Solids, Liquids and Phase Changes

Unit 10

ThermodynamicsThermodynamics- The study of energy

and the changes it undergoes

1st Law- the energy of the universe is constant.

Conservation of energy

2nd Law the entropy of the universe increases.

RememberEntropy (S) is disorder or randomness

Ssolid <Sliquid <<Sgasthere are many more ways for the molecules to

be arranged as a liquid than a solid.

Gases have a huge number of positions possible.

For exothermic processes DSsurr is

positive.

For endothermic processes DSsurr is

negative.

Consider this process

H2O(l)® H2O(g)

DSsys is positive

DSsurr is negative

Gibb's Free EnergyDG=DH-TDS at constant temperature

G = gibb’s free energy

H= enthalpy (heat exchange in a reaction)

T= temp in Kelvin

S= entropy

If DG is negative at constant T and P, the Process is spontaneous.

If DG is positive at constant T and P, the Process is non-spontaneous.

Third Law of ThermoThe entropy of a pure crystal at 0 K is 0.

All others must be>0.

Standard Entropies Sº ( at 298 K and 1 atm) of substances are listed.

Products - reactants to find DSº (a state function).

Free Energy in ReactionsDGº = standard free energy change.

Free energy change that will occur if reactants in their standard state turn to products in their standard state.

Can’t be measured directly, can be calculated from other measurements.

DGº=DHº-TDSº

Use Hess’s Law with known reactions.

Hess’s LawWe can add equations to come up with the

desired final product, and add the DH

Two rules

1. If the reaction is reversed the sign of DH is changed

2. If the reaction is multiplied, so is DH

Standard EnthalpyThe enthalpy change for a reaction at standard

conditions (25ºC, 1 atm , 1 M solutions)

Symbol DHº

When using Hess’s Law, work by adding the equations up to make it look like the answer.

The other parts will cancel out.

Example

O (g) + H (g) 2OH(g) 2 2 O (g) 2O(g)2 H (g) 2H(g)2

O(g) + H(g) OH(g)

Given

Calculate DHº for this reaction

DHº= +77.9kJDHº= +495 kJ

DHº= +435.9kJ

Since we can manipulate the equations

We can use heats of formation to figure out the heat of reaction.

Lets do it with this equation.

C2H5OH +3O2(g) ® 2CO2 + 3H2O

which leads us to this rule.

( H products) - ( H reactants) = Hfo

fo o

Enthalpy

Exothermic and Endothermic

CH + 2O CO + 2H O + Heat4 2 2 2

CH + 2O 4 2

CO + 2 H O 2 2

Pote

nti

al energ

y

Heat

N + O2 2

Pote

nti

al energ

y

Heat

2NO

N + O 2NO2 2 + heat

Direction Every energy measurement has three parts.

1. A unit ( Joules of calories).2. A number how many.3. and a sign to tell direction.

negative - exothermic

positive- endothermic

Some rules for Heat Heat is the exchange of energy

Heat = q

Heat given off is negative.

Heat absorbed is positive.

CalorimetryMeasuring heat.

Use a calorimeter.

Two kinds, but we will only focus on one

Constant pressure calorimeter (called a coffee cup calorimeter)

heat capacity for a material, C is calculated

C= heat absorbed/ DT = DH/ DT

specific heat capacity = C/mass

CalorimetryHeat capacity – (c) the amount of heat needed

to raise the temperature of a given quantity of a substance by one degree celsius

Specific heat – (s) the amount of heat energy required to raise the temperature of one gram of substance by one degree celsius

Calorimetrymolar heat capacity = C/moles

heat = specific heat x m x DT

heat = molar heat x moles x DT

Make the units work and you’ve done the problem right.

A coffee cup calorimeter measures DH.

An insulated cup, full of water.

The specific heat of water is 1 cal/gºC or 4.18 j/gºC

Heat of reaction= DH = s x mass x DT

Changes of StateMelting point – the temperature at which atomic

or molecular vibrations of a solid become so great that the particles break free from their fixed positions and start to slide past each other in a liquid state

Heating curve – a plot of temperature versus time for a substance where energy is added at a constant rate

Heating Curve for Water

TermsHeat of fusion – the amount of energy required

at the melting point temperature to cause the change of phase to occur

Heat of vaporization – the amount of heat needed to vaporize 1 gram of a liquid at constant temperature and pressure

Know these values for H2O

ExamplesWhat quantity of ice at 273K can be melted by

100 joules of heat?

How much heat is needed to change 100. grams of ice at 273K to steam at 373K?

Phase DiagramsWay to represent the phases of a substance as a

function of temperature and pressure

Triple point – the point at which all three states of a substance are present

Critical temperature – the temperature above which the vapor cannot be liquefied no matter what pressure is applied

Critical pressure – pressure required to produce liquefication at the critical temperature

Together, the critical temperature and critical pressure define the critical point

Phase Diagram for Water

Phase Change Terms

We must know the names of the phase changes.

Which ones do we know already?

How do we identify the phases?

http://www.teamonslaught.fsnet.co.uk/co2%20phase%20diagram.GIF

Phase Diagram for carbon dioxide

Kinetics of LiquidsMolecules of a cold sample of liquid have lower

kinetic energy than those in a warmer sample

If a particle near the surface has enough kinetic energy, it can overcome the attractive forces in a liquid and escape into the gaseous state

Known as a phase change

Distribution of Kinetic Energy of MoleculesN

um

ber

of

Mole

cule

s

ViscosityThe friction or resistance to motion that exists

between the molecules of a liquid when they move past one another

The stronger the attraction between the molecules in a liquid, the greater the resistance to flow

Liquids with large intermolecular forces tend to be highly viscous

Surface TensionThe resistance of a liquid to an increase in its

surface area

Which liquids will have high surface tensions and why?

Those with relatively large intermolecular forces

Because of decreased volume and increased molecular interaction, liquids expand and contract only very slightly with temperature change

Capillary ActionThe attraction of the surface of a liquid to the

surface of a solid

Liquids will rise very high in a narrow tube if a strong attraction exists between the liquid molecules and the molecules that make up the tubing

Pulls liquid up against force of gravity

Concave meniscus

Vapor PressureEvaporation (vaporization) – a process by which

the molecules of a liquid can escape the liquid’s surface and form a gas

Endothermic process

Heat of vaporization (enthalpy of vaporization) – energy required to vaporize one mole of a liquid at a pressure of 1 atm

Symbol: Δhvap

Vapor PressureCondensation – process by which vapor

molecules re-form a liquid

Phase EquilibriumEventually, enough

vapor molecules are present so that the rate of condensation equals the rate of evaporation

The system is at equilibrium

The pressure of the vapor present at equilibrium is called vapor pressure

Phase EquilibriumWhat will happen if the temperature is

increased?

The number of liquid molecules will be reduced

The number of gaseous molecules will be increased

The rates of evaporation and condensation will become equal again

This illustrates Le Châtelier’s Principle

Le Châtelier’s PrincipleWhen a system at equilibrium is disturbed by

the application of a stress, it reacts so as to minimize the stress and attain a new equilibrium position

Le Châtelier’s PrincipleChanging Concentration:

Adding a product to a reaction pushes a reversible reaction at equilibrium in the direction of the reactants.

Changing Temperature:

Increasing the temperature causes the equilbrium position of the reaction to shift in the direction that absorbs the heat.

Le Châtelier’s PrinciplePressure:

Change in pressure only affect gaseous equilibria that have an unequal number of moles of reactants and products

Increasing the pressure on a system results in a shift in the equilibrium position that favors the formation of product

Decreasing pressure will shift the equilibrium position to favor the reactants

Le Châtelier’s PrincipleExample:

Heat + H2O (l) H2O (g)

The equation will shift to the right until equilibrium is reached at the new temperature

Boiling PointThe point at which the liquid’s vapor pressure is

equal to the atmospheric pressure

Rapidly converting from liquid to the vapor phase within the liquid as well as at the surface

Intermolecular ForcesBoth solids and liquids are condensed states of

matter

Relatively weak forces which occur between molecules

*It is important to recognize that when a substance such as water changes from solid to liquid to gas, the molecules remain intact. The changes in state are due to changes in the forces among the molecules rather than within the molecules*

Dipole-dipole Forces

•The attractive force resulting when polar molecules line up so that the positive and negative ends are close to each other

•Try to maximize the + ----- interactions

•In the gas phase, these forces are unimportant

•Weaker than ionic or covalent bonds

Hydrogen Bonding

•Unusually strong dipole-dipole attractions that occur among molecules in which hydrogen is bonded to a highly electronegative atom

Physical PropertiesNonpolar tetrahedral hydrides show a steady

increase in boiling point

Polar tetrahedral hydrides, the lightest member has an unexpectedly high boiling point

This is due to hydrogen bonding that exist among the smallest molecule with the most polar X—H bond.

Boiling Points of Metal Hydrides

London Dispersion

Forces•Forces which exist among noble gas atoms and nonpolar molecules

•Involve an accidental dipole that induces a momentary dipole in a neighbor

The Liquid StateLow compressibility, lack of rigidity, and high density

when compared to gases

Surface tension – the resistance of a liquid to an increase in its surface area

Which liquids will have high surface tensions and why?

Those with relatively large intermolecular forces

Because of decreased volume and increased molecular interaction, liquids expand and contract only very slightly with temperature change

The Liquid StatePolar liquids exhibit capillary action

This is the spontaneous rising of a liquid in a narrow tube, due to:

Cohesive forces – the intermolecular forces among the molecules of the liquid

Adhesive forces – the forces between the liquid and its container

Which of these are stronger for water? Adhesive

The Solid StateCan be classified into very broad categories:

1. Crystalline solids – highly regular arrangement of components

2. Amorphous solids – have considerable disorder in their structure

3. Polycrystalline solid – an aggregate of a large number of small crystals in which the structure is regular but the crystals are arranged in random fashion

Crystalline Solid

Amorphous Solid

Polycrystalline Solid

Crystalline SolidsLattice structure – a 3D system of points

designating the positions of the components

Network SolidsAtomic solid containing strong directional

covalent bonds

Allotropes – forms of the same element that differ in crystalline structure

• Differ in properties because of differences in structure

Allotropes of Carbon