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Chem 109 at UW-Madison Midterm 3

by: y-chen9

Chem 109 at UW-Madison Midterm 3 Chem 109

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Study guide for chem 109 midterm #3 at UW-Madison. Covers reaction rates, enzyme kinetics, thermochemistry, enthalpy/entropy, Gibbs Free Energy, equilibrium, Le Chatelier's Principle
Advanced General Chemistry
Ive Hermans
Study Guide
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This 12 page Study Guide was uploaded by y-chen9 on Thursday November 12, 2015. The Study Guide belongs to Chem 109 at University of Wisconsin - Madison taught by Ive Hermans in Fall 2015. Since its upload, it has received 284 views. For similar materials see Advanced General Chemistry in Chemistry at University of Wisconsin - Madison.


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Date Created: 11/12/15
Chemistry 109: Midterm III Review By: Yang Chen Professor: Ive Hermans TA: Hangjian Zhao Textbook: “Chemistry: The Molecular Science” Moore & Stanitski (5e) _________________________________________________________________________ ________________________ CHAPTER 4-Review all) Energy and Chemical Reactions 4-1) The Nature of Energy  Thermodynamics: Physical science of heat and temperature in relationship to work and energy  Energy  Kinetic Energy  Potential Energy (ex. Gravitational, Electrostatic, Potential)  Calorie: Amount of E to raise 1 g water 1 degree Celcius o 1 cal = 4.184 J  Power: Energy per unit of time 4-2) Conservation of Energy  First Law of Thermodynamics/Law of Conservation of Energy: E can’t be created or destroyed  Heating is a process of transfer of kinetic E 4-3) Keeping Track of Energy Transfers  System: E in region concerned with vs Surroundings: things that exchange E with system  Internal Energy: Sum of all energy in the system o Depends on T, kind of particle and number of particles  Change in E = q (quantity of E transferred by heating the system) + w (quantity of E transferred by doing work on the system)  Energy transferred into a system (Positive ΔE)  Energy transferred out of a system (Negative ΔE) 4-4) Heat Capacity  Heat Capacity: Quantity of E required to raise T of a sample 1 degree  Specific Heat Capacity: Quantity of E needed to raise the T of 1 g of substance by 1 degree C. Used to distinguish one substance from another. o Specific heat capacity of water: 4.186 J/g * C o Spec. heat Cap. = Qaunt. of heat E transfer/sample mass * change in T  C = q/m * ΔT  Molar Heat Capacity: Quantity of E that needs to be transferred to raise the T of 1 mole of a substance 1 degree C o Given specific heat capcity of molar mass of a substance, we can calculate its molar heat capacity in J/mol* C 4-5) Energy and Enthalpy  “q”: Quantity of Heat E Transfer  Exothermic: Energy transferring out of system (-)  Endothermic: Energy transferring into system (+)  Enthalpy: Heat transfer at constant pressure o H = U (sum of internal E) + PV (Pressure * Volume) o Δ Enthalpy = ΔH = q  Fusion Enthalpy (Δ fusH Heat E transfer when melting, from solid to liquid  Freezing Enthalpy (Δ freezing Heat E transfer from liquid to solid  Vaporization Enthalpy (Δ vap): Heat E transfer from liquid to vapor gas  Condensation Enthalpy (Δ condensationHeat transfer from vapor gas to liquid 4-6) Reaction Enthalpies for Chemical Reactions  Thermochemical Expression: Balanced chemical equation with total change in reaction enthalpy  Standard Reaction Enthalpy (Δ H) (rnits- kJ/mol): Enthalpy of pure rxn product – rxn reactants at standard pressure of 1 bar and specific T 4-7) Where Does the Energy Come From?  Bond breaking – Endothermic  Bond making – Exothermic o Bond Enthalpy/Bond Energy  Almost always positive 4-8) Measuring Reaction Enthalpies: Calorimetry  Calorimeter: Device that measures heat transfers 4-9) Hess’s Law  Regardless of the multiple stages or steps of a reaction, the total enthalpy change for the reaction is the sum of all changes o Bc enthalpy is a state function (only depends on state at moment, T, volume, pressure, etc)  1) Examine target expression you want to calculate change in enthalpy for  2) Rearrange step equations in order to represent target expression  3) Assign coefficients to reactants and products in step equations so that they are the same as those in target expression. Change sign of enthalpy as needed  4) Add together to get total change in enthalpy 4-10) Standard Formation Enthalpies  Standard Formation Enthalpy: (Δ): fnthalpy for formation of one mole of a compound at standard state (1 bar atmosphere and specific T)  You can calculate standard reaction enthalpy from standard formation enthalpy o Δ f = Sum(coefficient of product x enthalpy change of product) – Sum(coefficient of reactant x enthalpy change of reactant) 4-11) Fuels for Society and Our Bodies  Chemical fuels: Coal, natural gas, petroleum  Bodily fuels: Carbohydrates, fats, protein CHAPTER 11-all) Chemical Kinetics: Rates of Reactions 11-1) Rates of Reactions  Chemical Kinetics: Study of speed of reaction and nanoscale arrangements of atoms/molecules as they go from reactants to products or vice versa  Things that influence the rate of a reaction o Nature of reactants—some are more reactive than others  Concentrations of reactions  Increases the number of molecules per unit volume o Temperature of reaction  Raising the temperature increases the fraction of molecules that are energetic enough to surmount the activation energy barrier. o Presence of catalyst  For heterogeneous reactions, greater surface area of a catalyst means greater reaction rate  Participates in the reaction mechanism  Definition of reaction rate: o General formula: Rate = change in concentration/change in time  However, we must take in stoichiometry—take coefficients  Reactants have a “minus sign” o Rate = 1/coeef * change in conc/change in time o Instantaneous Rate: Rate = 1/coeff * dc/dt  Average Reaction Rate and Instantaneous Reaction Rate Chapter 11-2) Rate Law Order of Reaction  Main Idea: Rate laws and reaction orders are only determined experimentally. Once these values are determined, and we know the rate constant and initial concentration of reactant or product we can use the integrated rate law to calculate the concentration of a reactant or product at any time after the reaction has begun.  Rate law for almost all homogenous reactions: Rate = k[A] [B] . m n o k: temperature o m/n: orders of reaction with respect to A and B  usually positive whole numbers but could be negative and or a fraction  sum of m and n and any other exponents gives the overall reaction orer o A/B: Could be reactants, products, or catalysts  Initial Rate Chapter 11-3) Rate Law and Order of Reaction  Integrated Rate Law: Experimental determination of rate law and rate constant using calculus principles   Zeroth Order: A reaction is 0 order if plotting [A] vs t gives a straight line  First Order: ..plotting ln[A] vs t gives a straight line  Second Order: ..plotting 1/[A] vs t gives a straight line  Half-Life (t ): Time required for the concentration of a reactant A 1/2 fall to one half of its initial value o Half-life is related to the 1 -order rate constant o Radioactive decay  Half-Life Equations:  Determining order: Reaction rates for different reactions show different dependencies on reactant concentrations. Generally, reaction rates increase with increasing concentration, but not always.  When reaction rate is directly proportional to reactant concentration, rate = constant × [reactant] the reaction is first-order.  When reaction rate is directly proportional to the square of reactant concentration, rate = constant × [reactant] 2 the reaction is second-order.  When reaction rate does not depend on reactant concentration, rate = constant the reaction is zero-order. CHAPTER 11-4) A Nanoscale View: Elementary Reactions Main Idea: Kinetic-molecular theory of matter and ideas about molecular structure help us understand how atoms and molecules move and how their bonds are made and broken. This understanding is known as elementary (nanoscale) reactions—They are the “building blocks” to more complicated large scale chemical reactions that take place in nature.  There are 2 important types of molecular transformations o Unimolecular reaction: One single particle rearranges structurally to form a different particle.  Breaking a bond to form 2 new molecules, rearrangement of one isomeric structure into another  Molecularity of 1  Ex: cis-2-butene  trans-2-butene  Transition state/activated complex  Higher energy barrier = slower the reaction. The minimum E required to surmound the barrier is called activation E (E a. o Bimolecular reaction: 2 particles collide and rearrange bonds to form products  New bonds are formed between reactants and existing bonds can be broken. 2 particles can come together to form a larger particle. 2 or more new particles can also be formed from 2 reactants  Molecularity of 2  Steric Factor: Limiting of collisions depending on 3D geometric shapes of reacting molecules  Reaction Rates speed up with temperature because higher temperatures lead to a greater fraction of reactant molecules have enough energy to pass the activation energy barrier.  A reaction is faster at a higher T bc its rate constant is bigger o Rate constants are only constants for a specific reaction at a given temperature 11-5) Temperature and Reaction Rate: The Arrhenius Equation  Arrhenius Equation: Can be used to calculate the rate constant at any temperature  k(T) = Ae -Ea/RT -1) o A: frequency factor (sec  Depends on (steric factor) how often molecules collide when all concentrations are 1 mol/L and on whether the molecules are properly oriented when the collide o e: base of natural log -1 o E a activation E (kJ mol ) o R: Ideal gas law constant  8.314 J mol K -1 -1 o T: Temperature in Kelvin Chapter 11-6) Rate Laws for Elementary Reactions  Rate Law for Unimolecular Reaction: o A  products:  Rate = k[A]  Rate Law for Bimolecular Reaction: o A + B  products:  Rate = k[A][B] o A + A  products 2  Rate = k[A] Chapter 11-7) Reaction Mechanisms  Reaction Mechanism: A set of equations for elementary reactions o Any valid mechanism must consist of a series of unimolecular or bimolecular elementary steps to the overall reaction and can correctly predict the experimentally observed rate law  Step 1: Rate-Limiting Step (fast)  The rate of the overall reaction is limited by, and equal to, the rate of the slowest step in the mechanism  Step 2: (slow) Chapter 11-8) Catalysts and Reaction Rate  Catalysts Cis-Trans: o Step 1: Dissociation o Step 2: Attachment of I atom to cis-2-butene o Step 3: Rotation around the C-C bond o Step 4: Loss of an I atom and reformation of C=C o Step 5: Regeneration Chapter 11-9) Enzymes: Biological Catalyst  Enzyme: Biological catalysts that speed up reactions by lowering activation E o Very specific o Usually proteins or other macromolecu9es 19ke RNA o Can increase reaction rates by 10 -10 o Some enzymes need cofactors which are organic or inorganic molecules or ions that allow the enzyme to be catalytic  Substrate: Reactant to be catalyzed by enzyme  Active Site: Part on enzyme that binds to substrate  Induced Fit: The new shape of the enzyme-substrate complex 11-10) Catalysis in Industry  Heterogeneous Catalysts: Catalysts in different phase from reactants o Acetic Acid Chapter 12.1-3, 6) Chemical Equilibrium 12-1) Characteristics of Chemical Equilibrium  Chemical Equilibrium: Concentrations of reactants and products are equal o During chemical equilibrium a reaction can be reactant- favored or product-favored o The equilibrium concentrations are reached whether you start with reactants or products  Dynamic Equilibrium: On going reaction at equilibrium as rate of reactants forming is equal to rate of products forming  Catalysts do not affect equilibrium concentrations. This is because catalysts speed up the forward and back reaction rates equally. A catalyst will bring a reaction to equilibrium faster. 12-2) The Equilibrium Constant  (Kc) = k forwardreverseQuotient of equilibrium concentrations of product and reactant substances that is constant at a specific T. K c can help answer: o Do products or reactants dominate @ equilibrium? o What are the [reactants] and [products]? o Given initial [reactants] and [products], in which direction will the reaction need to go to achieve equilibrium?  When writing an equilibrium constant expression: o K =c[products]/[reactants]  Only include those in gas state and in dilute solutions. Not pure liquids and solids  Coefficients (# of moles of a compound)  Gets raised to the power, as an exponent of the concentration of that compound  See page 531 for examples  “Whenever the stoichiometric coefficients of a balanced equation are multiplied by some factor, the equilibrium constant for the new equation (K )c2s the old equilibrium constant (K ) raic1d to the power of the multiplication” (Moore & Stanitski, 532)  “If two chemical equations can be summed up to give a third, the equilibrium constant for the overall equation equals the product of the two equilibrium constants for the equations that were summed” (Moore & Stanitski, 533)  Pressure Equilibrium Constant (K ): p   In some gas-phase equilibrium K =K c p  You can calculate K from K p c Chapter 12-3) Determining Equilibrium Constants  ICE Table- Initial concentration (M), Change in concentration (M), Equilibrium concentration (M) Chapter 12-4) The Meaning of the Equilibrium Constant  Kctells us if a reaction is reactant or product favoring at equilibrium o If K cs very large, then almost all the reactants have been converted to products at equilibrium o If Kcis very small, then almost no reactants are converted to products at equilibrium o If Kcis close to 1, then appreciable quantities of both reactants and products at equilibrium o If Chapter 12-6) Shifting a Chemical Equilibrium: Le Chatelier’s Principle  Le Chatelier’s Principle: If a system at equilibrium is disturbed (by change in T, pressure, concentration of reactants), the system will counteract that change to re-maintain equilibrium. Chapter 16.1-6) Thermodynamics Reactant-Favored and Product-Favored Processes  Reactant-Favored: When reactants dominate at equilibrium  Product-Favored: When products dominate at equilibrium o Also called spontaneous o Examples: bromine w aluminum, rusting of iron & combustion of gasoline o Energy spreads out (disperses) unless it is hindered from doing so  Dispersal of energy occurs bc the probability is much higher that energy will be spread over many particles than that it will be concentrated in a few’ Chapter 16-2) Chemical Reactions and Dispersal of Energy  More E spread out  more product-favoring  Energy will naturally spread out unless something prevents it from doing so  “Dispersal of E occurs because the probability is much higher that energy will be spread over many particles than that it will be concentrated in a few” (Moore & Stanitski, 697) o This is especially true with a large number of atoms/molecules  Energy wants to be dispersed over as many particles as possible  Greater volume = E more dispersed  more product-favoring Chapter 16-3) Measuring Dispersal of Energy: Entropy  Entropy: A measure of molecular randomness, disorder, nanoscale dispersal of energy (J/K) o Can be measured with a calorimeter  Change in Entropy o ΔS = S final Sinitiarev/T  Third Law of Thermodynamics: Entropy of a system approaches a constant value as the T in Kelvins approaches 0.  Entropy gets bigger as you go from solid to liquid to gas…more dispersal of E  Entropy is bigger for bigger molecules  Entropies of ionic solids with similar formulas: The one with the weaker ionic bond has greater entropy  Entropy increases when a solid or liquid dissolves in a solvent Chapter 16-4) Calculating Entropy Changes  Δ r = Sum(coefficient of product x entropy of product) – Sum(coefficient of reactant x entropy of reactant) Chapter 16-5) Entropy and the Second of Thermodynamics  Second Law of Thermodynamics: Entropy in the universe is always increasing  Δ S Δ S + Δ S r universer=system r surroundings  If exothermic and entropy of products is greater than entropy of reactants…definitely exothermic..and vice-versa for endothermic  Δ r ΔrS Product- Δ r Favored? - + Yes - (Spontaneous) - - Yes @ low T - (S) No @ high T + (Not Spontaneous) + + No @ low T + (NS) Yes at high T - (S) + - No + (NS) Chapter 16-6) Gibbs Free Energy  Gibb’s Free Energy: The amount of Energy in a system that is available to do useful work  ΔG = ΔH – TΔS o This defines a combined enthalpy/entropy function that determines spontaneity Temperature impacts ΔG  The negative temperature times the entropy change of the universe is. o Decrease/negative/exothermic in Gibbs is product favored o Predicts whether rxn is product favored Source: Moore, John W., and Conrad L. Stanitski. Chemistry: The Molecular Science. N.p.: Cengage Learning, n.d. Print.


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