Chem 130 Study Guide Exam 2!
Chem 130 Study Guide Exam 2! CHEM 130 - 003
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This 18 page Study Guide was uploaded by Izabella Nill Gomez on Friday September 25, 2015. The Study Guide belongs to CHEM 130 - 003 at University of Tennessee - Knoxville taught by Bin Zhao in Summer 2015. Since its upload, it has received 177 views. For similar materials see General Chemistry II in Chemistry at University of Tennessee - Knoxville.
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Chem 130 Exam 2 Study Guide Chapter 14 Reaction rate the speed at which a chemical reaction occurs Reaction mechanism step by step molecularlevel view of the pathway from reactants to products Chemical kinetics area of chemistry concerned with the speedrates of reactions 4 factors that allow us to change the rate at which a reaction occurs 1 Physical state of the reactants The more readily the reactants collide with one another the more rapidly they react Most are homogeneous Heterogeneous conditions limit the area of contact of reactants They tend to proceed faster if the surface area of the solid is increased 2 Reactant concentrations Most chemical reactions proceed faster if the concentration is increased 3 Reaction temperature Reaction rates generally increase as temperature increases As molecules move more rapidly they collide more frequently 4 Presence of a catalyst Catalysts are agents that increase the reaction rates without themselves being used up The speed of an event is de ned as the change that occurs in a given time interval The reaction rate of a chemical reaction is the change in the concentration of reactantsproducts per unit of time Ms changeEconcentration of B MB A B Average rate of appearance B Changee me At changeEconcentration of A AA Average disappearance of A ChangeEtlme T N Rates are always expressed in positive quantities It is typical for rates to decrease as a reaction proceeds because the concentration of reactants decreases Ex C4H9ClaqH20l gtC4H90HaqHClaq f C4H9Cl Measuring concentration 0 various times after t0 The resulting data is used to calculate the average rate of disappearance of C4H9Cl Graphs showing how concentration of a reactant or product changes with time allows us to evaluate the instantaneous rate of a reaction the rate at a particular instant during a reaction determined from the slop of the curve at a particular point in time datataken 017 042M 800 400s AC4H9Cl Instantaneous rate T 63105 l graphl 5 At t 01 instantaneous rate This is the initial rate of the reaction The rate of appearance of a compound equals the rate of disappearance of the other 0411901 6 Rate AZ Z In general for a reaction aAbBCCdD Lowercase letters are coef cients 1AA 1AlBl lmlcl mlDl Rate is given by Rate a b C At At At At Changing the initial concentration of either reactant changes the initial reaction rate Z aQN282H20l Ex NH Z 6 If NH doubled and N0 held constant the rate doubles The same if the rst compound is increased by a factor of 4 the reaction increases by 4 The way reaction concentrations are depended by the rate is expressed Ratek NH N0 Rate lawk A B wherekis the rate constant magnitude changes with temperature and determined how the temperature affects rate Product concentration does not appear in the rate aw rate law is for reactants There is a linear relationship between reaction rate and concentration 1 n Reaction orders m and n in the k reacmml reacmmz Ex Because the Z Z exponent of NH is 1 the rate is rst order in N0 Overall reaction order is the sum of the orders with respect to each reactant represented in the rate law 112 reaction is second order overall Exponents in a rate law can indicated how the rate is affected by each reaction concentration Ex rate depends on how many Z g powers NH is increased The same with NO 2 2 2 If a rate law is second order with respect to reactant A then 2 4 3 9 For any reaction the rate law must be determined experimentally 9 In general k 10 or higher is a fast reaction and k 10 or lower is a slow reaction 2 Units of rate units of rate constantunits of concentration 6 2 umts of concentration o 4 6 units of rate Z Units of rate constant for a reaction of second order overall In most reactions reaction orders mn are 012 If 0 the reactant has no effect on the reaction Rate of a reaction depends on the concentration but not the rate constant Rate constant is affected by temperature or a catalyst First order reaction one whose rate depends on the concentration of a single AlAl reactant raised to the rst power Ex rate M klA this expressed how rate depends on the concentration differential rate law 1nAt kt1nA0 Integrated rate law similar to ymxb Second order reaction one whose rate depends on either reactant concentration raised to the second power or concentrations of 2 reactants each raised to the rst 1 1 kz powequot W W Zero order reaction one in which the rate of disappearance of A is independent of A Ex gas in decomposition on the surface of a solid ALI ktA0 Overall Reaction Order Units for k 1 Zero MS 1 First S M 1 1 Second S I I I M1 0verallorder 1 The unIt of k for a reaction of any overall order IS 5 Halflife time required for the concentration of a reactant to reach half its initial value A 12lAlo Convenient to describe how fast it occurs especially for rst t12 693 orders Fast reactions have short half lifes t12 k For a rst order rate law t does not depend on initial concentration of any reactant Halflife is constant throughout the reaction In a rst order reaction the concentration of the reactant decreases by 12 in each of a series of regularly space time interval each equivalent to t12 For a second order rate law halflife depends on the initial concentration of the 1 M N NIH reactant the lower the initial concentration the longer the halflife The rates of most chemical reactions increase as temperature rises The faster rate is due to an increase in the rate constant with increasing temperature Rate constant k is temperature dependent Approximately the rate of reaction doubles for each 100 C rise The central idea for the collision model is that molecules must collide to react The greater the frequency the greater the reaction rate Collision Model based on the kineticmolecular level theory it accounts for both effects at the molecular level Molecules must collide to react The greater the number of collisions the bigger the increase in reaction rate Increasing temperature increases molecular speeds As molecules move faster they move more forcefully and more frequently increasing reaction rates Orientation factor molecules must be oriented in a certain way during collision for a reaction to occur The ineffective collision of molecules will not result in a reaction Molecules must possess a certain amount of energy to react and this minimum energy required to activate a chemical reaction is the activation energy Ea This is the energy barrier molecules must overcome for a reaction to occur The lower the activation energy the more molecules that can participate in a reaction During a reaction the chemical compound that is being reacted must twist the bond 1800 at the highest energy point where it is ready to break into its components After the climax chemicals change and form a new bond If it does not pass the energy barrier it returns to its original form A H lt 0 exothermic A H gt0 endothermic EaAE The activation energy for the reverse reaction is The rate depends on the magnitude of Ed generally the lower the value the faster the reaction At a lower temperature molecules have less energy to react At a higher temperature a larger amount of molecules have higher energy As temperature increases the fraction of molecules that can overcome the activation barrier increases The collision frequency also increases As a result reaction rate increases Activation energy minimum energy required to initiate a chemical reaction Ed is the difference between the energy of the starting molecule and the highest energy along the reaction pathway Activated complextransition state the molecule having the arrangement shown at the top of the barrier Rate depends on the magnitude of Ed generally the lower the value of Ed the faster the reaction The fraction of molecules that have energy equal or greater than Ed is given by fzeRT Where R is a gas constant and T is the absolute temperature Arrheniusfound that for most reactions to increase in rate with increasing temperature the graph is nonlinear Most reaction rate data obeyed an equation on a the fraction of molecules possessing Ea or greater b the number of collisions per second c fraction of collisions that have proper orientation Ea kAeRT Frequency factor A is constant as temperature is varied Related to frequency of collisions and probability that collisions are favorably oriented for a reaction As Ea increases k decreases Reaction rates decrease as Ed increases Ink 1nA RT E Graphically a can be determined as E where a is the slope klE 1 1 of the resultant line y mxb Nongraphically lnk2 T23 Cl where T is in Kelvin and R is a constant 8314 JmolK Reaction Mechanism the steps by which a reaction occurs Describes the order in which bonds are broken and formed and the changes in relative positions of the atoms in the course of the reaction Reactions may occur all at once of through several discrete steps Elementary Reactions reactions that occur in a single eventstep Ex N033033AN023023 This is a bimolecular reaction as it contains two molecules that participate as reactants Molecularity de nes the number of molecules that participate as reactants Termolecular 3 reactions are far less likely to occur than unimolecular 1 or bimolecular 2 The net change represented by a balanced chemical equation often occurs by a multistep mechanism consisting of a sequence of elementary reactions N02lglN02lgl gtN03lglN0lgl N03glCOggtN02lgl002g The chemical equations for the elementary reactions in a multistep mechanism must always add to give the chemical equation of the overall process 2N02lglN03gC0lg gtNOslglN02lglCOzlgl N02lglC0lgl gtN08C02lgl Both sides of the equation in the rst part must cancel if identical to achieve the resultant balanced equation N03 is neither an initial reactant nor a nal product so it is an intermediate produced after the rst reaction and consumed in the next If a reaction is elementary its rate law is based directly on its molecularity Aapmducm39 As number of molecules A increases the number that react in a given time increases RatekA For bimolecular AB pr0ducm the RatekAB Rate determininglimiting The slowest step in a reaction mechanism that limits the overall reaction rate has the highest EA Governs the rate law for the overall reactionit is the rate law for the overall reaction In general whenever a fast step precedes a slow one we can solve for the concentration of an intermediate by assuming that an equilibrium is established in the fast step Catalyst substance that changes the speed of a chemical reaction without undergoing a permanent chemical change itself A homogeneous catalyst is present in the same phase as the reactants in a reaction mixture Ex aq gtBr2aq2H20l Zaq H202aq2HZ 2Brz aQ02g Zaq2HZ aq gt2Brz Br2aqH2 02aq2HZ Z The catalyst Brquot is there at the start and end of the reaction and the intermediate is formed during the course of the reaction The catalyst affects the E numerical value of k determined by a and lowers it Heterogeneous catalyst exists in a phase different from the phase of the reactant molecules usually as a solid in contact with gas Composed usually of metalmetal oxides Initial step is usually absorption where molecules are taken up into the interior of a substance Adsorption binds molecules to a surfaceoccurs because atoms or ions of a surface of a solid are extremely reactive with unused bonding capacity Enzymes are biological catalysts The reaction any given enzyme catalyzes takes place at the active site Substances that react are called substrates and operate under the lockandkey model Combination of enzyme and substrate is enzyme substrate complex Chapter 15 Chemical equilibrium occurs when opposing reactions proceed at equal rates the rate at which the products form from reactants equals the rate at which the reactants form the products As a result concentrations cease to change making it appear stopped Equilibrium state mixture of reactants and products whose concentrations no longer change with time After compounds dissociate and reform what is left is an equilibrium mixture of both substancesif in a closed system equilibrium will eventually be reached N204 Equilibrium can be reached if the reaction is reversible ex can form N02 39 and N02 N204 can form N2048 2N023 colorless to brown 15 N204 f N204 Decomposition o is a forward reaction and the formation 0 is the reverse reaction N0 z N204g gt2N02g RatefkfN204 2 EX 2N02g gtN204gRaterer N0 2 Nazi 26 At equilibrium kfN204 k z 3 r At equilibrium concentrations no longer change however the equilibrium is dynamic so some compounds are always transforming At equilibrium the concentration of reactants and products no longer change with time For equilibrium to occur neither reactants nor products can escape from the system At equilibrium a particular ratio of concentration term equals a constant Ex N2g3H2g2NH3g Haber process is critical for production of fertilizers The above equation illustrates N2 and H2 that react at high temperature and pressure in the presence of a catalyst to form ammonia However in a closed system this does not lead to the complete consumption all 3 substances are present at equilibrium Law of mass action experiments carried out to discover how to analyze the gases in an equilibrium mixture Expresses for any reaction the relationship between concentrations of the reactants and products present at equilibrium aA bB lt gt dDeE This is the equilibrium constant expression Equilibrium constant Kc numerical value obtained when we substitute molar equilibrium concentrations into the equilibrium constant expression Concentrations are expressed in molarity the equilibrium constant expression depends only on the stoichiometry of the reaction not the mechanism Kc at any given temperature does not depend on the initial amounts of reactants and products Does not matter whether other substances are present as long as they do not react with a reactant or product When reactants and products in a chemical reaction are gases we can form the equilibrium constant expression in terms of partial pressures denoting Kc as p p for pressure PA is the partial pressure of A PB is the partial pressure of Betc For a given reaction the numerical value of K6 is different from KP RT 2 Since pVZnRTaszch where KP is the partial pressure equilibrium constant K6 is the equilibrium constant and R is a constant 0821 T is in Kelvin and Anzz moles of gasaqueous productmoles of gasaqueous reactant Ex N2048 2N02gAn2 11 Equilibrium constants are reported without units units present cancel Magnitude of the equilibrium constant for a reaction give important information about the composition of the equilibrium mixture If K 1 large Kequilibrium lies to the right products predominate If K 1 small Kequilibrium lies to the left reactants predominate The equilibrium constant expression for a reaction written in one direction is the reciprocal of the expression for the reaction written in reverse NOZLZ EX N204g2N02g 2N028N2048 N204L2 a If we multiply by 2 212g 6 Z 2N204g 4N028KCZZ You must relate each equilibrium constant you work with to a speci c balanced chemical equation Substances39 concentrations remain the same no matter how you write the equation but the Kc you calculate depends on how you write the reaction It is also possible to calculate Kc if we know the equilibrium constants for the other reactions 2N03rglt gt2N0gBr gK 2 014 EX 2 cl 39 Br LL2C1272 239 BrCl2 Br2gClzgltgt2BrClgKcz Z 3 Net sum ZNOBrgCl2g 2N0g239BrClgKc3KclKCZ Homogeneous equilibria substances that are all in the same phase Heterogeneous equilibria substances that are in different phases Whenever a pure solid or pure liquid is involved in a heterogeneous equilibrium its concentration is not included in the equilibrium constant expression 2 P19quot 3 zz EX ZCZL ZaqKcZ 2Zaq2ClZ Pbczzlglt gtPb 3 Concentration of pure solidliquid remains constant If we do not know the equilibrium concentrations of all species in an equilibrium mixture we can use stoichiometry of the reaction to deduce the equilibrium concentrations of others 1 Tabulate all known initial and equilibrium concentrations of species that appear in the equilibrium constant expression 2 For those species for which initial and equilibrium concentrations are known calculate the change in concentration that occurs as the system reaches equilibrium 3 Use stoichiometry of the reaction to calculate changes in concentration for all other species in the equilibrium constant expression 4 Use initial concentrations from step 1 and changes in concentration from step 3 to calculate any equilibrium constant expression 5 Determine the value of the equilibrium constant With the equilibrium constant you can predict the direction in which a reaction mixture achieves equilibrium and calculate equilibrium concentrations of reactants and products Reaction quotient Q is a number obtained by substituting reactant and products concentrations or partial pressures at any point during a reaction into an equilibrium constant expression d e aAbBlt gtdDeE chw A B Unlike Kc the reaction quotient varies as the reaction proceeds Q can tell us if a reaction is at equilibrium good for slow reactions QK reaction quotient equals the equilibrium constant only if the system is at equilibrium Q Z K concentration of products is too large and that of reactants too small Substances from right to left Q Z K concentration of products too small and of reactants too large Proceeds from left to right achieves equilibrium by forming more products Le Chatelier39s principle If a system at equilibrium is disturbed by a change in temperature pressure or a component concentration the system will shift its equilibrium position so as to counteract the effect of the disturbance If a substance is added to a system at equilibrium the system reacts to consume some of the substance If the substance is removed from the system the system reacts to produce more of the substance reducing the volume of a gaseous equilibrium mixture causes the system to shift in the direction that reduces the number of moles of gas If the temperature of a system at equilibrium is increased the system reacts as if we added a reactant to an endothermic reaction or a product to an exothermic reaction The equilibrium shifts in the direction that consumes the quotexcess reactantquot namely heat A system at dynamic equilibrium is in a state of balance When concentrations of species are altered the equilibrium shifts until a new balance is achieved Equilibrium constant remains the same If a chemical system is already at equilibrium and the concentration of any substance is increased either reactant or product the system reacts to consume some of that substance If the concentration is decreased in a substance the system reacts to produce some of that substance Ex Nzlgl3H2gltgt2NH33 H2 Adding causes the system to shift to reduce the concentration of H2 in the system Can occur only if H2 is consumed and simultaneously consumes N2 NH3 NH3 to form more Adding N2 also causes more to be formed NH3 NH3 NH3 Removing produces more and adding causes less to be made decomposing NzAHz If a system containing 1 or more gases is at equilibrium and the volume is decreased increasing pressure the system responds by reducing the pressure by reducing the number of gas molecules At constant temperature reducing the volume of a gaseous equilibrium mixture results in the system shifting in the direction that produces more gas molecules Ex N233H2392NH33 Increasing the pressure causes more formation of NH3 But with H2g12gHHIg increasing pressure does not affect the equilibrium Endothermic reactants heat 9 product Exothermic reactants H productheat When temperature of a system at equilibrium is increased the system reacts as if we added a reactant to an endothermic reaction or a product to an exothermic one The equilibrium shifts in direction that consumes the excess reactant or product namely heat Increasing T causes the equilibrium to shift to the right to more products K increases in an exothermic reaction the result is the opposite Cooling has an opposite effect The lower the temperature the equilibrium shifts to where there s heat endothermic to the left K decreases and exothermic to the right K increases
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