CHAPTER 1

3

STATE

S OF

MATTE

R; LIQ

UIDS A

ND SOLID

S

PHASE C

HANGES

SECTION 1

3.1

COMPARISON OF GASES, LIQUIDS AND SOLIDS Gases are compressible fluids. Their molecules are widely

separated. Liquids are relatively incompressible fluids. Their molecules are

more tightly packed. Solids are nearly incompressible and rigid. Their molecules or

ions are in close contact and do not move. Vapors term customarily used for the gasesous state of a

substance that exists naturally as a solid or liquid at 25 C and 1 atmosphere

States of Matter

ENERGY REQUIREMENTS FOR PHASE CHANGES

Water (solid) water (liquid) = 6 kJ/molWater (liquid) water (gas) = 41 kJ/mol

What does this mean????Takes more energy to convert water from a liquid to a gas then to convert water from a solid to a liquid.

WHY??Solid and liquid states of water are more similar

then the liquid and gas states

SECTION 13.1 WATER AND ITS PHASE CHANGES

FYI!!!WATER IS UNIQUE IN THAT AS IT

COOLS IT EXPANDSWHY???

EXPLAIN WHY ICE FLOATS?WHY PIPES BURST IN THE WINTER?

HOW POTHOLES FORM?

PHASE TRANSITIONSMelting: change of a solid to a

liquid.

Freezing: change a liquid to a solid.

Vaporization: change of a solid or liquid to a gas. Change of solid to vapor often called sublimation.

Condensation: change of a gas to a liquid or solid. Change of a gas to a solid often called deposition.

H2O(s) H2O(l)

H2O(l) H2O(s)

H2O(l) H2O(g)

H2O(s) H2O(g)

H2O(g) H2O(l) H2O(g) H2O(s)

ENERGY OF HEAT AND PHASE CHANGE

Temperature does not change during the change from one phase to another.

notice: there is either a temperature change OR a phase change.

You cannot have BOTH at the same time

Increase in temperature = increase in kinetic energy (after all the definition of temperature is average kinetic energy). There is no change in the potential energy

Increase in heat = increase in potential energy. There is no change in the kinetic energy because the temperature does not change!!!!

VAPORIZATION VS BOILING VS EVAPORATION

Boiling For water to boil, it must be 100 Celsuis. Boiling creates an actual gas. The substance changes phase.

Evaporation also involves liquids become gaseous. However, the body of liquid does not need to be at the boiling temperature. It occurs because the molecules of a liquid are not tightly bound together, and so some escape with time.

Vaporization is a blanket term referring to both boiling and evaporation. In the broadest sense, it is liquid becoming gas.

Boiling point the temperature at which the vapor pressure of a liquid is equal to the pressure of the external atmosphere.

Normal boiling point the temperature at which the vapor pressure of a liquid is equal to atmospheric pressure (1 atm).

Vapor Pressure the pressure of the vapor over a liquid at equilibrium in a closed container

BOILING POINT VS VAPOR PRESSURE

VAPOR PRESSUREIf a liquid is placed in a non-closed container, some of the molecules will escape the surface and evaporate.

In a sealed container, some of a liquid still evaporates but cannot “leave” the container. The molecules move back and forth between liquid and gas phases until they establish an equilibrium and thus a pressure in the vapor phase.

Vapor pressure: partial pressure of the vapor over the liquid measured at equilibrium and at some temperature.

TEMPERATURE DEPENDENCE OF VAPOR PRESSURES

The vapor pressure above the liquid varies exponentially with changes in the temperature.

What does this graph indicate about the relationship between vapor pressure and temperature??

There is a _______ relationship between vapor pressure and temperature. In other words as vapor goes UP the temperature goes _______ or as vapor pressure goes DOWN the temperature goes _______

PHASE D

IAGRAMS

SECTION 1

3.1

PHASE DIAGRAMS

Graph of pressure-temperature relationship; describes when 1,2,3 or more phases are present and/or in equilibrium with each other.

Lines indicate equilibrium state two phases.

Triple point- Temp. and press. where all three phases co-exist in equilibrium.

Critical temp.- Temp. where substance must always be gas, no matter what pressure.

• Critical pressure- vapor pressure at critical temp.• Critical point- point where system is at its critical pressure and

temp.

PHASE DIAGRAM

Phase Diagram of Water

ENDOTHERMIC

VS

EXOTH

ERMIC R

EACTIONS

SECTION 3

.6

TEMPERATURE, ENERGY AND HEATTemperature = average kinetic energy

Energy – the ability to do work

Heat – the transfer of energy

Energy and Reactions

Energy must be added to break bonds. Many forms of energy can be used to break bonds:

heat - electricity sound - light

Forming bonds releases energy. Example: When gasoline burns, energy in the form of heat and light is

released as the products of the isooctane-oxygen reaction and other gasoline reactions form.

Energy is conserved in chemical reactions. Chemical energy is the energy released when a chemical compound

reacts to produce new compounds. The total energy that exists before the reaction is equal to the total

energy of the products and their surroundings.

EXOTHERMIC VS ENDOTHERMIC REACTIONS

An exothermic reaction is a chemical reaction in which heat is released to the surroundings.

An endothermic reaction is a chemical reaction that absorbs heat.

The graphs to the right represent the changes in chemical energy for an exothermic reaction and an endothermic reaction.

CATALYSTS

Catalysts provide alternative pathways for a reaction, usually with a lower activation energy. With this lower energy threshold, more collisions will have enough energy to result in a reaction. An enzyme is a large organic molecule that folds into a unique shape by forming intermolecular bonds with itself. The enzyme’s shape allows it to hold a substrate molecule in the proper orientation to result in an effective collision. The rate of a chemical reaction is the change in the amount of reactants or products in a specific period of time. Increasing the probability or effectiveness of the collisions between the particles increases the rate of the reaction.

CALCULA

TING E

NERGY

CHANGES

SECTION 1

3.2 A

ND 3.7

Specific Heat Capacity: the amount of energy required to change the temperature of 1 gram of a substance 1 celsius degree (in j/g ◦C)Heat of vaporization: heat (energy) needed for the vaporization of a 1 mol of a liquid.

H2O(l) H2O(g) H = 40.6 kJ/molHeat of fusion: heat (energy) needed for the melting of a I mol of a solid substance.

H2O(s) H2O(l) H = 6.02 kJ/mol

EQUATION

Q = sm∆t

Q= energy required (joules)

s = specific heat capacity (given in table

page 70 in book)

M = mass of water (grams)

∆t = change in temperature (Celcius)

CALCULATING ENERGY CHANGES

MOLAR HEAT OF FUSION AND VAPORIZATION

Example: calculate the energy (in kJ) required to heat 25g of water from 25C to 100C and change it to steam at 100C. The specific heat capacity of water is 4.18J/g and the molar heat of vaporization of water is 40.6kJ/molStep 1: Q = s m ∆t = 4.18J x 25g x (100C – 25C) g= 7.8x103J Convert to kJ 7.8kJStep 2: now use the molar heat of vaporazation to calculate heat energy required to vaporize 25g of water at 100C. Heat of vaporation is given in mols to we must convert 25g water to mols of water25g H2O x 1 mol H2O = 1.4 mol H2O 18 g H2ONow calcualte energy need to vaporize the water

40.6 kJ x 1.4 mol H2O = 57kJ molStep 3: TOTAL THE ENEGY NEEDED: 7.8kJ + 57 kJ = 65kJ

FYI

One calorie is the amount of heat (or energy) needed to raise the temperature of 1 gram of water by 1°C.

4.184 joules = one calorie.

1,000 calories = 1 Calorie

INTE

RMOLECULA

R

FORCES

SECTION 1

3.3

INTRA VS. INTER MOLECULAR FORCES

Intramolecular Forces:• The attractive forces between atoms

and ions within a molecule• Ex: ionic bonds, covalent bonds• STRONGIntermolecular Forces• The attractive forces between

molecules• Ex: London dispersion forces , dipole-dipole forces

and hydrogen bonding

• WEAKI.e. much less energy to melt H2O (inter) than for it to decompose into H2 and O2 (intra)

COMPARISON OF ENERGIES FOR INTERMOLECULAR FORCES

Interaction Forces Approximate Energy within Bond or Attraction between molecules

Intermolecular

London 1 – 10 kJ

Dipole-dipole 3 – 4 kJ

Hydrogen bonding 10– 40 kJ

Chemical bonding

Ionic 100 – 1000 kJ

Covalent 100 – 1000 kJ

•Aka: van der waals forces

•Attractive forces between ALL molecules

•temporary attractive force that results when the electrons in two adjacent atoms occupy positions that make the atoms form temporary dipoles

•Exist with nobles gases and nonpolar molecules, H2, N2, I2

LONDON DISPERSION FORCES

Intermolecular ForcesIntermolecular Forces

London Dispersion ForcesLondon Dispersion ForcesInduced Dipole – Induced DipoleInduced Dipole – Induced Dipole

- Weakest of all intermolecular forces.- It is possible for two adjacent nonpolar molecules to

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

electrons of the adjacent molecule (or atom).- This attraction causes the electron clouds become

distorted.- In that instant a polar molecule (dipole) is formed

(called an instantaneous dipole).

DIPOLE-DIPOLE

• Forces of attraction between oppositely charge ends of polar molecules

• Line up + and - ends

HYDROGEN BONDING• Strong/Special type of dipole-dipole

force• IT IS NOT A BOND.• Between the positive H atom attached

to an N, O or F and the negative N, O, or F of another molecule

RELATIVE STRENGTH

London Dispersion < Dipole-Dipole < H-bonds

DETERMINING INTERMOLECULAR FORCESSTEP ONE

Is the compound polar or nonpolar???a.Nonpolar = ONLY london dispersion forcesb.Polar = go to step 2

STEP TWODoes the compound contain N-H, O-H, or F-H Bonds???

a.NO = dipole-dipole and london dispersion forces

b.YES = hydrogen bonding, dipole-dipole and london dispersion forces

EXAMPLES: WHAT TYPE(S) OF INTERMOLECULAR FORCES EXIST BETWEEN EACH OF THE FOLLOWING ?

HBr???HBr is a polar molecule: dipole-dipole forces. There are also london dispersion forces between HBr molecules.

CH4??? CH4 is nonpolar: london dispersion forces.

SO2??? SO2 is a polar molecule: dipole-dipole forces. There are also london dispersion forces between SO2 molecules.

DETERMINING RELATIVE BOILING POINTS

STEP ONEDetermine the types of intermolecular forces

a.Different forces =

H > D > L

Stronger force =

higher boiling point

b. Same forces = go to step 2

STEP TWOa. Look at the size

of the compounds

Larger molecules = higher boiling

point

THE E

ND