the kinetic molecular theory0

8/10/2019 The Kinetic Molecular Theory0

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States of MatterGases, Liquids and Solids


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The Kinetic Molecular Theory

Literal interpretation:

The theory of moving molecules


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Observations to support the theory:

Diffusion in gases and liquids

Movement of substances from an area of highconcentration to one of lower concentration

Ability of a gas to spread out and fill a container

Brownian MovementThe observable movement of particles due to collisions

with moving molecules.


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The theory explains these observations

The theory describes the differences between gas,liquids and solids

The theory explains the gas laws


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Major points: Supports the concept of an ideal gas

An ideal gas is one that perfectly fits all the assumptions ofthe kinetic-molecular theory.

Do not actually existin theory this is how they wouldbehave:


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1. Gases are made of tiny particles far apart relative to

their size:

Volume occupied by the molecules is inconsequential

Volume is mostly space

Explains why gases are compressible


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2. Gas particles are in continuous, rapid, random

motion

As a result there are collisions with othermolecules or with the wall of the container

Creates pressure

Increase in temperature increases themovement of the molecules and thus thepressure exerted by the gas


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3. There are no attractive forces between molecules

under normal conditions of temperature and

pressure

Gas molecules are moving too fast

Gas molecules are too far apart

Intermolecular forces are too weak


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4. Collisions between gas particles and betweenparticles and container walls are elastic collisions.

Collisions in which there is no net loss of total kineticenergy

Kinetic energy can be transferred between two particles

during collisions

Total kinetic energy remains the same as long astemperature remains the same


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5. All gases at the same temperature have the same average

kinetic energy. The energy is proportional to the

absolute temperature.

Absolute temperature = Kelvin temp scale

Ke= mv2

Ke = the kinetic energy

m = mass

v = the velocity


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1. Gases are made of tiny particles far apart relative to their size

2. Gas particles are in continuous, rapid, random motion

3. There are no attractive forces between molecules

under normal conditions of temperature and pressure

4. Collisions between gas particles and between particles andcontainer walls are elastic collisions.

5. All gases at the same temperature have the same averagekinetic energy. The energy is proportional to the

absolute temperature.


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Applies only to idealgases

Most gases behave likean ideal gas undernormal conditions

Gases with littleattraction betweenmoleculesHe/H2/N2

Real gases

Deviate from idealbehavior

Due to intermolecular

interaction (H2O, NH3) High pressure

Low temperature


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The Kinetic Theory and

Changes of State

GasesAttractions areinsignificant

LiquidsAttractions aremore important leadingto a more ordered state

Solids Attractions aremost important with anordered state


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Kinetic Molecular Theory and Changes of State

Attractions between particles in strength:

Least London dispersion forces

Dipole-dipole interaction

Hydrogen bonding

Greatest Metallic, Ionic and Covalent network


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Changes of state occurwith a change intemperature or pressure

Particles of a substanceovercome (or succumb)

to intermolecularattraction

Involvesenergy


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Solids, liquids and gasescan undergo variouschanges in processes thatare either endothermic

or exothermic


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Consider the evaporationof a liquid:

Temperature= theaverage kinetic energy

Some molecules havemore kinetic energy thanothers

These molecules escapeand become gasmolecules


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Evaporation will occur inclosed container alsoexcept

As the liquid evaporatesthe space above starts tofill with gas molecules

until it can hold no more

Gas will start tocondense.



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Eventually the rate ofevaporation will equalthe rate of condensation

Two processes will occursimultaneously with no

net change

State of Equilibrium

Vapor molecules abovethe liquid will collidewith each other and thecontainerand exert a pressure.

Equilibrium vaporpressure!!!


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Every liquid has a specificvapor pressure at a giventemperature.

Reflection of the strengthof the intermolecularbonding between

molecules

Vapor pressure alsoincreases with temperature


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Kinetic Molecular Theory and Changes of State Equilibrium vapor

pressure is (EVP) used todefine boiling point, BPt

Boiling point is thetemperature at which the

equilibrium vaporpressure equalsatmospheric pressure


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Boiling point of water is100 oC only at 760mm Hg

When atmosphericpressure is > 760 mm Hgthe boiling pt is > 100.

When atom0sphericpressure is


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Boiling requires acontinuous supply ofenergy..

Water boils at 100oC andthe temperature does notchange.even though

there is a continuoussupply of energy..

Where does the energy go?


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Same is true when icemelts

It melts (or freezes) at 00

C

At this temperature thereis a state of equilibrium

Temp will not change ifboth phases are present


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Energy is can be addedcontinuously, but thetemperature does not

change

Energy is used to change

the physical statethisrequires a lot of energy!!


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The amount of heatenergy required to meltone mole of a solid at the

solids melting point isthe solids molarenthalpy of fusion.

DHf

Energy absorbedrepresents potentialenergy

For water it is 6.009kJ/mol

Xj/g =6.009kJ/M x 1M/18g x 1000J/1kJ

= 333.8 j/g


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The amount of heatenergy required tovaporize one mole of a

liquid at the liquidsboiling point is theliquids molar enthalpyof vaporization.

D

Hv Energy absorbed

represents potentialenergy

For water it is 40.79kJ/mol

Xj/g =40.79J/M x 1M/18g x 1000J/1kJ

= 2266 j/g


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Compared to othersubstances these valuesare very high.

Water has very strongintermolecular bonding

Hydrogen bondsbetween highly polarmolecules


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Unique properties ofwater is related to thehydrogen bond

4-8 molecular groups inliquid water

Hexagonal arrangementin solid

--> Dipole w/ partial +/-

High boiling pt of

water

Solid is less dense..Ice

floats


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Changes of State are Shown in Phase Diagrams

Changes of phase aredepicted in phasediagrams

Show the relationshipbetween state of matter,

temperature andpressure


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Changes of State Shown in Phase Diagrams

Phase diagrams define:

Triple point=the T/P conditionsat which all three phases coexist

Critical point = Critical tempand press

Critical temp = temp abovewhich the substance cannot

exist as a liquid

Critical press= lowest pressure atwhich the substance can exist asa liquid at the criticaltemperature


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Phase Diagram of Water

Interesting points ADIce and vapor in

equilibrium

ACLiquid and vapor in

equilibrium ABIce and liquid in

equilibrium. Notean increase in pressurelowers melting point

nbp=normal boiling pt mp =melting point

Critial temp =373.99


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Phase Diagram of Carbon Dioxide

Note the following:

Very different temp andpressure compared towaters diagram

Liquid is only possible at

high pressureAt normal room

conditions CO2 onlyexists as a gas


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Phase Change vs Temperature change in a single phase Melting/Fusion

Molar heat of fusion

6.009 kJ/mol

Vaporizing

Molar hear of vaporization

40.79kJ/mol

Raising the temperature of ahomogeneous material

Specific heat


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Phase Change How much energy is absorbed when 47g of ice melts?

(at STP)

Energy =47g x1 mol x 6.009kJ

18g 1 mol

= 15.7 kJ


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Phase Change How much energy is absorbed when 47g of water

vaporizes? (at STP)

Energy =47g x1 mol x 40.79kJ

18g 1 mol

= 106 kJ (vs 15.7 kJgases have ahigher energy content)


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Phase ChangeWhat mass of steam is required to release 4.97 x 105kJ

of energy when it condenses?

grams =4.97 x 105kJ x 1mol x 18g

40.79kJ 1 mol

= 2.19 x 105g


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Temperature change in a single phase Quantity of energy

transferred as heat whilea temperature change

occurs depends on The nature of the

substance

The mass of the material

The size of thetemperature change.

Water has a high specificheat

Metals have low specificheat

Units = J/(g x oC)


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Temperature change in a single phase Specific heat of water (l) = 4.18 J/goC

Specific heat of water (s) = 2.06

Specific heat of water (g) = 1.87 Specific heat of ethanol (g) = 1.42

Specific heat of ethanol (l) = 2.44

Specific heat of mercury (l) = 0.140

Specific heat of copper (s) = 0.385

Specific heat of lead (s) = 0.129

Specific heat of aluminum (s) = 0.897


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Solids and the Kinetic Molecular theory Properties: Dominated by the fact that

Closely packed particles

Relatively fixed positions

Highest intermolecular or interatomic attractions

Properties are

Definite shape and volume Definite melting point

High density and incompressibility

Low rate of diffusion


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Solid structureSolids may be crystalline Solids may be amorphous

Crystals in which particles arearranged in a regular

repeating pattern

Particles are randomly arranged


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Solid structureCrystals

Total 3-D arrangement ofparticles is the crystal

structure CUBIC

BODY CENTERED CUBIC

TETRGONAL

HEXAGONAL TRIGONAL

MONO


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4-Classes of Crytsalline Solids

Ionic --IonsHard and Britle

Covalent Network

Network of molecules

Quartz (SiO)

Diamond

Metallic Crystals

Free moving e-

Covalent Molecular Crystals

Weak.

Water, dry ice


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Amorphous solidsWithout shape No regular pattern

Glasses Plastics


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the kinetic molecular theory0

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