solids, liquids and phase changes unit 10. thermodynamics thermodynamics- the study of energy and...
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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