*
Chapters 12 and 13
* More complicated than gases…
* particles are close together due to attractive
forces
* these attractive forces are mostly ignored when
dealing with gases
*
* directly related to properties such as melting
point and boiling point (energy needed to
overcome attractions)
* solubility of gases, liquids, and solids in various
solvents
* determines structures of biomolecules such as
DNA and proteins
*
* Forces between particles
* in ionic compounds, the ions are held together by
electrostatic attraction
* in molecular compounds, the intermolecular
forces (forces between particles) are based on
electrostatic attractions that are weaker than
ionic forces
*
* Polar Molecules mixed with Ionic Compounds
* Ions will be attracted to polar ends of molecules.
* Can be used to determine the enthalpy of
solvation (or hydration)
* Ion dipole forces
*
* Polar Molecules mixed with Ionic Compounds
(cont.)
* The force depends on:
* distance between ion and dipole.
* charge on ion.
* magnitude of dipole.
*
* Molecules with Permanent Dipoles
* molecules with dipoles interact by dipole-dipole
attraction
* these attractions influence endothermic
evaporation (ΔHvap) and exothermic condensation
* polar bonds are stronger and require more
energy to break bonds
*
*
*Solubility
*“like dissolves like” means that polar molecules are
more likely to dissolve in polar solvents, and nonpolar
molecules are more likely to dissolve in nonpolar
solvents
* A hydrogen bond is an attraction between the
hydrogen atom of an X-H bond and Y, where X
and Y are atoms of highly electronegative
elements and Y has a lone pair of electrons.
* These bonds are an extreme form of dipole-
dipole in which one atom is always H and the
other is often O, N, or F.
*
* due to large electronegativity differences,
these hydrogen bonds are very polar; partial
charges are formed
* hydrogen atom becomes a bridge between
electronegative elements
*
* WATER!!
* density in the solid state (most dense at about
4oC)
* lake turnover
* high heat capacity
* coastal climates
*
* Nonpolar Molecules
* polar molecules (like water) can induce a dipole in molecules without a
permanent dipole (such as oxygen gas and water)
* the force of attraction is called a dipole/induced dipole interaction
* The process of inducing a dipole is called polarization (a
molecule/atom has a certain polarizability)
* The higher the molar mass, the larger the cloud and greater the
polarizability of the molecule
*
* London Dispersion Forces
* when two nonpolar atoms approach each other,
attractions or repulsions between their electrons
and lead to distortions in their clouds; leading to
intermolecular attraction
* Arise between all molecules
* this force of attraction in nonpolar molecules is
an induced dipole/induced dipole force, or
London dispersion forces
*
* You mix the liquids water, CCl4, and hexane
(CH3CH2CH2CH2CH2CH3). For each pair of
compounds, what type of intermolecular forces
can exist between the compounds? If you mix
these three liquids, describe what observations
you might make.
*
* Decide what type of intermolecular force is
involved in (a) liquid O2, (b) liquid CH3OH, (c)
O2 dissolved in H2O. Place the interactions in
order of increasing strength.
*
* particles interact with each other like a solid,
but there is little order in their arrangement
* vaporization or evaporation is the process in
which a liquid becomes a gas; molecules
escape the liquid surface and enter the
gaseous state
*
* The heat energy required to vaporize a sample
is given as the standard molar enthalpy of
vaporization, ΔHovap. (kJ/mol)
* The opposite process is condensation, in which
a molecule may reenter the liquid phase. This
releases energy, which is why a steam burn is
much worse than one from boiling water!
*
* The molar enthalpy of vaporization of
methanol, CH3OH, is 35.2 kJ/mol at 64.6oC.
How much energy is required to evaporate 1.00
kg of this alcohol?
*
* liquid in a sealed flask will form a dynamic
equilibrium
* when this equilibrium has been established,
the pressure exerted by the vapor is the
equilibrium vapor pressure (a measure of the
tendency of molecules to escape to the vapor
phase)
* volatility
*
* If 0.50 g of pure water is sealed in an
evacuated 5.0L flask, and the whole assembly
is heated to 60oC, will the pressure be equal to
or less than the equilibrium vapor pressure of
water at this temperature? What if you use 2.0
g of water? Under either set of conditions is
any liquid water left in the flask, or does it all
evaporate?
*
* Relates equilibrium vapor pressure to the molar
enthalpy of vaporization at a specified
temperature
* You can calculate ΔHvapo for a liquid using the
Clausius-Clapeyron equation
ln(P2/P1) = -ΔHvapo/R[1/T2 – 1/T1]
*
* boiling point is the temperature at which a
liquid’s vapor pressure is equal to the external
pressure; at standard pressure this point is
called the normal boiling point
* Water boils at a lower temperature at higher
altitudes…why?
*
* When the interface between the liquid and
vapor disappears, this point is called the
critical point and has a Tc and Pc.
* At this point, the substance is called a
supercritical fluid, meaning that it is like a gas
under high pressure so that its density is like a
liquid but its viscosity is like a gas.
* supercritical CO2 used to decaffeinate coffee
*
* surface molecules are attracted to molecules below
them, makes the liquid behave as if it had a skin –
toughness of that “skin” is surface tension (water
striders)
* molecules may be attracted to adhesive forces
between two different substances in such a way
that overcomes the cohesive forces between the
molecules themselves – capillary action (paper)
* viscosity – resistance to flow (honey vs. water)
*
* After reading sections 12.1-12.4, you should be
able to do the following…
* P. 581 (2-24 even)
*
* molecules, atoms, or ions cannot move (although
they vibrate or rotate some)
* solids have regular, repeating patterns of atoms
or molecules within the structure
* attractive forces are maximized and repulsive
forces are minimized
* the unit cell within a crystalline solid is the
smallest repeating unit (such as a “repeat” in a
wallpaper pattern)
*
* a crystal lattice refers to a bunch of unit cells
all put together
* The lattice points defining each unit cell in solids
represent identical environments for the ions,
atoms, or molecules.
*
* Ionic compounds have high melting points due
to the strength of bonding in the lattice.
* Lattice energy is a measure of the strength of
ionic bonding
* Measured as lattice enthalpy!
*
* Using the ΔHf of compounds and enthalpy for ion formation in
gas phase, you can calculate the lattice enthalpy of a
compound.
* Add up the steps:
* Formation of solid sodium chloride = formation of each element
as a gas plus the formation of ions plus the lattice enthalpy
*
*
* low vapor pressure due to strong interactions
of positive and negative ions
* brittle due to repulsion of like charges caused
when one layer slides across another layer
* do not conduct electricity unless melted or in
solution
* do not dissolve in nonpolar solvents
*
* positive kernels consisting of nucleus and inner
electrons surrounded by a sea of mobile
valence electrons
* good conductors
* malleable and ductile
* pure substances or mixtures (alloys)
* interstitial alloys (steel)
* substitutional alloys (brass)
*
* Only formed from nonmetals: elemental
(diamond, graphite) or two nonmetals (silicon
dioxide, silicon carbide)
* Covalent network solids have high melting
points
* Generally for in the carbon group due to their
ability to form four covalent bonds
* Graphite is an allotrope of carbon that forms
sheets; high melting point due to covalent bonds,
but soft due to LDF layers
*
* Nonmetals, diatomic elements, two or more
nonmetals
* Nonconductors
* Low melting points due to weak IMF
* Sometimes very large molecules or polymers
*
* the melting point is when the lattice collapses
and the solid is converted to liquid
* melting requires energy; enthalpy of fusion
(ΔHfus)
* ionic compounds have higher lattice energies
and therefore higher melting points
*
*
*Molecules can escape directly from the solid to
the gas phase by sublimation, which is
endothermic.
*frost
* After reading sections 13.1-13.6, you should be
able to do the following…
* P. 611 (14-18 even, 22)
*