Transcript Solids and Liquids

Solids, Liquids and
Phase Changes
Unit 10
Thermodynamics
 Thermodynamics- 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.
Remember
Entropy (S) is disorder or randomness
Ssolid <Sliquid <<Sgas
 there 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 Energy
 DG=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 Thermo
 The 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 Reactions
 DGº = 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 Law
 We 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 Enthalpy
 The 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
Given
O 2 (g) + H 2 (g)  2OH(g) DHº= +77.9kJ
O 2 (g)  2O(g) DHº= +495 kJ
H 2 (g)  2H(g) DHº= +435.9kJ
Calculate DHº for this reaction
O(g) + H(g)  OH(g)
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.
o
o
o
( DH f products) - ( DH f reactants) = DH
Enthalpy
Exothermic and Endothermic
Potential energy
CH 4 + 2O 2  CO 2 + 2H 2 O + Heat
CH 4 + 2O 2
Heat
CO 2 + 2 H 2 O
N 2 + O 2 + heat  2NO
Potential energy
2NO
Heat
N2 + O2
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.

Calorimetry
 Measuring 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
Calorimetry
 Heat 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
Calorimetry
 molar 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 State
 Melting 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
Terms
 Heat 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
Examples
 What 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 Diagrams
 Way 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?
Phase Diagram for carbon dioxide
http://www.teamonslaught.fsnet.co.uk/co2
%20phase%20diagram.GIF
Kinetics of Liquids
 Molecules 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
Number of Molecules
Distribution of Kinetic Energy
of Molecules
Viscosity
 The 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 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
Capillary Action
 The 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 Pressure
 Evaporation (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 Pressure
 Condensation – process by which vapor molecules reform a liquid
Phase Equilibrium
 Eventually, 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 Equilibrium
 What 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 Principle
 When 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 Principle
Changing 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 Principle
Pressure:
 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 Principle
Example:
Heat + H2O (l)
H2O (g)
 The equation will shift to the right until equilibrium is
reached at the new temperature
Boiling Point
 The 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 Forces
 Both 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 Properties
 Nonpolar 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 State
 Low 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 State
 Polar 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 State
 Can 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 Solids
 Lattice structure – a 3D system of points designating
the positions of the components
Network Solids
 Atomic 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