Liquids and solids

They are similar u compared to gases. u They are incompressible. u Their density doesn’t change with temperature. u These similarities are due to the molecules being close together in solids and liquids and far apart in gases u What holds them close together?

Intermolecular forces u Inside molecules (intramolecular) the atoms are bonded to each other. u Intermolecular refers to the forces between the molecules. u These are what hold the molecules together in the condensed states.

Intermolecular forces u Strong covalent bonding ionic bonding u Weak Dipole dipole London dispersion forces u During phase changes the molecules stay intact. u Energy used to overcome forces.

Dipole - Dipole u Remember where the polar definition came from? u Molecules line up in the presence of a electric field. The opposite ends of the dipole can attract each other so the molecules stay close together. u 1% as strong as covalent bonds u Weaker with greater distance. u Small role in gases.

Hydrogen Bonding u Especially strong dipole-dipole forces when H is attached to F, O, or N u These three because- They have high electronegativity. They are small enough to get close. u Effects boiling point.

CH 4 SiH 4 GeH 4 SnH 4 PH 3 NH 3 SbH 3 AsH 3 H2OH2O H2SH2S H 2 Se H 2 Te HF HI HBr HCl Boiling Points 0ºC

Water ++ -- ++

London Dispersion Forces u Non - polar molecules also exert forces on each other. u Otherwise, no solids or liquids. u Electrons are not evenly distributed at every instant in time. u Have an instantaneous dipole. u Induces a dipole in the atom next to it. u Induced dipole- induced dipole interaction.

Example HH HH HH HH ++ ++ HH HH ++ -- ++ 

London Dispersion Forces u Weak, short lived. u Lasts longer at low temperature. u Eventually long enough to make liquids. u More electrons, more polarizable. u Bigger molecules, higher melting and boiling points. u Much, much weaker than other forces. u Also called Van der Waal’s forces.

Liquids u Many of the properties due to internal attraction of atoms. Beading Surface tension Capillary action u Stronger intermolecular forces cause each of these to increase.

Surface tension u Molecules in the middle are attracted in all directions. u Molecules at the the top are only pulled inside. u Minimizes surface area.

Capillary Action u Liquids spontaneously rise in a narrow tube. u Inter molecular forces are cohesive, connecting like things. u Adhesive forces connect to something else. u Glass is polar. u It attracts water molecules.

Beading u If a polar substance is placed on a non- polar surface. There are cohesive, But no adhesive forces. u And Visa Versa

Viscosity u How much a liquid resists flowing. u Large forces, more viscous. u Large molecules can get tangled up. u Cyclohexane has a lower viscosity than hexane. u Because it is a circle- more compact.

How much of these? u Stronger forces, bigger effect. Hydrogen bonding Polar bonding LDF

Model u Can’t see molecules so picture them as- u In motion but attracted to each other u With regions arranged like solids but with higher disorder. with fewer holes than a gas. Highly dynamic, regions changing between types.

Phases u The phase of a substance is determined by three things. u The temperature. u The pressure. u The strength of intermolecular forces.

Solids u Two major types. u Amorphous- those with much disorder in their structure. u Crystalline- have a regular arrangement of components in their structure.

Crystals u Lattice- a three dimensional grid that describes the locations of the pieces in a crystalline solid. u Unit Cell-The smallest repeating unit in of the lattice. u Three common types.

Cubic

Body-Centered Cubic

Face-Centered Cubic

Solids u There are many amorphous solids. u Like glass. u We tend to focus on crystalline solids. u two types. u Ionic solids have ions at the lattice points. u Molecular solids have molecules. u Sugar vs. Salt.

The book drones on about u Using diffraction patterns to identify crystal structures. u Talks about metals and the closest packing model. u It is interesting, but trivial. u We need to focus on metallic bonding. u Why do metal atoms stay together. u How there bonding effect their properties.

Metallic bonding 1s 2s 2p 3s 3p Filled Molecular Orbitals Empty Molecular Orbitals Magnesium Atoms

Filled Molecular Orbitals Empty Molecular Orbitals The 1s, 2s, and 2p electrons are close to nucleus, so they are not able to move around. 1s 2s 2p 3s 3p Magnesium Atoms

Filled Molecular Orbitals Empty Molecular Orbitals 1s 2s 2p 3s 3p Magnesium Atoms The 3s and 3p orbitals overlap and form molecular orbitals.

Filled Molecular Orbitals Empty Molecular Orbitals 1s 2s 2p 3s 3p Magnesium Atoms Electrons in these energy level can travel freely throughout the crystal.

Filled Molecular Orbitals Empty Molecular Orbitals 1s 2s 2p 3s 3p Magnesium Atoms This makes metals conductors Malleable because the bonds are flexible.

Carbon- A Special Atomic Solid u There are three types of solid carbon. u Amorphous- coal uninteresting. u Diamond- hardest natural substance on earth, insulates both heat and electricity. u Graphite- slippery, conducts electricity. u How the atoms in these network solids are connected explains why.

Diamond- each Carbon is sp3 hybridized, connected to four other carbons. u Carbon atoms are locked into tetrahedral shape.  Strong  bonds give the huge molecule its hardness.

Why is it an insulator? Empty MOsFilled MOs E The space between orbitals make it impossible for electrons to move around

u Each carbon is connected to three other carbons and sp 2 hybridized. u The molecule is flat with 120º angles in fused 6 member rings.  The  bonds extend above and below the plane. Graphite is different.

This  bond overlap forms a huge  bonding network. u Electrons are free to move through out these delocalized orbitals. u The layers slide by each other.

Molecular solids. u Molecules occupy the corners of the lattices. u Different molecules have different forces between them. u These forces depend on the size of the molecule. u They also depend on the strength and nature of dipole moments.

Those without dipoles. u Most are gases at 25ºC. u The only forces are London Dispersion Forces. u These depend on size of atom. u Large molecules (such as I 2 ) can be solids even without dipoles.

Those with dipoles. u Dipole-dipole forces are generally stronger than L.D.F. u Hydrogen bonding is stronger than Dipole-dipole forces. u No matter how strong the intermolecular force, it is always much, much weaker than the forces in bonds. u Stronger forces lead to higher melting and freezing points.

Water is special H O H   -- u Each molecule has two polar O-H bonds.

Water is special H O H   u Each molecule has two polar O-H bonds. u Each molecule has two lone pair on its oxygen.

Water is special u Each molecule has two polar O-H bonds. u Each molecule has two lone pair on its oxygen. u Each oxygen can interact with 4 hydrogen atoms. H O H  

Water is special H O H   H O H   H O H   u This gives water an especially high melting and boiling point.

Ionic Solids u The extremes in dipole dipole forces- atoms are actually held together by opposite charges. u Huge melting and boiling points. u Atoms are locked in lattice so hard and brittle. u Every electron is accounted for so they are poor conductors-good insulators.

Vapor Pressure u Vaporization - change from liquid to gas at boiling point. u Evaporation - change from liquid to gas below boiling point  Heat (or Enthalpy) of Vaporization (  H vap )- the energy required to vaporize 1 mol at 1 atm.

u Vaporization is an endothermic process - it requires heat. u Energy is required to overcome intermolecular forces. u Responsible for cool earth. u Why we sweat. (Never let them see you.)

Condensation u Change from gas to liquid. u Achieves a dynamic equilibrium with vaporization in a closed system. u What is a closed system? u A closed system means matter can’t go in or out. u Put a cork in it. u What the heck is a “dynamic equilibrium?”

Dynamic equilibrium u When first sealed the molecules gradually escape the surface of the liquid.

Dynamic equilibrium u When first sealed the molecules gradually escape the surface of the liquid. u As the molecules build up above the liquid some condense back to a liquid.

Dynamic equilibrium u When first sealed the molecules gradually escape the surface of the liquid. u As the molecules build up above the liquid some condense back to a liquid. u As time goes by the rate of vaporization remains constant but the rate of condensation increases because there are more molecules to condense.

Dynamic equilibrium u When first sealed the molecules gradually escape the surface of the liquid u As the molecules build up above the liquid some condense back to a liquid. u As time goes by the rate of vaporization remains constant but the rate of condensation increases because there are more molecules to condense. u Equilibrium is reached when

Rate of Vaporization = Rate of Condensation u Molecules are constantly changing phase “Dynamic” u The total amount of liquid and vapor remains constant “Equilibrium” Dynamic equilibrium

Vapor pressure u The pressure above the liquid at equilibrium. u Liquids with high vapor pressures evaporate easily. They are called volatile. u Decreases with increasing intermolecular forces. Bigger molecules (bigger LDF) More polar molecules (dipole-dipole)

Vapor pressure u Increases with increasing temperature. u Easily measured in a barometer.

Dish of Hg Vacuum P atm = 760 torr A barometer will hold a column of mercury 760 mm high at one atm

Dish of Hg Vacuum P atm = 760 torr A barometer will hold a column of mercury 760 mm high at one atm. If we inject a volatile liquid in the barometer it will rise to the top of the mercury.

Dish of Hg P atm = 760 torr A barometer will hold a column of mercury 760 mm high at one atm. If we inject a volatile liquid in the barometer it will rise to the top of the mercury. There it will vaporize and push the column of mercury down. Water

Dish of Hg 736 mm Hg Water Vapor u The mercury is pushed down by the vapor pressure. u P atm = P Hg + P vap u P atm - P Hg = P vap u = 24 torr

Temperature Effect Kinetic energy # of molecules T1T1 Energy needed to overcome intermolecular forces

Kinetic energy # of molecules T1T1 Energy needed to overcome intermolecular forces T1T1 T2T2 u At higher temperature more molecules have enough energy - higher vapor pressure. Energy needed to overcome intermolecular forces

Mathematical relationship u ln is the natural logarithm ln = Log base e e = Euler’s number an irrational number like    H vap is the heat of vaporization in J/mol

u R = J/K mol. u Surprisingly this is the graph of a straight line. (actually the proof is in the book) Mathematical relationship

Changes of state u The graph of temperature versus heat applied is called a heating curve. u The temperature a solid turns to a liquid is the melting point.  The energy required to accomplish this change is called the Heat (or Enthalpy) of Fusion  H fus

Heating Curve for Water Ice Water and Ice Water Water and Steam Steam

Heating Curve for Water Heat of Fusion Heat of Vaporization Slope is Heat Capacity

Melting Point u Melting point is determined by the vapor pressure of the solid and the liquid. u At the melting point the vapor pressure of the solid = vapor pressure of the liquid

Solid Water Liquid Water Water Vapor Vapor

Solid Water Liquid Water Water Vapor Vapor u If the vapor pressure of the solid is higher than that of the liquid the solid will release molecules to achieve equilibrium.

Solid Water Liquid Water Water Vapor Vapor u While the molecules of condense to a liquid.

u This can only happen if the temperature is above the freezing point since solid is turning to liquid. Solid Water Liquid Water Water Vapor Vapor

u If the vapor pressure of the liquid is higher than that of the solid, the liquid will release molecules to achieve equilibrium. Solid Water Liquid Water Water Vapor Vapor

Solid Water Liquid Water Water Vapor Vapor u While the molecules condense to a solid.

u The temperature must be above the freezing point since the liquid is turning to a solid. Solid Water Liquid Water Water Vapor Vapor

u If the vapor pressure of the solid and liquid are equal, the solid and liquid are vaporizing and condensing at the same rate. The Melting point. Solid Water Liquid Water Water Vapor Vapor

Boiling Point u Reached when the vapor pressure equals the external pressure. u Normal boiling point is the boiling point at 1 atm pressure. u Super heating - Heating above the boiling point. u Supercooling - Cooling below the freezing point.

Phase Diagrams. u A plot of temperature versus pressure for a closed system, with lines to indicate where there is a phase change.

Temperature Solid Liquid Gas 1 Atm A A B B C C D D D Pressure D

Solid Liquid Gas Triple Point Critical Point Temperature Pressure

Solid Liquid Gas u This is the phase diagram for water. u The density of liquid water is higer than solid water. Temperature Pressure

Solid Liquid Gas 1 Atm u This is the phase diagram for CO 2 u The solid is more dense than the liquid u The solid sublimes at 1 atm. Temperature Pressure