Chapters 13, 18 and 19 outline
Chapters 13, 18 and 19 outline Chem 0120
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This 25 page Study Guide was uploaded by Anna Perry on Thursday November 19, 2015. The Study Guide belongs to Chem 0120 at University of Pittsburgh taught by Dr. Vines in Winter2015. Since its upload, it has received 115 views. For similar materials see General Chemistry 2 in Chemistry at University of Pittsburgh.
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Chapter 18: Thermodynamics and Equilibrium 18.1 First Law of Thermodynamics: A Review a. Recall the statement of the first law of thermodynamics The first law of thermodynamics is the law of conservation of energy applied to thermodynamic systems. Energy remains constant. b. Know the definition of enthalpy The enthalpy of a system is the system’s internal energy plus pressure times volume; it is a state function H = U + PV c. Recall that the heat of reaction at constant pressure equals the enthalpy change The heat involved in a physical or chemical change for a given fixed pressure and temperature equals the change in enthalpy for the system q = ΔH d. How to calculate the enthalpy of reaction from standard enthalpies of formation ΔH° = nΔH° (froducts) - mΔH° (rfactants) 18.2 Entropy and the Second Law of Thermodynamics a. Define spontaneous process A spontaneous process is a physical or chemical change that occurs by itself. b. Define entropy Entropy, S, is a thermodynamic quantity that is a measure of how dispersed the energy of a system is among the different possible ways that system can contain energy c. Relate entropy to disorder in a molecular system (energy dispersal) When energy of a thermodynamic system is concentrated in a relatively few energy states, the entropy of the system is low. When that energy is dispersed, the entropy of the system is high. Entropy of a system plus its surroundings increases in a spontaneous process d. State the second law of thermodynamics in terms of system plus surroundings and in terms of the system only The second law of thermodynamics states that the total entropy of a system and its surroundings always increases for a spontaneous process e. Calculate the entropy change for a phase transition ∆ S= q T The heat of vaporization, ΔH of CCl is 39.4 KJ/mol vap 4 If 1 mol of liquid CC4 at 25°C has an entropy of 216 J/K, what is the entropy of 1 mol of the vapor in equilibrium with the liquid at this temperature? f. Describe how ΔH – TΔS functions as a criterion of a spontaneous reaction ΔH – TΔS < 0 18.3 Standard Entropies and the Third Law of Thermodynamics a. State the third law of thermodynamics The third law of thermodynamics states that a substance that is perfectly crystalline at 0 K has an entropy of 0 b. Define standard entropy Standard entropy, S°, is the entropy value for the standard state of the species c. State the situations in which the entropy usually increases Entropy increases when 1. A reaction in which a molecule is broken into two or more smaller molecules 2. A reaction in which there is an increase in moles of gas 3. A process in which a solid changes to liquid or gas or a liquid changes to a gas d. Predict the sign of the entropy change of a reaction Is ΔS° positive or negative? 2NH 3g) + CO (2) NH CO2H (aq)2+ H O(l) 2 e. Express the standard change of entropy of a reaction in terms of standard entropies of products and reactants ∆ S°= ∑ nS° (roducts)−∑ mS°(reactants) f. Calculate ΔS° for a reaction Calculate the ΔS° at 25°C // The standard entropy of NH 2ONH (a2) is 174 J(molK) 2NH (g) + CO (g) NH CONH (aq) + H O(l) 3 2 2 2 2 18.4 Free Energy and Spontaneity a. Define free energy, G. Free energy, G, gives a thermodynamic quantity defined by the equation G = H – TS As a reaction proceeds at a given temperature and pressure, reactants form products and the enthalpy H and entropy S change. b. Define the standard free-energy change The free-energy change that occurs when reactants in their standard states are converted to products in their standard states ΔG° = ΔH° - TΔS° c. Calculate ΔG° from ΔH° and ΔS° What is the ΔG° for the reaction at 25°C? N (g) + 3H (g) 2NH (g) 2 2 3 d. Define the standard free energy of formation, ΔG ° f Standard free energy of formation, ΔG °,fis the free-energy change that occurs when 1 mol of substance is formed from its elements in their reference forms at 1 atm and at a specified temperature e. Calculate ΔG° from standard free energies of formation ΔG°= ∑ nΔG° (roducts)−∑ mΔG° r(actants) Calculate ΔG° for C2H 5H(l) + 3O 2g) 2CO (2) + 3H O2g) f. State the rules for using ΔG° as a criteria for spontaneity The following rules are useful in judging spontaneity of a rxn 1. When ΔG° is a large negative number, the reaction is spontaneous as written, and reactants transform almost entirely to products when equilibrium is reached 2. When ΔG° is a large positive number, the reaction is nonspontaneous as written, and reactants do not give significant amounts of products at equilibrium 3. When ΔG° has a small negative or positive value, the reaction gives an equilibrium mixture with significant amounts of both reactants and products g. Interpret the sign of ΔG° A positive ΔG° means the reactants tend to go mostly to products A negative ΔG° means the reaction is nonspontaneous as written When ΔG° is close to zero, the reaction gives an equilibrium mixture 18.5 Interpretation of Free Energy a. Relate the free-energy change to maximum useful work The free-energy change is the maximum energy available, or free, to do useful work w max= ΔG b. Describe how the free energy changes during a chemical reaction Spontaneous: The free energy decreases as the reaction proceeds. At equilibrium, the free energy becomes a minimum, and the equilibrium mixture is mostly products Nonspontaneous: The free energy decreases until equilibrium, at which the minimum value is reached. There is very little reaction, however, because the equilibrium mixture is mostly reactants. To change reactants to mostly products is to undergo a nonspontaneous reaction 18.6 Relating ΔG° to The Equilibrium Constant a. Define the thermodynamic equilibrium constant, K The thermodynamic equilibrium constant, K, is the equilibrium constant in which the concentrations of gases are expressed in partial pressures in atmospheres, whereas the concentrations of solutes in liquid solutions are expressed in molarities. b. Write the expression for a thermodynamic equilibrium constant 2NH (3) + CO (g2 NH CO2H (aq)2+ H O(l) 2 c. Indicate how the free-energy change of a reaction and the reaction quotient are related ΔG = ΔG° + RTlnQ d. Relate the standard free-energy change to the thermodynamic equilibrium constant ΔG° + -RTlnK e. Calculate K from the standard free-energy change (molecular equation) Find the value of the equilibrium constant K at 25°C (298K) for the reaction 2NH 3g) + CO (2) NH CO2H (aq)2+ H O(l)2 ΔG° at 25°C equals -13.6kJ. f. Calculate K from the standard free-energy change (net ionic equation) Calculate Kspt 25°C for the reaction AgCl(s) Ag (aq) + Cl (aq) using standard free energies of formation 18.7 Change of Free Energy With Temperature a. Describe how ΔG° at a given temperature is approximately related to ΔH° and ΔS° at 25°C ΔG° T ΔH° - TΔS° b. Describe how the spontaneity or nonspontaneity of a reaction is related to each of the four possible combinations of signs of ΔH° and ΔS° ΔG° ΔH° ΔS° Description - - + Spontaneous + + - Nonspontaneous +/- - - Spontaneous at low T; Nonspontaneous at high T +/- + + Nonspontaneous at low T; Spontaneous at high T c. Calculate ΔG° and K at various temperatures (a) What is ΔG° at 1000°C for CaCO (3) CaO(s) + CO 2g) Is this reaction spontaneous at 1000°C and 1 atm? (b) What is the value of K pt 1000°C for this reaction? What is the partial pressure of CO 2 Chapter 19: Electrochemistry 19.1 Balancing Oxidation-Reduction Reactions in Acidic/Basic Solutions a. Learn the steps for balancing oxidation-reduction reactions in acidic solution using the half-reaction method 1. Assign oxidation #s to each atom, decide what’s being oxidized/reduced 2. Split the skeleton equation into two half-reactions 3. Balance each half-reaction a. Balance all atoms except O and H b. Balance O atoms by adding H Os to one side + c. Balance H atoms by adding H ions to one side d. Balance electric charge by adding electrons to the more positive side 4. Combine the two half-reactions to obtain the final balanced equation - a. Multiply each half-reaction by a factor such that the e s cancel b. Simplify the balanced equation by canceling species on both sides b. Balance equations by the half-reaction method (acidic solution) Zinc metal reacts with HNO . 3inc reduces nitrate to ammonium; zinc is oxidized to zinc ion. Write the net ionic equation: c. Learn the additional steps for balancing oxidation- reduction reactions in basic solution using the half- reaction method + 5. Note the number of H ions in the equation. Add this number of OH ions to both sides of the equation 6. Simplify by noting that H reacts with OH to give H O. Cancel 2 out waters occurring on both sides and reduce the equation to simplest terms d. Balance equations by the half-reaction method (basic solution) MnO (4q) + SO (aq3 MnO (s) + S2 (aq) 42- 19.2 Construction of Voltaic Cells a. Define electrochemical cell, voltaic (galvanic) cell, half cell, and electrolytic cell An electrochemical cell is a system consisting of electrodes that dip into an electrolyte and in which a chemical reaction either uses or generates an electric current A voltaic (galvanic) cell is an electrochemical cell in which a spontaneous reaction generates an electric current A half cell is the portion of an electrochemical cell in which a half-reaction takes place An electrolytic cell is an electrochemical cell in which an external energy source drives an otherwise nonspontaneous reaction b. Describe the function of the salt bridge in a voltaic cell The salt bridge is a tube of an electrolyte in a gel that is connected to the two half-cells of a voltaic cell; it allows the flow of ions but prevents the mixing of the different solutions that would allow direct reaction of the cell reactants c. State the reactions that occur at the anode and the cathode in an electrochemical cell Oxidation occurs at the anode Reduction occurs at the cathode d. Define cell reaction The net reaction that occurs in the voltaic cell is the cell reaction e. Sketch and label a voltaic cell 19.3 Notation for Voltaic Cells a. Write the cell r2+ction f3+m the ce2+ notation Zn(s)|Zn (aq)||Fe (aq), Fe (aq)|Pt 19.4 Cell Potential a. Define cell potential and volt The maximum potential difference between the electrodes of a voltaic cell is referred to as the cell potential or electromotive force (emf) of the cell The volt, V, is the SI unit of potential difference b. Define potential difference Potential difference is the difference in electric potential between two points c. Calculate the quantity of work from a give amount of cell reactant W max= -nFE cell The potential of a particular voltaic cell with the cell reaction Hg (aq) + H (g2 2Hg(l) + 2H (aq) Is 0.650 V. Calculate the maximum electrical work of this cell when 0.500 g H i2 consumed d. Know Faraday’s constant The Faraday constant, F, is the magnitude of charge on one mole of electrons F = 96,485 C/mole e _ w = -F x potential differences 19.5 Standard Cell Potentials and Standard Electrode Potentials a. Explain how the electrode potential of a cell is an intensive property The electrode potential is an intensive property, which means that its value is independent of the amount of species in the reaction b. Define standard cell potential and standard electrode potential Standard cell potential, E° cell the emf of a voltaic cell operating under standard-state conditions (1 M, 1 atm, 25°C). The standard electrode potential, E°, is the electrode potential when the concentrations of solutes are 1 M, the gas pressures are 1 atm, and the temperature is 25°C c. Interpret the table of standard reduction potentials Reduction half-reactions with more positive electrode potentials have the greater tendency to oxidize. T ~The strongest oxidizing agents are the ones with the most positive E° values ~The strongest reducing agents are the ones with the most negative E° values d. Determine the relative strengths of oxidizing and reducing agents Order the following oxidizing agents by increasing strength: Cl2(g), H2O2(aq), Fe (aq) e. Determine the direction of spontaneity from electrode potentials The stronger oxidizing agent will be on the reactant side when the equation is written as a spontaneous reaction. The stronger reducing agent will be on the reactant side of the spontaneous reaction Consider the reaction: 2+ 2+ 3+ Zn (aq) + 2Fe (aq) Zn(s) + 2Fe (aq) Does the reaction go spontaneously in the direction indicated, under standard conditions? f. Calculate cell potential from standard potentials E°cell E° cathode E°anode Calculate the standard cell potential of the voltaic cell at 25°C 3+ 2+ Al(s)|Al (aq)||Fe (aq)|Fe(s) 19.6 Equilibrium Constants from Cell Potentials a. Calculate the free-energy change from electrode potentials ΔG = w max ΔG° = -nFE° cell Using standard electrode potentials, calculate the standard free-energy change at 25°C for Zn(s) + 2Ag (aq) Zn (aq) + 2Ag(s) b. Calculate the cell potential from free-energy change Calculate the standard potential for this cell at 25°C from free-energies of formation Zn(s) + Cl 2g) Zn (aq) + 2Cl (aq) c. Know the relationships among K, ΔG°, and E° cell nFE° cell RTlnK E° cell (0.0592/n)logK d. Calculate the equilibrium constant from the cell potential The standard cell potential for the following voltaic cell is 1.10 V: Zn(s)|Zn (aq)||Cu (aq)|Cu(s) Calculate th2+equilibriu2+constant K for che reaction Zn(s) + Cu (aq) Zn (aq) + Cu(s) 19.7 Dependence of Cell Potential on Concentration a. Calculate the cell potential for nonstandard conditions What is the cell potential of this cell at 25°C? Zn(s)|Zn (1 x 10 M)||Cu 2+(0.100 M)|Cu(s) When the standard potential is 1.10 V b. Know the Nernst equation E cell E° cell (O.0592/n) logQ c. Describe how pH can be determined using a glass electrode Use test solution as the electrolyte for a hydrogen electrode and bubble in hydrogen gas at 1 atm According to the Nernst equation, the cell potential depends on the hydrogen-ion concentration and can be used to find Q. 19.8 Some Commercial Voltaic Cells I don’t think this matters idk tho 19.9 Electrolysis of Molten Salts a. Define electrolysis The process of producing a chemical change in an electrolytic cell is electrolysis 19.10 Aqueous Electrolysis a. Learn the half-reactions for water undergoing oxidation and reduction You can reduce water to H or ox2dize it to O 2 2H O2l) + 2e H (g)2+ 2OH (aq) - (reduction) + - 2H O2l) O (g2 + 4H (aq) + 4e (oxidation) b. Learn the half-reactions for electrolysis of sulfuric acid solutions + Sulfuric acid is a strong acid and ionizes completely into H and HSO . T4e HSO ion is4relatively strong and for the most part ionizes into H and SO . Therefore the species + - 4 you have to look at are H , SO , and4H O. 2 - At the cathode: + _ o 2H (aq) + 2e H (g) 2 o 2H O(2) + 2e H (g) +22OH (aq) - - At the anode: - 2- - o 2SO (aq4 S O (a2) 8 2e o 2H O(2) O (g)2+ 4H (aq) + 4e - The species whose oxidation half-reaction has the larger (less negative) oxidation potential is the more easily oxidized. c. Predict the half-reactions in an aqueous electrolysis What are the half-reactions in the electrolysis of aqueous copper(II) sulfate? 19.11 Stoichiometry of Electrolysis a. Understand units One faraday (9.6485 x 10 C) is equivalent to the charge on one mole of electrons Electric charge = electric current x time lapse So, Coulombs = amperes x seconds - The ampere (A) is the base unit of current. b. Calculate the amount of charge from the amount of product in an electrolysis When an aqueous solution of copper(II) sulfate, CuSO , is 4 electrolyzed, copper metal is deposited. Cu (aq) + 2e Cu(s) If a constant current was passed for 5.00 hours and 404 mg of Cu(s) was deposited, what was the current? c. Calculate the amount of product from the amount of charge in an electrolysis. When an aqueous solution of potassium iodide is electrolyzed using platinum electrodes, the half-reactions are 2I (aq) I2(aq) + 2e - 2H 2(l) + 2e H (g2 + 2OH (aq) - How many grams of iodine are produced when a current of 8.52 x 10 A flows through the cell for 10.0 min? Chapter 13: Rates of Reaction 13.1 Definition of a Reaction Rate a. Define reaction rate The reaction rate is the increase/decrease in molar concentration of product of a reaction per unit time b. Explain instantaneous rate and average rate of a reaction You obtain the instantaneous rate at a given time from the slope of the tangent at the point on the curve corresponding to that time You obtain the average rate is given by the change in concentration over change in time c. Explain how the different ways of expressing reaction rates are related For 2N 2 (5) 4NO (g2 + O (g)2 - Rate of formation of oxygen = Δ[O 2 Δt - Rate of decomposition of N O = −Δ N[O2 5] 2 5 Δt - Two moles of N O 2ec5mpose for each mole of oxygen Δ[O 2 −1 Δ[N O2]5 formed. Therefore, = Δt 2 Δt How is the rate of formation of NO F r2lated to the rate of reaction of Fluorine in 2NO 2g) + F (2) 2NO F(g2? d. Calculate average reaction rate Calculate the average rate of decomposition of N O , 2 5 −Δ N[O ] 2 5 , by the reaction 2N O 2g)5 4NO (g) 2 O (g) 2 Δt During the time interval from t=600s to t=1200s using the following data: TIME [N 2 ]5 600s 1.24 x 10 M-2 -2 1200s 0.93 x 10 M 13.2 Experimental Determination of Rate a. Describe how reaction rates may be experimentally determined To obtain the rate of a reaction, you must determine the concentration of a reactant or product during the course of the reaction 13.3 Dependence of Rate on Concentration a. Define and provide examples of a rate law and rate constant. A rate law is an equation that relates the rate of a reaction to the concentrations of reactants (and catalyst) raised to various powers - Rate = k[NO ][2 ] 2 The rate constant is a proportionality constant in the relationship between rate and concentrations rate - k= NO F [ 2][ 2 b. Define reaction order and overall order of a reaction. The reaction order equals the exponent of the concentration of that species in the rate law, as determined experimentally The overall order of a reaction equals the sum of the orders of the reactant species in the rate law. c. Determine the order of reaction from the rate law For 5Br (aq) + BrO (a3) + 6H (aq) 3Br (aq) +23H O(l) 2 The experimen-ally d-ter+ 2ed rate law is Rate = k[Br ][BrO ]3H ] What is the order of reaction with respect to each reactant species? What is the overall order of the reaction? d. Explain how exponents are determined experimentally Exponents are determined by doing a series of experiments in which the initial concentrations of reactants are varied. Then the initial rates are compared, from which the reaction orders can be deduced. Effect on Rate of Doubling the Initial Concentration of Reactant: m Rate is Multiplied by: -1 ½ 0 1 1 2 2 4 e. Determine the rate law from initial rates - + - For H 2 2aq) + 3I (aq) +2H (aq) I 3aq) + 2H O2l) - + Obtain the reaction orders with respect to H O2,2I , and H and k. Initial Conce-trations (mo+/L) Initial Rate H2O 2 I H [mol/(Ls)] Exp 1 0.010 0.010 0.00050 1.15 × 10 -6 -6 Exp 2 0.020 0.010 0.00050 2.30 × 10 Exp 3 0.010 0.020 0.00050 2.30 × 10 -6 Exp 4 0.010 0.010 0.00100 1.15 × 10 -6 13.4 Change of Concentration with Time a. Learn the integrated rate laws for first-order, second- order and zero-order reactions The integrated rate law is a mathematical relationship between concentration and time b. Learn the half-life equations for first-order, second-order and zero-order reactions [A]t First-order integrated rate law: ln =−kt [A] 0 1 1 Second-order integrated rate law: =kt+ [A]t [A]0 [A] =−kt+[A] Zero-order integrated rate law: t 0 c. Use an integrated rate law The decomposition of N O to 2O 5nd O is 2irst or2er with a rate constant of 4.80 × 10 /s at 45°C. (a)If the initial concentration is 1.65 × 10 mol/L, what is [N 2 5 after 825s? (b)How long would it take for the concentration of N O to 2 5 decrease to 1.00 × 10 mol/L from its initial value? d. Define half-life of a reaction The half-life, 1/2 of a reaction is the time it takes for the reactant concentration to decrease to one-half of its initial value. e. Relate the half-life of a reaction to the rate constant 0.693 First-order: t1/2 k 1 Second-order: t1/2 k[A]0 [A]0 Zero-order: t1/2 2k f. Plot kinetic data to determine the order of a reaction The order of the reaction is determined by which graph gives the best straight line fit to the experimental data. First-order: ln[A] vs t Second-order: 1/[A] vs t Zero-order: [A] vs t 13.5 Temperature and Rate; Collision and Transition-State Theories a. State the postulates of collision theory The collision theory of reaction rates is a theory that assumes that, for a reaction to occur, reactant molecules must collide with an energy greater than some minimum value and with the proper orientation. b. Explain activation energy (E ) a The activation energy of a reaction is the minimum energy of collision required for two molecules to react. c. Describe how temperature, activation energy, and molecular orientation influence reaction rates k = Zfp where Z = the collision frequency f = the fraction of collisions having energy greater than E a p = the fraction of collisions that occur with proper orientation Z depends on temperature. - As the temperature rises, gas molecules move faster and therefore collide more frequently. f is related to the activation energy - -Ea/RT f = e - f decreases with increasing value of E a - Because k depends on f, this means that reactions with large E aave a small k, and reactions with small E have a a large k. If the molecular orientation is wrong, the molecules will not bond and a reaction will not occur. d. State the transition-state theory The transition-state theory explains the reaction resulting from the collision of two molecules in terms of an activated complex. e. Define activated complex An activated complex (transition state) is an unstable grouping of atoms that can break up to form products. Draw: f. Describe and interpret potential-energy curves for endothermic and exothermic reactions Draw: 13.6 Arrhenius Equation a. Use the Arrhenius equation The Arrhenius equation expresses the dependence of the rate constant on temperature. –Ea/RT k = Ae If you plot lnk against 1/T, you should get a straight line with a slope of –E aR, from which you can obtain the E . a - The intercept is lnA. k2 Ea 1 1 Two point: ln = ( − ) k1 R T 1 T2 For H 2g) + I 2g) 2HI(g), k = 2.7 × 10 L/(mols) at 600K and -3 3.5 × 10 L/(mols) at 650K (a)Find the activation energy (b)Calculate the rate constant at 700K. b. Define frequency factor The frequency factor, which is assumed to be a constant, is the symbol A in the Arrhenius equation. 13.7 Elementary Reactions a. Define elementary reaction, reaction mechanism, and reaction intermediate. An elementary reaction is a single molecular event, such as a collision of molecules, resulting in a reaction A reaction mechanism is the set of elementary reactions whose overall effect is given by the net chemical equation. A reaction intermediate is a species produced during a reaction that does not appear in the net equation because it reacts in a subsequent step in the mechanism. b. Write the overall chemical equation from a mechanism Cl2 2Cl Cl + CHCl 3HCl + CCl 3 Cl + CCl CCl 3 4 c. Define molecularity The molecularity is the number of molecules on the reactant side of an elementary reaction. d. Give examples of unimolecular, bimolecular, and termolecular reactions A unimolecular reaction is an elementary reaction that involves one reactant molecule - Ex: O 3 O +2O A bimolecular reaction is an elementary reaction that involves two reactant molecules. - Most common - Ex: NO 2 NO N2 + NO 3 A termolecular reaction is an elementary reaction that involves two reactant molecules. - Ex: Br + Br + Ar Br +2Ar e. Determine the molecularity of an elementary reaction What is the molecularity of each step in the mechanism Cl2 2Cl Cl + CHCl 3HCl + CCl 3 Cl + CCl 3 CCl 4 f. Write the rate equation for an elementary reaction For an elementary reaction, the rate is proportional to the product of the concentration of each reactant molecule. A B + C - Rate = k[A] A + B C + D - Rate = k[A][B] A + B + C C + D - Rate = k[A][B][C] Write a rate equation for O 3 + NO O + 2O 2 13.8 The Rate Law and the Mechanism a. Explain the rate-determining step of a mechanism The rate-determining step is the slowest step in the reaction mechanism b. Determine the rate law from a mechanism with an initial slow step For O (g) + 2NO (g) O (g) + N O (g) 3 2 2 2 5 The proposed mechanism is O3+ NO 2O + O 3 2 (slow) NO 3 NO N2O 2 5 (fast) What is the rate law predicted by this mechanism? c. Determine the rate law from a mechanism with an initial fast, equilibrium step For 2NO(g) + 2H (g2 N (g)2+ 2H O(g)2 The proposed mechanism is 2NO N O2 2 (fast, equilibrium) N2O 2 H 2 O +2H O 2 (slow) N2O + H 2N + 2 O 2 (fast) What is the rate law predicted by this mechanism? 13.9 Catalysis a. Describe how a catalyst influences the rate of reaction Catalysis is the increase in rate of a reaction that results from the addition of a catalyst A catalyst speeds up a reaction without being consumed by it. They allow a reaction to occur with a reasonable rate at a much lower temperature (lower energy cost). b. Indicate how a catalyst changes the potential-energy curve of a reaction The catalyst must participate in at least one step of a reaction and be regenerated in a later step. The rate is increased either by increasing the frequency factor A or by decreasing the activation energy E .a c. Define homogeneous catalysis and heterogeneous catalysis. Homogeneous catalysis is the use of a catalyst in the same phase as the reacting species. Heterogeneous catalysis is the use of a catalyst that exists in a different phase from the reacting species, usually a solid catalyst in contact with a gaseous or liquid solution of reactants - Thought to occur by chemical absorption of the reactants onto the surface of the catalyst d. Explain enzyme catalysis The catalysts of biological organisms are protein molecules called enzymes. Each enzyme is highly specific, and acts only on a specific substance, catalyzing it to undergo a particular reaction.
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