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Chem 130 Final Exam Study Guide!

by: Izabella Nill Gomez

Chem 130 Final Exam Study Guide! CHEM 130 - 003

Izabella Nill Gomez
GPA 3.81
General Chemistry II
Bin Zhao

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Hey guys! It's finally the end of the semester and time for one last push! Here is an accumulation of everything we've worked on this entire year! It includes all of my past study guides along with...
General Chemistry II
Bin Zhao
Study Guide
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This page Study Guide was uploaded by Izabella Nill Gomez on Monday November 30, 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 133 views. For similar materials see General Chemistry II in Chemistry at University of Tennessee - Knoxville.


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Date Created: 11/30/15
Chemistry 130 Final Exam 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 changeEconcentmtian of B MB A B Average rate of appearance 8 Changeaime At ChangeE concentration of A MA Average disappearance of A ChangeEtlme T M 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 C4H9CllaqH20lgtC4H90HlaqHClaq Measuring concentration of C4H9Cl various times after tO 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 Ac4H9Cll Instantaneous rate At 39 5Z 2M6310S l gmph S S At t Oi 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 C4H9Cl 6 Rate AZ 2 In general for a reaction AJ39bBach39dD Lowercase letters are coef cients 1AA 71ABl lmlcl lAlDl Rate is given by Rate a C d At At Ar Ar Changing the initial concentration of either reactant changes the initial reaction rate amplaqlN2lgl2H20ltlgt Ex NH Z 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 i 1H 2 Reaction orders m and n In the k madam madam 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 2 g powers NH is increased The same with N0 2 2 2 If a rate law is second order with respect to reactant A then 2 4lt 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 0f concentratlon c 6 units of rate 2 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 klAl 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 kt powequot Mi Mi 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 ALZ ktAO Overall Reaction Order Units for k 1 Zero MS 1 First S M 1 1 Second 5 I I I M1 0verallorder 1 The unIt of k for a reaction of any overall order IS S 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 IM 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 tlZ 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 O 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 Ea 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 a1nA RT E Graphically a can be determined as E where a is the slope klE 1 1 of the resultant line y mxb Nongraphically kHz F T57 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 N038038 N02802ltggt 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 unimoecuar 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 N02lglN02lglN03lglN0lgl N03lglC0lglN02lglC02g The chemical equations for the elementary reactions in a multistep mechanism must always add to give the chemical equation of the overall process 2N02gN038C08N038N028C028Z N02lglC0lgl N0 glC02lgl 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 Aapmcmas39 As number of molecules A increases the number that react in a given time increases RatekA For bimolecular ABapr0duas 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 ZaqgtBr2aq2H20l ampaqH202aq2H 28 lt3laql 028 Zaq2Hz ZaqgtZBrz Br2aqH2 02aq2H 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 and N02 N204 can form N204ltggt92N02ltggt colorless to brown f N204 f N204 Decomposition o is a forward reaction and the formation 0 is the reverse reaction N0 a N204ltggt 2N02ltggt RatefzkfN204 2 EX 2N02g gtN204gRaterkr N0 2 N022 26 Z At equilibrium kfN204 k g 2 I 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 pvanTerzKCg 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 N0222 EX N204ltggt92N02ltggt 2N02gN204g N20422 c If we multiply by 2 2128 6 Z 2N2048H4N028Kcamp You must relate each equilibrium constant you work with to a speci c balanced chemical equation Substances 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 2NOBrglt gt2N0gBr gK 2 014 EX 2 cl 39 Br 222C1272 239 BrCl2 Br2gC12glt gtZBrClgKcz 3 Net sum 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 Pbquot 2 42 Ex 26 ZaqKCZ 2Zaq2Clz PbC12glt gtPbquot 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 QC 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 2 K concentration of products is too large and that of reactants too small Substances from right to left Q 2 K concentration of products too small and of reactants too large Proceeds from left to right achieves equilibrium by forming more products Le Chatelier s 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 Nzlgl3H2g92NH3gl 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 N2 NH3 to form more Adding 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 N2g3H2g92NH38 Increasing the pressure causes more formation of NH3 But with H2g12ggtHHlltggt increasing pressure does not affect the equilibrium Endothermic reactants heat H 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 Chemistry 130 Chapter 16 Notes Taste is one of the ve senses there are 5 basic tastes umami delicious bitter sour sweet salty The sour response of our taste buds is due to the presence of acids ex citric acid Bases have a bitter taste ex soap 4 Acid substance that when dissolved in water increases the concentration of Hquot ions Base substance that when dissolved in water increases the concentration of 6 0H ions Hydrogen chloride is an Arrhenius acid that is highly soluble in water which quot ZClquot produces J Ions H lt3a61 HClgH20lgtClquot The aqueous solution of HCl is hydrochloric acid About 37 HCl by mass if 2 39 3 concentrated NaOH is an Arrhenius basedissociates into 620 ions or 2 Acid base reactions involve the transfer of Hquot ions from one substance to Zaq another The interaction of a proton with water forms a hydronium ion H304 insert equation by hand 2 n dissolving in water HCI molecule transfers Hquot ion a proton to an H20 molecule Hilaql zH30quot HClgH20l gtClz Acid substance that donates a proton to another substance Base substance that accepts a proton Therefore HCI is a BronstedLowry acid and H20 is a BronstedLowry base lt3a61 Zaq0H NHglaqlH20llleNHZ Base acid Amphiprotic substances are capable of acting as a base when combine with something more acidic and acidic when combined with something more basic laq ZaqH3OZ HXaqH20leXquot amp H20 accepts a base X accepts the proton and is a HX donates an acid 2 base and H30 donates a proton and is an acid Conjugate acidbase pair acid and base that differ only in the presence or absence of a proton Every acid has a conjugate base formed by removing a proton from the acid Every base has a conjugate acid formed by adding a proton to the base The stronger the acid the weaker the conjugate base and vice versa 1 A strong acid completely transfers its protons to water leaving no undisassociated molecules in the solution conjugate base of a strong acid has negligible basicity 2 A weak acid only partially disassociates in an aqueous solution and therefore exists as a mixture of the acid and conjugate base conjugate base of a weak acid is a weak base 3 A substance with negligible acidity contains hydrogen but does not demonstrate any acidic behavior in water ts conjugate base is strong Z reacting completely with water abstracting protons to form 0H ions conjugate base of a substance with a negligible acidity is a strong base 6 Stronger acids react with water to produce Hquot ions and stronger bases react with water to produce 0H ionseveing effect laq ZaqH3OZ HXaqH20leXquot f H20 is a stronger base than Xquot H20 abstracts the proton from HX to 2 zH 0quot 1 produce X and the equilibrium lies to the right If X IS a stronger base the equilibrium lies to the left In every acidbase reaction equilibrium favore the transfer of the proton of the stronger acid to the stronger base to from the weaker acid and weaker base Auto ionization ability of water to donate a proton to another water molecule Because forward and reverse reactions are extremely rapid no water molecule remains ionized for long H Oz 3 lt3aql Zaq0H H20llt gtHquot 4 Ho Where f2 the solution is neutral OHquot 2 6 3 In most solutions Hquot and OHquot concentrations are not equal As concentration of one increases the other decreases The pH of a neutral solution is 7 Z Z OH Hquot 2 Z Z and Z pHZ log pHZ log increases 2 Ho pH decreases as the concentration of Z pHp0H14 25 6 4 EX the concentration of Hquot aq in a solution of pH 6 is 10 times the Hquot aq concentration in a solution of pH 7 The pH of a solution can be measured with a pH meter consisting of a pair of electrodes connected to a meter capable of measuring sma voltages that varies with pH Strong acids and bases are strong electrolytes Strong acids HCI HBr HI HNO3HClO4m0n0pr0tiCHZSO4dipr0tic lt3a61 iaCIN0 compete ionization HN03aqH20l gtH3Oquot Constant Strong bases ionic hydroxides of alkali metals NaOHKOH SMOHL Most acidic substances are weak acids that only partially ionize in an aqueous solution lt3a61 ZaqAquot HAaqH20lltgtH3OL 2 H30 3 ooooiicmcmo to K22 Ka is the acid dissociation constant Many weak acids are organic compounds made of HC and O The magnitude of Ka indicates the tendency of the acid to ionize in water The larger the Ka the stronger the acid concentration ionized 100 Percent Ionization originalconcemmtion The stronger the acid the greater the percent ionization For any acid the concentration of acid that ionizes equals the concentration of 6 Hquot thatforms assuming H20 autoionization is negligible H 2 2 z equilibrium I 2 Percent Ionization HA3 initial C As the concentration of a weak acid increases the equilibrium concentration of H Z 426 increases But percent ionization decreases as the concentration increases 2 Wag Concentration of H5 is not directly proportional to the concentration of the weak acid Ex doubling the concentration of a weak acid does not double the la61 concentration of H4 Polyprotic acid have more than one ionizable H atom EX ZaqKa117102 6aqHSO H2503aqltgtHz 2 ZaqKa264108 lt3laql50 ZaqltgtHZ H250 K612 refers to the equilibrium involving removing of the 2nOI proton The proton is more readily removed if K612 is small Finding pH from Pka using ICE table when can we make assumption that X is negligible As a general rule if X lt 5 of initial value X can be neglected If X gt 5 of initial value use quadratic formula If the acid is large and Ka is small it is more likely that X can be neglected Always check It is always easier to remove the rst proton from a polyprotic acid than to remove the second Also for 3 ionizable protons it is easier to remove the second proton than the third Weak bases react with water lt3a61 Zaq0Hz BaqH22llt gtHBZ pr10gKb Base dissociation constant always refers to the equilibrium in which a base reacts amp h H20 to form the corresponding conjugate acid and OH The greater the wit Kb the stronger the base KbgtiltKaKW prgtlltpKapKW14 Effect of anions and cations 1 If the salt contains an anion that does not react with water and a cation that does not react with water the pH is neutral 2 If the salt contains an anion that reacts with water to produce hydroxide ions and a cation that does not react with water the pH is basic 3 If the salt contains a cation that reacts with water to produce H and an anion does not react with water the pH is acidic 4 If the salts contains an anion and a cation both capable of reacting with water the pH depends on the relative abilities of the ions to react with water Factors affecting acid strength 1 Polarity of HX bond 2 Strength of HX bond 3 Stability of Conjugate base X When X is in the same group in the periodic table The strength of H X bond is the determining factor The strength of HX bond decreases as X increases in size The acidity of binary acids increases from top to bottom down a group When X is in the same period in the periodic table The polarity of H X bond is the determining factor Acidity increases as electronegativity of X increases The acidity of binary acids increases from left to right across a row Generally acidity increases with increasing electronegativity of E and with increasing number of electronegative atoms bonded to the central atom E Lewis acids are electronpair acceptors Atoms with an empty valence orbital can be Lewis acids Lewis bases are de ned as electronpair donors Anything that could be a Bronsted Lowry base is a Lewis base Lewis bases can interact with things other than protons however Chemistry Chapter 17 Notes C H3 COOH C H3 HOONa Consider both solutions contain 2 substances that share a common ion CH3C00quot lt3a61 WCH3COO A disociates completely CH3COONaaqgtNa quot I The weak electrolyte ionizes partially and the addition of CH3C00quot shifts the equilibrium reducing the number of H ions the presence of an added acetate ion causes the acetic acid to ionize less than it normally would Whenever a weak electrolyte and a strong electrolyte containing a common ion are together in a solution the weak electrolyte ionizes less than it would if it were alone in the solutioncommon ion effect onization of a weak base is also decreased by the addition of a common ion lt3a61l Zaq0H NH3 QH201 NH 6 Addition of NH shifts the equilibrium reducing the amount of OH ions Soutions with weak conjugate acidbase pairs resist drastic changes in pH when small amounts of strong acidbase are addedbuffered solutions buffers EX human blood resists changes in pH because of acid to neutralize added OH ions and a base to neutralize H ions Acid and base that make up the buffer must not consume each other through a neutralization reaction The requirements ful lled by a weak acidbase conjugate pair such as CH3C00H CH3C0 0quot or z NH3 Hquot NHfl The buffered solution example pH is determined by the value of a and the ratio of concentrations of conjugate 4 Ho 2 acidbase pairs 2 If OH ions are added to the buffer they react X6 2 446161 ZaqHXaq gtH20aqX OHZ X 2 Causes HX to decrease and 23 to increase 2 lt3aqHX aq If H ions are added WHY H6 4 Causes Xquot to decrease and HX to increase If the ratio of HXX is smathe pH is small 4 X6 2 4 X6 ogHpHlog Ka pKa pH z 2 Z lHXl pKa log H K 1 base HendersonHasselbach Equation p 19 a 0g acid used to calculate the pH of buffers Buffer capacity the amount of acid or base the buffer can neutralize before the pH begins to change to an appreciable degree Depends on the amount of acid and f CH3 COOH base used to prepare the buffer Ex pH of 1 M o and 1 M CH3C00Na is the same as a solution with the 1 M of each The second has greater buffering capacity because there s more of the compounds pH range pH range over which a buffer acts effectively Buffers effectively resist a change in pH in either direction when the concentrations of weak acid and pKa conjugate base are the same When pH the buffer pH is optimal Buffers pKa usually have a useabe range within or 1 pH unit of pKa pH or 1 Reactions between strong acids and weak bases proceed essentially to completion as do those between strong bases and weak acids As long as buffer capacity is not exceeded the strong acidbase is completely consumed To calculate how the pH of the buffer responds to the addition of strong acidbase 1 Consider acidbase neutralization reaction determine the effect of HX and X stoichiometry calculations 2 Use values HX and X with the Ka to calculate H Equilibrium calculation use HendersonHasselbach Strong acid conjugate base of buffer Z gtHX l lt ZH X6 gt i Buffer with HX and Xquot Calculate new HX and Xquot values use a to nd pH ZH20 gt T Z gtXquot T lt HX0Hquot Strong base conjugate acid of buffer Acidbase indicators can nd the equivalence point of titration point at which stoichiometrically equivalent quantities of acid and base have been brought together pH titration curve can monitor the progress of reaction The curve can determine equilibrium point and can also determine the Ka of a weak acid or Kb of a weak base 1 Strong acidstrong base titration 1 Initial pH pH of a solution before the addition of any base is through determination of concentration of the strong acid pH is low 2 Between initial pH and equilibrium point As NaOH base is added the pH increases slowly then rapidly pH levels are not yet neutralized 3 Equivalence point at the equivalence point there is an equal number of moles of base as acid reacted leaving only salt pH is 7 because cations and anions do not affect pH 4 After equivalence point pH is determined by the concentration of excess NaOH in the solution 2 Weak acidstrong base K 1 Initial pH Ex 50 mL of 1 M acetic acid 1M NaOH Use a pH289 ZH20gt 2 Between CH3COOHCH3COO use the neutralization reaction to nd CH 2 M3C00H and CH3 00 Calculate the pH of the buffer Z 6 ZH20gt Solution with weak acid 5 X calculate the new HX and Xuse Ka HX and gtHX0H X add strong base neutralization reaction HX decreases X increases to nd H 3Equivalence point is reached when 50 mL of the base is added to 50 mL of the acid 4 After equivalence point excess base pH is determined by the excess OH and pOH Weak acids have higher initial pH than strong acids pH change for rapid rise portion of the titration graph is smaller for weak acids than strong acids pH for equivalence point is above 7 for weak acid titration When weak acids have more than 1 ionizable H atom the reaction with OH occurs in steps laqlH20ll EX Zaq gtH2P0 H3P03aq0H 2 ZaqH20l Zaq gtHPO Zaq0Hz H2130 When neutralization steps of a polyprotic acid or polybasic base are suf ciently separated the titration has multiple equivalence points Phenophtaein is a good indicator for acidbase titrations Saturated solution one in which the solution is in contact with the undissolved solute 22 Z laCI Bfo 2ZaqSOl 2Z BaSO4Slt gtBa 50 K 2 Sp Solubility product constant indicates how soluble the solid is in water and is denoted Kw In general the solubility product Kw of a compound equals the product of the concentration of the ions involved in the equilibrium each raised to the power of its coef cient in the equilibrium equation Coef cient for each ion in the equilibrium equation also equals its subscript in the compound s chemical formula Molar solubility is the number of moles of solute that dissolve in forming 1 L of saturated solution of the solute molL Soubility of a slightly soluble salt is decreased by the presence of a second solute that furnishes a common ion pH of a solution affects the solubility of any substance whose anion is basic Soubility of a compound containing a basic anion anion of a weak acid increases as the solution becomes more acidic Solubility of a slightly soluble salt containing basic anions increases as H increases or as pH is lowered The more basic the anion the more solubility is in uenced by pH Solubility of salts with negligible basicity such as ClBr l and Z N0 is unaffected by pH changes Characteristic property of metal ions is their ability to act as Lewis acids toward H20 molecules which act as Lewis bases Bases other than water can also interact with metal ions such as transition metal ions Assembly of Lewis bases bonded to a metal ion is a complex ion Stability can be judged by size at equilibrium Constant for formation from a hydrated metal ion Kfz formation constant Solubility of metal salts increases in the presence of Lewis bases such as quot Z0Hquot NH3CNquot Amphoteric oxides and hydroxides soluble in strong acids and bases because they 28 2zSnquot 3Zan 6 Crz Al can behave as an acid or base Ex Z 2 zv0Hquot o react with acids 0 These dissolve because their anions The use of reaction quotient Q can be used to determine the direction in which a reaction must proceed to reach equilibrium lf Q gt Kspa precipitation occurs reducing the ion concentrations until Q KW lf Qlt Ks the solid dissolves increasing ion concentrations until Q Kw lf Q Kspa equilibrium exists saturated solution Ions can be separated from each other based on the solubilities of their salts selective precipitation Chemistry 130 Chapter 5 Notes Thermodynamics the study of energy and its transformations Thermochemistry relationships between chemical reactions and energy changes that involve heat Energy capacity to do work or transfer heat Work energy used to cause an object to move against a force Heat energy used to cause the temperature of an object to increase KE 1 2 Kinetic energy the energy of motion Tim V Potential energy an object has this relative to its position to other objects quotStoredquot energy that arises from the attractionsrepulsions one object experiences with another Electrostatic Potential Energy arises from interactions between charged particles proportional to the electric changes on the two interacting objects Q1 Q2 kQ1Q2 Eel d k is the constant of proportionality As d becomes in nitely large Eel goes to zero When Q1 Q2 have the same sign the particles repel each other and Eel is positive PE decreases as particles move farther apart and vice versa SI unit for energyJoule 1 cal4184J The energy that kinetic energy possesses results from the rising temperature in a reaction The potential energy is from chemical energy System portion of the universe used for study everything else is surroundings Open system is one where matter and energy is exchanged with surroundings Closed systems do not do this An isolated system is one in which neither energy nor matter can be exchanged with its surroundings Force any push or pull exerted on an object wFd Heat is transferred from a hotter object to a colder one Energy as heat is transferred from the hotter system just reacting substances to cooler surroundings everything else First Law of Thermodynamics energy lost by the system is gained by surroundings and vice versa energy is conserved Internal Energy E sum of all the kinetic and potential energies of the components of the system Generally the numerical value is not known Concerned with the change in AEEfinalEinitial Positive AE energy gained by the system and lost to by the surroundings Negative AE energy lost by the system and gained by the surroundings System may exchange energy with the surroundings as heat or work A1Z IJ39VV 1heat szmkgt When heat is added to a system or work is done on a system internal energy increases When work is done to the system w is positive The same with heat When heat is lost and work is done by the system internal energy is lowered Positive I Negative q heat gained heat lost w work done to the system work done by the system 3 net gain by system net loss by the system Endothermic process in which the system absorbs heat Heat ows into the system from the surroundings Ex ice melting Exothermic process in which the system loses heat Heat ows out of the system to the surroundings Ex combustion of gasoline State function property of a system that is determined by specifying the system s condition Value depends only on present state of the system not the path taken EX AEEPVAH Heat and work are not state functions Enthalpy Internal energy plus pressure times volume H E PV also state function Pressurevolume work work involved in expansion or compression of gases When WZ PAV AVZV nalV pressure is constant initial pressure is always positive or zero Because the expanding system does work on the surroundings w is negative volume decreases AHAEPAV AHqpw wqp where pressure is constant so where qp is heat when pressure is constant The change in enthalpy equal the heat qp gained or lost at constant Pressure AHZ Ip When qgtO the reaction is endothermic When AH is positive heat is gained endothermic state function When AH is negative heat is lost exothermic state function A H le OdItCIS Hreactants Enthalpy of reaction enthalpy change that accompanies a reaction aka heat of reaction Aern 1 Enthalpy is an extensive property 2 AH for a reaction is equal in magnitude and opposite in sign for the AH of the reverse reaction 3 Enthalpy change for a reaction depends on the states of the reactants and products if gas less heat is readily available Heat capacity amount of heat required to raise the temperature of a substance by 1 K or 1 deg C Molar heat capacity heat capacity of one or mole of a substance Cm Speci c heat capacity heat capacity of one gram of a substance Cs quantity of heat transferred Grams of a substance Temperature change When a substance absorbs heat the temperature increases qZCsmAT qsoln qrxn qsoln Csmm g of S01 qrxn If A Tgt O V am lt0 the reaction is exothermic Combustion reactions are studied using a bomb calorimeter which withstands high pressures sample ignited with electricity and heat absorbed by water which measures the change in temperature qrxnCcalAT Ccal is the speci c heat of calorimeter Carried out at constant value AE is measured Hess s Law states that if a reaction is carried out in a series of steps AH for the overall reaction equals the sum of the enthalpy changes f or the individual steps Overall enthalpy change is independent of the number of steps and the path by which the reaction is carried out consequence of being a state function AH For a particular set of reactant and products is the same whether the reaction takes place in one step or a series of steps Standard enthalpy change enthalpy change when all reactants and produces are in 0 their standard states AH 6 0 Standard enthalpy of formation AHf change of enthalpy for the reaction that forms one mole of compound from its elements with all substances in standard states 0 If elements standard statescompound 1 mol in standard then AHAHf Standard enthalpy of formation of the most stable form of any element is zero because no formation reaction is need when the element is already in its standard state Chemistry 130 Chapter 19 Notes Spontaneous process one that proceeds on its own without any outside assistance Occurs in one direction only and reverse process is nonspontaneous Ex gas spontaneously expands into a vacuum but the reverse process does not happen In general processes spontaneous in one direction are nonspontaneous in the other Depending on a temperature a process can be spontaneous or not Ex spontaneous ice melting at TgtO deg C Thermodynamics tells of the direction and extent of a reaction but not the speed Reversible process speci c way in which a system changes state by exactly reversing the state the system and surroundings can be restored to the original state no net change Irreversible process one in which the system and surroundings cannot simply be reversed to restore to original state Ex heat ow to colder objects Reversible change produces the maximum amount of work that can be done by a system on its surroundings Isothermal process expansion of ideal gas at a constant temperature Gas expands spontaneously with no external pressure no PV work on surroundings wO Any spontaneous process is irreversible Entropy associated with the extent of randomnessdisorder in a system or the extent to which energy is distributed among the various motions of the molecules of the system is a state function Aszs nalSinitial AS h For an isothermal process T where delta S is for any isothermal process qrev is heat for the reversible process but applies to other processes as well and T is the absolute temperature Melting of a substance at the melting point and vaporization of a substance at the qrev A Hfusion AS fuszon T T qZAH boiling point are isothermal processes fusion For a spontaneous process the total change in entropy sum of entropy change of the system plus the surroundings is greater than zero Generay when a solid is dissolved in a solvent entropy increases The dissolving of a salt in water involves both a disordering process the ions become less con ned and an ordering process some water molecules become more con ned The disordering processes are usually dominant Second law of thermodynamics any reversible process results in an increase in total entropy whereas any reversible process results in no overall change in entropy A S universe A S system A S Surroundings O Reversible AS AS AS universe system surroundings gt IrreverSIble O The entropy of the universe increases in any spontaneous process Microstate Single possible arrangement of the positions and kinetic energies of the gas molecules when the gas is in a speci c thermodynamic state It consists of a particular microscopic arrangement of atoms or molecules of the system that corresponds to its state Number of microstates in a system is W Sk In W Entropy is a measure of how many microstates are associated with a W ASk1n W initial particular microscopic state Entropy increases with the number of microstates in a system Entropy increases in a gas as volume increases Also increases with increasing temperature Molecules exhibit translational motion movement of a molecule from one place to another vibrational motion periodic motion of atoms within a molecule rotational motion rotation of a molecule about an axis the number of microstates for a system increases with an increase in volume temperature or an increase in the number of molecules because any of these changes increases the possible positions and kinetic energies of the molecules making up the system The number of microstates also increases with the complexity of molecules more vibrational motions are available Entropy increases for 1 Gas formation from solids or liquids 2 Liquids or solutions formation from solids 3 The number of gas molecules that increase during a chemical reaction lf thermal energy is decreased by lower temperatures the energy of motion decreases and entropy decreases The third law of Thermodynamics entropy of a pure crystalline substance at an absolute zero is zero SO KO The order is perfect among molecules no thermal motion only one microstate WS k In 10 0 Standard molar entropies S 1 At 298 K it is not zero 2 For gases it is greater than when liquid or solid 3 Generally increases with increasing molar mass 4 Generally increases with an increasing number of atoms in the formula of a substance ASo nSOpr0ducts mSOreactants 8 is for sum qsystem system For isothermal processes ASsurroundings T V T Spontaneous processes that result in a decrease of entropy are always exothermic Gibbs free energy GzH TAS AH AS AS AS ASSystem T system gt universe system surroundings lf AGltO reaction is spontaneous in the forward direction If AGZO reaction is at equilibrium lf AGgtO reaction in the forward direction is nonspontaneous but spontaneous in the reverse reaction AG It is more convenient to use as a criterion for spontaneity instead of AS AG universe because refers to the system alone In any spontaneous process carried out at constant temperature and pressure free energy always decreases 0 Standard free energies of formation A Gf A Gm 8n G0 products 8m G0 reactants gt eacUon 5 pontaneous at all T 6 Nonspontaneous at all T Spontaneous at low T nonspontaneous 2 2 2 z z at high T Spontaneous at high T or nonspontaneous at owT Chemistry Chapter 20 Notes Electrochemistry study of the relationships between electricity and chemical reactions Includes spontaneous and nonspontaneous processes By keeping track of oxidation numbers oxidation states one can determine whether a given reaction is oxidation reduction Elemental forms of molecules have an oxidation number of 0 oxygen always has 2uness it is in peroxide hydrogen is 1 and so forth By determining the change in oxidation state in a chemical reaction oxidizing agents oxidants and reducing agents reductants can be determined Oxidants are the substances that make it possible for other substances to be oxidized Reductants are the substances that make it possible for other substances to be reduced Reducing agent is therefore oxidized in the process and vice versa Whenever balancing a chemical equation one must obey the law of conservation of mass Gains and losses must also be balanced If a substance loses a certain amount of electrons during the process of a reaction another gains them 2 laq 4Zaq2Fez EX 3ZaqgtSnz 2Zaq2Fez Snz Separate the substances and their electron lossgain to make things easier Z 4 326 2Zaq gtSnquot Oxidation Sn 26ltaql Z gt2Fe 3 ltaqgt2 Reduction 2 Fe 2 Halfreactions illustrate either oxidationreduction alone To balance an equation via halfreaction 1 Divide the equation into 1 oxidation and 1 reduction halfreaction 2 Balance each half reaction 3 Multiply halfreactions by integers as needed to make the numbers of elections lost in the oxidation halfreaction equal to the number gained by the reduction 4 Add halfreactions and simplify if possible 5 Check to make sure atoms and charges are balanced If a redox reaction occurs in an acidic solution 1 Assign oxidation numbers to determine what is oxidized and what is reduced 2 Write the oxidation and reduction halfreactions 3 Balance each halfreaction a Balance elements other than H and O b Balance 0 by adding H20 c Balance H by adding H d Balance charge by adding electrons 4 Multiply the halfreactions by integers so that the electrons gained and lost are the same 5 Add the halfreactions subtracting things that appear on both sides 6 Check to make sure that atoms and charges are balanced If a redox reaction occurs in a basic solution the equation must be balanced by using OH and H20 rather than H and H20 One way is to balance the half reaction like in an acidic solution and count the number of H in each half reaction and add the same number of OH to each side for the halfreaction This way it is also mass balanced This way protons are neutralized to form water In spontaneous oxidationreduction redox reactions electrons are transferred and energy is released A voltaic cell is a device in which the transfer of electrons takes place through an external pathway rather than directly between reactants present in the same 2ZaqCuS reaction vessel Ex when zinc is reacted with copper 2 2504Z 5 Zn Cuquot Electrons owing through a wire and ions moving in solution both constitute electrical current The ow of electrical charge can be used to complete electrical work The two solid materials connected by an external circuit are electrodes in a cell The anode negative is where oxidation occurs and the cathode positive is where reduction occurs Each compartment of the voltaic cell is called a hafce One half cell is the site of oxidation the other for reduction Ex 2 22 2 2 anode oxidation halfreaction 616 f ZnSgtZn Z gtCus cathode reduction halfreaction 2 Z 1226 It Electrons become available as zinc is oxidized at the anode ow through external circuit to the cathode where copper is reduced For a voltaic cell to work the solutions in the two halfcells must remain electrically neutral To allow zinc cations to move out of the solution and anions to migrate in there must be a means to connect the two solutions For copper cations to leave the solution and excess anions to remain there must be something connecting the solution as well A salt bridge is composed of a Ushaped tube with an electrolyte solution Allows for anions to always migrate toward the anode and the cations toward the cathode Electrons ow from the anode through the external circuit to the cathode This is why the anode has a negative sign and the cathode has a positive sign Potential energy of electrons is higher in the anode than in the cathode allowing for spontaneous ow toward the electrode with a more positive electron potential The difference between potential energy per electric charge is measured in volts J0 ule 1 V Coulomb Cell potential is the potential difference between two electrodes of a voltaic ce AKA Electromotive force emf pushes electrons through the external circuit AKA voltage Cell potential of any voltaic cell is positive Magnitude depends on the reactions that occur at the cathode and anode concentrations of reactants and products and the temperature standard conditions25 deg C Standard cell E0 1139 E0 potentialemf ce 6 IS the cell potential at standard conditions a standard reduction potentials E36 E2Cath0de E an0de The greater the difference between oxidizing and reducing agents the greater the voltage of the cell The values of standard reduction potentials are referenced to a standard hydrogen electrode SHE Consists of a platinum wire connected to a piece of platinum foil covered with nely divided platinum that serves as an inert surface for the reaction and can operate as either the anode or cathode Whenever we assign an electrical potential to a halfreaction we write the reaction as a reduction Standard reduction potentials are intensive properties The more 0 positive the value of EA the greater the tendency for reduction under standard conditions AG for a redox reaction can be found by using the equation AG nFE n the number of moles of electrons transferred F Faraday constant 1 F 96485 Cmol 96485 JV mol E emf A positive value of E and a negative value of AG both indicate that a reaction is spontaneous Under standard conditions AG nFE 0 A G RTnK 0 nF E RTnK 0 RT E 7 an Finding the emf under nonstandard conditions use Nernst equation RT EE0 ln nF Q Nernst equation implies that cell could be created that has the same substance for anode and cathode as long as concentrations are different Electrolytic cell an electrochemical cell in which electric power is used to cause a nonspontaneous redox reaction to occur Ex 2 NaCll 2 Na Cl2g electrolysis Since AG WW and AG nFE The maximum useful electrical work obtained from a voltaic cell is Wmax nFE When an external potential Em is applied to a cell the amount of work performed is WanEext


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