Unit 11 – Intermolecular Forces/Solids, Liquids and
Solutions
IQ #11) What is bond polarity?
2) What determines the type of bond that exists between two atoms?
3) List the 3 major bonds types and the difference in electronegativity that exists between the atoms.
Definition of IMF
• Attractive forces between molecules.
• Much weaker than chemical bonds within molecules.
• a.k.a. van der Waals forces
Attractions between molecules• They are what make solid and liquid
molecular compounds possible.• The weakest are called van der Waal’s
forces - there are two kinds:1. Dispersion forces
weakest of all, caused by motion of e-
increases as # e- increases- halogens start as gases; bromine is
liquid; iodine is solid – all in Group 7A
Dispersion Forces
• London Dispersion Forces
View animation online.
Dipole interactions
• Occurs when polar molecules are attracted to each other.
2.Dipole interaction happens in water–Figure 8.25, page 240 –positive region of one molecule attracts the negative region of another molecule.
Dipole interactions
• Occur when polar molecules are attracted to each other.
• Slightly stronger than dispersion forces.
• Opposites attract, but not completely hooked like in ionic solids.
H Fδ+ δ-
H Fδ+ δ-
Dipole Interactions
•Dipole-Dipole Forces
+ -
View animation online.
3. Hydrogen bonding• …is the attractive force caused by
hydrogen bonded to N, O, F, or Cl• N, O, F, and Cl are very
electronegative, so this is a very strong dipole.
• The hydrogen partially share with the lone pair in the molecule next to it.
• This is the strongest of the intermolecular forces.
Hydrogen bonding defined:• When a hydrogen atom is:
a) covalently bonded to a highly electronegative atom, AND b) is also weakly bonded to an unshared electron pair of a nearby highly electronegative atom.– The hydrogen is left very electron
deficient, thus it shares with something nearby
– Hydrogen is also the ONLY element with no shielding for its nucleus when involved in a covalent bond!
Hydrogen Bonding
HH
O+ -
+
H HO+-
+
Hydrogen bonding
HH
O H HO
HH
O
H
H
OH
HO
H
HO HH
O
Hydrogen Bonding
Types of IMF
Determining IMF
• NF3
– polar = dispersion, dipole-dipole
• CH4
– nonpolar = dispersion• HF
– H-F bond = dispersion, dipole-dipole, hydrogen bonding
Examples:1. Explain, in terms of intermolecular forces, why(a) the boiling point of O2 (-183oC) is higher than that of N2 (-196oC).
(b) the boiling point of NO is higher than either N2 or O2.
(a) Both O2 & N2 are non-polar molecules it is based on molar mass. As molar mass increases, so does the dispersion force resulting in stronger bonds in turn a higher boiling pt. O2 has a higher BP, because it has a greater molar mass in turn a greater dispersion force.
Ion
ic c
har
acte
r
0
Non – Polar Covalent
0.5
Polar Covalent
1.7
Ionic
4.0
(b) Both O2 & N2 are non-polar molecules, but NO is a polar molecule. NO has stronger intermolecular forces in turn a higher boiling pt.
What types of intermolecular forces are present in H2? CCl4? OCS? NH3?
H2
=dispersionCCl4
=dispersion
SO C••
••
••
••OCS =
NH3
=
NH
HH
••
dispersion, dipole-dipole
dipole-dipole, hydrogen bonding
IQ #2
Attractions and properties•Why are some chemicals gases,
some liquids, some solids?
–Depends on the type of IMF!
–Table 8.4, page 244
Kinetic Molecular Theory
•KMT
–Particles of matter are always in motion.
–The kinetic energy (speed) of these particles increases as temperature increases.
Forces and Phases- Substances with very little
intermolecular attraction exist as gases.
- Substances with strong intermolecular attraction exist as liquids.
- Substances with very strong intermolecular (or ionic) attraction exist as solids.
Phase Differences
SolidSolid – definite volume and shape; particles packed in fixed positions; particles are not free to moveLiquidLiquid – definite volume but indefinite shape; particles close together but not in fixed positions; particles are free to move
GasGas – neither definite volume nor definite shape; particles are at great distances from one another; particles are free to move
Three Phases of Matter
Liquid Properties
• Surface Tension– attractive force between
particles in a liquid that minimizes surface area.
Liquid Properties• Capillary Action
– attractive force between the surface of a liquid and the surface of a solid.
water mercury
Applications:
1) Blood tests (finger)
2) Plants: absorb subsurface ______ with tiny tubes in the _____. This can lift water about a maximum of __ ft. or ___ cm. Plants taller than one foot must use _______ (_____________).3) Paper: ______________4) Sponges, towels, diapers, etc: ____________
waterroots
130
xylem Active transportCellulose
fibersCotton fibers
Viscosity
Definition: The _________ of a liquid to flow.- Examples of viscous liquids: - Cause? The more the molecules ______
each other, the ______ the viscosity.- Effect of temperature: As the temperature
increases, the viscosity _________.Examples: Fudge, syrup, • motor oil (summer: ______ viscosity vs.
winter: ______ viscosity)
resistance
Malasses, oil, & honey
attracthigher
decreases
highlow
The Solid State 1. Types of Solidsa) Crystalline: A solid in which the particles are
arranged in an orderly, ____ repeating pattern.Example: _____
Seven types of crystals: cubic, orthorhombic, tetragonal, monoclinic, triclinic, hexagonal, rhombohedral.
b) Amorphous: Without ______. A non-crystalline solid whose particles are in a ______ arrangement.
Example: _______
3-DNaCl
shaperandom
glass
http://www.emporia.edu/
Phase Changes
Which has a higher m.p.?• polar or nonpolar?• covalent or ionic?
Phase Changes
• Melting Point– equal to freezing point
polarionic
IMF m.p.
Phase Changes
• EvaporationEvaporation– molecules at the surface gain
enough energy to overcome IMF.
• VolatilityVolatility– measure of evaporation rate– depends on temp & IMF.
Phase Changes
• EquilibriumEquilibrium– trapped molecules reach a
balance between evaporation & condensation
Phase Changes
• Vapor PressureVapor Pressure– pressure of vapor above
a liquid at equilibrium
IMF v.p.temp v.p.
•depends on temp & IMF•directly related to
volatility
p.478
temp
v.p
.
Phase Changes
• Boiling Point– temp at which v.p. of liquid
equals external pressure.
IMF b.p.Patm b.p.
•depends on Patm & IMF
•Normal B.P. = b.p. at 1 atm
Think About It!
Example: Which substance would have a higher vapor pressure at 25°C: O
║ or H2O?H3C—C—CH3 (acetone)
Dipole-Dipole
Hydrogen Bonding
Effect of Pressure on Boiling PointBoiling Point of Water at Various Locations
Location Feet above sea
level
Patm (kPa) Boiling Point (C)
Top of Mt. Everest, Tibet
29,028 32 70
Top of Mt. Denali, Alaska
20,320 45.3 79
Top of Mt. Whitney, California
14,494 57.3 85
Leadville, Colorado 10,150 68 89
Top of Mt. Washington, N.H.
6,293 78.6 93
Boulder, Colorado 5,430 81.3 94
Madison, Wisconsin 900 97.3 99
New York City, New York
10 101.3 100
Death Valley, California
-282 102.6 100.3
Think About It! 1) If you place a glass of water in a
bell jar and turn on the vacuum pump, what will happen to boiling point?
2) Can you cook an egg faster if you turn up the flame under a pan of boiling water? Explain.
Patm , B.P.
No, the temperature remains constant at the boiling point. High energy molecules escape, which cools the liquid. Thus, continuing to heat the water just maintains the temperature.
3) Does it take more or less time to boil an egg on Mt. Everest or here in Fullerton? Explain.
4) Does food cook faster in a pan with a lid on it?
Explain. 5) How does a pressure cooker work?
More time. The atmospheric pressure on Mt. Everest is only 240 mmHg. , the water boils at 70 oC, and the food would take longer to cook at the lower temperature.
Yes. The lid traps the high energy molecules, which keeps the heat from escaping.
The pressure cooker increases the pressure, which increases the boiling point of water to ~ 150 oC –200 oC. , more heat-faster cooking.
IQ #31) What is primarily responsible in
determining the state of a compound or element? Explain.
2) Define: surface tension, capillary action, and viscosity.
3) Explain in terms of intermolecular forces why: (a) NaCl has a higher melting point than Br2.
(b) C2H6 has a higher boiling point than CH4
IQ #3 cont.
4) Define: Volatility, Boiling Point, Vapor Pressure, Melting Point.
5) What relationship does IMF have with all of these?
6) Will increasing the elevation lower or raise your boiling point?
Solution Chemistry- Definitions
Solution - - homogeneous mixture
Solvent - present in greater amount
Solute - substance being dissolved
Definitions
Solute Solute - KMnO4 Solvent Solvent - H2O
Concentrated vs. Dilute
Solvents
Solvents at the hardware store
Solvation
Solvation – – the process of dissolving
solute particles are separated and pulled into solution
solute particles are surrounded by solvent particles
Dissolution of sodium Chloride
Solvation
NONPOLAR
NONPOLAR
POLAR
POLAR
““Like Dissolves LikeLike Dissolves Like””““Like Dissolves LikeLike Dissolves Like””
Solvation
• Soap/DetergentSoap/Detergent– polar “head” with long nonpolar “tail”– dissolves nonpolar grease in polar water
Solubility
•Solubility–maximum grams of solute that will dissolve in 100 g of solvent at a given temperature
–varies with temp
Saturation of Solutions• A solution that contains the maximum
amount of solute that may be dissolved under existing conditions is saturated.
• A solution that contains less solute than a saturated solution under existing conditions is unsaturated.
• A solution that contains more dissolved solute than a saturated solution under the same conditions is supersaturated.
Solubility
SATURATED SOLUTIONno more solute
dissolves
UNSATURATED SOLUTION
more solute dissolves
SUPERSATURATED SOLUTION
becomes unstable,
crystals form
concentration
Solubility• Solids are more soluble at...Solids are more soluble at...
– high temperatures.– Increasing surface area of the solid
• Gases are more soluble Gases are more soluble at...at...
– low temperatures &– high pressures (Henry’s
Law).– EX: nitrogen narcosis,
the “bends,” soda
Therefore…Solids tend to dissolve best when:
• Heated• Stirred• Ground into small particles
Liquids tend to dissolve best when:• The solution is cold
• Pressure is high
Solubility Chart
Calculations of Solution Concentration
Concentration - A measure of the amount of solute in a given amount of solvent or solution
Molality - moles of solute divided by the mass of solvent in kilograms
Parts per million – the ratio of parts (mass) of solute to one million parts (mass) of solution
Grams per liter - the mass of solute divided by the volume of solution, in liters
Molarity - moles of solute divided by the volume of solution in liters
Molarity• Concentration of a solution.
solution of liters
solute of moles(M)Molarity
total combined volume
substance being dissolved
Molarity
2M HCl
L
molM
L 1
HCl mol 2HCl 2M
What does this mean?
Molarity Calculations
molar mass
(g/mol)
6.02 1023
(particles/mol)
MASS
IN
GRAMS
MOLESNUMBER
OF
PARTICLES
LITERSOF
SOLUTION
Molarity(mol/L)
Sample Molarity Calculations
1. How many grams of NaCl are required to make 0.500L of 0.25M NaCl?
0.500 L
0.25 mol
1 L
= 7.3 g NaCl
58.44 g
1 mol
L 1
mol0.25 0.25M
Molarity Calculations2. Find the molarity of a 250 mL
solution containing 10.0 g of NaF.
10.0 g 1 mol
41.99 g = 0.238 mol NaF
0.238 mol
0.25 L
M =
= 0.95M NaF
L
molM