close

Вход

Забыли?

вход по аккаунту

?

Chapter 11: Intermolecular Forces

код для вставкиСкачать
Intermolecular Forces, Liquids
and Solids
CHAPTER 11
CHEM 160
A Molecular Comparison
of Liquids and Solids
• Physical properties of substances understood in terms of
kinetic molecular theory:
– Gases are highly compressible, assumes shape and volume of
container:
• Gas molecules are far apart and do not interact much with each
other.
– Liquids are almost incompressible, assume the shape but not the
volume of container:
• Liquids molecules are held closer together than gas molecules, but
not so rigidly that the molecules cannot slide past each other.
A Molecular Comparison
of Liquids and Solids
– Solids are incompressible and have a definite shape and
volume:
• Solid molecules are packed closely together. The molecules are so
rigidly packed that they cannot easily slide past each other.
A Molecular Comparison
of Liquids and Solids
A Molecular Comparison
of Liquids and Solids
A Molecular Comparison
of Liquids and Solids
• Converting a gas into a liquid or solid requires the
molecules to get closer to each other:
– cool or compress.
• Converting a solid into a liquid or gas requires the
molecules to move further apart:
– heat or reduce pressure.
• The forces holding solids and liquids together are called
intermolecular forces.
Intermolecular Forces
• The covalent bond holding a molecule together is an
intramolecular forces.
• The attraction between molecules is an intermolecular
force.
• Intermolecular forces are much weaker than
intramolecular forces (e.g. 16 kJ/mol vs. 431 kJ/mol for
HCl).
• When a substance melts or boils the intermolecular forces
are broken (not the covalent bonds).
Intermolecular Forces
Intermolecular Forces
Ion-Dipole Forces
• Interaction between an ion and a dipole (e.g. water).
• Strongest of all intermolecular forces.
Intermolecular Forces
•
•
•
•
•
Dipole-Dipole Forces
Dipole-dipole forces exist between neutral polar
molecules.
Polar molecules need to be close together.
Weaker than ion-dipole forces.
There is a mix of attractive and repulsive dipole-dipole
forces as the molecules tumble.
If two molecules have about the same mass and size, then
dipole-dipole forces increase with increasing polarity.
Intermolecular Forces
Dipole-Dipole Forces
Intermolecular Forces
Dipole-Dipole Forces
Intermolecular Forces
•
•
•
•
•
London Dispersion Forces
Weakest of all intermolecular forces.
It is possible for two adjacent neutral molecules to affect
each other.
The nucleus of one molecule (or atom) attracts the
electrons of the adjacent molecule (or atom).
For an instant, the electron clouds become distorted.
In that instant a dipole is formed (called an instantaneous
dipole).
Intermolecular Forces
London Dispersion Forces
• One instantaneous dipole can induce another
instantaneous dipole in an adjacent molecule (or atom).
• The forces between instantaneous dipoles are called
London dispersion forces.
Intermolecular Forces
•
•
•
•
•
London Dispersion Forces
Polarizability is the ease with which an electron cloud
can be deformed.
The larger the molecule (the greater the number of
electrons) the more polarizable.
London dispersion forces increase as molecular weight
increases.
London dispersion forces exist between all molecules.
London dispersion forces depend on the shape of the
molecule.
Intermolecular Forces
London Dispersion Forces
• The greater the surface area available for contact, the
greater the dispersion forces.
• London dispersion forces between spherical molecules
are lower than between sausage-like molecules.
Intermolecular Forces
London Dispersion
Forces
Intermolecular Forces
London Dispersion Forces
Intermolecular Forces
Hydrogen Bonding
• Special case of dipole-dipole forces.
• By experiments: boiling points of compounds with H-F,
H-O, and H-N bonds are abnormally high.
• Intermolecular forces are abnormally strong.
Intermolecular Forces
Hydrogen Bonding
• H-bonding requires H bonded to an electronegative
element (most important for compounds of F, O, and N).
– Electrons in the H-X (X = electronegative element) lie much
closer to X than H.
– H has only one electron, so in the H-X bond, the + H presents
an almost bare proton to the пЃ¤- X.
– Therefore, H-bonds are strong.
Hydrogen Bonding
Hydrogen Bonding
Intermolecular Forces
Hydrogen Bonding
• Hydrogen bonds are responsible for:
– Ice Floating
•
•
•
•
•
•
•
•
Solids are usually more closely packed than liquids;
Therefore, solids are more dense than liquids.
Ice is ordered with an open structure to optimize H-bonding.
Therefore, ice is less dense than water.
In water the H-O bond length is 1.0 Г….
The O…H hydrogen bond length is 1.8 Å.
Ice has waters arranged in an open, regular hexagon.
Each пЃ¤+ H points towards a lone pair on O.
Intermolecular Forces
Hydrogen Bonding
Intermolecular Forces
Some Properties of Liquids
Viscosity
• Viscosity is the resistance of a liquid to flow.
• A liquid flows by sliding molecules over each other.
• The stronger the intermolecular forces, the higher the
viscosity.
Surface Tension
• Bulk molecules (those in the liquid) are equally attracted
to their neighbors.
Some Properties of Liquids
Viscosity
Surface Tension
Some Properties of Liquids
Surface Tension
• Surface molecules are only attracted inwards towards the
bulk molecules.
– Therefore, surface molecules are packed more closely than bulk
molecules.
• Surface tension is the amount of energy required to
increase the surface area of a liquid.
• Cohesive forces bind molecules to each other.
• Adhesive forces bind molecules to a surface.
Some Properties of Liquids
Surface Tension
• Meniscus is the shape of the liquid surface.
– If adhesive forces are greater than cohesive forces, the liquid
surface is attracted to its container more than the bulk
molecules. Therefore, the meniscus is U-shaped (e.g. water in
glass).
– If cohesive forces are greater than adhesive forces, the meniscus
is curved downwards.
• Capillary Action: When a narrow glass tube is placed in
water, the meniscus pulls the water up the tube.
Phase Changes
• Surface molecules are only attracted inwards towards the
bulk molecules.
• Sublimation: solid  gas.
• Vaporization: liquid  gas.
• Melting or fusion: solid  liquid.
• Deposition: gas  solid.
• Condensation: gas  liquid.
• Freezing: liquid  solid.
Phase Changes
Phase Changes
•
•
•
•
•
•
Energy Changes Accompanying
Phase Changes
Sublimation: пЃ„Hsub > 0 (endothermic).
Vaporization: пЃ„Hvap > 0 (endothermic).
Melting or Fusion: пЃ„Hfus > 0 (endothermic).
Deposition: пЃ„Hdep < 0 (exothermic).
Condensation: пЃ„Hcon < 0 (exothermic).
Freezing: пЃ„Hfre < 0 (exothermic).
Phase Changes
Energy Changes Accompanying
Phase Changes
• Generally heat of fusion (enthalpy of fusion) is less than
heat of vaporization:
– it takes more energy to completely separate molecules, than
partially separate them.
Phase Changes
Phase Changes
Energy Changes Accompanying
Phase Changes
• All phase changes are possible under the right conditions.
• The sequence
heat solid п‚® melt п‚® heat liquid п‚® boil п‚® heat gas
is endothermic.
• The sequence
cool gas п‚® condense п‚® cool liquid п‚® freeze п‚® cool solid
is exothermic.
Phase Changes
Heating Curves
• Plot of temperature change versus heat added is a heating
curve.
• During a phase change, adding heat causes no
temperature change.
– These points are used to calculate Hfus and Hvap.
• Supercooling: When a liquid is cooled below its melting
point and it still remains a liquid.
• Achieved by keeping the temperature low and increasing
kinetic energy to break intermolecular forces.
Phase Changes
Critical Temperature and Pressure
• Gases liquefied by increasing pressure at some
temperature.
• Critical temperature: the minimum temperature for
liquefaction of a gas using pressure.
• Critical pressure: pressure required for liquefaction.
Phase Changes
Critical Temperature and Pressure
Vapor Pressure
•
•
•
•
Explaining Vapor Pressure on the
Molecular Level
Some of the molecules on the surface of a liquid have
enough energy to escape the attraction of the bulk liquid.
These molecules move into the gas phase.
As the number of molecules in the gas phase increases,
some of the gas phase molecules strike the surface and
return to the liquid.
After some time the pressure of the gas will be constant
at the vapor pressure.
Vapor Pressure
Explaining Vapor Pressure on the
Molecular Level
Vapor Pressure
•
•
•
•
Explaining Vapor Pressure on the
Molecular Level
Dynamic Equilibrium: the point when as many molecules
escape the surface as strike the surface.
Vapor pressure is the pressure exerted when the liquid
and vapor are in dynamic equilibrium.
Volatility, Vapor Pressure, and Temperature
If equilibrium is never established then the liquid
evaporates.
Volatile substances evaporate rapidly.
Vapor Pressure
Volatility, Vapor Pressure, and
Temperature
• The higher the temperature, the higher the average kinetic
energy, the faster the liquid evaporates.
Vapor Pressure
Volatility, Vapor Pressure, and Temperature
Vapor Pressure
Vapor Pressure and Boiling Point
• Liquids boil when the external pressure equals the vapor
pressure.
• Temperature of boiling point increases as pressure
increases.
Vapor Pressure
Vapor Pressure and Boiling Point
• Two ways to get a liquid to boil: increase temperature or
decrease pressure.
– Pressure cookers operate at high pressure. At high pressure the
boiling point of water is higher than at 1 atm. Therefore, there
is a higher temperature at which the food is cooked, reducing
the cooking time required.
• Normal boiling point is the boiling point at 760 mmHg (1
atm).
Phase Diagrams
• Phase diagram: plot of pressure vs. Temperature
summarizing all equilibria between phases.
• Given a temperature and pressure, phase diagrams tell us
which phase will exist.
• Any temperature and pressure combination not on a
curve represents a single phase.
Phase Diagrams
• Features of a phase diagram:
– Triple point: temperature and pressure at which all three phases
are in equilibrium.
– Vapor-pressure curve: generally as pressure increases,
temperature increases.
– Critical point: critical temperature and pressure for the gas.
– Melting point curve: as pressure increases, the solid phase is
favored if the solid is more dense than the liquid.
– Normal melting point: melting point at 1 atm.
Phase Diagrams
Phase Diagrams
The Phase Diagrams of H2O and CO2
Phase Diagrams
The Phase Diagrams of H2O and CO2
• Water:
– The melting point curve slopes to the left because ice is less
dense than water.
– Triple point occurs at 0.0098C and 4.58 mmHg.
– Normal melting (freezing) point is 0C.
– Normal boiling point is 100C.
– Critical point is 374C and 218 atm.
Phase Diagrams
The Phase Diagrams of H2O and CO2
• Carbon Dioxide:
– Triple point occurs at -56.4C and 5.11 atm.
– Normal sublimation point is -78.5C. (At 1 atm CO2 sublimes
it does not melt.)
– Critical point occurs at 31.1C and 73 atm.
Structures of Solids
•
•
•
•
•
Unit Cells
Crystalline solid: well-ordered, definite arrangements of
molecules, atoms or ions.
Crystals have an ordered, repeated structure.
The smallest repeating unit in a crystal is a unit cell.
Unit cell is the smallest unit with all the symmetry of the
entire crystal.
Three-dimensional stacking of unit cells is the crystal
lattice.
Structures of Solids
Unit Cells
Structures of Solids
Unit Cells
• Three common types of unit cell.
– Primitive cubic, atoms at the corners of a simple cube,
• each atom shared by 8 unit cells;
– Body-centered cubic (bcc), atoms at the corners of a cube plus
one in the center of the body of the cube,
• corner atoms shared by 8 unit cells, center atom completely enclosed
in one unit cell;
– Face-centered cubic (fcc), atoms at the corners of a cube plus
one atom in the center of each face of the cube,
• corner atoms shared by 8 unit cells, face atoms shared by 2 unit cells.
Unit Cells
Unit Cells
Structures of Solids
Unit Cells
Structures of Solids
The Crystal Structure of Sodium
Chloride
• Two equivalent ways of defining unit cell:
– Cl- (larger) ions at the corners of the cell, or
– Na+ (smaller) ions at the corners of the cell.
• The cation to anion ratio in a unit cell is the same for the
crystal. In NaCl each unit cell contains same number of
Na+ and Cl- ions.
• Note the unit cell for CaCl2 needs twice as many Cl- ions
as Ca2+ ions.
Structures of Solids
The Crystal Structure of Sodium
Chloride
Structures of Solids
The Crystal Structure of Sodium
Chloride
Structures of Solids
•
•
•
•
•
Close Packing of Spheres
Solids have maximum intermolecular forces.
Molecules can be modeled by spheres.
Atoms and ions are spheres.
Molecular crystals are formed by close packing of the
molecules.
We rationalize maximum intermolecular force in a crystal
by the close packing of spheres.
Structures of Solids
•
•
•
•
Close Packing of Spheres
When spheres are packed as closely as possible, there are
small spaces between adjacent spheres.
The spaces are called interstitial holes.
A crystal is built up by placing close packed layers of
spheres on top of each other.
There is only one place for the second layer of spheres.
Structures of Solids
Close Packing of Spheres
• There are two choices for the third layer of spheres:
– Third layer eclipses the first (ABAB arrangement). This is
called hexagonal close packing (hcp);
– Third layer is in a different position relative to the first
(ABCABC arrangement). This is called cubic close packing
(ccp).
Structures of Solids
Close Packing of Spheres
Structures of Solids
•
•
•
•
Close Packing of Spheres
Each sphere is surrounded by 12 other spheres (6 in one
plane, 3 above and 3 below).
Coordination number: the number of spheres directly
surrounding a central sphere.
Hexagonal and cubic close packing are different from the
cubic unit cells.
If unequally sized spheres are used, the smaller spheres
are placed in the interstitial holes.
Bonding in Solids
• There are four types of solid:
– Molecular (formed from molecules) - usually soft with low
melting points and poor conductivity.
– Covalent network (formed from atoms) - very hard with very
high melting points and poor conductivity.
– Ions (formed form ions) - hard, brittle, high melting points and
poor conductivity.
– Metallic (formed from metal atoms) - soft or hard, high melting
points, good conductivity, malleable and ductile.
Bonding in Solids
Bonding in Solids
•
•
•
•
Molecular Solids
Intermolecular forces: dipole-dipole, London dispersion
and H-bonds.
Weak intermolecular forces give rise to low melting
points.
Room temperature gases and liquids usually form
molecular solids and low temperature.
Efficient packing of molecules is important (since they
are not regular spheres).
Bonding in Solids
•
•
•
•
Covalent-Network Solids
Intermolecular forces: dipole-dipole, London dispersion
and H-bonds.
Atoms held together in large networks.
Examples: diamond, graphite, quartz (SiO2), silicon
carbide (SiC), and boron nitride (BN).
In diamond:
– each C atom has a coordination number of 4; each C atom is
tetrahedral; there is a three-dimensional array of atoms.
– Diamond is hard, and has a high melting point (3550 C).
Bonding in Solids
Covalent-Network Solids
Bonding in Solids
Covalent-Network Solids
• In graphite
– each C atom is arranged in a planar hexagonal ring;
– layers of interconnected rings are placed on top of each other;
– the distance between C atoms is close to benzene (1.42 Å vs.
1.395 Г… in benzene);
– the distance between layers is large (3.41 Å);
– electrons move in delocalized orbitals (good conductor).
Bonding in Solids
Ionic Solids
• Ions (spherical) held together by electrostatic forces of
attraction.
• There are some simple classifications for ionic lattice
types.
Ionic Solids
Bonding in Solids
Ionic Solids
• NaCl Structure
• Each ion has a coordination number of 6.
• Face-centered cubic lattice.
• Cation to anion ratio is 1:1.
• Examples: LiF, KCl, AgCl and CaO.
• CsCl Structure
• Cs+ has a coordination number of 8.
• Different from the NaCl structure (Cs+ is larger than Na+).
• Cation to anion ratio is 1:1.
Bonding in Solids
Ionic Solids
• Zinc Blende Structure
•
•
•
•
•
Typical example ZnS.
S2- ions adopt a fcc arrangement.
Zn2+ ions have a coordination number of 4.
The S2- ions are placed in a tetrahedron around the Zn2+ ions.
Example: CuCl.
Bonding in Solids
Ionic Solids
• Fluorite Structure
•
•
•
•
Typical example CaF2.
Ca2+ ions in a fcc arrangement.
There are twice as many F- per Ca2+ ions in each unit cell.
Examples: BaCl2, PbF2.
Bonding in Solids
•
•
•
•
•
Metallic Solids
Metallic solids have metal atoms in hcp, fcc or bcc
arrangements.
Coordination number for each atom is either 8 or 12.
Problem: the bonding is too strong for London dispersion
and there are not enough electrons for covalent bonds.
Resolution: the metal nuclei float in a sea of electrons.
Metals conduct because the electrons are delocalized and
are mobile.
End of Chapter 11
Intermolecular Forces, Liquids
and Solids
Документ
Категория
Презентации
Просмотров
13
Размер файла
6 514 Кб
Теги
1/--страниц
Пожаловаться на содержимое документа