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CH 104: CHEMICAL KINETICS. Chemical kinetics is the study of the rates of reactions. The rate of a reaction is the change in concentration per unit of time. Some reactions are very fast. For example, H 3 O + (aq) + OH – (aq) → 2H 2 O (l) is completed in about 0.0000001 seconds.

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slide1

CH 104: CHEMICAL KINETICS

  • Chemical kinetics is the study of the rates of reactions. The rate of a reaction is the change in concentration per unit of time.
  • Some reactions are very fast. For example, H3O+(aq) + OH–(aq) → 2H2O(l) is completed in about 0.0000001 seconds.
  • Some reactions are very slow. For example, 2H2(g) + O2(g) → 2H2O(l) is completed in about 1,000,000,000 years.
  • Which is these reactions is faster, (a) Na(s) and Br2(l), or (b) the rusting of Fe(s)?
  • (a) Na(s) and Br2(l)
slide2

CHEMICAL KINETICS

  • Factors affecting the rates of chemical reactions:
    • Nature of reactants
    • Presence or absence of catalysts
    • Solvent
    • Concentration of reactants
    • Temperature
  • In Part A of today’s experiment you will measure the affect of the concentration of reactants on rate.
  • In Part B of today’s experiment you will measure the affect of temperature on rate.
slide3

CHEMICAL KINETICS

  • Given the following general reaction:
  • aA + bB + cC + … → dD + eE + fF + …
  • The rate equation equals:
  • This rate has been arbitrarily defined as the disappearance of A
  • (–Δ[A]/Δt). However, it could have been defined as the disappearance of any reactant, or the appearance of any product.
  • m need not equal a, n need not equal b, etc.
  • m is “the order in A”, n is “the order in B”, etc.
  • m + n + p + … is “the overall order”
  • m, n, p, etc. usually equals 0, 1, or 2; however, they may also equal 1/2, 3/2, etc.
  • k is the specific rate constant. It is a constant for any given reaction in a specific solvent and at a specific temperature.
  • What does k equal when all the concentrations are 1 M?
  • Rate = k[1]m[1]n[1]p
  • Rate = k
slide4

CHEMICAL KINETICS AND CONCENTRATION

  • S2O82–(aq) + 3I–(aq) → 2SO42–(aq) + I3–(aq)
  • The method of initial rates is used to measure the orders of a reaction. For example, the order in S2O82–(aq) is measured as follows.
  • Step #1: To find the order in S2O82–(aq), select the experiments with different initial concentrations of S2O82–(aq) and equal concentrations of I–(aq). What are these experiments?
  • Experiments 1 and 2. In Part A of today’s experiment you must assign the initial concentrations 3 different of reactants (CH3COCH3, I2, and H+). How will you do this so that you can measure the order of each reactant?
  • Step #2: Use the ratio of these rate equations to solve for the order in S2O82–(aq).
  • 2 = 2m
  • m = 1
  • Therefore, the order in S2O82–(aq) is 1.
slide5

CHEMICAL KINETICS AND CONCENTRATION

  • S2O82–(aq) + 3I–(aq) → 2SO42–(aq) + I3–(aq)
  • What is the order in I–(aq)?
  • Step #1: To find the order in I–(aq), select the experiments with different initial concentrations of I–(aq) and equal concentrations of S2O82–(aq). What are these experiments?
  • Experiments 2 and 3.
  • Step #2: Use the ratio of these rate equations to solve for the order in I–(aq).
  • 2 = 2n
  • n = 1
  • Therefore, the order in I–(aq) is also 1.
slide6

CHEMICAL KINETICS AND CONCENTRATION

  • S2O82–(aq) + 3I–(aq) → 2SO42–(aq) + I3–(aq)
  • What is the overall order?
  • The Order in S2O82–(aq) + The Order in I–(aq) = 1 + 1 = 2
  • Therefore, the overall order is 2.
slide7

CHEMICAL KINETICS AND CONCENTRATION

  • S2O82–(aq) + 3I–(aq) → 2SO42–(aq) + I3–(aq)
  • What is the rate constant (k) for this reaction?
  • Rate = k[S2O82–]1[I–]1
  • 1.4 x 10–5 = k[0.038][0.060]
slide8

CHEMICAL KINETICS AND CONCENTRATION

  • S2O82–(aq) + 3I–(aq) → 2SO42–(aq) + I3–(aq)
  • What is the rate of this reaction when [S2O82–] = 0.050 M and [I–] = 0.025 M?
  • Rate = (6.1 x 10–3 L mol–1 s–1)[S2O82–]1[I–]1
  • Rate = (6.1 x 10–3 L mol–1 s–1)[0.050][0.025]
  • Rate = 7.7 x 10–6 mol L–1 s–1
slide9

CHEMICAL KINETICS AND TEMPERATURE

  • In Part B of today’s experiment you will measure the affect of temperature on rate.
  • Experience tells us that the rates of reactions increase with temperature.
  • For example, fuels such as gasoline, oil, and coal are relatively inert at room temperature; however, they rapidly burn at elevated temperatures.
  • In addition, many foods last almost indefinitely in a freezer; however, they spoil quickly at room temperature.
slide10

CHEMICAL KINETICS AND TEMPERATURE

  • The activation energy (Ea) is the minimum energy that is needed for molecules to react.
  • In other words, Ea is the height of the energy barrier between reactants and products.
slide11

CHEMICAL KINETICS AND TEMPERATURE

  • Svante Arrhenius noted that the temperature dependence of the specific rate constant is mathematically similar to the Boltzmann distribution of energies.
slide12

CHEMICAL KINETICS AND TEMPERATURE

  • The Arrhenius equation describes the relationship between the specific rate constant (k), the activation energy (Ea), and the absolute temperature (T). A graph of ln k versus 1/T is called an Arrhenius plot. It is a straight line with slope of m = –Ea/R and a y-intercept of b = ln A.
  • k is the specific rate constant.
  • Ea is the activation energy.
  • R is the gas constant, 8.314 J mol–1 K–1.
  • T is the temperature in Kelvin.
  • A is a constant for a given reaction.
slide13

CHEMICAL KINETICS AND TEMPERATURE

  • Calculate the Ea for this reaction.
  • 2HI(g) → H2(g) + I2(g)
  • Step #1: Complete this table.

–14.860

556

0.00180

–10.408

629

0.00159

–8.426

666

0.00150

–6.759

700

0.00143

–3.231

781

0.00128

slide14

CHEMICAL KINETICS AND TEMPERATURE

  • Step #2: Use Excel to plot ln k versus 1/T. Then calculate the slope (–Ea/R) of this Arrhenius plot.
slide15

CHEMICAL KINETICS AND TEMPERATURE

  • Step #3: Calculate Ea.
  • Slope = (–Ea/R)
  • –Ea = (Slope)R
  • Ea = –(Slope)R
  • Ea = –(–2.24 x 104 K)( 8.314 J mol–1 K–1)
  • Ea = 1.86 x 105 J mol–1
  • Ea = 186 kJ mol–1
slide16

SAFETY

  • Give at least 1 safety concern for the following procedure.
  • Using acetone (CH3COCH3), hydrochloric acid (HCl), and iodine (I2).
  • These are irritants. Wear your goggles at all times. Immediately clean all spills. If you do get either of these in your eye, immediately flush with water.
  • Acetone is extremely flammable. Never use it near a flame or spark.
  • Your laboratory manual has an extensive list of safety procedures. Read and understand this section.
  • Ask your instructor if you ever have any questions about safety.
slide17

SOURCES

  • Barnes, D.S., J.A. Chandler. 1982. Chemistry 111-112 Workbook and Laboratory Manual. Amherst, MA: University of Massachusetts.
  • McMurry, J., R.C. Fay. 2004. Chemistry, 4th ed. Upper Saddle River, NJ: Prentice Hall.
  • Petrucci, R.H. 1985. General Chemistry Principles and Modern Applications, 4th ed. New York, NY: Macmillan Publishing Company.