final exam wednesday,december 11, at 10:15 a.m. – 12:15 p.m. in the ic building, room 421

FINAL EXAM

Wednesday,December 11, at 10:15 a.m. – 12:15 p.m. in the IC

building, Room 421

Prentice Hall © 2003 Chapter 11

• The forces holding solids and liquids together are called intermolecular forces.

• The covalent bond holding a molecule together is an intramolecular forces.

• The attraction between molecules is an intermolecular force.

• Intermolecular forces are much weaker than intramolecular forces

• When a substance melts or boils the intermolecular forces are broken (not the covalent bonds).

Intermolecular ForcesIntermolecular Forces


Ion-Dipole Forces• Interaction between an ion and a dipole (e.g. water).• Strongest of all intermolecular forces.



Dipole-Dipole Forces• Exist between neutral polar molecules.• Polar molecules need to be close together.• Weaker than ion-dipole forces.• There is a mix of attractive and repulsive dipole-dipole

forces as the molecules tumble.• If two molecules have about the same mass and size, then

dipole-dipole forces increase with increasing polarity.


• Sublimation: solid gas.• Vaporization: liquid gas.• Melting or fusion: solid liquid.• Deposition: gas solid.• Condensation: gas liquid.• Freezing: liquid solid.

Phase ChangesPhase Changes

.

• ENERGY ASSOCIATED WITH HEATING CURVES

Topics

• Vapor Pressure

• Normal Boiling Point

• Normal Freezing

• Specific Heat

• Enthalpy (Heat) of Vaporization

• Enthalpy (Heat) of Fusion

Vapor Pressure

• THE PRESSURE OF A VAPOR IN EQUILIBRIUM WITH ITS LIQUID (OR ITS SOLID)

NORMAL BOILING POINT & FREEZING POINTS

• NORMAL BOILING PT. - THE TEMPERATURE @WHICH VAPOR PRESSURE = 1 atm

• NORMAL FREEZING PT. – THE TEMPERATURE @ WHICH THE VAPOR PRESSURE OF THE SOLID AND THE LIQUID ARE THE SAME

Heat Capacity aka Specific Heat (C)

• Specific Heat (C) = the amount of energy required to raise the temperature of 1 gram of substance 1 degree celcius

Specific Heat (C) aka Heat Capacity

• Units for: specific heat (C) = J/g-oC

where J = joulesoC = temperature in oC

g = mass in grams

Specific Heat (C) Values(aka Heat Capacity)

• Example: Water

• LIQUID: CLiq = 4.18 J/ (oC . g)

• LIQUID: Csol = 2.09 J/ (oC . g)

• LIQUID: Cgas = 1.84 J/ (oC . g)

Use of Specific Heat

• q = mCT

• q = gm substance x specific heat x T

• where:

• M = mass of substance in grams

• q = amount of heat (energy)

• C = specific heat

• And T = change in temperature

Enthalpy of Vaporizationaka heat of vaporization (Hvap)

• Is the amount of heat needed to convert a liquid to a vapor at its normal boiling point

Enthalpy of Fusion aka heat of fusion (Hfus)

• Is the amount of heat needed to convert a solid to a liquid at its normal melting (freezing) point

Units for Hvap, Hfus and heat(q)

Heat of fusion Hfus = kJ/mol

Heat of vaporization Hvap = kJ/mol

Heat (q) = Joules

Therefore:

• To come up with Joules which is the unit of heat, if:

H is given, then:qvap = Hvap x moles

and

qfus = Hfus x moles

(2) Specific heat (C) is given, then: q = mCT

Sample Problem

• Calculate the enthalpy change upon converting 1 mole of ice at -25 oC to steam at 125 oC under a constant pressure of 1 atm? The specific heats are of ice, water and steam 2.09 J/g-K for ice, 4.18 J/g-K for water and 1.84 J/g-K for steam. For water, Hfus= 6.01 kJ/mol, and Hvap = 40.67kJ/mol.

• Note: The total enthalpy change is the sum of the changes of the individual steps.

Energy Changes Accompanying Phase Changes

• All phase changes are possible under the right conditions.• The sequence

heat solid melt heat liquid boil heat gas

is endothermic.• The sequence

cool gas condense cool liquid freeze cool solid

is exothermic.


Heating Curves• Plot of temperature change versus heat added is a heating

curve.• During a phase change, adding heat causes no

temperature change.• Supercooling: When a liquid is cooled below its melting

point and it still remains a liquid.• Achieved by keeping the temperature low and increasing

kinetic energy to break intermolecular forces.


HEATING CURVES

• ENERGY ASSOCIATED WITH HEATING CURVES

• During a phase change, adding heat causes no temperature change.

Critical Temperature and Pressure• Gases liquefied by increasing pressure at some

temperature.• Critical temperature: the minimum temperature for

liquefaction of a gas using pressure.• Critical pressure: pressure required for liquefaction.


• Phase diagram: plot of pressure vs. Temperature summarizing all equilibria between phases.

• Given a temperature and pressure, phase diagrams tell us which phase will exist.

• Any temperature and pressure combination not on a curve represents a single phase.

Phase DiagramsPhase Diagrams

• Features of a phase diagram:– Triple point: temperature and pressure at which all three

phases are in equilibrium.

– Vapor-pressure curve: generally as pressure increases, temperature increases.

– Critical point: critical temperature and pressure for the gas.

– Melting point curve: as pressure increases, the solid phase is favored if the solid is more dense than the liquid.

– Normal melting point: melting point at 1 atm.


The Phase Diagrams of H2O and CO2


3 Things Learned

• Reading a phase diagram

• Determining triple point on phase diagram

• Determining critical point on phase diagram


• Water:– The melting point curve slopes to the left because ice is less

dense than water.

– Triple point occurs at 0.0098C and 4.58 mmHg.

– Normal melting (freezing) point is 0C.

– Normal boiling point is 100C.

– Critical point is 374C and 218 atm.



• Carbon Dioxide:– Triple point occurs at -56.4C and 5.11 atm.

– Normal sublimation point is -78.5C. (At 1 atm CO2 sublimes it does not melt.)

– Critical point occurs at 31.1C and 73 atm.



• The forces holding solids and liquids together are called intermolecular forces.

• The covalent bond holding a molecule together is an intramolecular forces.

• The attraction between molecules is an intermolecular force.

• Intermolecular forces are much weaker than intramolecular forces

• When a substance melts or boils the intermolecular forces are broken (not the covalent bonds).



Ion-Dipole Forces• Interaction between an ion and a dipole (e.g. water).• Strongest of all intermolecular forces.



Dipole-Dipole Forces• Exist between neutral polar molecules.• Polar molecules need to be close together.• Weaker than ion-dipole forces.• There is a mix of attractive and repulsive dipole-dipole

forces as the molecules tumble.• If two molecules have about the same mass and size, then

dipole-dipole forces increase with increasing polarity.



Dipole-Dipole Forces



London Dispersion Forces• Weakest of all intermolecular forces.• It is possible for two adjacent neutral molecules to affect

each other.• The nucleus of one molecule (or atom) attracts the

electrons of the adjacent molecule (or atom).• For an instant, the electron clouds become distorted.• In that instant a dipole is formed (called an instantaneous

dipole).



London Dispersion Forces• Polarizability is the ease with which an electron cloud can be

deformed.

• The larger the molecule (the greater the number of electrons) the more polarizable.

• London dispersion forces increase as molecular weight increases.

• London dispersion forces exist between all molecules.

• London dispersion forces depend on the shape of the molecule.



London Dispersion Forces• The greater the surface area available for contact, the

greater the dispersion forces.• London dispersion forces between spherical molecules

are lower than between sausage-like molecules.



London Dispersion Forces



Hydrogen Bonding• Special case of dipole-dipole forces.• By experiments: boiling points of compounds with H-F,

H-O, and H-N bonds are abnormally high.• Intermolecular forces are abnormally strong.



Hydrogen Bonding

Problems Chapter 11

2, 3, 15-19, 27, 49, 51, 53, 55, 67, 85, 87, 89, 90, 101, 102, 104

final exam wednesday,december 11, at 10:15 a.m. – 12:15 p.m. in the ic building, room 421

Documents

intramolecular forces

repulsive dipoledipole

iondipole forces interaction

prentice hall

instantaneous dipole

phase changes slide

hydrogen bonding slide

spherical molecules