States of Matter; Liquids and Solids

1
Chapter 13

• States of Matter Liquids and Solids

2
Phase changes

• Section 13.1

3
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

4
Energy Requirements for Phase Changes
Water (solid) water (liquid) 6
kJ/mol Water (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
5
Section 13.1 Water and its phase changes

• FYI!!!
• water is unique in that as it cools it EXPANDS
• WHY???
• Explain why ice floats?
• Why pipes burst in the Winter?
• How potholes form?

6
Phase Transitions
H2O(s) ? H2O(l) H2O(l) ? H2O(s) H2O(l) ? H2O(g)
H2O(s) ? H2O(g) H2O(g) ? H2O(l) H2O(g) ?
H2O(s)

• Melting 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.

7
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!!!!

8
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.

9
Boiling point vs vapor pressure

• 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

10
Vapor Pressure

• If 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.

11
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
  _______

12
Phase Diagrams

• Section 13.1

13
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.

14
(No Transcript)
15
Phase Diagram
16
(No Transcript)
17
Phase Diagram of Water
18
Endothermic vs Exothermic Reactions

• Section 3.6

19
Temperature, energy and heat

• Temperature 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.

20
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.

21
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 enzymes
  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.

22
Calculating Energy Changes

• Section 13.2 and 3.7

23
Calculating energy changes

• 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) DH 40.6 kJ/mol
• Heat of fusion heat (energy) needed for the
  melting of a I mol of a solid substance.
• H2O(s) ?H2O(l) DH 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)

24
Molar heat of fusion and vaporization
25
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/mol
Step 1 Q s m
?t
4.18J x 25g x
(100C 25C) g 7.8x103J
Convert to kJ 7.8kJ Step
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
water 25g H2O x 1 mol H2O 1.4 mol H2O
18 g H2O Now calcualte energy
need to vaporize the water 40.6 kJ x 1.4
mol H2O 57kJ mol Step 3
TOTAL THE ENEGY NEEDED 7.8kJ 57 kJ
65kJ
26
FYI

• One calorie is the amount of heat (or energy)
  needed to raise the temperature of 1 gram of
  water by 1C.
• 4.184 joules one calorie.
• 1,000 calories 1 Calorie

27
Intermolecular Forces

• Section 13.3

28
Intra vs. inter molecular forces

• Intramolecular Forces
• The attractive forces between atoms and ions
  within a molecule
• Ex ionic bonds, covalent bonds
• STRONG
• Intermolecular Forces
• The attractive forces between molecules
• Ex London dispersion forces , dipole-dipole
  forces and hydrogen bonding
• WEAK
• I.e. much less energy to melt H2O (inter) than
  for it to decompose into H2 and O2 (intra)

29
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
30
London dispersion forces

• 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

31
Intermolecular Forces

• London Dispersion ForcesInduced 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).

32
Dipole-dipole

• Forces of attraction between oppositely charge
  ends of polar molecules
• Line up and - ends

33
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

34
Relative strength

• London Dispersion lt Dipole-Dipole lt H-bonds

35
Determining Intermolecular forces

• STEP ONE

• STEP TWO

• Does the compound contain N-H, O-H, or F-H
  Bonds???
• NO dipole-dipole and london dispersion forces
• YES hydrogen bonding, dipole-dipole and london
  dispersion forces

• Is the compound polar or nonpolar???
• Nonpolar ONLY london dispersion forces
• Polar go to step 2

36
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.
37
Determining Relative Boiling points

• Step ONE

• STEP TWO

• Determine the types of intermolecular forces
• Different forces
• H gt D gt L
• Stronger force
• higher boiling point
• b. Same forces go to step 2

• Look at the size of the compounds
• Larger molecules higher boiling point

38
THE END