Chapter 10 Liquids and Solids Problems: 14, 15, 18, 21-23, 29, 31-35, 37, 39, 41, 43, 46, 81-83, 87, 88, 90-93, 99, 104-106, 113
Recall: Intermolecular vs. Intramolecular Forces Intramolecular: bonds between atoms that make molecules. e.g. Covalent, polar covalent and ionic Intermolecular: bonds between molecules holds solids and liquids together
Types of Intermolecular Bonds 1. Dipole Dipole: Fairly strong intermolecular bond. Forms between polar molecules. Typical energy is 5 25 kj to break 1 mole of these bonds. 2. Ion Dipole: Either a cation or an anion bonded to a polar molecule (dipole molecule) Typical energy is 40 600 kj/mole 3. Hydrogen Bond: A special type of dipole dipole bond between: N-H O-H N-X O-X Typical energy is 10 40 kj/mole F-H F-X Why is H-Cl H-Cl not considered a H bond?
Polarizability the ease in which the electron cloud distribution of an atom can be distorted. The larger the cloud (atom) the easier an atom is to polarize. Induced polarity is when a polar structure causes a distortion in another electron cloud 4. Induced Dipole Intermolecular Bond caused by 1 molecule which has a permanent charge separation (dipole) inducing a dipole in another non-polar molecule. 2 Types exist: a) ion induced dipole b) dipole induced dipole
5. Instantaneous Dipole (aka London Dispersion Force or Van der Waals Force): A temporary uneven distribution of electrons around the atom causing an induced dipole in a neighboring molecule or atom, and a temporary bond.
Capillary Action: spontaneous rising of a liquid up into a narrow tube. 2 Forces involved: 1. Cohesion: Attraction between like molecule. 2. Adhesion: Attraction between unlike molecules.
Surface Tension: Energy required to stretch the surface of a liquid by a given area. Viscosity: Measure of a fluid s resistance to flow. Viscous liquids are caused by: strong intermolecular bonding long molecule strands that tangle with each other. Homework: 35, 37, 39, 41, 43, 46
Solids Lattice: internal 3-D arrangement of particles in a solid 2 types of solids: 1. Crystalline internal atomic order or arrangement (ordered lattice) 2. Amorphous lacking internal atomic order (unordered lattice)
Crystalline Solids 5 kinds 1. Ionic solid or ionic crystal ordered lattice of cations and anions Strong ionic bonds cause: a) high melting and boiling points b) hard c) poor conductors in the solid state (no mobile electrons)
2. Molecular Solid or Crystal molecules such as H 2 O, sugar, I 2, P 4, that contain only non-metallic elements Have relatively strong intra-molecular bonds, but relatively weak intermolecular bonds, such as dipole-dipole or H-Bonds, or London Forces. Low melting and boiling points. Not conductive 3. Covalent Network Solid A network of covalent bonds, no intermolecular bonds.
Carbon Diamond sp 3 hybridized, tetrahedral bonding -non conductor of electricity -very hard Graphite sp 2 hybridized with 1 unhybridized p orbital. Overlapping p orbitals are weak bonds. These p orbitals form delocalized molecular bonds, which make graphite conductive.
Graphite sp 2 hybridized with 1 unhybridized p orbital, form trigonal planar plates of carbon. Overlapping p orbitals are weak bonds. These p orbitals form delocalized molecular bonds, which make graphite conductive.
Silicon: Unlike CO 2, SiO 2 is a strong solid network. Si is too large to form pi bonds. Quartz SiO 2 empirical formula, forms networks of SiO 4 4- tetrahedra These tetrahedrons bond to form various anions: Si 2 O 7-6 Si 3 O 9-6 Silicate minerals comprise most rocks e.g. KAlSi 3 O 8
4. Metallic Solids Bonding can be described 1 of 2 ways A) Electron Sea Model: Ordered array of metal cations in a sea of valence electrons Mobile electrons cause metals to be conductive to heat and electricity. Metal cations can slide past each other. Explains why metals are malleable and ductile. B) too complicated.
Metal Alloys: mixture of elements that has metallic properties 2 Types: a. Substitutional Alloy: Elements of similar sizes mix. e.g. Brass Copper and Zinc Sterling Silver Silver and Copper Solder 95% Tin and 5% Antimony b. Interstitial Alloy: Smaller atoms of 1 element fit in the holes between a larger element. Steel Iron (larger) and Carbon (smaller)
5. Group 18 Solids Very weak bonding (London Dispersion Force) therefore low melting and boiling points.
Amorphous Solids Lacking internal order a) Glass quickly cooled quartz Very viscous liquid which eventually freezes into quartz snowflake obsidian To alter properties of glass: Add B 2 O 3 expands and contracts less when heated or cooled. Called Pyrex glass Add K 2 O Makes glass much harder
b) Ceramics - amorphous
Vapor Pressure the pressure of a gas evaporated from a liquid or solid. Equilibrium Vapor Pressure:
Vapor Pressure of a gas is dependent on temperature. The higher the temperature, the higher the vapor pressure. Line represents the minimum KE needed to become a gas.
Vapor pressure is a function of 1) molecular weight 2) intermolecular bonding Low Vapor Pressure substances have: a) high molecular weight b) strong intermolecular bonding
Vapor Pressure as a function of temperature can be calculated by the Clausius-Clapeyron equation. ln (P 1 ) = ΔH vap (1 1 ) Where R = 8.31 J (P 2 ) R (T 2 T 1 ) K. mole and ΔH vap (water) = 40.79 kj/mole at 25 o C Problem: The vapor pressure of water at 20. o C is 17.535 torr. What is it at 23 o C?
Changes in State of Matter Liquid Gas boil: requires energy from the surroundings Gas Liquid condense: releases heat into the surroundings Heat of Vaporization (ΔH) Water = 40.79 kj/mole Argon = 6.3 kj/mole Ethanol = 39.3 kj/mole Hg = 59.0 kj/mole
iquids boil when the atmospheric pressure (external pressure) = vapor pressure. The vapor pressure of water at 100 o C is 760 torr Boil: Evaporate: External Pressure = Vapor Pressure External Pressure > Vapor Pressure
Problem: Calculate the energy required to boil 25 g of liquid 100 o C water, changing it to 100 o C gaseous water? When steam condenses back to liquid, 40.79 kj/mole is given off.
Solid Liquid Melt (fuse): Freeze: requires energy from surroundings releases energy to surroundings ΔH f : Heat of fusion Water Carbon Tetrachloride Sodium Chloride = 6.02 kj/mole = 2.51 kj/mole = 30.2 kj/mole
Why does ice melt at the temperature that it does? 10 P (torr) solid extrapolated liquid 0 0 o C
Supercooled a liquid below its freezing temperature. Occurs because the liquid molecules have not attained the degree of order necessary to form a solid. Superheated a liquid above its boiling temperature. Bumping a hot liquid above its boiling point all of a sudden forms 1 large bubble of gas, and the liquid accelerates out of the test tube.
Solid Vapor Sublimation Deposition
Heating curve for water Specific Heat Ice = 2.03 J/g o C Water = 4.184 J/g o C Steam = 1.99 J/g o C Problem: How much energy is needed to heat 36 g ice at -27 o C to steam at 120. o C?
Phase Diagrams Show states of matter relative to temperature and pressure. Critical Point Fluid forms. Temp. is too high to be a liquid, pressure is too high to be a gas. Called a supercritical fluid. Triple Point Where a solid, liquid and gas all co-exist.
Phase diagram of water with negative solid/liquid slope. Phase diagram of carbon dioxide with positive solid/liquid slope. Homework: Problems: 14, 15, 18, 21-23, 29, 31-34, 43, 46, 81-83, 87, 88, 90-93, 99, 104-106, 113
B) Molecular Orbital Model or Band Model: valence electrons form molecular orbitals around the metal. Conduction Band Valence Band Energy Gap Insulator Conductor