Unit 11 – Intermolecular Forces/Solids, Liquids and Solutions

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