Intermolecular forces
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Intermolecular Forces. Kinetic Molecular Theory. Describes the behavior of subatomic particles Liquids, solids, and gases are composed of small particles that have mass. Particles are in constant, random, rapid motion. Particles have collisions.

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Kinetic molecular theory
Kinetic Molecular Theory

  • Describes the behavior of subatomic particles

    • Liquids, solids, and gases are composed of small particles that have mass.

    • Particles are in constant, random, rapid motion.

    • Particles have collisions.

    • Particles have an avg. KE directly related to temperature.

    • The state of a substance at room temperature depends on the strength of the attractions between its particles.


Definition of imf
Definition of IMF

  • Attractive forces between molecules.

  • Much weaker than chemical bonds within molecules.


Definition of imf1
Definition of IMF

  • Intramolecular forces:

    Covalent Bonding

  • Much stronger than chemical bonds between atoms.

  • Examples : nonpolar, polar sharing


Definition of imf2
Definition of IMF

  • Intermolecular Forces

  • Attractive forces between molecules.

  • Much weaker than chemical bonds within molecules.


Intermolecular forces1
Intermolecular Forces

  • Attractive forces between molecules or particles (ions, metal atoms, etc…)

  • Examples:

    • dispersion, (London /Vander Waals);

    • dipole-dipole,

    • dipole-ion,

    • hydrogen “bonding”,

    • metallic bonding,

    • ion-ion


Intermolecular forces2
Intermolecular Forces

  • Relative Strength:

Weakest

Strongest

  • Examples:

    • dispersion, (London /Vander Waals);

    • dipole-dipole,

    • dipole-ion,

    • hydrogen “bonding”,

    • metallic bonding,

    • ion-ion



Types of imf1
Types of IMF

  • London Dispersion Forces

View animation online.


Types of imf2

+

-

Types of IMF

  • Dipole-Dipole Forces

View animation online.


Types of imf3
Types of IMF

  • Hydrogen Bonding


Determining imf
Determining IMF

  • NCl3

    • polar = dispersion, dipole-dipole

  • CH4

    • nonpolar = dispersion

  • HF

    • H-F bond = dispersion, dipole-dipole, hydrogen bonding


Liquids solids

Liquids & Solids

Physical Properties


Liquids vs solids

LIQUIDS

Stronger than in gases

Y

high

N

slower than in gases

SOLIDS

Very strong

N

high

N

extremely slow

Liquids vs. Solids

IMF Strength

Fluid

Density

Compressible

Diffusion


Liquid properties
Liquid Properties

  • Surface Tension

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


Liquid properties1

water

mercury

Liquid Properties

  • Capillary Action

    • attractive force between the surface of a liquid and the surface of a solid


Types of solids

decreasing

m.p.

Types of Solids

  • Crystalline - repeating geometric pattern

    • covalent network

    • metallic

    • ionic

    • covalent molecular

  • Amorphous - no geometric pattern


Types of solids1
Types of Solids

Ionic

(NaCl)

Metallic


Types of solids2
Types of Solids

Covalent

Molecular

(H2O)

Covalent

Network

(SiO2 - quartz)

Amorphous

(SiO2 - glass)


Liquids solids1

Liquids & Solids

Changes of State



Phase changes1
Phase Changes

  • Evaporation

    • molecules at the surface gain enough energy to overcome IMF

  • Volatility

    • measure of evaporation rate

    • depends on temp & IMF


Phase changes2

# of Particles

temp

volatility

IMF

volatility

Kinetic Energy

Phase Changes

Boltzmann Distribution

p. 477


Phase changes3
Phase Changes

  • Equilibrium

    • trapped molecules reach a balance between evaporation & condensation


Phase changes4

temp

v.p.

IMF

v.p.

Phase Changes

p.478

  • Vapor Pressure

    • pressure of vapor above a liquid at equilibrium

v.p.

  • depends on temp & IMF

  • directly related to volatility

temp


Phase changes5

Patm

b.p.

IMF

b.p.

Phase Changes

  • Boiling Point

    • temp at which v.p. of liquid equals external pressure

  • depends on Patm & IMF

  • Normal B.P. - b.p. at 1 atm


Phase changes6

IMF

m.p.

Phase Changes

  • Melting Point

    • equal to freezing point

  • Which has a higher m.p.?

    • polar or nonpolar?

    • covalent or ionic?

polar

ionic


Phase changes7
Phase Changes

  • Sublimation

    • solid  gas

    • v.p. of solid equals external pressure

  • EX: dry ice, mothballs, solid air fresheners


Heating curves

Gas - KE

Boiling - PE 

Liquid - KE 

Melting - PE 

Solid - KE 

Heating Curves


Heating curves1
Heating Curves

  • Temperature Change

    • change in KE (molecular motion)

    • depends on heat capacity

  • Heat Capacity

    • energy required to raise the temp of 1 gram of a substance by 1°C

    • “Volcano” clip -

water has a very high heat capacity


Heating curves2
Heating Curves

  • Phase Change

    • change in PE (molecular arrangement)

    • temp remains constant

  • Heat of Fusion (Hfus)

    • energy required to melt 1 gram of a substance at its m.p.


Heating curves3
Heating Curves

  • Heat of Vaporization (Hvap)

    • energy required to boil 1 gram of a substance at its b.p.

    • usually larger than Hfus…why?

  • EX: sweating, steam burns, the drinking bird


Phase diagrams
Phase Diagrams

  • Show the phases of a substance at different temps and pressures.


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