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CHEM 1040 Exam 3 Study Guide

by: Hannah B.

CHEM 1040 Exam 3 Study Guide CHEM 1040

Marketplace > Auburn University > Chemistry > CHEM 1040 > CHEM 1040 Exam 3 Study Guide
Hannah B.
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Dr. Yngards Fundamentals of Chemistry II. Study guide for exam three, covering chapters 17-19.4. Good luck!
Fundamental Chemistry II
Ria Astrid Yngard
Study Guide
Auburn University, Chemistry, Chem, chem1040, Yngard, exam, exam3, au
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This 7 page Study Guide was uploaded by Hannah B. on Thursday April 21, 2016. The Study Guide belongs to CHEM 1040 at Auburn University taught by Ria Astrid Yngard in Winter 2016. Since its upload, it has received 67 views. For similar materials see Fundamental Chemistry II in Chemistry at Auburn University.


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Date Created: 04/21/16
CHEM 1040 YNGARD EXAM 3 REVIEW Important Term Important PrincipleImportant Concept Important Equation Important Shit CHAPTER 17: ACID-BASE EQUILIBRIA AND SOLUBILITY EQUILIBRIA CHAPTER 17.1: THE COMMON ION EFFECT -An aqueous solution of a weak electrolyte contains both the weak electrolyte and its ionization products, which are ions -If a soluble salt that contains one of these ions is added, the equilibrium shifts to the left, thereby suppressing the ionization of the weak electrolyte -common ion effect: when a compound containing an ion in common with a dissolved substance is added to a solution at equilibrium, the equilibrium shifts to the left CHAPTER 17.2: BUFFER SOLUTIONS -a solution that contains a weak acid and its conjugate base (or a weak base and its conjugate acid) is a buffer, or a buffer solution -buffer solutions resist changes in pH upon addition of small amounts of either an acid or base by converting it from strong to weak -Henderson-Hasselbalch equation pKa = -log[H ] -effective buffers 1. 10 >/= (conj. base / weak acid) </= 1 2. pH = pKa +/- 1 (buffer range) CHAPTER 17.3: ACID-BASE TITRATIONS -strong acid - strong base titrations OH-(aq) + H+(aq) —> H2O(l) -initial pH determined by strong acid -between initial pH and equivalence point: excess of acid -at equivalence point: acid neutralized (pH = 7.0) -after equivalence point: excess of base -weak acid - strong base titrations CH3COOH(aq) + OH-(aq) —> CH3COO-(aq) + H2O(l) -higher initial pH, more gradual change, shorter vertical region near equivalence point -initial pH: use Ka for acid since weak -equivalence point: greater than 7.0 -can use Henderson-Hasselbalch to solve -strong acid - weak base titrations H+(aq) + NH3(aq) —> NH4+(aq) -equivalence point: less than 7.0 -the equivalence point is the point at which the acid has been neutralized completely by the added base -an acid-base indicator is usually a weak organic acid or base for which the ionized and un- ionized forms are different colors -the endpoint of a titration is the point at which the color of the indicator changes -phenolphthalein changes color between 8.3 - 10.0 -methyl red changes color between 4.2 - 6.3 Page 1 of 7 CHAPTER 17.4: SOLUBILITY EQUILIBRIA AgCl(s) <—> Ag+(aq) + Cl-(aq) Ksp = [Ag+][Cl-] -Ksp is the solubility product constant -the smaller the Ksp value, the less soluble the compound -molar solubility: the number of moles of solute in 1L of a saturated solution (mol/L) -solubility: the number of grams of solute in 1L of a saturated solution (g/L) -for predicting precipitation reactions use reaction quotient (Q) -if Q = Ksp —> solution saturated -if Q < Ksp —> no precipitate -if Q > Ksp —> precipitate CHAPTER 17.5: FACTORS AFFECTING SOLUBILITY -common ion effect -ion is less soluble with common ion effect than when in pure water solution -pH -adding OH- ions (increasing the pH) shifts the equilibrium left, decreasing the solubility -adding H+ ions (decreasing the pH) shifts the equilibrium right, increasing the solubility -complex ion formation -a complex ion is an ion containing a central metal cation bonded to one or more molecules or ions -a measure of the tendency of a metal ion to become complex is given by the formation constant (Kf) (“stability constant”), which is the equilibrium constant for complex ion formation -the larger the Kf, the more stable the complex ion is, the more soluble CHAPTER 17.6: SEPARATION OF IONS USING DIFFERENCES IN SOLUBILITY -even when both products are insoluble, we can still achieve some degree of separation by choosing the proper reagent to bring about precipitation -fractional precipitation is a technique that separates ions from solution based on their different solubilities -qualitative analysis is the practice of using principle of selective precipitation to identify the types of ions present in a solution Page 2 of 7 CHAPTER 18: ELECTROCHEMISTRY CHAPTER 18.1: BALANCING REDOX REACTIONS -electrochemistry is the study of the relationships between electrical energy and chemical reactions (redox reactions) -redox reaction: reaction of electron transfer -gain of electrons: reduction the oxidizing agent (oxidant) is reduced -loss of electrons: oxidation -the reducing agent (reductant) is oxidized -balancing redox reactions in acidic media 1. separate the reaction into half-reactions (an oxidation or reduction that occurs as part of the overall redox reaction) 2. balance with regards to atoms other than O and H 3. balance for O by adding H2O 4. balance for H by adding H+ 5. balance for charge by adding electrons 6. match up number of electrons in each half-reaction 7. add balanced half-reactions back together -balancing redox reactions in basic media -use OH- to balance O instead of H2O CHAPTER 18.2: GALVANIC CELLS -the experimental apparatus for generating electricity through the use of a spontaneous reaction is called a galvanic cell -the zinc and copper bars are called electrodes -the anode is the electrode at which oxidation occurs (anions move towards anode) -the cathode is the electrode at which reduction occurs (cations move towards cathode) -each combination of container, electrode, and solution is called a half-cell -salt bridge: an inverted U tube containing an innert electrolyte solution -electrical current flows from the anode to the cathode because of a difference in the gravitational potential energy -the difference in electrical potential between the anode and cathode is measured by a voltmeter and the reading is called the cell potential (Ecell) -depends on nature of electrodes and ions in the solution, and concentrations of the ions, and the temperature operated CHAPTER 18.3: STANDARD REDUCTION POTENTIALS -standard half-cell potentials are referencedOto a standard hydrogen electrode (SHE) 2H+ (1M) + 2e- —> H2 (1 atm) E = 0V -standard cell potential, E cell equation O -a +E means that the redox rxn will favor the formation of the products at equilibrium -a -E means that reactants will be favored at equilibrium -the strongest oxidants have the most positive reduction potentials; the strongest reductants have the most negative reduction potentials -standard reduction potential is an intensive property (does not depend on the amount of substance involved) Page 3 of 7 CHAPTER 18.4: SPONTANEITY OF REDOX REACTIONS (STANDARD STATE CONDITIONS) electrical energy (J) = cell potential (V) x total electrical charge (C) total charge = nF n: number of moles of electrons that pass through the circuit F: Faraday constant, the electrical charge contained in one mole of electrons 1F = 96,500 C/mol e- 1F = 96,500 J/V x mol e- Wmax = Welectrical = -nFEcell = G G = —RTlnK —> -nFE cell = RTlnK -by converting to the base-10 logarithm of K, we get: O O G < 0, E > 0, K > 1 G > 0, E < 0, K < 1 CHAPTER 18.5: SPONTANEITY OF REDOX REACTIONS (NOT STANDARD STATE CONDITIONS) 2 Nernst Equation -as Q increases, E decreases; until they reach equilibrium and E=0 & Q=K -a concentration cell is a galvanic cell from two half-cells composed of the same material but differing in ion concentrations Page 4 of 7 CHAPTER 18.6: BATTERIES -a battery is a galvanic cell, or a series of connected galvanic cells, that can be used as a portable, self-contained source of direct electric current -dry cell battery: -alkaline battery: -lead storage battery: (“the battery of the future”) -lithium-ion battery: -fuel cells: -fuel cell: a galvanic cell that requires a continuous supply of reactants to keep functioning and gets contaminated very easily Page 5 of 7 CHAPTER 18.7: ELECTROLYSIS -the use of electric energy to drive a nonspontaneous rxn is electrolysis -an electrolytic cell is one used to carry out electrolysis -under ordinary atmospheric conditions (1atm & 25 C), water will not spontaneously decompose to form hydrogen and oxygen gas because the standard free-energy charge for the reaction is a large positive quantity -the overvoltage is the difference between the calculated voltage and the actual voltage required to cause electrolysis -Faraday: “for any half-reaction, the amount of substance reduced/oxidized is directly proportional to the number of electrons passed into the cell” 1 charge (C - coulomb) = 1 current (A - amperes) x 1 time (s - seconds) -a coulomb is the quantity of electric charge passing any point in the circuit in 1s when the current is 1A -charge (coulombs) / Faraday constant (96500) = number of mols of electrons -mols substance / mols e- = moles of substance reduced or oxidized -mols of substance x molar mass of substance = grams of product CHAPTER 18.8: CORROSION -corrosion refers to the deterioration of a metal by an electrochemical process -redox reaction for rust: + 2+ 2Fe(2+ + O2(g) + 4H (aq) —> 2Fe (g) + 2H2O(l) + 4Fe (aq) + O2(g) + (4+2x)H2O(l) —> 2Fe2O3(xH2O)(s) + 8H (aq) -prevention of corrosion: coat surface with paint, coat surface with oxide layer (passivation) or other metals (cathodic protection) -galvanization, an example of cathodic protection, involving zinc-plating CHAPTER 19.1-19.4: CHEMICAL KINETICS CHAPTER 19.1: REACTION RATES -chemical kinetics is the study of how fast reactions take place -factors that increase reaction rate: increased reactant concentration, increased temperature, increased surface area of a solid reactant, and the presence of a catalyst CHAPTER 19.2: COLLISION THEORY OF CHEMICAL REACTIONS -general chemical reaction equation: reactants —> products -reactions generally occur as a result of collisions between reacting molecules -a greater frequency of collisions usually leads to a higher reaction rate -according to the collision theory of chemical kinetics, the reaction rate is directly proportional to the number of molecular collisions per second: -a collision that does result in a reaction is called an effective collision which is effected by… -the activation energy (Ea) is the minimum amount of energy required to initiate a chemical reaction -molecules must also be oriented in a way that favors reaction Page 6 of 7 -when an effective collision occurs between molecules, they form an activated complex (a state also known as the transition state), a temporary species formed by the reactant molecules as a result of the collision CHAPTER 19.3: MEASURING RXN PROGRESS AND EXPRESSING RXN RATE -reaction rate: the change in concentration of a reactant or product with time -average reaction rate equation: -negative sign on [A] shows that the concentration is decreasing over time -do not make your value negative, the sign just shows whats happening -the instantaneous rate is the rate for a specific instant in time; equal to the slope of a tangent to the curve at any particular time -rate = k[Br2] —> k = rate/[Br2] -k is called the rate constant, and is constant at a constant temperature -stoichiometry and reaction rate -in general: aA + bB —> cC + dD -the rate is given by: CHAPTER 19.4: DEPENDENCE OF RXN RATE ON REACTANT CONCENTRATION -rate = k[Br2] is an example of a rate law (an equation that relates the rate of reaction to the concentrations of reactants -for the reaction aA + bB —> cC + dD, the rate law is: -k is the rate constant -x and y are numbers that must be determined experimentally -the sum of x and y is called the overall reaction order -experimental determination of the rate law: -the initial rate is the instantaneous rate at the beginning of the reaction -check to see if rates change proportionally -if yes, exponent is 1 -if no, exponent is determined based on change in experimental data Page 7 of 7


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