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Chemical Kinetics. BLB 11 th Chapter 14. Chemical Reactions. Will the reaction occur? Ch. 5, 19 How fast will the reaction occur? Ch. 14 How far will the reaction proceed? Ch. 15. Expectations. Review: concentration units, graphing (line, slope) Work with reaction rates.
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Chemical Kinetics BLB 11th Chapter 14
Chemical Reactions • Will the reaction occur? Ch. 5, 19 • How fast will the reaction occur? Ch. 14 • How far will the reaction proceed? Ch. 15
Expectations • Review: concentration units, graphing (line, slope) • Work with reaction rates. • Determine a rate law for a given reaction from experimental data. • Work with rate laws. • Understand the factors that affect the rate of a chemical reaction. • Calculate the activation energy.
Thermochemistry review… C(s, diamond) C(s, graphite) ΔH °rxn = Is the reaction favorable?
Kinetics … • studies the rates at which chemical reactions occur. • gives information about how the reaction occur, that is, the reaction mechanism • Determining the reaction mechanism is the overall goal of kinetic studies. (sect. 6)
14.1 Factors that Affect Reaction Rates • Physical state of the reactants – states that promote contact have faster rates; homogeneous vs. heterogeneous • Concentration of the reactants: conc. ↑, rate ↑ (or pressure for gases) • Temperature: temp. ↑, rate ↑ due to higher molecular energy and speed (section 5) • Catalysts: rate ↑ by changing the mechanism and reaction energy (section 7) • Other physical things like stirring and grinding solid reactants.
14.2 Reaction Rates • Rate – change in some variable per unit time • Reaction rate – change in concentration per unit time; M/s or mol/L·s • Rates are determined by monitoring concentration as a function of time. • Rates are positive quantities; for reactant A:
C4H9Cl(aq) + H2O(l) → C4H9OH(aq) + HCl(aq) In this reaction, the concentration of butyl chloride, C4H9Cl, was measured at various times.
Reaction Rates cont. • Rates change over time: • reactant rates decrease • product rates increase • Instantaneous rate – rate at a specific time • Average rate – Δ[A] over a specific time interval • Initial rate – instantaneous rate at t = 0 • Note: Rates and rate laws are not based on stoichiometry!! They must be determined experimentally.
C4H9Cl(aq) + H2O(l) → C4H9OH(aq) + HCl(aq) (Fig. 14.4, p. 577)
Reaction Rates and Stoichiometry • The molar ratios between reactants and products correspond to the relative rates of the reaction. • Relative rates – relationship between rates of reactant disappearance and product appearance at a given time. 2 HI(g) → H2(g) + I2(g)
2 NO2(g) → 2 NO(g) + O2(g) Rate = 2.4 x 10-5 M/s Rate = 8.6 x 10-5 M/s Rate = 4.3 x 10-5 M/s
Relative Rates PracticeAt a given time, the rate of C2H4 reaction is 0.23 M/s. What are the rates of the other reaction components?C2H4(g) + 3 O2(g) → 2 CO2(g) + 2 H2O(g)0.23 M/s ? ? ?
Types of Rate Laws • Differential rate law or rate law (section 3) • Shows how the reaction rate changes with concentration • Must be specific in how defined • Integrated rate law (section 4) • Shows how concentration changes with time • Graphical determination of the order
14.3 Concentration & Rate aA + bB → cC + dD • General form of rate law: [A], [B] – conc. in M or P rate = k[A]m[B]nk – rate constant; units vary m, n – reaction orders • Reaction orders and, thus, rate laws must be determined EXPERIMENTALLY!!! • Note: m ≠ a and n ≠ b • Overall order = sum of individual orders • Rate constant is independent of concentration.
Reaction OrdersFor the reaction: A →B, the rate law is: rate = k[A]m
2 NO(g) + 2 H2(g) → N2(g) + 2 H2O(g)rate = k[NO]2[H2] • What is the order with respect to NO? • What is the order with respect to H2? • What is the overall order? • If [NO] is doubled, what is the effect on the reaction rate? • If [H2] is halved, what is the effect on the reaction rate? • What are the units of k?
PtCl2(NH3)2 + H2O → PtCl(H2O)(NH3)2 + Cl¯rate = k[PtCl2(NH3)2] k = 0.0090 h-1 • Calculate the rate of reaction when the concentration of PtCl2(NH3)2 is 0.020M. • What is the rate of Cl¯ production under these conditions?
Practice problem 26 2 NO(g) + 2 H2(g) → N2(g) + 2 H2O(g)rate = k[NO]2[H2] k = 6.0 x 104M-2s-1 @1000K (b) Calculate rate when [NO] = 0.025 M and [H2] = 0.015 M.
Determining Rate Laws Initial Rates Method • Find two experiments in which all but one reactant’s concentration is constant. • Observe the relationship between concentration change and rate change to determine the order for that reactant. • Repeat for other reactant(s).
NH4+(aq) + NO2¯(aq) → N2(g) + 2 H2O(l) Determine the rate law and calculatek.
2 NO(g) + Cl2(g) → 2 NOCl(g) Determine the rate law and calculatek.
BrO3¯(aq) + 5 Br¯(aq) + 6 H+(aq) → Br2(l) + 2 H2O(l) Determine the rate law and calculatek.
14.4 The Change of Concentration with TimeaA→ products • 1st order y = mx + b Plot: ln[A] vs. t slope = −k integrate
14.4 The Change of Concentration with TimeaA → products • 2nd order y = mx + b Plot: 1/[A] vs. t slope = k integrate
14.4 The Change of Concentration with TimeaA → products • zero order y = mx + b Plot: [A] vs. t slope = −k integrate
Integrated rate law analysis (BLB 42)CH3NC(g) → CH3CN(g) - isomerization
Integrated rate law analysis (BLB 46)C12H22O11(aq) + H2O(l) → 2 C6H12O6(aq)
Half-life, t1/2 • t1/2 – time required for the concentration of a reactant to decrease by half of its initial value
1st order half-life • All half-lives same length of time • Independent of initial concentration Note: Radioactive decay follows 1st order kinetics.
2nd order half-life • All half-lives different length of time • Dependent on initial concentration • Zero half-life • All half-lives different length of time • Dependent on initial concentration
14.5 Temperature and Rate • Generally, as temperature increases, so does reaction rate. • This is because k is temperature dependent.
14.5 Temperature and Rate Collision theory • In a chemical reaction, bonds are broken and new bonds are formed. In order for molecules to react, they must collide. • Collisions are either effective or ineffective due to orientation of molecules. • Collisions must have enough energy to overcome the barrier to reaction, the activation energy. • Temperature affects the number of collisions.
Activation Energy, Ea • Energy barrier (hump) that must be overcome for a chemical reaction to proceed • Activated complex or transition state – arrangement of atoms at the top of the barrier
Activation Energy, Ea • Energy difference between the reactant and the highest energy along the reaction pathway • Reaction specific • Rate of reaction is dependent upon the magnitude of Ea; Ea ↓, rate↑ (generally) • Temperature independent
Temperature Effects Maxwell-Boltzman Distribution • At higher temperatures, more molecules will have adequate energy to react. • This increases the reaction rate.
Arrhenius Equation • Svante Arrhenius developed an equation for the mathematical relationship between k and Ea. • A is the frequency factor, which represents the number of effective collisions.
Arrhenius Equation y = m x + b