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CH 104: CHEMICAL KINETICS

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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CH 104: CHEMICAL KINETICS

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  1. 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)

  2. 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.

  3. 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

  4. 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.

  5. 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.

  6. 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.

  7. 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]

  8. 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

  9. 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.

  10. 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.

  11. 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.

  12. 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.

  13. 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

  14. 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.

  15. 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

  16. 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.

  17. 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.

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