FSU CHM 1046 - CH 13.2 Rate of a Chemical Reaction
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CH 13.2 Rate of a Chemical ReactionAverage rate of reaction aA+bB →cC +dDRate¿−(1a)Δ[A]Δt=−(1b)Δ[B]Δt=+(1c)Δ[C]Δt=+(1d)Δ[D]Δt The Rate LawA productsRate = k[A]n k=rate constant n = reaction orderDetermining the Order of a Reaction The order of a reaction can be determined only by experiment. A common way to determine reaction order is by the method of initial rates. Example Aproducts[A] (M) Initial Rate (M/s).1 .015.2 .030.4 .060When the [A] doubles, the rate doubles. This means the initial rate is directly proportional to the initial concentration, therefore it is first order. Second Order[A] (M) Initial Rate (M/s).1 .015.2 .060.4 .240Rate quadruples when doubling concentration Zero Order[A] (M) Initial Rate (M/s).1 .015.2 .015.4 .015Rate is unaffected by concentration changeIf unsure or numbers not obvious rate2rate1=k [ A]n2k [ A ]n1 Reaction Order for Multiple ReactantsaA + bB  cC + dDRate = k[A]m[B]n m= reaction order with respect to A n=reaction order for Bm & n are experimentally determined.Overall Order = m + n13.4: Integrated Rate Law: Dependence of Concentration on TimeRelationship between the concentrations of the reactants and time.First order integrated rate lawA  Productsln[A]t=−kt +ln[A]0y = mx + borln[ A ]t[ A ]0=−kt[A]t is concentration at anytime, [A]0 is initial concentration Second order integrated rate law1[ A ]t=kt+1[ A]0y = mx + bZero order integrated rate law[ A ]t=−kt+[ A ]0y = mx + bHalf Life of a Reaction (t1/2)First Order Half-Lifet12=.693kt1/2 is independent of the initial concentration Second Order Half-life t1 /2=1k [ A ]0Zero Order Half-lifet1 /2=[ A ]02 k13.6 Reaction Mechanisms *******************reaction mechanism is a series of individual chemical steps by which an overall chemical reaction occursH2(g) + 2ICl(g)  2HCl(g) + I2(g)• Step 1: H2(g) +ICl(g)  HI(g) + HCl(g)• Step 2: HI(g) + ICl(g)  HCl(g) + I2(g)Each step is called an Elementary Step, cannot be broken down into smaller stepsReaction intermediates – things that are formed in one step, and consumed in another. 14.3 The Equilibrium ConstantaA + bB  cC + dDK=[C ]c[D ]d[ A]a[B]b• large K = forward reaction favored. (products) • Small K = reaction favors reverse reaction (reactants)• K~0 = neither reaction is favoredKoverall = K1 * K2 14.4: Equilibrium Constant in Terms of PressureaA + bB  cC + dDKp=Kc(RT )ΔnΔn = c + d – (a + b) 14.5: Hetergeneous Equilibria: solids and liquidssolids and liquids are left out of Kc expressions14.6: Calculating Equilibrium Constant from Measured Equilibrium ConcentrationsICE tables ****************DO EXAMPLES********14.7: The Reaction Quotient: Predicting the Direction of Changeis like Kc but at any point in the reaction, not just at equilibrium.aA + bB  cC + dDQc=[C]c[D]d[ A ]a[ B]b• Only reactants, Qc = 0• Only products, Qc = infiniteQ<K – reaction towards productsQ=K – reaction at equilibriumQ>K – reaction towards reactions 14.8: Finding Equilibrium ConcentrationsFinding equilibrium concentrations when given K and all but one concentrationExample 14.8 on paperFinding equilibrium concentrations when given K and initial concentrations or pressuresExamples 14.9 & 14.10Finding equilibrium partial pressures when you are given the equilibrium constant and initial partial pressuresExample 14.11More practice Examples 14.12 & 14.1314.9: Le Châtelier’s Principle: How a system at equilibrium responds to disturbancesChanging Concentrations• Increasing concentration of reactants – Q<K favors products• Increasing concentration of products – Q>K favors reactants• Decreasing concentration of reactants – Q>K favors reactants• Decreasing concentration of products – Q<K favors productsChanges in concentrations DO NOT matter in solids and liquidsChanging Pressure• Increase in pressure/decrease in volume – shifts to side with fewer moles of gas• Decrease in pressure/increase in volume – shifts to side with more moles of gas When equal moles are present on each side, no shift takes place.Also when adding inert gas at a fixed volume, no shift takes placeEffect of Temperature Change on EquilibriumExothermic – A+B  C + D + Heat• Add Heat – shifts to the left, K decreases• Remove Heat – shifts to the right, K increasesEndothermic – A + B + Heat  C + D• Add Heat – shifts right, K increases• Remove Heat – Shifts left, K decreases15.3 Definitions of Acids and BasesArrhenius Definition• Acid: a substance that produces H+ ions in aqueous solutions• Base: a substance that produces OH- ions in aqueous solutionsHCl(aq)  H+(aq) + Cl-(aq)HCl is an acid under Arrhenius because it produces H+ ion in solutionNaOH(aq)  Na+(aq) + OH-(aq)NaOH is a base under Arrhenius because it produces OH- ion in solutionBrønsted-Lowry Definition• Acid: proton(H+ ion) donor• Base: proton (H+ ion) acceptor• Amphoteric: substances that can act as acids and bases HCl(aq) + H2O(l)  H3O+(aq) + Cl-(aq)• HCl is and acid under Brønsted because it donates a proton to H2O in solutionNH3(aq) + H2O(l)   NH4+(aq) + OH-(aq)• NH3 is a base under Brønsted because is accepts a proton from H2O in solution15.4 Acid Strength and the Acid Ionization Constant Ka• strong acid completely ionizes in solution o HCl + H2O  H3O+ + Cl- • weak acid only partially ionizes.o Cl + H2O  H3O+ + Cl- (double arrows)Acids can be classified according to their number of protons:1. Monoprotic Acid: HA2. Diprotic Acid: 3. Triprotic Acid: H3AAcid strength depends on attraction between H+ and the anion of the acid.The stronger the acid, the weaker the conjugate base, visa versa. Acid Ionization Constant: Ka• The equilibrium constant for the ionization reaction of a weak acid.HA(aq) + H2O(l)  H3O+(aq) + A-(aq)• Ka = [H3O+][A-] / [HA]Smaller Ka = weaker acidLarger Ka = stronger acid15.5 Autoionization of water and pH:Water is amphoteric, it can act as either acid or base. Water acts as an acid and a base with itself in pure water, this is called autoionization.• H2O(l) + H2O(l)  H3O(aq) + OH-(aq)The equilibrium constant is called the ion product constant: Kw. • Kw = [H3O+][OH-]• Kw = 1*10-14 at 25°C in pure water. An acidic solution contains an acid the creates additional H3O+ ions, causing [H3O+] to increase. However Kw will always be the same,

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# FSU CHM 1046 - CH 13.2 Rate of a Chemical Reaction

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