = V2 x 293 K• V2 = 2.1 L

CHARLES LAW: CHARLES LAW: ΔΔT and T and ΔΔVV

CHARLES in charge was on CHARLES in charge was on TVTV

Ideal Gas

3. Temperature 3. Temperature vsvs PressurePressure• Gay-Lussac’s Law• For a given mass of gas at a constant

volume, the pressure of a gas varies directly with its Kelvin temperature.

• Temperature increases, pressure increases• As temperature (speed of particles)

increases, the particles collide (pressure) more often in a set volume (space).

• P1 /T1 =P2 /T2 or P1 T2 =P2 T1

Ostrich Egg in Ostrich Egg in MicrowaveMicrowave

Combined Gas LawCombined Gas Law• Combines all three gas laws into one

expression.

Practice ProblemPractice Problem• You have a 2.0 liter balloon that was at

20.00C and 1.5 atm. If you take this balloon and place it in a room at STP conditions, what volume would the balloon occupy?

• P1 V1 T2 =P2 V2 T1 (Remember KELVIN)• 1.5 atm x 2.0 L x 273 K = 1atm x V2 x 293 K

• 819 = V2 x 293 K• V2 = 2.8 L

4. Moles (amount) 4. Moles (amount) vsvs TempTemp• moles increases, temp. decreases• Inverse relationship• In a set volume (space), adding more

moles of a gas (amount), will cause the particles to slow down (temp.) in order to maintain a constant pressure.

• Compress tanks become colder as you fill them

Ideal Gas

5. Moles (amount) 5. Moles (amount) vsvs pressurepressure• moles increases, pressure increases• In a set volume (space), adding more

moles of a gas (amount), will cause more collisions (pressure) between gas particles.

• Think of a super soaker or simply filling your tire

6. Avogodro6. Avogodro’’s Law (ch19.1)s Law (ch19.1)• Amount (moles) is directly proportional

to the space occupied.• The greater the moles of a gas

(amount), the more volume (space) the particles will need in order to maintain constant pressure (particles collide)

• 1 mole of gas at STP= 22.4 liters of any gas

Practice ProblemPractice Problem• How many liters of Hydrogen are in

6.2 grams of H2 at STP?• Molar of mass of is H2 2 gram/mole• 6.2 grams H2 x 22.4 L H2 / 2 gram of H2

• 69 L of H2

7. Ideal Gas Law (ch19.1)7. Ideal Gas Law (ch19.1)•PV = nRT•R = constant

0.08206 (L*atm)/(mol*K)8.31 (L*kPa)/(mol*K)

•n = represents the number of moles.

•Can be used in determining densities of different gases.

Application for Ideal Gas LawApplication for Ideal Gas Law

•Tire companies have to manage the amount of moles a tire can hold under different temperatures and pressures in order to properly engineer tires.

•When you fill your tire, you are applying the Ideal Gas Law

Practice ProblemPractice Problem• A propane tank that holds 3000. g of

C3 H8 . How much larger a container would be needed to hold the same amount of propane if the gas is at 250C and a pressure of 3.0 atm?

SolutionSolution• PV=nRT.• First solve for n• 3000. g of C3 H8 . x 1 mole =

44 grams C3 H8

• V*3.0 atm= 68.18 moles x 0.0821x298K• V = 560 L

68.18 moles

Density and Molar Mass can Density and Molar Mass can the Ideal Gasthe Ideal Gas

• D=P (MM) / RT (MM=molar mass)• MM = mRT/PV (m= mass)• What is the molecular mass of an

unknown gas if 8.11 g of it occupy 2.38 L at 109.1 kPa and 10.00C?

• MM = mRT/PV• MM = 8.11 g x 8.31 x 283K /109.1 kpa x 2.38 L

• MM= 73.5 g/ mol

Practice ProblemPractice Problem

•2.0 grams of N2 is kept under a pressure of 0.95 atm, and a temperature of 30.00C. What is the density of the gas under these conditions?

SolutionSolution• PV=nRT.• First solve for n• 2.0 g of N2 x 1 mole =

28 grams N2

• V*0.95 atm= 0.071 moles x 0.0821x303K• D=m/v• D =2.0 g / 1.9 L = 1.1g/L

0.071 moles

8. Graham8. Graham’’s Law of Diffusion s Law of Diffusion (ch18.2)(ch18.2)

•Diffusion is the random scattering of gas molecules.

•The longer they diffuse the more evenly distributed they will become in the container.

•The heavier the gas the slower the rate of diffusion.

LAST LAW!LAST LAW!I PROMISEI PROMISE

(ch18)(ch18)

Gas Law Gas Law Review RapReview Rap

Ch. 17.1Ch. 17.1Changing StatesChanging States

liquid gassolid

Changes of StateChanges of State

Exothermic Process

Endothermic Process

Melting and FreezingMelting and Freezing• Melting Point is the temperature at

which a solid becomes a liquid• Melting and freezing take place at

the same threshold temperature. • According to Kinetic Theory, almost

all solids and liquids expand and become disordered when the temperature is raised.

Evaporation Evaporation • Conversion of a liquid to a gas or

vapor below its boiling point.• It occurs only at the surface.• Remember the difference between a

vapor and gas.–Vapor is normally a liquid or solid at

room temperature

Vapor PressureVapor Pressure• The pressure exerted by a vapor in

equilibrium with its liquid state.• Vapor pressure measures how easily

a liquid changes into vapor• Liquids with high vapor pressures

turn into vapors very easily. (Volatile liquids)

Ex. Gasoline, perfume

Once equilibrium is reached, the vapor particles will begin

to condense back to a liquid at the same rate they change into a vapor.

Dynamic EquilibriumDynamic Equilibrium

Vapor Equilibrium reached

SublimationSublimation• Process where solid goes directly to a

gas (vapor), because the vapor pressure is so high, liquid phase does not exist.

• Ex. Iodine, Dry Ice

Boiling PointBoiling Point• The temperature at which the vapor

pressure of the liquid equals the atmospheric pressure

• The entire liquid is changing state, not just the surface.

• Liquids with low boiling points are considered volatile

Super Heated Super Heated WaterWater

DistillationDistillation• A method of separating substance

with different boiling points.• Used in desalinating sea water.

LiquefactionLiquefaction• Changing a gas into a liquid. • A gas can be changed into a liquid by

two methods:–must be placed under tremendous

pressure (compressing)–Placed in really cold temperature

conditions

Intermolecular ForcesIntermolecular Forces• The forces holding molecules to each

other.• What phase is strongest?• Solids• What phase is the weakest?• Gases (vapor)

STRONG INTERMOLECULAR FORCESSTRONG INTERMOLECULAR FORCES

• Don’t change phase easily• High melting points• Low vapor pressure• Nonvolatile Substances• High boiling points• High viscosity• High surface tension

WEAK INTERMOLECULAR FORCESWEAK INTERMOLECULAR FORCES

• Do change phase easily• Low melting points• High vapor pressure• Volatile Substances• Low boiling points• Low viscosity• Low surface tension

Phase DiagramPhase Diagram• Shows how states of matter of a

substance are affected by changes in temperature and pressure.

• Triple Point- point where all three states of matter meet.

• Critical Point- Point where only the vapor can exist.

Heating CurveHeating Curve• Used to show how much enthalpy

energy (Heat transfer) is needed to change phase.

• Enthalpy (heat) of Fusion- energy required to change from solid to liquid

• Enthalpy (heat) of Vaporization- energy required to change from liquid to vapor

Heating Curve

Enthalpy of Fusion

Enthalpy of Vaporization

Ch. 17.2 LiquidsCh. 17.2 Liquids• Definite volume• no definite shape(takes shape of container)• Difficult to compress• disorderly arrangement on particles• Flowing motion of particles

Liquid PropertiesLiquid Properties

• Viscosity- the resistance of a fluid to flow

•Thick fluids have high viscosity

•Ex. Syrup

http://photos.runic.com/photos/syrup1.jpg

Liquid PropertiesLiquid Properties

• Surface Tension- Ability of liquid molecules to hold on to each other.

• Apparent “skin” affect• Ex. Over filling a liquid in a glass

with out the liquid spilling

Hg

Liquid PropertiesLiquid Properties

• Capillary Rise- the tendency of a liquid to rise in a small diameter tube due to the surface tension of the liquid.

• Used to measure surface tension of a liquid

Hydrogen BondingHydrogen Bonding• Causes water to be very polar.• This bonding also causes water to decrease in

density and expand as it freezes (increases space between molecules)

Ch. 16 SolidsCh. 16 Solids•Definite shape•Definite volume•Difficult to compress•Orderly arrangement of particles

•Smallest amount of movement of particles.

Solid StructuresSolid Structures• Metal Solids• Crystal structure (repeating patterns)• Allotropes (different forms of same

element, ex. Carbon)• Amorphous (no crystal structure)

–Glasses, rubber, plastics