11/17/2010 1 Chapter 10 Liquids and Solids The Three States (Phases) of Matter Changes of State The Phase Changes of Water Evaporation and Condensation Enthalpy (Heat) of Vaporization, ∆H vap The energy needed to vaporize 1 mol of a liquid at 1 atm pressure Vaporization is endothermic since energy is required to overcome the intermolecular forces in the liquid Example: Enthalpy of vaporization of water ∆H vap = 40.7 kJmol -1 The large enthalpy of vaporization of water (due to hydrogen bonding) helps cool the surface of the Earth as well as the body through perspiration
Welcome message from author
This document is posted to help you gain knowledge. Please leave a comment to let me know what you think about it! Share it to your friends and learn new things together.
Transcript
11/17/2010
1
Chapter 10Liquids and Solids
The Three States (Phases) of Matter
Changes of State
The Phase Changes of Water
Evaporation and Condensation
Enthalpy (Heat) of Vaporization, ∆HvapThe energy needed to vaporize 1 mol of a liquid at 1 atm
pressure
Vaporization is endothermic since energy is required to overcome the intermolecular forces in the liquid
Example:
Enthalpy of vaporization of water∆Hvap= 40.7 kJmol-1
The large enthalpy of vaporization of water (due to hydrogen bonding) helps cool the surface of the Earth as well as the
body through perspiration
11/17/2010
2
Vapor Pressure of a Liquid in a Closed System
(a) Initially molecules evaporate (at a constant rate at a given temperature) and the amount of liquid decreases
(b) As the amount of vapor increases the vapor starts to condense back into liquid and condensation rate increases
Eventually the rate of vaporization equals the rate of condensation and the system reaches equilibrium
The vapor pressure (Pvap) at equilibrium is called the equilibrium vapor pressure or just the vapor pressure of the liquid
closed system
Measurement of Vapor Pressure
Pvapor = Patmopshere - PHg
Vapor Pressure and Intermolecular Forces
The vapor pressure of a liquid is determined by the size of the intermolecular forces between the molecules of the liquid
The stronger the forces, the lower the vapor pressure
Polar molecules that interact via stronger hydrogen bonding and dipole-dipole forces have low vapor pressures
Small non-polar molecules that interact via weak dispersion forces have high vapor pressures and are said to be volatile
However, non-polar molecules with high molecular weights have low vapor pressures because dispersion forces become significant
Solids also have vapor pressures, but they are typically much lower than liquids
Place the following in order of increasing vapor pressure:
CH4
H2O
NaCl
C10H22
He
NH3
Vapor Pressure and Temperature
Vapor pressure increases rapidly with temperature since more molecules have the kinetic energy required to escape the liquid
Normal Boiling Point (nbpt)
Temperature at which the vapor pressure of a liquid equals
1 atm or 760 mmHg
Example:
The normal boiling point of water is 100 C
11/17/2010
3
Quantitative Dependence of Pvap on Temperature
1ln( ) vap
vap
HP C
R T
This is a simple linear equation of the form y = mx + b so if you measure Pvap at different temperatures a plot of ln(Pvap) vs. 1/T will yield a straight line with a slope of –∆Hvap/R and an intercept of C (a
constant characteristic of a given liquid)
The slopes are always negative consistent with ∆Hvap being positive (endothermic)
The smaller the slope, the lower ∆Hvap and the more volatile the liquid
Water has a large ∆Hvap due to the strong hydrogen bonding between molecules
Vapor Pressure of Liquid Nitric Acid
1/T (K-1)
0.0026 0.0028 0.0030 0.0032 0.0034 0.0036 0.0038
ln(P
vap)
2
3
4
5
6
7
ln(Pvap) = -4640(1/T) + 19.68 If we know ∆Hvap and Pvap at one temperature we can calculate Pvap at a given temperature or the temperature at a given Pvap
Since the constant, C does not depend on temperature we can solve for C at two temperatures T1 and T2:
Melting and Freezing
11/17/2010
4
Enthalpy (Heat) of Fusion, ∆HfusThe energy needed to melt 1 mol of a solid at 1 atm pressure
Melting is endothermic since energy is required to overcome the intermolecular forces in the solid
Example:
Enthalpy of fusion of ice∆Hfus= 6.02 kJmol-1
Normal Melting Point (nmpt)
Temperature at which a solid melts at pressure of a 1 atm or
760 mmHg
Example:
The normal melting point of water is 0 C
At the normal melting point of a substance, the vapor pressures of the solid and the liquid are equal to 1 atm:
Note that the vapor pressure of ice increases with temperature at a faster rate than the vapor pressure of water
A solid’s enthalpy of fusion and melting point is related to the strength of the intermolecular or interatomic forces in the solid:
Changes of state do not always occur exactly at the melting or boiling points!
Supercooling
This occurs when a liquid is rapidly cooled below it melting point at 1 atm but still remains a liquid for some time due its inability to
immediately reorganize its structure into the solid
However, eventually the solid does form, releasing energy and bringing the temperature back up to melting point where the
remainder of the liquid freezes
11/17/2010
5
Superheating
This occurs when a liquid is rapidly heated above its boiling point at 1 atm but still remains as a liquid
Bubbles of hot vapor form in the liquid which rapidly expand and burst before reaching the surface, blowing the liquid out of the out
the container
This is called bumping and can be controlled by using boiling chips which prevent the formation of large bubbles
Sublimation and Deposition
Solid turns directly into a gas or a gas turns directly into a solid without passing through the liquid state
Examples:
carbon dioxide‘dry ice’
iodine
Enthalpy (Heat) of Sublimation, ∆HsubThe energy needed to sublime 1 mol of a solid at 1 atm
pressure
Sublimation is endothermic since energy is required to overcome the intermolecular forces in the solid
Example:
Enthalpy of sublimation of iodine∆Hvap= 28.7 kJmol-1
The enthalpy of sublimation is the sum of the enthalpy of fusion and the enthalpy of vaporization:
∆Hsub= ∆Hfus + ∆Hvap
Heating Curves
11/17/2010
6
It is possible to calculate the total amount of energy required to convert at solid at some initial temperature to another phase (either
a liquid or a gas) at a given temperature by summing up the energies required at each stage
This requires calculating the amount of energy required to heat a particular phase to a given temperature using its specific heat
capacity:
q = s x m x ∆T
where:
s = specific heat capacity (Jg-1ºC-1)q = energy required (J)m = mass of sample (g)
∆T = temperature change (ºC) = Tfinal - Tinitial
Cooling Curves
Energy is released when a vapor condenses at the boiling point:
The energy released when 1 mol of a gas condenses into liquid at 1 atm pressure = -∆Hvap
Energy is released when a liquid freezes at the melting point:
The energy released when 1 mol of a liquid freezes into a solid at 1 atm pressure = -∆Hfus
Energy is released when a solid deposits from a gas at the sublimation point:
The energy released when 1 mol of a gas sublimes into a solid at 1 atm pressure = -∆Hsub
11/17/2010
7
Phase Diagrams
Show which states exist for a closed system as a function of temperature and pressure
supercritical fluid region
The Phase Diagram of Water
Tm = normal melting point (0 C, 1 atm)
Tb = normal boiling point (100 C, 1 atm)
T3, P3 = triple point where all three states can coexist simultaneously
(0.0098 C, 0.0060 atm)Tc = critical temperature above which the vapor cannot be
liquefied at any pressure (374 C)Pc = critical pressure required to produce liquefaction at the critical
temperature (218 atm)Tc, Pc = critical point (374 C, 218 atm)
Supercritical Fluid Region
At temperatures and pressures higher than the critical point, a substances exists as a supercritical fluid, which has properties in
between that of a liquid and a gas
For example, it can diffuse through solids like a gas, and dissolve materials like a liquid!
Other Features
The solid/liquid boundary line has a negative slope indicating that the melting point of water decreases as the external pressure
increases due to ice being less dense than water
A liquid boils at the temperature where the vapor pressure of the liquid equals the external pressure so as we go higher in altitude,
the external pressure goes down and so does the boiling point
The Phase Diagram of Carbon Dioxide
The solid/liquid line has a positive slope since the solid is denser than
the liquid
T3, P3 = triple point (-56.6 C, 5.1 atm)
Tc, Pc = critical point (31 C, 72.8 atm)
At 1 atm the solid sublimes at (-78 C) leading to it commonly being referred to as “dry ice”
11/17/2010
8
Astronomical Applications
Carbon dioxide ice also sublimes on the surface of the planet Mars which has a maximum surface temperature of 20 C, a minimum
surface temperature of -140 C and an average surface temperature of -53 C and a surface pressure of 0.0063 atm
The Planet Mars has Polar Caps like the Earth
11/17/2010
9
The North Polar Cap of Mars during Winter The outer carbon dioxide cap sublimes into the atmosphere during Spring
The Residual North Polar Cap of Water Ice during Summer
Titan the largest moon of Saturn
Has a cold, nitrogen atmosphere containing hydrocarbons (mostly methane, CH4 and ethane, C2H6
Under the surface conditions on Titan (pressure 1.6 atm, temperature 94 K), the methane phase diagram indicates that liquid methane
should be present on the surface
Phase Diagram of CH4
Not to scale!
11/17/2010
10
Liquid Erosion Features Hydrocarbon lakes?
Cassini Orbiter deploying Huygens Probe Parachuted into Titan’s Atmosphere
Jan 14th 2005 – The Surface of Titan
A slushy mixture of water ice and liquid hydrocarbons
Related Documents
