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Chemistry Exam 1 study guide

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by: Audrey Notetaker

Chemistry Exam 1 study guide 111001

Marketplace > Boston College > Chemistry > 111001 > Chemistry Exam 1 study guide
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Chapter 11: Intermolecular forces in liquids, vapor pressure, phase diagrams, unit cells Chapter 12: Intro to solutions, heat of dissolution concentration terms, colligative properties, colloids a...
General Chemistry II
Neil Wolfman
Study Guide
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This 10 page Study Guide was uploaded by Audrey Notetaker on Saturday February 13, 2016. The Study Guide belongs to 111001 at Boston College taught by Neil Wolfman in Fall 2016. Since its upload, it has received 84 views. For similar materials see General Chemistry II in Chemistry at Boston College.

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Date Created: 02/13/16
Chemistry Exam 1 Study Guide Chapter 11, 12 Chapter 11: Intro to liquids Gases are compressible fluids Liquids are NOT compressible because the molecules are more closely packed, but they can still  move around a bit Liquids have a definite volume but indefinite shape Solids are NOT compressible, have a fixed volume and shape. Atoms, ion, molecules cannot  move freely but can vibrate or oscillate – most closely packed and dense Solid  liquid Melting/fusion Liquid  solid Freezing Liquid  gas Vaporization Gas  liquid Condensation Solid  gas Sublimation Gas  solid deposition Intermolecular forces among liquid molecules ­ Force between molecules is weaker than the force between atoms in a molecule ­ Bond length is shorter than distance between molecules Types of intermolecular forces 1. Dipole­Dipole: the positive end of one molecule is attracted to the negative end of  another ­  higher dipole­moment = higher boiling point ­ Solubility o Polar solutes will dissolve in polar solvent o Non­polar solutes will dissolve in non­polar solvents o CANNOT dissolve polar in non­polar 2. Ion­Dipole 3. London Dispersion 4. Hydrogen Bonding Chapter 11: Intermolecular Forces in Liquids Ion­Dipole – the positive end to the negative of a ionic compound in a polar solvent London Dispersion – electron density is not distributed evenly, causing a brief dipole­moment/  polarity; all molecules have this ­ Larger atoms = stronger the force ­ Larger molecules = stronger the force *Stronger intermolecular force = higher boiling point *Isomers affect dispersion forces because the different arrangement = different distribution of  electrons and different surface area; more surface area = stronger force Hydrogen Bonding – occurs in polar molecules; only between O­H / N­H / F­H Surface Tension – amount of energy it takes to increase the surface area (mJ/m^2). The  molecules towards the middle of a solution have no net force  For a given volume, a sphere has the lowest surface area = lowest boiling point Greater intermolecular force = greater the surface area Viscosity – resistance to flow. Measured in poise (P) or (g/cm*s) Greater forces = greater viscosity Greater molecular weight = greater viscosity Greater SA of molecule = greater viscosity High temperature = low viscosity Chapter 11: Vapor Pressure of Liquids  Capillary Action: the ability of liquid to flow up a narrow tube against gravity Adhesive – molecules stick to the wall Cohesive – molecules stick to each other ­ if the adhesive force is greater than the cohesive, then the liquid will creep up the side  (i.e. water) ­ if the cohesive force is greater than the adhesive, then the liquid will bulge (i.e. mercury) Dynamic Equilibrium: rate of evaporation = rate of condensation ­ the rate of evaporation is a constant and only dependent on temperature ­ vapor pressure is a constant and only dependent on temperature *liquids with a high vapor pressure are volatile Normal Boiling Point: the temperature at which pressure is at 1 atm. A liquid has reached its  boiling point when the vapor pressure of liquid = the ambient pressure (1 atm). Freezing point and melting point does not depend on pressure because solids and liquids  cannot be compressed *temperature does not change during a phase change ∆ H vap = how much heat it takes to heat a substance over a certain amount of time. Always  ∆ H positive and larger than  fus  Chapter 11: Phase Diagrams P2 −∆ H 1 1 ln( ) = ( − ) R= 8.314 J/mol*K P1 R T2 T 1 *pressure, delta H, and temperature should never be negative  *vapor pressure goes up if the temperature goes up Phase Diagram: the lines represent equilibrium ­ if the line representing the change from solid to liquid leans to the right, the solid is more  dense than the liquid ­ if the line representing the change from liquid to solid leans to the left, the liquid is more  dense than the solid (i.e. water) ­ all three phases are at equilibrium at the triple point ­ there is no liquid below the triple point Supercritical Fluid: there is no differentiation between liquid and gas. Takes place above the  critical point. Critical Point: point at which a substance can no longer exist as a liquid Critical Pressure: minimal pressure needed to liquefy a gas at critical temperature There are different forms of ice at high pressures *water moderates temperature *water is a universal solvent because of its polarity Solids: essentially incompressible Intermolecular forces within solids: molecular, non­binding atomic, metallic atomic,  ionic, covalent Molecular: weak intermolecular forces – low melting points H2O, CO2 Non­binding atomic: weak London forces – low melting points Noble gases Metallic atomic: Transition metals Chapter 11: Unit Cells and Cubic Systems Ionic Solids: cations and anions NaBr, KOH, MgF2 Covalent Solids: atoms, molecules held together in networks or chains Graphite, diamond, quartz *highest melting point *Stronger forces = higher melting point ­ melting point instead of boiling because it’s a solid Hardness: how easily structural units can be moved Molecular solids – soft ionic solids – hard, brittle 3D covalent networks – hardest substance metallic solid­ not brittle, malleable, conducts electricity Amorphous: no long­range order (i.e. glass, plastic) Crystalline: ordered (NaCl, sucrose) Crystal: 3D ordered arrangement “unit cell” – brick/building block when replicated in all three dimensions can give entire  lattice structure *different unit cell shape is based on edge length and angle Types of crystals: Cubic System Simple cubic ­ (1/8 *8) = 1 atom per unit cell  o 8 other unit cells share the same corner/make up the atom ­ Coordination number = 6 ­ Packing efficiency = 52% ­ Edge length (l) = 2*r(radius) Body­centered cubic ­ (1/8 *8) + 1 = 2 atoms per unit cell ­ Coordination number = 8 ­ Packing efficiency = 68% ­ Edge length (l) = 4*r Face­centered cubic ­ (1/8 *8) + (1/2 *6) = 4 atoms per unit cell ­ Coordination number = 12 ­ Packing efficiency = 74% ­ Edge length (l) = 4*r (diagonal across the face) Chapter 11: Calculations with Unit Cells Polymorphic – substance can exist in more than one structure X­ray Diffraction crystallography: determine the structure; finds where each atom sits in 3D ­ Crystals are highly ordered Constructive Interference: brighter, larger amplitude Destructive Interference: dimmer; smaller amplitude *If the difference in distance travelled by wavelength is an integer, then the waves are  constructive so the light will be lighter nλ=2dsinθ nλ=2a Bragg’s Law:                n = integer (n = 1: first order diffraction; n = 2: second order diffraction, etc.) λ  = wavelength (m)  d = distance of separation between layers (m) θ  = angle of incidence  Chapter 12: Introduction to Solutions Alloy: dissolving two metals Solution: homogeneous mixture (atoms, molecules, ions, solids, liquids, gas) ­ Completely miscible liquids can be completely dissolved ­ Solute: solid or gas dissolved; whichever is the smaller amount ­ Solvent: liquid is dissolved *unreactive gases are completely miscible to each other Solubility: how much of a compound dissolves in a certain amount of solvent ­ Like dissolves like (polar and polar / non­polar and non­polar) ­ Group 1A, ammonium ions, and salts are soluble in water Unsaturated: more solute can be added and it will still dissolve Saturated: greater than or equal to the amount at which the solute will stop dissolving or  it has completely dissolved Supersaturated: not in equilibrium; very unstable *things have a natural tendency towards disorder or entropy which leads to mixing ­ Want the state of lowest energy o If solute/solvent forces are greater than the solute then there is no mixing o If solute/solvent forces are equal to the solute: solvent then there is mixing Molecular Solutions ­ Gases: entropy predominates (is stronger) so there is mixing ­ Non polar liquids: entropy is stronger so there is mixing ­ Polar and non polar: entropy is weak because it takes too much energy to dissolve so  there is no mixing *longer hydro­carbon chain becomes less polar and less likely to be miscible in water Solubility depends on: ­ Ion­dipole force favors dissociation: bringing ions into solution ­ Interactions of ions (lattice energy): favors order Δ H +  is endothermic: needs/absorbs heat; cold Δ H ­  is exothermic: releases heat; hot Dissolving a Solid Δ H 1. Break the crystal lattice – opposite of the lattice energy ­ + 2. Break the hydrogen bonds (separate water) ­ + Δ H 3. Mixing – ion dipole forces ­ ­ Δ H Chapter 12: Heat of Dissolution Concentration Terms ΔH (solution)=ΔH (solute)+ΔH (solvent+ΔH(mixing) ΔH solution)=ΔH (solute)+ΔH(hydration) *most ionic solids get more soluble the higher the temperature *most gases become less soluble as temperature increases *change in pressure has no effect on solubility of liquids or solids (because they are  incompressible) in water but has a HUGE effect on gas Henry’s Law: S(gas) =K(h)P(gas) ­ Solubility will change according to the temperature Coligative Property: only depends on quantity not quality; number of moles of whatever is added not what is actually added Chapter 12: Colligative Properties of Solutions masssolute(mol) Molality:  masssolvent(kg) *adding a nonvolatile solute to a volatile solvent will lower the vapor pressure ­ More solute pressure will cause less liquid to evaporate  ­ Vapor pressure of solution is less than the vapor pressure of solvent Rault’s Law: vapor pressure of solution = mol fraction of solvent x vapor pressure of solvent Psolution= (Xsolvent )Psolvent ) ∆ P=(Xsolute)(Psolvent) *change in vapor pressure will always be lowering Ideal Solution: solution obeys Rault’s law at high and low mol fraction and the solute and  solvent contribute to the solution vapor pressure *liquid with higher pressure is more volatile and has a lower boiling point *more solute dissolved = higher boiling point ∆ Tb=m xKb                            will be an increase change ∆ Tf=m xKf        will be a decreased change *Solutions will boil at higher temperatures and freeze at lower temperatures  Osmotic pressure: amount of pressure required to stop osmotic flow, dependent upon molality of  solute Π=MRT Van Hoff factor (i): actual number of particles in solution Something that doesn’t ionize will have a factor of i=1 Ion pairing: reduces the effect so factor is less than predicted Hyperosmotic: solutions have higher osmotic pressure than body fluids Hyposmotic: solutions have lower osmotic pressure than body fluids Isosmotic: solutions have equal osmotic pressure with body fluids Colloids: dispersions of particles of oen substance (dispersed phase) in another substance  (continuous phase) ­ Determined by the size of particles (if particles are between 1 nm and 1000 nm then it’s a colloid) Chapter 12: Colloids and Detergents Colloid: dispersed particles are larger than true solution  Tyndall Effect: light will scatter through colloid Hydrophilic (water loving): dispersed particles have strong attraction to water so they interact  with water Hydrophobic (water fearing): dispersed particles have no attraction to water so they do not  interact with water  ­ Unstable and tend to aggregate/clump over time Coagulation: process in which dispersed phase of colloids aggregate/clump Soaps/detergents: molecules that have a polar and non polar end Association colloid: colloidal size particles form micelles Micelle number: number of monomers in a micelle Critical micelle concentration (cmc): minimum concentration of soap/detergent for micelle to  form; measured in weight/volume Anionic: ionic head is negatively charged Cationic: ionic head is positively charged Zuidionic: ionic head is both positively and negatively charged 1. Does reaction occur?  thermodynamics (free energy) 2. If reaction occurs where does it end?  equilibrium (K) 3. How fast is the reaction?  kinetics (k): speed and mechanism a. Rate? b. What affects rate? c. Molecular level?


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