final study guide
final study guide CHEM 1040 - 002
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This 11 page Study Guide was uploaded by Joerdan Notetaker on Tuesday December 8, 2015. The Study Guide belongs to CHEM 1040 - 002 at Auburn University taught by Ria Astrid Yngard in Fall 2015. Since its upload, it has received 34 views. For similar materials see Fundamental Chemistry II in Chemistry at Auburn University.
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Date Created: 12/08/15
Dr. Yngard Chem 1040-‐002 Final Exam Study Guide For additional studying I suggest re-‐working old exam problems and problems that went over in class. Ch. 7.3, 13, 14 Ch. 7.3 Intermolecular Forces o Types of Forces o London Dispersion(weakest of them all): coulombic(particles w/ a charge) attractions between instantaneous dipoles of molecules § The higher the boiling point, the more energy needed-‐ attraction is stronger § Larger the molecule, more polarizable, stronger the force o Dipole-‐Dipole Interactions: attraction between the positive end of one molecule and the negative end of the neighboring molecule § The dipole moment measures how polar a compound is o Hydrogen Bonding: attraction between hydrogen atom and electronegative/nonbonding electron pair on F, O, or N atom § Example-‐ CH CO3H (draw out lewis structure to see O-‐H bond) Chapter 13 Physical Properties of Solutions o Solution: homogenous mixture of two or more pure substances o Composed of a solute & solvent o Solute is uniformly dispersed throughout the solvent o Unsaturated Solution: less solute than solvent, has capacity to dissolve @ specific temperature o Saturated Solution: maximum amount of solute that will dissolve in solvent at specific temp. o Solubility: amount of solute dissolved in a given volume of a saturated solution at a specific temp. o Example-‐ NaCl @ 20°C is 36g/100ml of water o Supersaturated Solution: has more dissolved solute and is generally unstable o Examples-‐ § 19g NaCl dissolved in 100ml water @ 20°C UNSATURATED § 45g NaCl dissolved in 100ml water @ 20°C SUPERSATURATED § 45g NaCl dissolved in 100ml water @ 20°C but not all has dissolved SATURATED o Solvation: solute molecules are separated & surrounded by solvent molecules o ΔH 1 0 (solute/solute) o ΔH 2 0 (solvent/solvent) o ΔH 3 0 (solute/solvent) usually the case… o ΔH soln can either be endo/exothermic o ΔH soln= ΔH1 + 2 + 3 o Solution process depends on if.. o Enthalpy is decreased o Entropy(disorder) is increased o Solute/Solvent interactions: “like dissolves like” o In general… § Polar dissolves polar § Nonpolar dissolves nonpolar o Miscible-‐ mixing in all proportions o Methanol is miscible in water and hexanol is almost insoluble in water: why? o B/c hexanol has more carbon thus making it more nonpolar and since water is polar it is harder to mix the two; methanol is simply polar o Molarity= M= moles of solute/liters of solution o Temp. dependent o Mole Fraction A= X = A moles A/sum of moles of all components o Percent by Mass= mass of solute/mass of solute+mass of solvent x 100% o Temp. independent o Calculate percentage by mass of NaCl in a solution containing 3.50g of NaCl in 75.0g of water %mass= (3.50gNaCl)/(3.50gNaCl + 75.0gWater) x 100%= 4.46% 6 o Parts per million= ppm= mass of solute/total mass of solution x 10 o Parts per billion= ppb= mass of solute/total mass of solution x 10 o Factors affecting solubility… o Temperature § Gases in water: solubility decreases, temp. increases § Solid solutes in water: solubility increases, temp. increases o Pressure § Solubility of solids/liquids is hardly affected § Solubility of gases increases w/ pressure § Henry’s Law: solubility of a gas in a liquid is proportional to the pressure of the gas over the solution o C= kP § C= molar concentration(mol/L) § k= constant(mol/L x atm) § P= pressure of gas over solution(atm) o Colligative Properties: properties that depend on the number of solute particles in a solution o Vapor Pressure Lowering § Raoult’s Law: solution containing a non-‐volatile solute • P 1 X1P° 1 • P 1 partial pressure of solvent • X 1 mole fraction • P° 1 vapor pressure o Boiling Point Elevation § ΔT =b T b -‐b T°b= K m • K =b constant(°C/m) • m= molality(mol/kg) o Freezing Point Depression § ΔT = f T° f – f T = f K m • K =f constant(°C/m) • m= molality(mol/kg) o Osmotic Pressure: π= MRT § M= molarity(mol/L) § R= ideal gas constant(L x atm/mol x K) § T= absolute temp.(K) § Osmosis: selective passage of solvent molecule thru a porous membrane from a more dilute solution to a more concentrated one o van’t Hoff Factor (i) § i= actual # of particles after dissociation/# of formulas units initially dissolved in solution § nonelectrolytes usually have a van’t Hoff factor of 1 § electrolytes van’t Hoff factor is different based on what electrolyte is being used o Colloids: a dispersion of particles of one substance throughout another substance § Particles much larger than the normal solute molecules, but too small to be settled by gravity § Hydrophilic-‐ water loving § Hydrophobic-‐ water hating(nonpolar) Chapter 14 Entropy & Free Energy o 1 Law of Thermodynamics: energy cannot be created nor destroyed but can be transferred or transformed o spontaneous processes: does occur under a specific set of conditions without ongoing outside intervention o CH (g) 4 + 2O (g) 2 à CO (g)2 + 2H O(2) ΔH°= -‐890.4 kj/mol § Exothermic; will release heat § An exothermic process is often spontaneous o H O(s)2 à H O(l)2 ΔH°= 6.01 kj/mol § Endothermic process o Entropy(S): a measure of how dispersed the systems energy is o Types of Motion o Translational-‐ movement of entire molecule from one place to another o Rotational-‐ rotation of the molecule about an axis or sigma(σ) bonds o Vibrational-‐ periodic motion of atoms within a molecule o Entropy changes in a system ΔS sys o ΔS = sys final-‐ initial o ΔS = sysW final § S°: absolute entropy of a substance of 1atm o Calculating reaction standard entropy o ΔS° = rxn°(products) -‐ Σm S°(reactants) o Example: § 3O (2)à 2O (g) 3 ΔSrxn ? § ΔS° =rxn [2(237.6 J/K x mol)] – [3(205.0 J/K x mol)]= -‐139.8 J/K x mol o Qualitative ΔS predictions § Temperature increases, entropy increases § Volume increases, entropy increases § Phase changes: increasing entropy solid< liquid< gas • Melting, vaporization, sublimation § Chemical reaction: rxn resulting in a greater # of gas molecules, entropy increases • Examples-‐ o CO (s) 2 à CO (g)2 ΔS= + o CaO(s) + CO (g)2 à CaCO (s) 3 ΔS= -‐ Ch. 15-‐17.3 § Equilibrium: the rate of the forward reaction is equal to the rate of the reverse reaction o Concentrations of reactants/products are not changing with time § The reaction quotient and equilibrium contstant o aA + bB ⇌ cC + dD o Q c K c = [C] [D] /[A] [B] b § Magnitude of K c o When the K cs large, the rxn favors the products o When the K cs small, the rxn favors the reactants o If in between, the equilibrium mixture contains comparable amounts of both reactants and products § Heterogeneous equilibrium: species in a reversible rxn are in different phases o Concentrations of solids and liquids don’t appear in the equilibrium expression § Manipulating the equilibrium expression o If the equation is reversed, invert the equilibrium constant o If coefficients in the rxn are multiplied by a factor, the equilibrium constant is raised to the power equal to that same factor o [Hess’s Law] equilibrium constant for a net rxn made up of two or more rxns is the product of the eq. constants for the individual rxns § Gaseous equilibrium-‐ set up like the reaction quotient and the eq. constant Δn o K p K cRT) o K and p ace only equal when the Δn=0 § Chemical Equilibrium and Free Energy o A rxn with a very large K has negative ΔG° (spontaneous) o A rxn with a very small K has a positive ΔG° § ΔG/ΔG° relationship o ΔG = ΔG° + RTlnQ § ΔG°/K relationship o ΔG° = -‐RTlnK § Le Chatelier’s Principle o When you apply stress to a system at eq. the system will respond by shifting in the direction that minimizes the effect of the stress o Stress can be... § Addition/removal of a reactant or product § Volume or pressure change § Temperature change o When the volume is decreased, eq. is driven towards the side with the smallest number of moles of gas o The addition of an inert gas that is neither reactant/product will not change the equilibrium o Temperature changes.. § Endothermic; reactants + heat ⇌ products § Exothermic; reactants ⇌ products + heat § Acids and Bases § Arrhenius definition o Acid-‐ substance, when dissolved in water produces H ions o Base-‐ substance, when dissolved in water produces OH ions -‐ § Bronsted definition o Acid-‐ proton donor o Base-‐ proton acceptor § Conjugate acid/base pair only differ in one proton § Factors affecting acid strength: bond strength and bond polarity § Hydrohalic Acids-‐ hydrogen/halogens o HI < HBr < HCl < HF increasing bond strength o HI > HBr > HCl > HF decreasing acid strength § Oxyacids o Increased electronegativity causes increased acidity o More oxygen atoms increases acidity § Carboxylic Acids-‐ acid strength depends on nature of the R group § Ion-‐Product Constant of Water: K w o Amphoteric: species that can act as an acid or base o K w [H O3][OH ] -‐ o @ 25°C K w= 1.0 x 10 -‐7 o in pure water K w= 1.0 x 10 § pH and pOH scale o pH = -‐log[H O3] o pH + pOH = 14.00 o pH = 7 NEUTRAL o pH > 7 BASIC o pH < 7 ACIDIC § Strong Acids(NEED TO KNOW): HCl, HBr, HI, HNO , HCl3 , HClO4, H SO3 2 4 o Everything else considered to be weak § Strong Bases: LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH) , Sr(OH2 , Ba(OH)2 2 o Hydroxides of alkali metals and heaviest alkaline earth metals § Weak Acids and K a o K c= Ka o Use the K a and an ICE chart to find [H ] and then calculate the pH o If ionization is less than five percent, then simplifying is allowed § % ionization = [H ]eq./[weak acid]initial x 100% § Weak Bases and K b o K c= Kb o Use the K b and an ICE chart to find the [OH ] then calculate pOH then calculate pH § Conjugate Acids and Bases o Strong acid has relatively weak conjugate base o Weak acid has relatively strong conjugate base o Strong base has relatively weak conjugate acid o Weak base has relatively strong conjugate acid o K w= (K a(K ) b § Polyprotic Acids-‐ if a1 is 1000x greater than a2K , the acid can be treated monoprotic when determining the pH § Oxides of Metals o Metal oxide + water à metal hydroxide (basic solution) o Metal oxide + acid à salt + water § Oxides of Nonmetals o Nonmetal oxide + water à acidic solution o Nonmetal oxide + base à salt + water § Lewis Acids and Bases o Acid: electron pair acceptor o Base: electron pair donor § Common Ion Effect: extent of ionization of a weak acid decreases in the presence of a strong electrolyte that shares a common ion (same holds for a weak base) § Buffers: solution of a weak acid and conjugate base or a weak base and its conjugate acid § Henderson-‐Hasselbach Equation o pH = pK a + log[conjugate base]/[weak acid] § Strong Acid-‐ Strong Base Titration o HCl(aq) + NaOH(aq) à NaCl(aq) + H O(l) 2 o Endpoint-‐ point where a change in color takes place o Equivalence point-‐ before the acid is in excess 1. Initial pH (pH= -‐log[H ] + 2. Between initial and eq. point (pH is excess of H ) 3. Eq. point acid complete titration (pH neutral) 4. After eq. point, base will be in excess (pH on basis of OH excess) § Weak Acid-‐ Strong Bases Titration 1. Initial pH (use the Ka) 2. Between initial pH and eq. point (buffer) 3. Eq. point (conjugate base, pH>7, Kb -‐> pOH -‐> pH) 4. After eq. point (to get pH use excess OH) Ch. 17.4-‐19 o Titrations-‐ understand how to work out titration problems from the end of chapter seventeen o Solubility Equilibria-‐ K sp • BaSO (s) <-‐> Ba (aq) + SO (aq) 4 2+ 2-‐ 4 o K sp [Ba ][SO ] 4 o The smaller the K ,sp the less soluble it is o Precipitation formation-‐ know the solubility rules • If Q sp K no precipitate forms • If Q sp K a precipitate forms • If Q sp K the solution is just saturated o Factors Affecting Solubility • Common ions-‐ the extent of ionization of a weak acid decreases in the presence of a strong electrolyte that shares a common ion • pH • Complexing agents o Complex ion: central metal cation bonded to one or more molecules or ions o Electrochemistry-‐ study of relationships between electrical energy and chemical reactions o Redox reactions-‐ reaction of electron transfer • “OIL RIG” o Oxidation Is Loss of electrons o Reduction Is Gain of electrons • The oxidizing agent is reduced • The reducing agent is oxidized o Steps for Balancing Redox Rxns in Acidic Media I. Assign oxidation numbers (all elements in elemental form have an oxidation number of zero) II. Write oxidation and reduction half reactions III. Balance each reaction a. Balance elements other than H and O b. Balance O by adding H 2 c. Balance H by adding H d. Balance charge by adding electrons IV. Multiply half reactions by integers if needed V. Add the half reactions and cancel what appears on both sides o Balancing Redox Rxns in Basic Media o Balance as if the reaction happens in acidic media -‐ + o Use OH to neutralize H o Galvanic cell-‐ spontaneous reaction used to produce energy o Electrolytic cell-‐ energy is used to drive a reaction • Anode: oxidation (anions move toward the anode) • Cathode: reduction (cations move toward the cathode) o Standard Reduction Potentials, E° • Cell potential (Ecell‐ potential difference between the anode and cathode in a cell; measured in volts • Standard Hydrogen Electrode (SHE) o 2H + 2e > 2 H E°= 0 V • A positive E° means the redox rxn will favor the products • A negative E° means the redox rxn will favor the reactants • Under standard conditions… E° = cell° cathode -anodeE • Oxidizing and Reducing Agents Ø Strongest oxidants have most positive reduction potentials Ø Strongest reductants have the most negative reduction potentials Ø The greater the difference between the two, the greater the cell voltage Ø Changing the stoichiometry does not affect the value of reduction potential • Electric energy (J) = cell potential (V) x total electric charge (C) -‐19 o 1.60 x 10 C = charge one electron • total electric charge = nF (F = Faraday’s constant = 96.500 C/mol ▯ e or 96,500 J/V ▯ mol ▯ e ) o Spontaneity: standard state conditions • ΔG° = -‐nFcell or cell = ΔG°/-‐nF • @ equilibrium ΔG° = -‐RTlnK o E° cell -‐(RT/nF)lnK o Spontaneity: non-‐standard state conditions • ΔG = ΔG° + RTlnQ • -‐nFE = -‐nFE° + RTlnQ • (Nernst Equation) E = E° -‐ (RT/nF)lnQ o Concentration cells: a cell with two half-‐cells containing the same components but with different ion concentrations o Battery: a galvanic cell, a series of galvanic cells, that can be used as a portable direct electric current (all rxns written out in ch. Eighteen powerpoint slides) • Dry cell battery-‐ has no fluid component (used in flashlights) o Has a zinc container (anode) and manganese dioxide with an electrolyte (ammonium chloride) • Alkaline battery-‐ also based on reduction of manganese dioxide and oxidation of zinc but takes place in basic medium • Lead storage battery-‐ six identical cells joined together each with a lead anode and a cathode made of lead dioxide; both immersed in sulfuric acid • Fuel cells-‐ requires continuous supply of reactants to keep functioning; hydrogen & oxygen gases bubbled through anode and cathode compartments o Electrolysis: process in which electrical energy is used to drive a non-‐ spontaneous chemical reaction (electrolytic cell carries this out) • Faraday-‐ for any half-‐reaction, the amount of substance reduced or oxidized is directly proportional to the # electrons passed into cell • Coulomb (electrical charge)-‐ amount of electrical charge transported in one (s) by a steady current of one ampere (A) o Corrosion: deterioration of a metal by an electrochemical process; formation of rust requires oxygen & water • Prevention-‐ coat surface with paint or coat with oxide layers o Chemical Kinetics: rate, or speed, at which a chemical process occurs o Collision Theory of Chemical Kinetics § Rate “related to” number of collisions/second o Effective collision-‐ § Activation energy § Proper orientation o Factors that affect reaction rates § Temperature § Concentration of reactants § Physical state of reactants § Catalysts o Average reaction rate-‐ the change in concentration of a reactant or product with time § A à B § Avg. rate of appearance B = Δ[B]/Δt § Avg. rate of disappearance A = -‐ Δ[A]/Δt o Instantaneous rate-‐ rate for a specific moment in time; slope of tangent to the curve at any particular time § Rate= k[Br ] 2 • k= 3.5e-‐3 s (rate constant) • constant at constant temperature o Reaction rate and stoichiometry § aA + bB > cC + dD § rate= (-‐1/a) Δ[A]/Δt = (-‐1/b) Δ[B]/Δt = (1/c) Δ[C]/Δt = (1/d) Δ[D]/Δt o Rate Law-‐ relationship between the rate of the reaction and the concentrations of the reactants § aA + bB à products § rate= k[A] [B] y § x is the order of the rxn with respect to A § y is the order of the rxn with respect to B § x & y can only be determined experimentally § overall reaction order= sum of exponents o Rate Law: Zero-‐Order § A à products § Rate = k[A] or rate = -‐Δ[A]/Δt § -‐Δ[A]/Δt = k § Integrated rate law: [A] t = -‐kt +0 [A] § Graph: [A] vs. time is a straight line • slope is –k • [A] is y-‐intercept 0 § Half life 1/2 ): the time required for one half of a reactant to drop half its initial value • t1/2 = 0[A] /2k o Rate Law: First-‐Order § A à products § Rate = k[A] § -‐Δ[A]/Δt = k[A] § Integrated rate law: ln[A] t= -‐kt + 0n[A] § Graph: ln[A] vs. time is a straight line • Slope = -‐k • Y-‐intercept = ln[A] 0 § t1/2 0.693/k o Rate Law: Second-‐Order § A à products 2 § Rate = k[A] § -‐Δ[A]/Δt = k[A] § Integrated rate law: 1/[A] t kt + 1/[A0 § Graph: 1/[A] vs. time is a straight line • Slope = k • Y-‐intercept = 1/[A] 0 § t1/2 = 1/k[A0 o Effect of temp. on reaction rate o Collision: orientation and energy § Activation energy (Ea)-‐ minimum energy required to initiate a chemical rxn § Activated complex or transition state-‐ a temporary, high energy, unstable chemical species formed § Arrhenius equation: k = Ae -‐Ea/RT • A = frequency factor • k = rate constant • Ea = activation energy (kJ/mol) • R = gas constant (8.314 J/mol K) • T = absolute temperature o Reaction Mechanism: the sequence of steps that sum to give the overall rxn o Elementary reaction-‐ each step in rxn mechanism involving a single collision of the reactant molecules o Molecularity of an elementary reaction-‐ number of reactant molecules involved in collision § Unimoleculer • A à products rate = k[A] first order § Bimolecular • A + Bà products rate = k[A][B] second order 2 • A + A à products rate = k[A] second order o Two Step Mechanism 2 § NO + NO à 2 N2 rate = k[NO] § N 2 2 + 2 à 22O 2 2 2 rate = k[N O ][O ] § Intermediate: N O 2 2 § Slowest step in mechanism is rate determining step
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