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Chapter 11 – Intermolecular Forces, Liquids and Solids
Homework:13, 16, 18, 19, 23, 43, 45, 47,
48, 49, 50, 51, 54, 55, 56
11.2 – Intermolecular Forces Strengths of intermolecular forces of
different substances varies Generally weaker than ionic or covalent
bonds Less energy needed to vaporize (or
evaporate) a liquid or melt a solid than to break the covalent bonds in molecules
In other words: Compound stay intact when melting or boiling, just breaking intermolecular forces
Many properties of liquids reflect strength of the intermolecular forces Such as boiling point Example:
Because the forces between HCl molecules are so weak, HCL boils at very low temperature
-85ºC at atmospheric pressure
Boiling/Melting A liquid boils when bubbles of its vapor
form within the liquid Molecules in the liquid must overcome
their attractive forces to do vaporize The stronger the forces, the higher the
temperature at which the liquid boils Same general principle applies to
melting
Types of Intermolecular Forces Three types exist between neutral
molecules dipole-dipole forces London dispersion forces hydrogen-bonding forces
a.k.a. van der Waals forces There is one other force that mostly
applies to solutions ion-dipole force
All of these forces tend to be less than 15% as strong as covalent or ionic bonds
Dipole-Dipole Forces Neutral POLAR molecules attract
each other when the positive end of one molecule is near the negative end of another These dipole-dipole forces only work
when the polar molecules are very close together
Weaker than ion-dipole forces
In liquids polar molecules are free to move with respect to one another Different configurations create
orientations that are attractive, and orientations that are repulsive.
Ion-Dipole Forces Pretty simple Exists between an ion and a partial
charge on a polar molecule. Increases as either the charge of
the ion increases, or the magnitude of the dipole moment increases
Rules for Dispersion Forces
1. When the molecules of two substances have similar weight and shape, dispersion forces are approximately equal
Differences in the attraction forces are due to differences in strengths of the dipole-dipole forces
More polar molecules have the stronger attractions
2. When the molecules of two substances differ in molecular weights, dispersion forces tend to be decisive in finding which substances has the stronger attractions
Differences in the magnitudes of the forces usually because of differences in atomic weight
More massive molecules having stronger attractions
Hydrogen Bonding Hydrogen bonding is a special
type of intermolecular force Always between a H atom in a polar
bond and an unshared electron pair on a nearby, small, electronegative ion or atom
Usually a H-F, H-O or H-N bond Usually a F, O or N atom in another
molecule
Example A hydrogen bond exists between
the H atom in HF molecule and the F atom of a nearby HF molecule F-HÅÅÅÅF-H dots represent the hydrogen bond
The hydrogen bond is a unique form of a dipole-dipole attraction Because F, N and O are so electronegative,
bonds are VERY polar H on positive end
H has no inner core of electrons Positive side has partially exposed nucleus Gives a large dipole effect (since not just electron
density, but actual nuclear charge) Also, since H is so small, it can approach an
electronegative atom very closely
Hydrogen bonds still weaker than ordinary bonds
Stronger than dipole-dipole or dispersion forces
Comparing Intermolecular Forces Dispersion forces are found in all
substances Strengths increase with increasing
molecular weight Strength increases with longer
molecules
Dipole-dipole forces adds to dispersion forces Found only in polar molecules
Hydrogen bonds require H atoms bonded to F, O or N Also adds to dispersion forces
11.5 – Vapor Pressure Molecules can escape from the surface
of a liquid to the gas phase by evaporation.
Suppose we place an amount of ethanol (C2H5OH) in an evacuated, closed container. Ethanol will quickly begin to evaporate So pressure exerted by the vapor above the
liquid will increase Eventually, the pressure of the vapor will
become a constant value This is called the vapor pressure of the
substance
Explaining Vapor Pressure Molecules of a liquid move at various
speeds At any instant, some of the molecules at
the surface of the liquid get enough kinetic energy to overcome the attractive forces of their neighbors Thus escaping into the gas phase
The weaker the attractive forces, the more particles that can escape, and therefore, more vapor pressure
At any given temperature the movement of molecules from liquid to gas goes on continuously.
As the number of gaseous molecules increase, the probability increases that a molecule in the gas phase will hit the liquid surface and be recaptured by the liquid
Eventually, rate at which molecules return to liquid equals the rate at which they escape
So we get a steady number of molecules in the gas phase
Dynamic Equilibrium The condition when two opposing processes are
occurring at the same time, and at the same rate, is called dynamic equilibrium
Often referred to simply as equilibrium A liquid and vapor are in equilibrium when
evaporation and condensation occurs at equal rates Often appears like nothing is happening, but no net
change But particles are constantly changing from liquid gas
and from gas liquid The vapor pressure of a liquid is the pressure
exerted by its vapor when vapor and liquid are in equilibrium.
Volatility, Vapor Pressure and Temperature Substances with high vapor pressure
evaporate more quickly than substances with low vapor pressure
Liquids that evaporate easily are said to be volatile
As temperature increases, particles move more, and will evaporate more. Vapor pressure will increase as temperature
increases Non-linear progression
Vapor Pressure and Boiling Point A liquid boils when vapor pressure
equals the external pressure acting on the surface of the liquid The temperature at which a liquid boils
increases with increasing external pressure
The boiling point of a liquid at 1 atm pressure is called its normal boiling point
11.6 – Phase Diagrams The equilibrium between liquid and
vapor is not the only dynamic equilibrium we deal with.
Under the right conditions (of pressure and temperature) we can have other dynamic equilibriums solid / liquid solid / vapor
A phase diagram is a graphical way to summarize the conditions which cause the various equilibrium to exist. Allow us to predict the phase of matter
given a temperature and pressure The pressure shown in the diagram is either
the pressure applied to the system or the pressure generated by the substance itself
Three main curves on the graph
1. The line from A to B Equilibrium between gas and
liquid phase The pressure at 1 atm on this
graph represents the normal boiling point of the substance
The curve ends at the critical point
Critical Point? The critical point is the critical
temperature and critical pressure of the substance Critical temperature is the temperature where
beyond it, the substance will never again be a liquid
Intermolecular forces too weak, no matter what the pressure
Critical pressure is the pressure that is needed to create a liquid at the critical temperature
2. The line AC represents the variation in the vapor pressure of the solid as it sublimes at different temperatures
3. The line AD represents the change in melting point of the solid with increasing temperature
Usually slopes right as pressure increases
Because solid is usually denser than liquid
An increase in pressure usually favors the more compact solid phase, so higher temperature needed
Melting and freezing point are the same
The melting point at 1 atm is normal melting point
Point A, where the lines meet, is known as the triple point All three phases are in equilibrium at
this temperature and pressure
Practice Interpreting Phase Diagrams - Water
Carbon Dioxide