Energy and phases

All matter can undergo changes in its state.

These changes have to do with the amount of energy in the particles of matter.

Kinetic theory of matter-

1.All matter is made of particles

2.These particles are in constant motion.

More energy causes the particles to move faster.

At 100°C, water becomes water vapor, a gas. Molecules can move randomly over large distances.

Below 0°C, water solidifies to become ice. In the solid state, water molecules are held together in a rigid structure.

Between 0°C and 100 °C, water is a liquid. In the liquid state, water molecules are close together, but can move about freely.

Matter has five states or phases

Solid : A definite shape and volume

Lower energy

Liquid: A definite volume but it takes the shape of its container

Higher energy Gas : No definite

shape or volume Yet even higher

energy

Plasma : No definite shape or volume and whose particles have broken apart

Bose-Einstein Condensate: Gases near absolute zero forming a super fluid

Plasma is by far the most common form of matter. Plasma in the stars and in the tenuous space between them makes up over 99% of the visible universe and perhaps most of that which is not visible.

Liquids: Are not very compressible

Useful in hydraulics

Viscosity: Liquids resistance to flow

Surface tension: holds the liquid together

Gases: Fill their container and can change pressure

http://www.stolaf.edu/people/giannini/flashanimat/transport/project.swf

Diffusion: Spreading of particles through an area until they are uniformly distributed

Changes of state: When matter changes form, it is a physical change and has to do with the energy of the material

Ex. Boiling, melting, freezing, condensing –all require an energy change

Heat of fusion: Energy required to go from solid state to liquid state (For water 334kJ/kg)

Melting point – different for substances

helium−269 hydrogen−253 Iron-2887 Graphite (carbon)3900 Diamond

(carbon)4827 Tungsten-5660 Gold-3080

Heat of Vaporization: Energy required to go from liquid to gas. (For water 2260kJ/kg)

Condensation: A gas changes to a liquid when cooled to or below its boiling point

Vaporization is at boiling point or below

Evaporation: A liquid changes to a gas without reaching its boiling point

Sublimation: Changing from a solid to a gas without existing as a liquid

Deposition : changing from a gas to a solid without being a liquid

Latent heat: heat absorbed without a change in temperature (stored until a phase change)

Is water a solid or liquid at 0C? Why?

Honor only

Intermolecular Forces

A phase is a homogeneous part of the system in contact with other parts of the system but separated from them by a well-defined boundary. 2 Phases

Solid phase - ice

Liquid phase - water

11.1

11.2

Intermolecular forces are attractive forces between molecules.

Intramolecular forces hold atoms together in a molecule.

Intermolecular vs Intramolecular

• 41 kJ to vaporize 1 mole of water (inter)

• 930 kJ to break all O-H bonds in 1 mole of water (intra)

Generally, intermolecular forces are much weaker than intramolecular forces.

“Measure” of intermolecular force

boiling point

melting point

Hvap

Hfus

Hsub

Intermolecular forces are feeble; but without them, life as we know it would be impossible. Water would not condense from vapor into solid or liquid forms if its molecules didn't attract each other. Intermolecular forces are responsible for many properties of molecular compounds, including crystal structures (e. g. the shapes of snowflakes), melting points, boiling points, heats of fusion and vaporization, surface tension, and densities. Intermolecular forces pin gigantic molecules like enzymes, proteins, and DNA into the shapes required for biological activity.

http://www.nationmaster.com/encyclopedia/Image:Myoglobin.png

Intermolecular Forces

1. London Forces (Dispersion Forces)

2. Dipole-Dipole Interactions

3. Ion-Dipole Interactions (Salt dissolving in solution)

4. Hydrogen Bonding

Dispersion ForcesOccur between every compound and arise from the net attractive forcesamount molecules which is produced from induced charge imbalances

The magnitude of the Dispersion Forces is dependent upon how easily itis to distort the electron cloud. The larger the molecule the greater it’s Dispersion Forces are.

Figure 10-8 Olmsted Williams

Figure 10-9 Olmsted Williams

Olmsted Williams Fig 10-10 Pg 437

The boiling point of alkanes increase with the length of the carbon chain. Long-chain alkanes have larger dispersion forces because of the increased polarizability of their larger electron cloud.

How molecular shape affects the strength of the dispersion forces

The shapes of the molecules also matter. Long thin molecules can develop bigger temporary dipoles due to electron movement than short fat ones containing the same numbers of electrons.Long thin molecules can also lie closer together - these attractions are at their most effective if the molecules are really close.For example, the hydrocarbon molecules butane and 2-methylpropane both have a molecular formula C4H10, but the atoms are arranged differently. In

butane the carbon atoms are arranged in a single chain, but 2-methylpropane is a shorter chain with a branch.

                                                                                                                                                            Butane has a higher boiling point because the dispersion forces are greater. The molecules are longer (and so set up bigger temporary dipoles) and can lie closer together than the shorter, fatter 2-methylpropane molecules.

http://www.chemguide.co.uk/atoms/bonding/vdw.html

Polarizability

11.2

the ease with which the electron distribution in the atom or molecule can be distorted.

Polarizability increases with:

• greater number of electrons

• more diffuse electron cloud

Dispersion forces usually increase with molar mass.

Is the Molecule Polar?

We have already talked about diatomic molecules. The moreElectronegative atom will pull the electron density of the bond Closer to itself giving it a partial negative charge leaving the otherAtom with a partially positive charge. Thus giving the molecule A dipole moment.

But what about molecules made up of more than two molecules?

Dipole-Dipole Forces

Attractive forces between polar molecules

Orientation of Polar Molecules in a Solid

11.2

Figure 10-11

Dipole Forces occur between molecules containing a dipole moment. The positive end of the dipole moment on one mole is attracted to theNegative end of the dipole moment on a nearby molecule.

Consider 2-methyl propane(left) and acetone (right) Both compounds are aboutEqual in size and shape therbyHaving similar dispersion forces,But Acetone contains anOxygen (red) and causes theMolecule to have a dipole Moment allowing it to haveDipole forces and thus a Higher boiling point

Olmsted Williams

Ion-Dipole Forces

Attractive forces between an ion and a polar molecule

11.2

Ion-Dipole Interaction

The larger the charge the stronger the force

Fig 10-34

A molecular picture showing the ion-dipole Interaction that helps a solid ionic crystal dissolve in water. The arrows indicate ion-dipole interactions.

Olmsted Williams

SO

O

What type(s) of intermolecular forces exist between each of the following molecules?

HBrHBr is a polar molecule: dipole-dipole forces. There are also dispersion forces between HBr molecules.

CH4

CH4 is nonpolar: dispersion forces.

SO2

SO2 is a polar molecule: dipole-dipole forces. There are also dispersion forces between SO2 molecules.

11.2

Hydrogen Bond

11.2

The hydrogen bond is a special dipole-dipole interaction between they hydrogen atom in a polar N-H, O-H, or F-H bond and an electronegative O, N, or F atom.

A H…B A H…Aor

A & B are N, O, or F

Intermolecular Forces

1. London Forces (Dispersion Forces)

2. Dipole-Dipole Interactions

3. Ion-Dipole Interactions (Salt dissolving in solution)

4. Hydrogen Bonding

These forces affect how molecules will interact with each other andAs a general rule as the strength of the force increases the boiling Point of the compound increases

Surface tension is the amount of energy required to stretch or increase the surface of a liquid by a unit area.

Strong intermolecul

ar forces

High surface tension

11.3

Liquids and Surface Tension

Properties of Liquids

Cohesion is the intermolecular attraction between like molecules

11.3

Adhesion is an attraction between unlike molecules

Adhesion

Cohesion

Evapora

tion

GreatestOrder

LeastOrder

11.8C

on

den

sati

on

T2 > T1

The equilibrium vapor pressure is the vapor pressure measured when a dynamic equilibrium exists between condensation and evaporation

H2O (l) H2O (g)

Rate ofcondensation

Rate ofevaporation=

Dynamic Equilibrium

11.8

A substance with a high Vapor pressure is consideredTo be volitile therefore, the lowerThe boiling point the higher the Vapor pressure and the weakerThe intermolecular forces

The boiling point is the temperature at which the (equilibrium) vapor pressure of a liquid is equal to the external pressure.

The normal boiling point is the temperature at which a liquid boils when the external pressure is 1 atm.

11.8

Melt

ing

11.8Fr

eezi

ng

H2O (s) H2O (l)

The melting point of a solid or the freezing point of a liquid is the temperature at which the solid and liquid phases coexist in equilibrium

11.8

Sub

limati

on

11.8

Deposi

tion

H2O (s) H2O (g)

Molar heat of sublimation (Hsub) is the energy required to sublime 1 mole of a solid.

Hsub = Hfus + Hvap

( Hess’s Law)

A phase diagram summarizes the conditions at which a substance exists as a solid, liquid, or gas.

Phase Diagram of Water

11.9

11.9