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Chemistry II Study Guide I

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by: Annika Coley

Chemistry II Study Guide I CHEM 1120

Marketplace > University of Memphis > Chemistry > CHEM 1120 > Chemistry II Study Guide I
Annika Coley
University of Memphis
GPA 3.99

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About this Document

This study guide covers what is going to be on the exam
Dr. Brewster
Study Guide
50 ?




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"Yes YES!! Thank you for these. I'm such a bad notetaker :/ will definitely be looking forward to these"
Ms. Andy Tromp


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This 4 page Study Guide was uploaded by Annika Coley on Monday February 15, 2016. The Study Guide belongs to CHEM 1120 at University of Memphis taught by Dr. Brewster in Spring 2016. Since its upload, it has received 128 views. For similar materials see GENERAL CHEMISTRY II in Chemistry at University of Memphis.


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Yes YES!! Thank you for these. I'm such a bad notetaker :/ will definitely be looking forward to these

-Ms. Andy Tromp


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Date Created: 02/15/16
Chemistry II Study Guide  Review how and when to use the following equations: o Boyle’s Law  P1V1=P 2 2 o Charles’ Law  V1/T1=V 2T 2 o Avogadro’s Law  V1/n1=V 2n 2 o Ideal Gas Law  PV=nRT o Average kinetic energy  KE=1/2mv 2 o Clausius-Clapeyron equation  Ln(P2/P1)(-Δ vapR)(1/t2-1/1 ) o Energy required for a change in temperature  Q=mCΔT o Energy required for a phase change  Q=nΔ vap o Heat of a solution  ΔH solutionH soluteΔH solventΔH mix o Henry’s Law  Sgas= kH· P pressure o Boiling Point Elevation/Freezing Point Depression  BP -BP =ΔT =m·k solution solute b b o Reaction Rate  Rate=–Δ[H ]2Δt o Rate Law n m  rate = k [A] [B] o  Review the following concepts o Intermolecular forces  The interactions and attractions between all molecules and atoms  The structure of a molecule determines the properties and strength of its intermolecular forces  The strength of the intermolecular forces compared to the amount of thermal energy or movement of the molecules determines the state of matter o Phase Changes  Caused by heating, cooling, or pressure  The amount of thermal energy in the particles is altered o Dispersion Forces  Temporary polarity in an atom due to temporary dipoles  Caused by the majority of electrons in an atom being on one side of the atom producing a negative charge at one end and a positive charge at the other o Dipole dipole attractions  Exists only between polar molecules  Acts in addition to dispersion forces  The partial negative charge of the polar molecule is strongly attracted to the partial positive charge of another polar molecule o Hydrogen Bonds  Strongest intermolecular force  Occurs when very electronegative atoms bond with hydrogen  Fluorine, Oxygen, and Nitrogen o Ion Dipole Attractions  Exist in a mixture  Determines solubility in water  The positive ions are attracted to the negative partial charge of the polar molecule and the negative ions are attracted to the positive partial charge of the molecule o Endothermic process  Process requiring energy o Exothermic process  Process that gives off energy o Vapor Pressure  The pressure exerted by a gas when it is in dynamic equilibrium with its liquid  When vapor pressure equals external pressure the liquid reaches boiling point  As temperature increases, vapor pressure increases o Phase diagram  Show the states of matter present for different combinations of temperature and pressure  Can be navigated horizontally or vertically to determine state of matter resulting from changes in heat or pressure o Crystal lattice  Arrangement of particles in a crystalline solid  Formed when a liquid is heated and then allowed to cool very slowly o Solution  Homogeneous mixture  Comprised of a solvent and solute  Solvent is what dissolves the solute (ex: water)  Solute is dissolved (ex: salt) o Solubility  Ability of a solvent to dissolve a solute  Changes depending on temperature and pressure  Heating things can make some things more soluble  Heating gases makes them less soluble o Exothermic o Solution Concentration  Amount of solute in a given amount of solution  Dilute o Relatively small amount of solute  Concentrated o Relatively large amount of solute o Molarity (M)  Moles of solute per 1 mL of solution  Measure of solution concentration o Molality (m)  Moles per kilogram of solvent  Compares moles and mass o Parts solute in parts solution  Parts can be measured by mass or volume  Percentage  Parts of solute in every 100 parts solution  Parts per million  Parts of solute in every 1 million parts solution o Colligative Properties  Properties that depend on the concentration of the solution  Vapor pressure  Freezing and boiling point  Osmotic pressure o Mole Fraction  Moles of a substance/total moles o Raoult’s Law  Two liquids mixed will both affect the vapor pressure  Add together both liquid’s individual vapor pressure o Osmosis  Flow of water from low concentration to high concentration of solution  Separated by semipermeable membrane and only the solvent moves  Often thought of as the flow of water against gravity o Osmotic pressure  Amount of pressure needed to keep osmotic flow from taking place  Proportional to molarity of solute particles  Meaning if there are moles of solute form the water to flow towards then the osmotic pressure will be higher  Also based off of Van’t Hoff factors  Larger Van’t Hoff factors equal larger osmotic pressure  П=MRT o Van’t Hoff Factor  Ratio of moles of solute particles to moles of formula units dissolved  Always less than theoretical values because the ions may dissociate but they may be still be close enough to each other to not be considered fully dissociated o Reaction rate  Speed of a chemical reaction  Important to be able to control the speed of a chemical reaction  3 things that effect reaction rate  Concentration of reactants o Greater reaction rate with greater concentration o Reaction happens when particles collide productively  Increased temperature o The particles move faster at a higher temperature so they are more likely to collide o If they are moving faster they are also more likely to have high enough activation energy to cause a reaction if they do collide o Gases are the exception – they are less likely to react if they are at a higher temperature  Reactant orientation o Can cause a change in what is produced when they collide o B + A-A = B-A-A o B + A-X = B-A-X or B-X-A o Average rate  Change in measured concentrations in any particular time period  As the reaction continues, the concentration of reactant decreases causing the reaction to slow down o Instantaneous Rate  Change in concentration at any one particular time  Slope at one point on the cure – tangent line o Reaction order  Coefficient in front of the element  Sum of the exponents on the reactants  Overall order of the reaction o Zero order  One reactant decomposing  Rate of reaction will not change with concentration change o First order  Directly proportional to concentration o Second orde2  K[A]  Quadruples the rate of the reaction  Concentration changed by a factor of 2 and the rate changes by a factor of 4 o


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