States of Matter

Kinetic Molecular Theory

Particles are always in motion.Temperature is a measure of average kinetic energy of particles.Intermolecular forces hold particles together. Stronger forces require more energy (higher temp.) to overcome.

Intermolecular Forces

Always weaker than chemical bondAffect structure and state of matter

Dipole-Dipole Forces

Positive and negative ends of polar molecules attract each other.About 1% as strong as covalent or ionic bondsWeaken as distance between molecules increases

Hydrogen BondingEspecially strong dipole-dipole forceOccurs when H bonds to a strongly electronegative atom—O, N, or FVery strong because 1) molecule is very polar & 2) small size of H

H BondingExample—waterMore pronounced in molecules formed from small atoms (dipoles can come closer)High boiling point

London Dispersion Forces

Forces that exist in all atoms and molecules but that are significant only among Noble gases and nonpolar moleculesResult from temporary dipoles formed when electrons distribute themselves unevenly—can induce a dipole in a neighboring atomVERY WEAK

London Dispersion Forces

Stronger in larger atoms or molecules due to the greater chance of the formation of instantaneous dipoles.

Physical PropertiesMelting and boiling points are higher when IM attractions are strongerMore energy required to separate molecules

Which would have the higher boiling point &

Why?Cl2 or F2?

H2O or H2S?

SiBr2 or SBr2?

CH4 or C10H22?

O2 or NO?

States of MatterGases—weak IM forces (like London Dispersion Forces)Liquids—intermediate IM forcesSolids—strong IM forces

Liquids & IM ForcesSurface tension—result of IM forces that resist an increase in surface areaCapillary action—result of cohesive forces within liquid and adhesive forces between liquid and tube

Liquids (cont’d)Viscosity—the ability of a liquid to resist flow (resist change in shape)All effects are higher with more polar molecules.

SolidsAmorphous—without definite structureCrystalline—definite structures

Solids (Crystals)Ionic solids—made of charged particles; ions at lattice pointsMolecular solids—made of neutral particles; molecules at lattice pointsAtomic solids—made of neutral particles; atoms at lattice points; 3 types

An atomic solid, an ionic solid & a molecular solid

Ionic SolidsIons at lattice pointsClosest packed spheresArranged to minimize repulsions and maximize attractionsConducts only when melted

Molecular SolidsLattice positions occupied by moleculesInternal covalent bonds are strong, but intermolecular forces are weakIM force: dipole/dipole if polar covalent bond; London dispersion forces (larger in larger molecules)

Atomic Solids1. Network—directional covalent

bonds; forms giant molecules (diamond, graphite,and silicon); highest melting points

2. Metallic—delocalized covalent bonds; atoms have closest packing structure; high melting points

Atomic Solids3. Group 8A—Noble gases

—London dispersion forces only; low melting points.

Network Atomic Solids

Strong, directional bondsForm giant “molecules”Typically brittle & poor conductorsExamples—carbon and silicon

Carbon NetworkFollows a molecular orbital (not atomic orbital) model

DiamondTetrahedral--sp3 hybridized bonds stabilize structureLarge gaps exist between filled and unfilled molecular orbitals—hard for electrons to move—no conductivity

GraphiteFused carbon rings form sheetsTrigonal planar—sp2 hybridized (1 p orbital remains unhybridized)Delocalized electrons in orbital causes graphite to be conductive

Figure 10.22: The structures of diamond and graphite. In each case only a small part of the entire structure is shown.

Closest Packed Solids

aba pattern—alternating layers—atoms in 3rd layer lie directly above atoms in 1st layer—hexagonal unit cell—body centeredabca pattern—atoms in 1st and 4th layers are in line; 2nd & 5th layer; 3rd & 6th layer—face-centered cubic cell

aba Packing

abca Packing

Density of Closest Packed Solids

To calculate density, you need to know:

MASSVOLUME

MassFigure out how many atoms in one unit cellMultiply by molar massDivide by Avogadro’s numberYou now know a mass in grams

Face-Centered Cubic Unit Cell

If these are atoms of calcium, what is the mass of the cell?

VolumeDetermine the length of one side of the cube by using the atomic radius (varies depending on type of unit cell)Cube the side length.

Simple Cubic--aaa

If the atomic radius of this atom is 122 pm, what is the volume?

Body-Centered Cubic--aba

If the atomic radius of the atom is 246 pm, what is the volume of the cell?

Face-Centered Cubic--abca

If the atomic radius is 291 pm, what is the volume of the cell?

Sample ProblemSilver crystallizes in a face-centered cubic closest packed structure. The radius of the silver atom is 144 pm. Calculate the density of silver.

Bonding in MetalsStrong, non-directional bondsAtoms are hard to separate but easy to move.“Electron sea” modelMobile electrons carry heat or electricity easily

Band Model or Molecular Orbital

ModelElectrons travel around metal crystal in a molecular (instead of atomic) orbitalsResult is a continuum of levels that eventually merge to form a band.

Figure 10.19: The

molecular orbital

energy levels produced

when various numbers of

atomic orbitals

interact.

Band or MO ModelEmpty orbitals close in energy exist.Electrons are very mobile into and out of these similar-energy orbitals--CONDUCTIVITY.

SemiconductorsSome electrons can cross the “energy gap” between molecular orbitals—somewhat conductiveHigher temperatures result in more electrons’ being able to reach conductive bands

DopingAdding other elements with one more or one less electron than a semiconductor can increase conductivity

n-type semiconductor

An element with one more valence electron is addedMore valence electrons are available to move into conduction bandsWhat could be used to dope Si?

p-type semiconductor

An element with one less valence electron is addedThe absence of a valence electron creates a hole through which electrons can travel

AlloysIntroducing other elements into metallic structure is easy.Substitutional—some “host” atoms are replaced by atoms of similar size (Brass = copper + zinc)Interstitial—small atoms occupy spaces between metal atoms (such as carbon used to harden steel)

Phase Changes

Phase ChangesCondensation/Evaporation

Boiling point

Freezing/MeltingMelting point

Deposition/Sublimation

Energy of Phase ChangesHeating/Cooling Lab

Energy of Phase ChangesHeating Curve

Temperature changes when only one phase is presentNo temperature change during phase change

For WaterHeat of fusion—334 J/gHeat of vaporization—2260 J/gSpecific heat

1.84 J/goC for gas4.18 J/goC for liquid2.1 J/goC for solid

What amount of energy is needed to completely melt 45 g

of ice at its melting point?

What amount of energy is needed to completely vaporize

45 g of water at its boiling point?

Calculate the amount of energy required to raise the

temperature of 125 g of ice from -12oC to 75oC.

What is the final temperature/state of 95 g of ice at -15oC if 65 kJ of energy are

added?

The heat of vaporization for carbon dioxide is 571 J/g. How

much energy would be required to change 6.92 g of dry ice to

carbon dioxide gas?

Vapor PressureIn a closed system,

# of liquid molecules decreases as they enter the gas phase. Eventually, an equilibrium is reached—constant number of molecules in both phases.

Equilibrium

Vapor PressureAt equilibrium evaporation rate is exactly the same as condensation rate.Vapor pressure = the pressure of the vapor present at equilibrium

Measuring Vapor Pressure

Difference between atmospheric pressure with a vacuum above

and with vapor above.

Effect of IM ForcesLow IM forces lead to high vapor pressure; less attraction allows molecules to vaporize more easilyHigh IM forces lead to low vapor pressure; molecules are tightly held in liquid

Enthalpy (H) of Vaporization

Change in energy when a liquid vaporizesVapor pressure varies with temperature. (Why??)Note: Function is not linear

Linear FunctionGraphing ln of vapor pressure and 1/T gives a linear graph.

Can use slope-intercept form of equation to solve for points.

Line Equation

ln(Pvap) =- vap (1)

+ C

R

(T)

Format is y = mx + bR is the universal gas constant.C is a constant (unique for each

substance)

Clausius-Clapeyron Equation

Since C is temperature independent, by measuring vapor pressure at several different temperatures, heat of vaporization can be calculated.

Problem 2Ethanol’s vapor pressure at 30.o C is 100. torr. At 62o C the vapor pressure has increased to 400. torr. What is H for ethanol?

In Your Scientist’s Notebook

What is ethanol’s vapor pressure at 15o C?

Heating CurveChanges of state are evident on a heating curve:

Heating CurveAs heat is added temperature typically increases.At the plateaus, all energy is going into the phase change, so temperature is constant.

Enthalpy (H) of Fusion

Enthalpy change that occurs at the melting point when a solid melts or freezes

ProblemThe enthalpy of fusion for water is 6.02 kJ per mole. How much energy would be required to change 43.2 g of ice to liquid water?

What is the enthalpy change when 52.8 g of CCl4 freezes?

Enthalpy of fusion is 2.51 kJ/mol.

Boiling PointThe temperature at which the vapor pressure of a liquid is equal to the atmospheric pressureVaries with pressure (Think about a pressure cooker.)

Phase Change Enthalpy

The amount of heat gained or lost—HEndothermic: boiling (evaporating), melting, sublimatingExothermic: condensing, freezing, depositing

Phase Change Entropy

The amount of disorder in a systemIncreasing entropy: boiling, melting, sublimatingDecreasing entropy: condensing, freezing, depositing

Phase DiagramLines show equilibrium boundaries for phase changesEquilibrium—the system shows no net changeDynamic equilibrium—two processes of change occur at exactly the same rate

Typical Phase Diagram

Points of InterestCritical temperature—the temp. above which a gas cannot be liquefied at any pressureCritical pressure—the pressure required to liquefy the gas at critical temperature

Points of InterestTriple point—the point at which solid, liquid & gas have equal vapor pressure and exist in equilibrium“Normal” boiling or melting point—read at 1 atm pressure

What is the enthalpy / entropy change

when…Evaporation occurs?Sublimation occurs?Freezing occurs?Melting occurs?Deposition occurs?

Calculations with Enthalpy

To calculate total energy change, consider each stage in the process.Ice at -10o C to steam at 110o C:

Heating ice from -10o to 0o CMeltingHeating water from 0o to 100o CBoilingHeating steam from 100o to 110o C

Finding the NumbersMelting or Freezing: heat of fusion

(Hfus)

Vaporizing or Condensing: heat of vaporization (Hvap)

Temperature changes: specific heats (q = mcT)Sign depends on whether heat is gained or lost in the process.

ExampleWhat is the total energy required to change 45.0 g of ice at -5o C to water at 50o C?Specific heat of ice: 2.1 J/g oCSpecific heat of water: 4.21 J/g oCHfus: 6.02 kJ/mol

Hvap: 40.7 kJ/mol