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Intermolecular Forces, Liquids, and
SolidsChapter 11 – Brown & LeMay
Temperature Review Measure of kinetic energy What can you say about the KE of salt
particles, water molecules, and oxygen particles at room temperature?
State determined by strength of forces that keep particles together
Strength Compare energy needed for phase change
vs. decomposition in HCl(l) Intermolecular (called weak) because they
are weaker than ionic or covalent Boiling point reflects strength of bonds in
liquid Melting point reflects strength of bonds in
solids
Kinds of Intermolecular Forces Three major kinds: dipole-dipole, London
dispersion, and hydrogen bonding In solutions, ion-dipole All are electrostatic in nature Approximately 15% of covalent or ionic
strength
Ion - dipole When? Ionic solid + polar liquid Increases with increasing charge of ion or
polarity of solvent Determines solubility
Dipole-Dipole forces Weaker than previous + end of one attracts – end of another If size is equal, more polar has stronger
dipole attractions. (NH3 vs H2O) If polarity is the same but masses differ,
than smallest is stronger. (Able to orient better)
London Dispersion Forces All molecules have this Only attraction in nonpolar molecules How can Iodine be a solid? Temporary lopsided charge builds up from
random motion of electrons - 1930 Increases with mass – we say it has greater
polarizability Straight molecule is more polarizable than a
curled up molecule – why? Halogen Family is a great essay
Hydrogen Bond Strongest of all “weak” forces Is caused when H is bonded to F, O, or N These are so electronegative that the H is a
“naked nucleus” or bare proton Very attractive! Will bond to nearby electron pairs
Importance of Hydrogen Bonding Biological systems
– DNA, proteins Water chemistry
(MP, BP, specific heat, surface tension)
Density of ice
Density Most solids are more dense than
liquid Water is less dense because of
hydrogen bonding At 4°C, water becomes less
dense Important for life in winter Causes lake turnover Alum example
Practice Look at Flow Chart
Properties of Liquids Viscosity “Slower than….. Resistance of a liquid to flow Time it as it goes through a small tube with gravity
acting upon it. Poise – 1g/cm-s Trends – same substance – decreases with
increasing temperatureseries (same structure) – increases with increasing mass
Surface Tension How many drops on a penny? Uneven forces at surface Acts like pond scum Definition – energy needed to increase the surface area of
a liquid by a certain amount Water is high – why?
Called “cohesive” force – together Water moving up a stem – adhesive force Capillary acion – rise up a thin tube Meniscus!
Phase Changes Solid to Liquid is called Heat of Fusion Hfus
For water, 6 kJ/mol
Liquid to Gas is called Heat of Vaporization
Hvap
For water, 40.7 kJ/mol
Hsub is sum of each
Heating Curve Try a problem Remember - flat during phase change,
temperature change when heating a single phase
Cooling is opposite
Supercooling Happens with some liquids - remove heat
and it doesn’t freeze when it should Very unstable May happen during hibernation
Critical Temperature Highest temperature at which a liquid can
form from a gas when pressure is applied. Above this, the substance is called a
supercritical fluid. Gas just becomes more compressed. Critical pressure - pressure at the critical
temperature
Vapor pressure Vapor pressure forms above any liquid if
container is closed – why? Equilibrium is reached This is vapor pressure Higher if forces holding liquid together are
weak - called a volatile (fleeing) liquid
Boiling Point Temperature at which the VP equals
atmospheric pressure Normal BP - boiling point at 1 atm Everest? Autoclave?
Phase Diagram
Handout Look at lines Look at slope
of AB Freeze-drying
- library book example
Water vs. CO2
Structure of Solids Amorphous (rubber, plastics) - large or
mixtures - no true structure Crystalline - highly ordered structure Crystalline solids have true melting points
Unit Cell Repeating unit of a
solid 7 types – (6-sided
parallelograms) Ni, Na, NaCl Array of points in the
crystal lattice
3 cubic unit cells
Total Atoms for each unit cell
Packing Spheres naturally pack hexagonally Animation
Bonding Shown by x-ray diffraction Molecular - low MP If unit packs well, mp can be high Covalent Network Solid - very strong Many covalent bonds in 3-D Diamond, graphite, SiO2, SiC, BN Ionic - greater charge, greater MP Metallic solids - hexagonal close packed, mp
varies