1. CHAPTER 10 AND 11
AP CHEMISTRY

2. INTERMOLECULAR FORCES
Molecules tend to be
Nonconductors of electricity when pure substances
Most molecules when placed in water won’t conduct
EXCEPTION- highly polar molecules i.e. HCl
Insoluble in water but soluble in nonpolar solvents (iodine dissolves in alcohol) i.e. toluene, CCl4
A few are very soluble in water i.e. ethyl alcohol(C2H5OH)
Low melting and boiling points
Many are gases at room temperature
Many boiling points are directly related to the strength of the intermolecular forces

3. DIPOLE FORCES Ion-dipole Dipole-dipole What are dipole forces?
Ions attracted to a dipole substance i.e. NaCl in water
Dipole-dipole
What are dipole forces?
Attraction between polar molecules
As the dipole moment increases the boiling point increases
London dispersion forces (Van der Waal)
What is a dispersion force?
Electrical temporary dipoles (POLARIZABILITY)
All molecules have dispersion forces, the strength depend on two factors
Number of electrons in the molecule
Ease with which electrons are dispersed to form temporary dipoles
As the molar mass increases the strength will increase

4. HYDROGEN BONDS Polarity has a small effect on boiling point
Compounds have to have hydrogen attached to nitrogen, oxygen or fluorine
Hydrogen in one compound is attracted to a lone pair on the N, O, or F of a different compound
Two reasons hydrogen bonding is stronger than other dipole forces
Electronegative difference is large between hydrogen(2.2) and nitrogen(3.0), oxygen(3.5), or fluorine(4.0)
Small size of hydrogen allows the unpaired electron pair of F, O, or N atom to approach the H atom closely
The molar mass of H2S is double that of water, both are polar molecules, but the boiling point of water is 167 °C higher than H2S. Why?
Hydrogen bonding
What types of intermolecular forces are present in N2, SiCl4, OCCl2, N2H4

5. PROPERTIES OF LIQUIDS Viscosity Resistance to flow Surface tension
Energy required to increase surface area of liquid by a unit amount
Water increases surface tension because of hydrogen bonding

6. PHASE CHANGES Vapor pressure Pvapor = Patm - PHg column page 460-461
Heat of fusion
Enthalpy change associated with melting
Ice = 6.01 kJ/mol
Heat of vaporization
Heat required to vaporize a liquid
Water = kJ/mol
Calculate the enthalpy change associated with converting 1.00 mol of ice at - 25 oC to water vapor at 125 oC at 1 atm pressure. Heat capacities of ice, water, and steam are 2.09 J/g·C, 4.18 J/g·C, 1.84 J/g·C. The heat of fusion of ice is 6.01 kJ/mol, heat of vaporization is kJ/mol.
1 mol of water has 18 g/mol, use the specific heat formula
ΔH = (18 g/mol)(2.09 J/g·C)( ) = 940J = .94kJ
ΔH = (18 g/mol)(4.18 J/g·C)( ) = 7524J = 7.52 kJ
ΔH = (18 g/mol)(1.84 J/g·C)( ) = 828J = .83kJ
55.97 kJ/mol

7. PROBLEM
A 428 mL sample of hydrogen gas was collected by displacement of water in Salt Lake City a 25 oC at a pressure of 647 mm Hg. The vapor pressure of water at 25 oC is 23.8 mm Hg. Calculate the partial pressure of hydrogen gas and the number of moles of hydrogen gas formed.
PH2 = Patm -PH2O
623 mm Hg
PV = nRT
1.43 X 102- mol H2

8. CRITICAL TEMPERATURE
Highest temperature at which a substance can exist as a liquid
Critical pressure
Pressure required to bring about liquidefaction
Sublimation
When a substance goes from the solid phase to the gaseous phase without going through the liquid phase
Deposition
When a substance goes from the gaseous phase to the solid phase without going through the liquid phase

9. VAPOR PRESSURE Dynamic equilibrium Volatile Phase diagram
Opposing processes are occurring at the same rate
Evaporation and condensation
Volatile
Liquids with high vapor pressure evaporate quickly
Phase diagram
On the overhead

10. STRUCTURE OF SOLIDS Crystalline solid Amorphous solid Crystal Lattice
Geometry shaped to their arrangement of ions
Amorphous solid
No orderly arrangement
They do not melt at a specific temperature
Glass
Crystal Lattice
Unit cell repeated, page 447
If the unit cell of NaCl is 6.54 Å on an edge, calculate the density of NaCl
(6.54 Å)3(10-8cm/1Å)3
2.78 X cm3
8 atoms X 58.5 g/mol X 1mol/6.02 X 1023 atoms
7.78 X g/ 2.78 X cm3
2.76 g/cm3

11. CONTINUE Close packing of ions cause three types of unit cells
Simple Body centered

12. X-RAY DIFFRACTION
Scattering of x-rays by a regular arrangement of atoms or ions
Constructive interference
Higher intensity if waves are the same
Destructive interference
Cancels each other out
Bonding in solids
Table 10.3 page 451
Graphite and diamonds are allotropes of carbon (same number of carbons (60) but in different forms
You can convert graphite into diamonds
Read pages 454 to 470 take notes

13. PROBLEM
Silver crystallizes in a cubic closest packed structure. The radius of a silver atom is 1.44 Å (1 Å = 10-8cm). Calculate the density of solid silver.
This is face-centered so you would have one half (2r) and two fourths (2r). Using Pythagorean theorem d2 + d2 = (4r)2
d= r8
d= (1.44 Å)( 8)
(4.07 Å)3│ (10-8 cm)3
│ ( 1 Å )
6.74 X cm3
(4 atoms)(107.9 g ) (1 mol )
(1 mol )(6.02 X 1023 atom)(6.74 X cm3)
10.6 g/cm3

14. SOLUTION COMPOSTION Molarity = moles of solute L of solution
Percent mass = mass of solute X 100
Mass of solution
Mole fraction = A = XA = nA
nA + nB
nA = moles of solute
nB = moles of solution or other substance
Molality = mole of solute
kg of solvent

15. CONTINUE
A solution is prepared by mixing 1.00 g ethanol(C2H5OH) with g water to give a final volume of 101 mL. Calculate molarity, mass percent, mole fraction and molality of ethanol.
1.00g/ 1 mol
46 g
0.0217mol
0.0217mol/.101L
.215M
1.00 g X 100
100.0g water g
0.990%
0.0217mol mol
.0217mol/.1000kg
.217m

16. NORMALITY Equivalents/L solution Acid-base Oxidation-reduction 5
Equivalent mass of acid or base that can get 1 mole of protons
Equivalent mass of sulfuric acid is ½ it’s molar mass because you get 2 protons
Oxidation-reduction
Equivalent amount of oxidizing or reducing agent to accept or furnish 1 mole of electrons
KMnO4
MnO4- + 5e- + 8H+ ----> Mn2+ + H2O, so MMKMnO4 to find normality 5

17. ENERGY SOLUTION FORMATION
Enthalpy Hsol’n = H1 + H2 + H3
Fig 11.2 page 505 and table 11.3 page 507
Read and take notes
Vapor pressure
Nonvolatile solute lowers the vapor pressure of a solvent
Raoult’s law
Psol’n =Xsolvent·Posolvent
Po = vapor pressure of solvent
X = mole fraction of solvent
Calculate the expected vapor pressure of a solution that contains 215 g urea, NH2CONH2, dissolved in 435g of water (Vp pure water = mmHg)
P = 24.2mole water/( mol urea) X 23.76
20.7mmHg

18. CONTINUE Ptotal = XAPoA + XBPoB
Liquid-liquid solution that obey Raoult’s law is an ideal solution
Table 11.4 page 520

19. COLLIGATIVE PROPERTIES
Nonvolatile solute raises the boiling point and lowers the freezing point
ΔTf = kfm m = mol solute/kg solvent
ΔTb = kbm ΔT = change in temperature
kf or kb = boiling point or freezing pt. constant
Calculate the boiling point of a solution that contains 50.0g of glucose (MM = g/mol) in 400. g of water. The molal constant is 0.52 °C/m
ΔTb = (0.52)(50.0g/ 1 mol / )
( / g / .400 kg )
.36 °C °C

20. CONTINUE
A student dissolves 1.00 g of a new compound in 25.0g cyclohexane. The solution has a freezing point of 2.0 °C and that of cyclohexane is 6.5 °C. Cyclohexane has a kf of 20.2 °C/m. Using this data, calculate the molar mass of the compound.
=20.2°C/m (Xmol of unknown solute)
(0.0250kg of cyclohexane)
mol of unknown solute
1.00g/0.0056mol
180 g/mol unknown

21. OSMOTIC PRESSURE Semipermeable membrane
Solvents can pass through but not the solute molecule
The volume of solution increases and the solvent decreases, this flow is called OSMOSIS
Osmotic pressure
When you have an excess pressure on the solution in comparison to the pure solvent

22. CONTINUE Reverse osmosis Colligative properties of electrolytes
If an external pressure greater than the osmotic pressure is applied the solvent particles will go from the solution to the pure solvent
Colligative properties of electrolytes
Van’t Hoff factor i:
i = moles of particles in solution
Moles of solute dissolved
Colloids
Suspension of tiny particles in some medium
Blood
Coagulation - destroy a colloid - heating or adding an electrolyte (ions)