Ch. 12 – States of MatterCh. 12 – States of MatterCh. 12 – States of MatterCh. 12 – States of Matter

I. Intermolecular ForcesI. Intermolecular Forces

A. Definition of IMFA. Definition of IMF

Attractive forces between molecules.

Much weaker than chemical bonds within molecules.

a.k.a. van der Waals forces

C. Johannesson

B. Types of IMFB. Types of IMF

B. Types of IMFB. Types of IMF

London Dispersion Forces

View animation online.

B. Types of IMFB. Types of IMF

Dipole-Dipole Forces

+ -

View animation online.

B. Types of IMFB. Types of IMF

Hydrogen Bonding

C. Determining IMFC. Determining IMF

NCl3• polar = dispersion, dipole-dipole

CH4

• nonpolar = dispersionHF

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

II. Physical PropertiesII. Physical Properties

Ch. 12 - Liquids & SolidsCh. 12 - Liquids & SolidsCh. 12 - Liquids & SolidsCh. 12 - Liquids & Solids

A. Liquids vs. SolidsA. Liquids vs. Solids

LIQUIDS

Stronger than in gases

Y

high

N

slower than in gases

SOLIDS

Very strong

N

high

N

extremely slow

IMF Strength

Fluid

Density

Compressible

Diffusion

B. Liquid PropertiesB. Liquid Properties

Surface Tension• attractive force between particles in a

liquid that minimizes surface area

B. Liquid PropertiesB. Liquid Properties

Capillary Action• attractive force between the surface of

a liquid and the surface of a solid

B. Liquid PropertiesB. Liquid Properties

Viscosity• Measure of the resistance of a liquid to

flow

B. Liquid PropertiesB. Liquid Properties

Cohesion• Force of attraction between

identical moleculesAdhesion

• Force of attraction

between different

molecules water mercury

C. Types of SolidsC. Types of Solids

Crystalline - repeating geometric pattern• covalent network• metallic• ionic• covalent molecular

Amorphous - no geometric pattern

decreasingm.p.

C. Types of SolidsC. Types of Solids

Ionic(NaCl)

Metallic

C. Types of SolidsC. Types of Solids

CovalentMolecular

(H2O)

CovalentNetwork

(SiO2 - quartz)

Amorphous(SiO2 - glass)

C. Types of SolidsC. Types of Solids

Allotrope• In a covalent network, different forms in the same

state

Carbon

allotropeswater mercury

C. Types of SolidsC. Types of Solids

Amorphous Solid• No geometric pattern

water mercury

Ch. 12 - Liquids & SolidsCh. 12 - Liquids & SolidsCh. 12 - Liquids & SolidsCh. 12 - Liquids & Solids

III. Changes of StateIII. Changes of State

A. Phase ChangesA. Phase Changes

A. Phase ChangesA. Phase Changes

EvaporationEvaporation• molecules at the surface gain enough

energy to overcome IMF

VolatilityVolatility• measure of evaporation rate• depends on temp & IMF

A. Phase ChangesA. Phase Changes

Kinetic Energy

# o

f P

art

icle

s

Boltzmann Distribution

temp

volatility

IMF

volatility

A. Phase ChangesA. Phase Changes

EquilibriumEquilibrium• trapped molecules reach a balance

between evaporation & condensation

A. Phase ChangesA. 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

temp

v.p

.

A. Phase ChangesA. 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

Which has a higher m.p.?• polar or nonpolar?• covalent or ionic?

A. Phase ChangesA. Phase Changes

Melting Point• equal to freezing point

polar

ionic

IMF m.p.

A. Phase ChangesA. Phase Changes

Sublimation

• solid gas

• v.p. of solid equals external pressure

EX: dry ice, mothballs, solid air fresheners

B. Heating CurvesB. Heating Curves

Melting - PE

Solid - KE

Liquid - KE

Boiling - PE

Gas - KE

B. Heating CurvesB. 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

B. Heating CurvesB. 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.

B. Heating CurvesB. Heating Curves

Heat of Vaporization (Hvap)

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

EX: sweating, steam burns, the drinking bird

C. Phase DiagramsC. Phase Diagrams

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

The following slides… same information, different explanation and examples

Phase Changes

Why do liquids and solids form at all?

• KMT postulates– A gas is a collection of small particles traveling in

straight-line motion and obeying Newton's Laws. – The molecules in a gas occupy no volume. – Collisions between molecules are perfectly elastic

• no energy is gained or lost during the collision

– There are no attractive or repulsive forces between the molecules.

– Kinetic energy is proportional to temperatureX

What is a phase?• Region of matter that is:

– chemically uniform– physically distinct– mechanically separable.

• Often synonymous with (same meaning as) “state of matter”

Changing phases

Distinguish liquid vs. solid?

Properties of …• Gas phase

– Like/unlike soccer players on field

• Liquid phase– Like/unlike crowd at a rally– Like/unlike gases

• Solid phase– Like/unlike movie theatre

• Compare intermolecular interactions – gases vs. liquids vs. solids

Phases of matter

Add

mor

e en

ergy

Phase change vocabulary

Phase changes and IMF’s• As InterMolecular Forces increase, melting

and boiling temperatures _________? (increase or decrease)?

Metallic bonding

Network covalent bonding

Larger sphere, higher

melting point

Strength of interactions

• Which simulation has stronger intermolecular interactions? A or B– How do you know?

A BSame temperature

Ranking of intermolecular interactions • Water• Wood• Iron• Air• Gold• Mercury• Carbon dioxide• Oxygen• Gasoline• Lead

vs.

Why?

Why is water special?

• Periodic trends– Boiling and melting points of hydrides

KMT, energy and phase

changes

Total energy

Water phases present?

Heating curves

Why are b and d flat?

Boiling vs. melting

Which takes more energy?(same mass)

Consider liquid gasEvaporation removes energy

Evaporative Cooling

Evaporation and volatility

• Evaporation– molecules at the surface gain enough energy

to overcome attractive intermolecular forces (IMF)

• Volatility– measure of evaporation rate– depends on temperature and IMF

Volatility

Kinetic Energy

# o

f P

art

icle

s

Boltzmann Distribution

temp

volatility

IMF

volatility

Boiling – Fixed temperature

Why do liquids boil?• Boiling Point

–T at which liquid vapor pressure = atmospheric pressure• depends on Patm & IMF

–Normal boiling point - b.p. at 1 atm

IMF b.p.Patm b.p.

Temperature/pressure phase diagram for water

Transition to Supercritical CO2

Interpret phase diagram for water

Draw a phase diagram for NH3

(used as a refrigerant in RV’s and solar cooling systems)

Ammonia phase diagram

Phase Changes Need to know…

• Know why liquids and solids exist at all– Failure of KMT postulate

• Know phase names and changes – molecular views of phases and changes– distinguish S vs. L on molecular level

• Explain heating curve origins and parts– Why no change in T during phase change?

• Explain differences and similarities between evaporation, volatility and boiling

• Explain evaporative cooling on a molecular level

• Interpret temperature-pressure phase diagrams