INTERMOLECULAR FORCES• Intermolecular Forces are the forces of attraction between

molecules.

• They are the GLUE

that holds separated molecules together in a

liquid or solid state.

• These forces determine

whether something is a

solid, liquid, or gas at

room temperature.

• Intermolecular forces include hydrogen bonding, dipole-dipole attraction, and London dispersion (van der Waals) forces.

• They are much weaker than the forces that hold atoms together within a molecule (intramolecular forces = covalent, ionic, & metallic Bonding)

Intermolecular forces all depend on the PARTIAL CHARGES.

• Polar bonds form between elements with very different electronegativities.

• Non-polar bonds form between elements with similar electronegativities.

Difference in Electronegativity 0 .4 = nonpolar covalent bond 0.5 < 2.0 = polar covalent bond > 2.0 = ionic bond

Bond PolarityC. Johannesson

Nonpolar

Polar

Ionic

Polar BOND = Unequal Sharing between atoms creates a partial positive (+) & a partial negative (-).

Molecular PolarityPolar Molecule = One end of the whole

molecule is slightly negative, and one end is slightly positive. Polar molecules result when electrons are distributed unequally.

A polar molecule has two poles, so it is called a dipolar molecule. It has a dipole moment.

H Cl+ -

Nonpolar Molecules = No Permanent Dipole

Nonpolar Molecule: If a molecule has• all nonpolar bonds• polar bonds that cancel each other

out

Symmetrical shapes = Tetrahedral, trigonal planar, & Linear w/ 3 atoms

BF3

Polar Molecules• Polar Molecules

• Dipole moments are asymmetrical and don’t cancel .

netdipolemoment

http://www.youtube.com/watch?v=aPVwOtDEp5k

Determining Molecular Polarity• Therefore, polar molecules have...

• Polar bonds and an asymmetrical shape (lone pairs) • Asymmetrical Shapes = Bent & Pyramidal

CHCl3

H

Cl ClCl

netdipolemoment

Intermolecular Forces

•Forces that attract molecules to other molecules. These include:•Hydrogen Bonding (strongest)•Dipole-dipole attraction•London Dispersion (van der Waals) Forces (weakest)

Hydrogen BondingHydrogen Bonding

Bonds between hydrogen and N, O, F, or Cl have particularly large differences in electronegativities (=very polar bonds), so the forces of attraction between these molecules is particularly strong.H-Bonding in Water:

Hydrogen Bonding in DNAHydrogen Bonding in DNA

Base pairing in DNA by hydrogen bonding

• This is because F, O, Cl, and N strongly attract the electrons that they share with Hydrogen.

• These bonds are VERY POLAR.• The large partial charges result in an unusually strong

dipole-dipole attraction.

Dipole-Dipole AttractionDipole-Dipole Attraction

Attraction between oppositely charged regions of neighboring polar molecules.

*****Add this to your notes*********LD Forces•Caused by e- distributions becoming asymmetrical “instantaneous dipole”, which induces a dipole in a neighboring molecule.

F2

London (Dispersion) Forces The weakest of intermolecular forces, these forces are proportional to the size of the molecule These are the only forces of attraction between completely nonpolar molecules.

Large nonpolar molecules may have substantial dispersion forces, resulting in relatively high boiling points Small nonpolar molecules have weak dispersion forces and exist almost exclusively as gases

London Forces in HydrocarbonsLondon Forces in Hydrocarbons

B. Liquid Properties

•Surface Tension•attractive force between particles in a liquid that minimizes surface area

If molecules have stronger intermolecular forces, than they have

• Higher boiling points• Higher melting points• Higher heats of fusion & vaporization • i.e. they are solids or liquids at room temperature. • Higher surface tension

Type of Intermolecular Forces

I.D. whether nonpolar or polar molecule (VSEPR Sturcture)

All Bonds are Nonpolar

Nonpolar Molecule

London Dispersion

Has Polar Bonds

*Tetrahedral* Linear *Trigonal planar(Symmetrical)

NOT<-----

(Asymmetrical)

Polar Molecule

All Bonds have Equal Polarity(Same Diff. in electroneg.)

Nonpolar Molecule

London Dispersion

Bonds Do NOT have Equal

Polarity

Has H-N; H-O; H-F; or H-Cl

Does Not

H-Bond Dipole-Dipole

Look for H-bonds

Determining the Intermolecular Forces

• NCl3• Polar molecule = dispersion, dipole-dipole

• CH4

• nonpolar molecule= dispersion• HF

• Polar molecule = dispersion, dipole-dipole, hydrogen bonding

LIQUIDS & SOLUBILITY

The Nature of Liquids• Particles in a liquid are attracted to each other• Disruptive forces and attractions between particles determine the properties of each liquid.

• Definite Volume• Indefinite Shape• Much denser than gases• Evaporation – Liquid gas w/o boiling• Vaporization – liquid vapor (gas)

• Particles w/ enough kinetic energy escape• Add heat = more particles have enough energy to escape

• Vapor Pressure = force exerted by a gas above a liquid in a sealed container.• Dynamic Equilibrium rate of evaporation = rate of condensation• T↑VP↑ T↑VP↑• Volatile = how easily a liquid evaporates.

• Boiling Point – • Vapor pressure of liquid = external pressure ( LG)• BP = Dec. at Higher altitudes b/c Pressure dec.• Normal BP = BP at 1 atm or 101.3KPa

(Standard Pressure)

Effect of Pressure on Boiling Point

Ch. 13.4 – Changes of StateStates of Matter• Solid, Liquid , Gas (Vapor)• Plasma = high temp., atoms torn apart • Ex) stars, sun, fluorescent lights• Freezing, Melting (S –L)• Condensation ,Vaporization (L-G)• Sublimation = SG

• Occurs when VP exceeds Atmospheric Pressure

• Phase Diagram• Triple Point = Set of conditions at which all three

phases exist

1

Pressure

(atm)

Temperature (C) 0 100

Represents phases as a function of temperature and pressure.

Phase changes by NamePhase changes by Name

Critical temperature: temperature above which the vapor can not be liquefied.Critical pressure: pressure required to liquefy AT the critical temperature.

Carbon dioxideCarbon dioxidePhase Phase

Diagram for Diagram for CarbonCarbondioxidedioxide

CarbonCarbonPhase Phase

Diagram for Diagram for CarbonCarbon

Phase Diagram for SulfurPhase Diagram for Sulfur

PHASE CHANGE CALCULATIONS

• Calorimetry = measurement of the heat change during a chemical process (done in a calorimeter).

• Energy is measured in Joules (SI) or Calories• Calorie = the amount of energy required to heat 1 gram of water 1 C.

• Heat required to raise the temperature of any substance

• Heat (energy) = mCT

= (mass )(specific heat)(change in Temp)

• Specific heat of water = 4.18 J/gC or 1 cal/gC

• Heat (energy) required to melt a substance = S L phase change.

• Heat = m (Hfus)

• = (mass)(Heat of Fusion )(from a table)

• Heat (energy) required to vaporize a substance =L G phase change.

• Heat = m (Hvap)

• = (mass)(Heat of Vaporization)(from a table)

 • When cooling, heat is lost rather than absorbed but you

use the same formulas to calculate it.

• Hvap = - Hcond & Hfus

= -Hmelting

Heating Curve of Water

Energy

Te

mp

era

ture

(C

)

120

100

80

60

40

20

0

-20

Temperature remains __________ during a phase change.

Ex: Given the following values, calculate the amount of heat required to convert 25.0 g of ice initially at –15.0 C to steam at 125 C.

• Hfus _334 J/g Hvap

_2260 J/g_ Cice _2.09 J/gC_ Cliquid _4.18 J/gC Cgas _2.01 J/gC

1.Calculate the heat to raise temp of ice from –15.0 C to its melting point of 0 C.

2.Use the Hfus to calculate the amount of heat necessary to melt 25.0 g of ice.

 

3. Calc. the heat required to raise the temp. of water from 0C to its boiling point (100 C).

4. Use the Hvap to calculate the amount of heat energy necessary to vaporize 25.0 g of water.

 

 

5. Calc. the heat required to heat steam from 100C -125C.

 

6. Add the heat from each step to find the total heat needed.

784 J + 8350 J + 10500 J + 56500 J + 1260 J = 77400 J or 77.4 kJ

Ch. 15.2&15.3: Mixtures

• Homogeneous = (same) uniform• Solutions = homogeneous mixture, small particles that do not settle out

• Negative Tyndall Effect, particles cannot be seen distinctly

• *made of solute (that which is dissolved) and a solvent (the dissolving medium)

• *EX: salt water

• Heterogeneous = (different) not uniform• Colloid = heterogeneous mixture, medium particles that do not settle out, + Tyndall Effect Ex: milk

• Suspension = heterogeneous mixture, may appear uniform while being stirred but separates in different phases Ex: muddy water, thick tea

• Tyndall Effect = distinguishes between colloids and solutions. Colloid particles are large enough that they will reflect light when it is shined though. Solution particles are too small to reflect light. You see a beam of light through a colloid.

Electrolyte VS Nonelectrolyte• Nonelectrolyte = does not conduct electricity

(molecular compounds)• Weak electrolyte =conducts electricity poorly• Strong electrolyte = conducts electricity well

(ionic compounds – separate ions can carry e-)

Solutions & Solubility

• Solubility = the amount of a substance that dissolves in a solvent at a particular temperature and pressure

• Miscible = liquids or gases that will dissolve in each other. • Ex.) antifreeze and water; vinegar and water; olive

oil and vegetable oil; CO2 and water – club soda

• Immiscible = liquid or gases that will not dissolve in each other. • Ex.) vegetable oil and water; I2 and water

Water dissolves ionic compounds

Attack, I’ll use my negative end to get the positive ions!

We got this negative ion surrounded by our positive ends, I feel it easing away from its buddies!!

http://www.youtube.com/watch?v=7PHhBBg-6X0&NR=1

Immiscible

Oil

H H

O

H H

O

H H

O

H H

O

H H

O

H H

O

H H

O

Oil

Oil

Oil

Oil

Oil

Oil

Oil

Oil

Oil

Hey water, let me in, I’m not attracted to other oils

Forget it, oil! I don’t dislike you, but I’m seriously attracted to other water molecules.

The Polar Clique

http://www.youtube.com/watch?v=ek6CVVJk4OQ

1. Composition of the solvent and the solute determine whether a substance will dissolve.

Polar/Nonpolar• Polar = negative charge concentrated in 1 area (e-

concentrated) • All ionic compounds (charged ions)

• Some molecules (H20) are polar, due to unequal sharing of electrons. The more electronegative atoms (F is the most) hogs the electrons creating partial charges. EX: water or any salt.

• Nonpolar = negative charge is spread out• Equal sharing of electrons due to small differences in

electronegativity (no partial charges) or perfectly symmetrical molecules.

• ex) propane and vegetable oil

Factors Which Affect Solubility:

General Rule:

•“Like dissolves like.” •Polar dissolves polar•Nonpolar dissolves nonpolar

• Emulsifying Agent : a substance that causes two liquids that would not normally dissolve in each other to mix.• Usually has a polar and a nonpolar end. • EX: soap and detergent, egg yolk in mayonnaise

• Agitation - Shaking or stirring increases solubility• Particle size – the smaller a particle is the faster it will

dissolve. • Temperature – • The solubility of liquids and solids usually increases as

temperature increases. EX: sugar dissolves better in warm water than ice water.

• The solubility of gases usually decreases as temperature increases. EX: CO2 dissolves better in cold soda than in warm soda. That is why cold soda is more bubbly than warm soda.

• Pressure – solubility of gases increases as pressure increases. EX: Soda is bottled under high pressures, so that the soda contains more dissolved CO2.

•  

Factors Which Affect Solubility:

• Unsaturated = solution contains less than the amount that can normally dissolve, at a particular temperature

• Saturated = solution contains the maximum amount that can normally dissolve, at a particular temperature

• Supersaturated = solution contains more than the maximum amount that normally dissolves, at a particular temperature

Types of Solutions

How To Make A Supersaturated Solution:Make a saturated solution at a higher temperatureCool it

The Solute Will Fall Out Of The Solution If:

A seed crystal is addedDisturbing the solution

http://www.youtube.com/watch?v=HnSg2cl09PI&NR=1&feature=fvwphttp://www.youtube.com/watch?v=XSGvy2FPfCw

• Solution - Solution - homogeneous mixture

C. Johannesson

Solvent Solvent - present in greater amount

Solute Solute - substance being dissolved

Molarity (M)• Moles of solute per L of solution • Concentration = how much solute is dissolved in

solvent

• M = moles L

• EX: 2.0 M, 500 mL NaCl

• How much NaCl?

mol

M L÷

X

Dilution• Adding water to make a solution less concentrated

M1V1 = M2V2

Volume units don’t matter as long as they are the same.

EX:* need 0.4 M CuSO4

Colligative Properties• = properties that depend on the number of solute

particles dissolved in a particular amount of solvent. • * A pure solvent will have different properties than a

solution.

Vapor Pressure Lowering

• Vapor pressure = pressure exerted by a vapor that is in equilibrium with its liquid in a closed system

• Solution with a nonvolatile (does not vaporize readily) has a lower VP than pure solvent

• Solvent molecules surround solute decreasing the amount that can escape the liquid dec. VP

• EX: Salt water has a lower vapor pressure than pure water

Freezing Point Depression• The FP of a solution is lower than the FP of the pure

solvent• Solute particles disrupts the formation of the orderly

pattern of a solid dec. FP• Ex) 58.5 g NaCl added to 1000 g H2O• FP = -3.72 °C Pure H2O FP = 0 °C

• http://www.youtube.com/watch?v=A6lC68nb58U&feature=related•  • http://www.youtube.com/watch?v=yxlQRD1aM_Y

• The water molecules arrange into a uniform, crystalline pattern.

Imagine taking a bucket of golf balls ('liquid' form, randomly arranged) and stacking them up in a pyramid ('solidifying' them). This is relatively easy to do, right?

Now imagine the bucket of golf balls, but with a number of tennis balls mixed in (a solution). Try stacking the golf balls - freezing them - but having to randomly insert tennis balls in the stack. This will not be as easy.

The tennis balls (solute) will interfere with crystal formation of the golf balls (solvent). See the following diagram:

Freezing Point Depression

View Flash animation.

Freezing Point Depression

Boiling Point Elevation• The BP of a solution is higher than the BP of a pure

solvent.• Adding solute decreases VP. BP is the point where VP =

Atm P. Extra Energy needs to be added to get the lower VP up to Atm P.

• Ex.) Add to boiling H2O. H2O BP = 100 °C

• 58.5 NaCl to 1000 g H2O BP = 100.52 °C

C. Johannesson

Boiling Point Elevation

Solute particles weaken IMF in the solvent.

Boiling Point Elevation

C. Johannesson

Colligative Properties

•Applications• salting icy roads• making ice cream• antifreeze

• cars (-64°C to 136°C)• fish & insects