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Chemistry Exam 2 Study Guide

by: Anzlee

Chemistry Exam 2 Study Guide CHEM 1120

Marketplace > Middle Tennessee State University > Chemistry > CHEM 1120 > Chemistry Exam 2 Study Guide

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These notes cover the material for our test. It has review from the past section of notes as well as some practice problems. For additional review, refer to our smartwork assignments.
General Chemistry II
Ngee S Chong
Study Guide
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This 11 page Study Guide was uploaded by Anzlee on Thursday October 6, 2016. The Study Guide belongs to CHEM 1120 at Middle Tennessee State University taught by Ngee S Chong in Fall 2016. Since its upload, it has received 25 views. For similar materials see General Chemistry II in Chemistry at Middle Tennessee State University.


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Date Created: 10/06/16
EXAM 2 STUDY GUIDE Prerequisites to Review: *Remember how to do Lewis Dot Structure *Review of Group Names on Period Table: 1A/1- Alkali Metals 2A/2- Alkali Earth Metals 6A/16- Chalcogens 7A/17- Halogens (Diatomic) 8A/18- Noble/Inert Gases *States at Room Temperature: Gas: Mg and Br Liquid: Fr, Cs, Gr, and Rb Solid- all others *Solubility: -Solute- what’s going into solvent; smaller number of moles -Solvent- dissolves solute; larger number of moles -Solubility- maximum amount of substance it can dissolve at a given volume -Miscible- liquids are mutually soluble in any proportion Chapter 6 Intermolecular Forces •   Interactions between nonpolar molecules o   Dispersion- momentary shift in electron density o   Dispersion (London) forces- caused by presence of temporary dipoles in molecules (weakest bond) §   Larger molecules usually solid, then liquid, then gas o   Temporary (induced) dipole- separation of charge produced in atom or molecule by a temporary uneven distribution of electrons (middle strength) §   May be caused by a reaction between a polar and nonpolar molecule o   Polarizability- the ease that an electron cloud in a molecule, ion, or atom can be distorted, inducing a temporary dipole §   Larger molecules/elements have higher polarizability because intermolecular forces are weaker Factors Affecting Strength of Dispersion: •   Size of atom/molecule- larger are more polarizable because the outer valence electrons are being weakly pulled by the positive nucleus; dispersion increases with polarizability o   Molar mass size is directly correlated with boiling point/melting point o   RVP: reid vapor pressure of gasoline regulated by Environmental Protection Agency •   Shape of molecules- increased surface area causes increased interactions and stronger interactions, which affects physical and chemical properties; linear molecules have higher dispersion than branched molecules with similar molecule weight o   Constitutional isomers (structural isomers)- molecules that have the same formulas, but different connections between atoms; differently arranged o   Larger surface = larger dispersion = higher boiling point o   Viscosity- measure of resistance of a fluid to flow (cP- centipoise unit) §   factors: molecular shape, molar mass, and temperature §   higher molar mass= higher viscosity Interactions Involving Polar Molecules: •   Dipole-dipole attraction- the force between polar molecules •   Hydrogen bond- strongest dipole-dipole interaction o   occurs between hydrogen bonded to a small and highly electronegative element (F, O, N) and an atom of O or N in another molecule o   Ex. between complementary sites on double stranded DNA- between A and T or C and D o   All alcohols; all amino acids can have hydrogen bonds o   More hydrogen bonding = higher boiling point •   Boiling points of Binary Hydrides- when boiling points are related to attractive forces, be able to determine boiling point order •   Ion-Dipole- force between an ion and a molecule with a permanent dipole •   Sphere of hydration- water molecules surrounding ion in aqueous medium Solubility: •   Depends on relative strength of interactions between molecules •   Ionic/polar solutes are soluble in polar solvents o   No solution if different forces •   Nonpolar solutes are soluble in nonpolar solvents o   No solution if different forces •   More than one force may need to be examined o   Solubility decreases as hydrogen bonding energy decreases and dispersion increases •   Hydrophobic- repels water; lowers solubility •   Hydrophilic- attracts water; heightens solubility Physical States: •   Factors that affect state: intermolecular forces, temperature, pressure •   Phase diagram o   Graphical representation of substance’s states depending on temperature and pressure o   Lines represent points where two states on either side coexist in equilibrium o   Triple point- where all phases exist at same time o   Critical point- where liquid and gas have same density o   Supercritical fluid- substance above critical temperature and pressure Properties of Water: •   Surface tension- energy needed to separate molecules of the liquid’s surface; directly correlates with strength of intermolecular forces •   Cohesion- interactions between same particles •   Adhesion- interactions between different particles •   ex. meniscus §   concave- adhesive forces greater than cohesive forces §   convex- cohesive forces greater than adhesive forces •   Capillary action- liquid can spontaneously flow against gravity; involves adhesive and cohesive forces •   Density of water- decreases as a solid o   Important for aquatic ecosystems Chapter 11 Properties of Solutions •   Enthalpy of Solution depends on: o   Energies that are holding the solute ions in the crystal lattice o   Attractive force holding solvent together o   Interactions between solute ions and solvent molecules •   Lattice energy- energy released when 1 mole of the ionic compound forms from free ions in gas phase; energy released when crystal lattice forms (positive) •   ΔH ion-ionenergy required to remove ions from crystal lattice (negative) •   Born-Haber cycle- series of steps with change of energy (ΔH) that describe the formation of an ionic solid from its elements (sublimation, bond breaking, ionization, electron affinity, formation) Isoelectronic- same electron configuration when electrons are lost or gained Na is isoelectronic with Ne Vapor Pressure: •   Pressure exerted by gas in equilibrium with liquid; evaporation and condensation rates are equal •   Normal boiling point- ambient pressure must be standard pressure •   Factors affecting vapor pressure: o   Temperature/surface area; overcoming intermolecular forces o   Presence of nonvolatile solute- affects rate; decreases vapor pressure o   Clausius-Clapeyron Equation: related vapor pressure with temperature of substance to its heat of evaporation •   Fractional Distillation- separate a mixture of compounds based on their differing boiling points o   More volatile components have enriched vapor pressure •   Raoult’s Law- total of overall pressure is equal to the sum of each individual pressures of the components, in an ideal solution •   Real vs. Ideal solutions- due to differences in solvent-solvent and solvent-solute interactions Osmosis: •   Fluid flows across a semipermeable membrane to balance the concentration of that fluid on both sides •   Osmotic pressure- pressure applied to membrane to stop flow of solvent o   Formula: iMRT •   Hypertonic state- water migrates out •   Hypotonic state- water migrates in •   Isotonic state- water concentration is at equilibrium o   Cells within body should be isotonic •   Reverse osmosis- solvent pushed through membrane, leaving higher concentrated solution behind o   Ex. water distillation/purification Chapter 11 Online Homework Samples: 8, 16, 18                                             Test 2 Material Thermodynamics Spontaneous Reactions: •   Most are exothermic- ∆H < 0 o   ex. combustion •   Endothermic spontaneous reactions- ∆H > 0 o   particles in product are more spread out than starting reactants o   ex. cold packs •   Gas à liquid à solid o   Heat is released •   Solidà liquidà gas o   Heat is put in Entropy: •   Entropy (S)- how dispersed energy is in a system at a temperature o   Effected from molecular motion and volume o   Increases from: increasing temperature, volume, complexity of structure, and independent particles o   Can never be negative (only zero if K is zero) •   Second Law of Thermodynamics- entropy of universe is equal to the entropy of the system and the surroundings (greater than zero) •   Entropy and microstates o   Motion of molecules in quantized- different molecular states are separated by energies o   Energy state/energy level- allowed value of energy o   Microstate- distribution of particles within energy level •   Boltzmann Equation: -23 k Bn W o   kB= 1.38*10 J/K o   W = number of microstates •   Third Law of Thermodynamics- perfect crystal provides baseline for entropy because at absolute zero (0 K or -273 C), the entropy is zero 0 •   Standard Molar Entropy (S )- absolute energy of 1 mole in standard state (298 K and 1 bar of pressure) o   Found from how much heat energy a substance must have to change temperature by one value •   Types of Molecular Motion (increases with increasing temperature): o   Translational- movement through space o   Rotational- perpendicular spinning motion around axis o   Vibrational- movement of atoms to and away from each other •   Entropy Change: ∆S rxn= ∆S sys S final Sinitial •   Isothermal Process o   ∆S = q rev o   qrev= flow of heat for reversible process •   Gibbs Free Energy (G)- maximum energy released that is available to do work o   constant temperature and pressure o   ∆G = ∆H -T∆S o   ∆H = ∆G +T∆S o   efficiency = work done divided by energy produced Chapter 13: Chemical Kinetics •   Kinetics- rate of the change of concentrations of substances in chemical reactions o   Reaction Rate- how quickly reaction occurs; concentration change over time o   Reactant molecules must collide with sufficient energy (above activation energy) to form products o   Factors effecting rate- physical state, concentration, temperature, and catalysts o   Rate constant = k o   Rate Law- defime relntionship between rate of reaction and concentrations §   k (A) (B) •   m and n are reaction order with respect to A and B (usually 1 or 2; never 3 or above) •   m + n is the overall order of reaction o   Reaction order- number defining dependence of rate on concentration •   Photochemical Smog- internal combustion engines produce compounds that interact with sunlight, producing a mixture of gases Half-Life (t 1/2 •   Time of a chemical reaction where the concentration of the reactant decreases by half o   ln A/t =0-0.693 = -kt o   t1/2= 0.693/k §   Example: An isotope has a half life of 14.3 days. How long does it take for 95% of the sample to decay? Activation Energy (E ) a •   Energy of molecular collisions reactants must have to break bonds and therefore form products •   Arrhenius Equation o   k = Ae -Ea/RT §   A- collision frequency factor §   R- universal gas constant (note units) §   T- temperature in kelvin Reaction Mechanisms •   Steps that show how reaction occurs (molecular level) •   Needs to be consistent with the rate law o   Elementary Step- single process o   Intermediate- species produced in one step and used in the next step o   Molecularity- number of ions, atoms, or molecules in an elementary step


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