Intermolecular forces (IMFs) CONDENSED STATES OF MATTER

Intermolecular forces (IMFs) CONDENSED STATES OF MATTER

States of Matter: - composed of particles packed closely together with little space between them. Solids maintain a. - any substance that flows. (A fluid) - particles are free to slide past one another and continual change their positions. Particles are in - are fluids composed of particles in. Gases are not touching most of the time.

Phase Diagrams show the phase of matter at a variety of P and T. Substances can be almost any phase, given the right P and T.

H2O vs. CO2

is less dense in the solid state (ice has lots of space in it), water has a negative slope between the solid & liquid on the phase diagram. Water is densest at 4 o C. goes straight from a solid to a gas at atmospheric pressure.

Intramolecular vs. Intermolecular forces are the forces that hold atoms together in a molecule (within a molecule) Ex. covalent bonds. forces are the attractions that molecules have for one another (molecules next to each other) Ex. H-bonds

Intramolecular vs. Intermolecular Ex. The forces in H 2 O In water, the 2 hydrogens are attached to the oxygen through intramolecular forces. The reason why water molecules are attached to one another (and not separated like a gas) is because of the intermolecular forces: The partially negative O of one H 2 O molecule is attracted to the partially positive H of a different H 2 O molecule.

Kinetic-Molecular Description of Liquids & Solids Solids & liquids are atoms, ions, molecules are close to one another highly incompressible Liquids & gases are easily flow in liquids & solids are strong in gases are weak

IMF types: Ion ion attractions are the strongest. makes up ionic bonding tend to be crystalline solids (hard, but brittle) very high melting points Na + Cl - Na + Cl - Na + Cl - Na + Cl - Na + Cl - Na + Cl - Cl - Na + Cl - Na + Cl - Na + Cl - Na + Cl - Na + Cl - Na + Na + Cl - Na + Cl - Na + Cl - Na + Cl - Na + Cl - Na + Cl - ion-ion attractions

IMF types: Dipole Dipole attractions (fairly strong) happens with polar molecules b/c they have permanent dipole moments H Cl H Cl δ + δ - δ + δ -

IMF types: Hydrogen bonding a special dipole attraction, stronger than normal dipole-dipole attractions - very strong attraction ~ 2 conditions: must have a bonded to a H and at least one lone pair of electrons on N, O, F δ + δ - H.. l H N H δ - :N H δ + δ + l δ + l H H δ + δ +

IMF types: Induced Dipole (the weakest), also called attractions occurs in nonpolar molecules temporary dipole caused by interaction with another molecule boil and melt very easily Increases with molecular size (aka. The heavier the molecule the more they exhibit induced dipoles) δ + δ - F F F F δ - δ +

Mixing of IMF Ions and molecules can interact with each other in a variety of ways. ion-dipole Ions with polar molecules ion-induced dipole Ions with nonpolar molecules dipole-induced dipole Polar and nonpolar molecules

Mixing of IMF : An ion-dipole force is an attractive force that results from the electrostatic attraction between an ion and a molecule that has a dipole. Most commonly found in solutions. Especially important for solutions of ionic compounds in polar liquids. A positive ion (cation) attracts the partially negative end of a neutral polar molecule. A negative ion (anion) attracts the partially positive end of a neutral polar molecule. Ion-dipole attractions become stronger as either the charge on the ion increases, or as the magnitude of the dipole of the polar molecule increases

Ion-Dipole

Ion-Induced Dipole : An ioninduced dipole attraction is a weak attraction that results when the approach of an ion induces a dipole in an atom or in a nonpolar molecule by disturbing the arrangement of electrons in the nonpolar species.

Ion-Induced Dipole

Dipole Induced Dipole : A dipole-induced dipole attraction is a weak attraction that results when a polar molecule induces a dipole in an atom or in a nonpolar molecule by the arrangement of electrons in the nonpolar species.

Dipole Induced Dipole

IMF Flow Chart

Properties of liquids Properties of liquids at constant temperature ~ no definite ~ definite ~ have surface tension, diffuse, medium density, viscosity, evaporation, capillary action, and vapor pressure. of each depends on.

liquids diffuse into one another they are in each other for example: water/alcohol gasoline/motor oil liquids do not diffuse into each other they are in each other for example: water/oil water/cyclohexane

1. Surface tension Surface Tension - measure of the that occur at the surface of a liquid molecules at surface of a liquid are only attracted in a down direction. Denser on top At surface, molecules are attracted downward, thus liquid is denser on top water bugs

2. Viscosity how easily it flows ~ stronger IMF the liquid is ~ geometry of molecule affects viscosity (more complex shapes = more viscous) ~ very long chains more viscous b/c longer chains get tangled C C C C C C C C C C C C C C C more viscous less viscous

3. Capillary action tendency of a liquid to be or by a very narrow tube - Stronger IMF more cohesion ~ when a molecule has attraction for itself, it s called ~ when a molecule has attraction for other molecules, it is capillary rise implies (water) capillary fall implies (mercury)

4. Evaporation when a liquid changes to the vapor phase at a temp. that is less than it s boiling point. Why? If the molecule can gain enough, they break through the liquid and go into the atmosphere

5. Vapor pressure the pressure of a gas that exists over its solid or liquid state. Does not depend on how much liquid/solid you have. P vap depends on the temp. and type of substance. You have vapor pressure as long as there is evaporation of a liquid. Higher Temperature = Higher vapor pressure Boiling occurs when the P vap of liquid = P atm

The Liquid State Vapor Pressure (High = low )

Boiling Points & Distillation ~ All liquids have boiling points: based on IMF. The higher IMF, the higher the normal boiling point. ~ can separate liquids on the basis of their b.p. (distillation) CH 3 OH has a lower boiling point than C 2 H 5 OH, so it changes to a gas first. This means it has lower IMF

Trends in boiling points of Liquids Gas MW BP( o C) He 4-269 Ne 20-246 Ar 40-186 Kr 84-153 Xe 131-107 Rn 222-62 The boiling point increases in response to molecular size

Compound MW(amu) B.P.( o C) CH 4 16-161 C 2 H 6 30-88 C 3 H 8 44-42 n-c 4 H 10 58-0.6 n-c 5 H 12 72 +36 The boiling point increases in response to molecular size

In the Liquid State Compound MW(amu) B.P.( C) HF 20 19.5 HCl 37-85.0 HBr 81-67.0 HI 128-34.0 HF has the highest B.P. b/c of Hydrogen bonding. The rest increases in response to molecular size. o

In the Liquid State Compound MW(amu) B.P.( C) H O 18 100 2 H S 34-61 2 H Se 81-42 2 H Te 130-2 2 Water has the highest B.P. because of Hydrogen bonding. The rest increases b/c of increase in molecular size. o

Various boiling points Arrange the following substances in order of increasing boiling points. C 2 H 6, NH 3, Ar, NaCl, AsH 3

Amorphous & Crystalline Solids Amorphous solids have a well ordered structure. Particles are irregularly arranged so IMF vary in strength within a sample Ex. paraffin, glasses 35 Crystalline solids have well defined structures that consist of extended array of repeating units. Have defined IMF. give X-ray difraction patterns ~ see Bragg equation in book Ex. Ice, salt

Structure of Crystals unit cell - smallest repeating unit of a crystal Ex. bricks are repeating units for buildings 7 basic crystal systems We do not need to learn these 7 just an FYI for your future 36

Types of Solids 4 Types of : Covalent Ionic Metallic Molecular

Covalent Solids Covalent solids: Also known as or individual atoms are covalently bonded to other atoms and those atoms are bonded to other atoms, etc. In a network solid there are no individual molecules and the entire crystal is considered. This makes covalent solids very hard with very high melting points. Most are nonconductors. Formulas for network solids are simple ratios of the component atoms represented by a formula unit (just like ionic compounds)

Examples of Bonding in Solids Covalent Solids atoms that are covalently bonded to one another Very strong bonds and hard to break examples: SiO 2 (sand), diamond, graphite, SiC 39

Ionic Solids Ionic: positive and negative ions arranged in a specific lattice structure. are strong. Also known as a lattice of positive and negative ions held together by electrostatic forces All of the ionic compounds you are familiar with

Examples of Bonding in Solids Ionic Solids ions occupy the unit cell examples: CsCl, NaCl, ZnS 41

Metallic Solids Metallic: metals where each valence electron is thought to belong to the entire structure. It has a closely packed lattice with delocalized electrons throughout. Metals are seen as a positive nuclei with a of electrons. The mobility of electrons helps explain the electrical conductivity of metals. This occurs in all of the metals on the periodic table

Examples of Bonding in Solids Metallic Solids positively charged nuclei surrounded by a sea of electrons 43 positive ions occupy lattice positions examples: Na, Li, Au, Ag,..

Molecular Solids Molecular solids: are solids made up of molecules that are next to each other in unit cells held together by the. The attractive forces between individual molecules are. dipole forces are weaker than covalent or ionic bonds, so molecular solids are soft They are volatile and insulators. Have relatively low melting temperature Simple covalent compounds usually form molecular solids

Examples of Bonding in Solids 45 Molecular Solids molecules occupy unit cells low melting points,volatile & insulators examples: water, sugar, carbon dioxide, benzene

Type of Solid Covalent Network Interactions Properties Examples Covalent bonding High melting point, hard, nonconducting Ionic Ionic High melting point, conductors in molten or aqueous form, brittle, hard Metallic Metallic bonding Conducting, variable hardness and melting point (depends upon strength of metallic bonding) Molecular Hydrogen bonding, Dipole-dipole, London Dispersion Low melting points, (can easily turn into liquids or gases), nonconducting C (diamond) SiO 2 (quartz) NaCl, MgO, etc Fe, Mg, Cu, etc H 2 O, CO 2

Bonding in Solids - Variations in Melting Points Covalent Solids Melt at high temps (most > 1500 o C ) because the attractive forces between the individual particles are very strong. Substance Melting Point ( o C) sand, SiO2 1713 carborundum, SiC ~2700 diamond >3550 graphite 3652-3697

Bonding in Solids - Variations in Melting Points Ionic Solids Melt at fairly high temps b/c the attraction between ions are much stronger than in molecular solids but weaker than in covalent solids. Attractive forces increase as charges on ions increase & their radii decrease. Coulomb s Law Compound Melting Point ( o C) LiF 842 LiCl 614 LiBr 547 LiI 450 CaF2 1360 CaCl2 772 CaBr2 750 CaI2 740

Bonding in Solids - Variations in Melting Points Metallic Solids Melting points vary widely b/c there are large variations in the strengths of metallic bonding. Most metals have fairly high melting points but Mercury is a liquid at room temp. Metal Melting Point ( o C) Na 98 Pb 328 Al 660 Cu 1083 Fe 1535 W 3410

Bonding in Solids - Variations in Melting Points Molecular Solids Have low melting points (most < 300 o C ) because the attractive forces between the molecules are rather weak. Compound Melting Point ( o C) ice 0 ammonia -77.7 benzene, C6H6 5.5 napthalene, C10H8 80.6 benzoic acid, C6H5CO2H 122.4

Alloys are not bonds! An is a physical mixture made up of two or more metals. Ex. Brass, Bronze, pewter, steel Silver and gold used in jewelry is actually an alloy of that element mixed with other metals to make it stronger.

Brief summary Intermolecular attractions from strongest to weakest Ion-Ion: ionic compounds (metal/nonmetal) Hydrogen bonding: H attached to a N, O, or F and lone pair of e - on the central atom Dipole Dipole: polar compounds London Dispersion Forces (induced dipole): all compounds exhibit this, but it is most important with non-polar compounds.

Effects of intermolecular attractions The compound that has the highest boiling pt, melting pt, and heat of vaporization corresponds to the compound with the strongest IMF. The highest vapor pressure corresponds to the lowest intermolecular attractions. If two compounds are nonpolar, the one with the greatest molecular mass has the greater London Forces. If two compounds are ionic, the one with the greatest charge ions has the greater IMF. If same charge, the smallest ions have the greatest IMF. *Coulomb s Law

For many years, the world s record for flying gliders was 60,000 ft. It was set by a Texan who flew into an updraft in front of an approaching storm. The pilot had to fly out of the updraft and land, not because he was out of air (there was still plenty of air in his compressed air bottle) but because he was not wearing a pressurized suit. What would have happened to the pilot s blood if he had continued to fly higher?