Foundations of College Chemistry, 14th Ed.
Morris Hein and Susan Arena
Liquid water provides the basis for our bodies as well as recreational sports like windsurfing.
Liquids: intermediate between gases and solids. Contain particles close to one another but have
fluidity (can assume the shape of a container).
Significant attractive forces exist between particles in a liquid.
© 2014 John Wiley & Sons, Inc. All rights reserved.
Gases Liquids Solids Variable shape and
volume Variable shape, fixed volume Fixed shape and volume May expand or
compress May flow, not compressible Non-compressible crystalline solids Low densities High density High density
Mix to form
Evaporation or Vaporization:
Escape of molecules from the liquid to the gas phase.
Molecules in the liquid state have different kinetic energies (KEs).
Those with higher KEs can overcome attractive forces between particles
and escape to the gas phase.
Liquid Vapor
Evaporation
Phase change from the solid to gas phase that bypasses the liquid state.
Examples
Solid Vapor
CO2 (s) CO2 (g)
liquid vapor
condensation evaporation
In a closed container, an equilibrium develops between molecules evaporating and condensing.
Molecules from the liquid phase can escape to the vapor phase through evaporation.
Molecules in the gas phase can strike the surface of the liquid and return to the liquid phase.
This process is called condensation.
Vapor Pressure
Independent of the quantity of liquid or its surface area.
Increases with increasing temperature.
Depends on the strength of attraction between molecules in the liquid state.
Volatile liquids: very weak attractive forces
between molecules. Evaporate very rapidly at ambient temperature. Have high vapor pressures as a result.
Vapor pressure: pressure exerted by a vapor in equilibrium with its liquid phase.
Resistance of a liquid to an increase in surface area.
Molecules on a liquid surface are strongly attracted by molecules within the liquid.
Surface tension increases with increasing attractive interactions between molecules.
Mercury droplets form spheres due to surface tension.
Surface Tension
•
tendency to minimize surface area
•
Adhesion – Forces that bind a substance to a
surface
Spontaneous rise of a liquid in a narrow tube.
Cohesive forces exist between water molecules in a liquid.
Adhesive forces exist between water molecules and the walls of the container.
When the cohesive forces between molecules are less than the adhesive forces between liquid and container,
Capillary Action in Action
Shape of the meniscus reflects the relative strength of cohesive forces within the liquid and adhesive
forces between the liquid and the tube.
If convex:
adhesive forces < cohesive forces
If concave:
adhesive forces > cohesive forces
Hg H2O
© 2014 John Wiley & Sons, Inc. All rights reserved.
Temperature at which the vapor pressure of a liquid is equal to the external pressure above the liquid.
Where is the boiling point of a liquid higher, at or above sea level?
Normal boiling point:
temperature when the vapor pressure is 1 atm
Normal boiling points: Water: 100 ºC
Ether: 35 ºC
Ethyl Alcohol: 78 ºC
Vapor Pressure Curve
© 2014 John Wiley & Sons, Inc. All rights reserved.
The vapor pressure curve for water is given below. What is the boiling point of water at 300 mmHg?
Freezing/melting point: the temperature at which the solid phase of a substance is in equilibrium
with its liquid phase.
solid
melting
liquid
freezing
While both phases are present, the temperature remains constant.
The energy is used to change the solid to the liquid phase.
Freezing Point or Melting Point
Heat of vaporization
Heat of fusion
Melting
Boiling
Heat of fusion: energy required to change 1 g of a solid at its melting point to a liquid.
Heat of vaporization: energy required to change 1 g of a liquid to vapor at its normal boiling point.
Heat of fusion: energy required to change 1 g of a solid at its melting point to a liquid.
The heat of fusion for water is 335 J/g. Calculate the amount of heat needed to melt 25.0 g of water.
= 8380 J 25.0 g × 335 J
1 g
Use the heat of fusion as a conversion factor!
Energy and Phase Changes
How many kilojoules are released when 31.2 g of water at 0 °C freezes?
335 J/g
22
Heat of Vaporization
The heat of vaporization is the amount of heat
• absorbed to change 1 g of
liquid to gas at the boiling point
• released when 1 g of gas
changes to liquid at the boiling point
boiling point of H2O = 100 °C heat of vaporization (water)
24
Learning Check
How many kilojoules (kJ) are released when 50.0 g of steam from a volcano condenses at 100 °C?
Calculate the energy needed to convert 25.0 g of ice at 0 ºC to steam at 100 ºC?
1. ice melts (total energy = mass x heat of fusion)
2. liquid water is warmed from 0° to 100°C (energy = mass x specific heat x ΔT) 3. water evaporates (energy = mass x heat of vaporization)
Plan
The conversion of ice to steam is a three step process:
The overall energy required for the process is the sum of the 3 steps.
Given: heat of fusion = 335 J/g heat of vaporization = 2259 J/g
specific heat of liquid water = 4.184 J/gºC
Phase Change Practice
100 ºC
Calculate the energy needed to convert 25.0 g of ice at 0 ºC to steam at 100 ºC?
Calculate
= 8375 J 25.0 g × 335 J
1 g
= 10,460 J 4.184 J
1 g ºC
25.0 g × ×
2. warm water
3. evaporate water 1. ice melts
= 56,475 J 25.0 g × 2259 J
1 g Total energy required:
E = Step 1 + Step 2 + Step 3 = 8375 J + 10,460 J + 56475 J = 75,300 J
© 2014 John Wiley & Sons, Inc. All rights reserved.
How many joules of energy are needed to change 10.0 g of ice at 0.00 ºC to water at 20.0 ºC?
Given: heat of fusion = 335 J/g
specific heat of liquid water = 4.184 J/gºC
ANS 4.19 x 103 J
Plan
The conversion of ice to liquid water is a two step process.
Example
•
When a volcano erupts, 175 g of steam at 100
Attractive forces between molecules.
These forces allow for formation of liquids and solids.
The degree of intermolecular forces correlates with a compound’s physical properties.
The stronger the interaction between molecules in a liquid, the higher the boiling point and the lower
the vapor pressure. Example:
Dipole-Dipole Attractions
In covalent molecules, due to different atoms having different electronegativities, molecules are polar.
When polar molecules are put together, they will align to permit interaction between oppositely polarized portions of the molecules.
These interactions between dipoles in different molecules are called dipole-dipole forces.
The interaction of two polar H2O molecules.
Types of Intermolecular Forces
Water has very high melting and boiling points, and heats of fusion and vaporization.
These anomalous properties are due to strong attraction between water molecules due to hydrogen bonding,
a special type of dipole-dipole attraction.
The Hydrogen Bond
Hydrogen bonds: one type of strong intermolecular force/attraction between molecules.
Hydrogen bonds are much weaker than ionic or covalent bonds which are intramolecular forces.
Hydrogen Bonding between H2O molecules.
To form hydrogen bonds, a compound must have covalent bonds between hydrogen and F, O or N
(a very electronegative element).
Can hydrogen bond.
Cannot hydrogen bond.
(No H attached to oxygen).
Which of the following molecules would be expected to participate in hydrogen bonding?
The Hydrogen Bond Practice
Molecules without dipoles can also interact with one another.
These interactions between nonpolar molecules and noble gases are called London dispersion forces.
London forces arise from uneven, instantaneous charge distributions due to electron movement
in nonpolar molecules.
London forces are very weak forces.
Generally become more important as the size of the molecule increases. Larger sizes provide more
possible electrons to provide dipoles.
This instantaneous dipole can then induce a dipole in a neighboring nonpolar molecule, resulting in a small
attraction between particles.
© 2014 John Wiley & Sons, Inc. All rights reserved.
