Intermolecular ForcesLACC Chem101

States of MatterSolids

High density Little translational motion (if any) Rotations/vibrations give temperature Intermolecular forces cause 3D structures

Liquids High density Translational motion limited by frequent collisions High degree of intermolecular forces

Gases Low density Very high speed translational motion Little intermolecular forces

Often approximated as no interactions (ideal gases)

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Phase TransitionsMelting Point

Melting: Solid Liquid Freezing: Liquid Solid

Boiling Point Condensation: Gas Liquid Vaporization: Liquid Gas

Sublimation Point Sublimation: Solid Gas Deposition: Gas Solid

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Phase DiagramsShow equilibrium lines between phases

Phases defined by both temperature and pressure

Key Parts Triple point Critical point Normal boiling point

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Latent HeatHeat transfer of a constant temperature process

Heat energy added/lost, but no associated temperature change

Energy used to change structure of the material

Energy transfer changes the internal energy of the substance, and is therefore an enthalpy

We define enthalpies phase changes:

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Heating/Cooling CurveGraphical representation of the temperature/energy of a

substance

Example: Water moving from ice to gas

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Water example Calculate the amount of energy required to change 15.0g of ice at -5.00C to

steam at 125.0C.

The first step is to design a pathway: q1 = msDT for ice from -5.0 to 0.0 oC, the specific heat of ice is 4.213 J/g oC

q2 = DHfus for ice to liquid at 0.0oC

q3 = msDT for liquid 0.0oC to 100.0 oC

q4 = DHvap for liquid to steam at 100.0oC

q5 = msDT for steam 100.0 to 125.0 oC; the specific heat of steam is 1.900 J/g oC

qT = q1 + q2 + q3 + q4 + q5

The next step is to calculate each q: q1= (15.0 g) (4.213 J/g oC) (0.0 - (-5.0) oC) = 316 J

q2 = (15.0 g) (335 J / g) = 5025 J

q3= (15.0 g) (4.184 J/g oC) (100.0 - (0.0) oC) = 6276 J

q4 = (15.0 g) (2260 J / g) = 33900 J

q5= (15.0 g) (1.900 J/g oC) (110 - 100 oC) = 285 J

qT = 316 J + 5025 J + 6276 J + 33900 J + 285 J = 45.8 kJLACC Chem101

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Intermolecular ForcesNot to be confused with intramolecular forces

These are forces between atoms within molecules Covalent/Ionic bonds

Intermolecular Forces occur between molecules

Types of forces depend on compound type Neutral molecules

Dipole-Dipole forces London dispersion Hydrogen bonding

Ionic compounds Ion-dipole forces

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Intermolecular Forces Intermolecular forces affect boiling and melting points

Stronger force requires more energy to break the molecules apart

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Type of Interaction Energy Range (kJ/mol)

Intermolecular

Van der Waals 0.01 - 10

Hydrogen bond 10 - 40

Chemical Bond

Ionic 100 - 1000

Covalent 100 - 1000

Intermolecular ForcesIon-Dipole

between ions and polar moleculesstrength is dependent on charge of the ions or polarity of the bonds usually involved with salts & H20

Dipole-Dipolebetween neutral polar molecules weaker force than ion-dipolepositive dipole attracted to negative dipolemolecules should be relatively close togetherstrength is dependent on polarity of bonds

London dispersionall molecules and compoundsinvolves instantaneous dipolesstrength is dependent on Molar Mass (size)contributes more than dipole-dipoleshape contributes to strength

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Hydrogen BondingExists between hydrogen atom in a polar bond and a lone pair

of electrons on a nearby electronegative species Oxygen, Fluorine, and Nitrogen

Special case of dipole-dipole interaction

Stronger than dipole-dipole and London dispersion

Accounts for water’s notable properties High boiling point for small size Solid expands from liquid volume

Less dense than the liquid “universal” solvent High heat capacity

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Flowchart of Intermolecular Forces

Interacting molecules or ions

Are polar Are ions Are polar molecules involved? molecules

involved? and ions both

present? Are hydrogen

atoms bonded to N,

O, or F atoms?

London forces Dipole-dipole hydrogen bonding Ion-dipole Ionic only (induced forces forces Bondingdipoles)Examples: Examples: Examples Example: Examples:Ar(l), I2(s) H2S, CH3Cl liquid and solid KBr in NaCl,

H2O, NH3, HF H2O NH4NO3

NO NO YES NO

YES

Yes

NOYES

Van der Waals forcesLACC Chem101

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Properties of LiquidsViscosity

Resistance of a liquid to flow Depends on attractive forces between molecules May also be caused by structural features (entanglement)

Surface tension Energy required to increase the surface area of a liquid

(Energy/Area) Due to interactions between molecules

Also lack of interactions at an interface

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Vapor PressurePressure exerted by a vapor in equilibrium with its liquid or

solid state

Changes with intermolecular forces

Involves an equilibrium between liquid and gas Volatile Nonvolatile

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Clausius-Clapeyron EquationHigher temperature causes a weakening of intermolecular

forces This causes higher vapor pressure

Relationship is a differential equation:

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Clausius-Clapeyron ExampleThe vapor pressure of ethanol at 34.9C is 100.0mmHg. The

normal boiling point of ethanol is 78.5C. Calculate the heat of vaporization.

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Crystalline SolidComposed of crystal lattices

Unit cell: Geometric arrangement of lattice points smallest repeating unit in cell structure Edge lengths and angles describe unit cell

Many types of unit cells Metals and salts are usually cubic

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Unit cell exampleDetermine the number of ions in the lithium fluoride unit cell.

The structure is a face centered cube.

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Molecular SolidsSolid composed of molecules held together by van der Waals

forces

Require large number of atoms surrounding center for maximum attraction

Close-packing arrangement variations Hexagonal close-packed Cubic close-packed

Similar to face-centered cubic

Coordination number Number of nearest neighbors Highest is 12

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Metallic SolidsSea of delocalized electrons

Usually cubic or hexagonal close-packed

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Covalent NetworkDirectional covalent bonds

Hybridization affects structure Structure gives physical properties

Examples Tetrahedral structures: diamond, Si, Ge, Sn

sp3 hybridized Face-centered cubic cells

Hexagonal Sheets: graphite, carbon nanotubes sp2 hybridized Electrical properties

delocalized electrons

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CRYSTALLINE SOLIDS

Type of solid lattice site Type of force properties of examples particle type between particles solids

IONIC positive & electrostatic high M.P. NaCl negative ions attraction nonvolatile Ca(NO3)2

hard & brittle poor conductor

POLAR polar dipole-dipole & moderate M.P. Sucrose,MOLECULAR molecules London Dispersion moderate C12H22O11

forces volatility Ice, H2ONONPOLAR Nonpolar London Dispersion low M.P., Argon, Ar,MOLECULAR molecules & forces volatile Dry Ice, CO2

atomsMACRO- atoms covalent bonds extremely high Diamond, CMOLECULAR between atoms M.P. nonvolatile Quartz, SiO2

Covalent- Arranged in Very HardNetwork Network Poor conductorMETALLIC metal atoms attraction between variable M.P. Cu, Fe

outer electrons low volatility Al, Wand positive good conductor

atomic centersLACC Chem101

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TYPE OF MELTING POINT HARDNESS ELECTRICAL SOLID OF SOLID & BRITTLENESS CONDUCTIVITY

Molecular Low soft & brittle Nonconducting

Metallic Variable Variable hardness, conducting malleable

Ionic High to very hard & brittle Nonconducting high solid

(conducting liquid)

Covalent Very high Very hard UsuallyNetwork nonconducting

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X-ray diffractionMethod for determining atomic/molecular structure of a crystal

Atoms reflect x-rays directionally Angles and intensities of diffracted light rays give 3D picture of

electron densities Atom positions then deduced

Planes of atoms act as reflecting surfaces

X-rays reflect and make diffraction pattern on photographs Constructive interference give more intense waves

Only at angles where X-rays are in-phase Destructive interference cancels

Type of unit cell and size can be determined

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X-Ray Diffraction If molecular, the position of atoms can be determined using

the Bragg equation

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