Description At equilibrium at 1.0 L vessel contains 17.00 mol of H2, 20.00 mol of CO2, 13.00 mol of H2O, and 6.000 mol of CO at 427?C. CO2(g) + H2(g) ?? CO(g) + H2O(g) What is the value of K at this temperature for the reaction?
Description What is the percent yield of water if 0.90 g of water is obtained when 29.0 g of butane is burned in excess oxygen? The balanced chemical equation is... (a) 0.02% (b) 2% (c) 10% (d) 36%
Description If 16 grams of O2 react with excess C2H6, how many grams of CO2 will be formed? The formula mass of O2 = 32 amu and the formula mass of CO2 = 44 amu. The balanced chemical equation is... (a) 22 grams (b) 13 grams (c) 9 grams (d) 7 grams
Description NH4+ (aq) + NO2- (aq) ? N2 (g) + H2O (l) Experiment [NH4+]i [NO2-]i Initial rate (M/s) 1 0.24 0.10 7.2 x 10-4 2 0.12 0.10 3.6 x 10-4 3 0.12 0.15 5.4 x 10-4 4 0.12 0.12 4.3 x 10-4 First determine the rate law and rate constant. Under the same initial conditions as in Experiment 4, calculate [NH4+] at 284 seconds after the start of the reaction. In this experiment, both reactants are present at the same initial concentration. The units should be M, and should be calculated to three significant figures. 2. NO2(g) + CO(g)> NO(g) + CO2(g) The rate of the above reaction depends only on the concentration of nitrogen dioxide at temperatures below 225°C. At a temperature below 225°C the following data were obtained: Time (s) [NO2] (mol/L) 0 0.520 1.10×103 0.481 2.75×103 0.433 4.12×103 0.400 8.25×103 0.325 1.65×104 0.236 Which of the following expressions for the rate law (either differential or integrated) are completely consistent with the above experimental data. k[NO2]3 = -d[CO]/dt ln([NO2]/[NO2]0) = -kt k[NO2]2 = d[NO]/dt -d[CO]/dt = k[NO2]2 1/[NO2] - 1/[NO2]0 = kt 3.The data below was collected for the reaction: 2N2O5(g) ? 4NO2(g) + O2(g) at a temperature of 25°C: time (hr) [N2O5](M) 0.0 0.374 0.9 0.337 1.8 0.299 2.9 0.262 4.2 0.224 5.7 0.187 7.5 0.150 9.9 0.112 13.2 0.075 18.9 0.037 Using the data, Find t½. Plot [N2O5], ln[N2O5] and 1/[N2O5] as a function of time. You can cut and paste the data from the table into an Excell worksheet. From the graphs, determine the reaction order and the reaction rate constant. Reaction order: Rate constant: The effect of temperature on the rate of a reaction was studied and the following data obtained: k (s-1) T (°C) 1.91×10-4 5 2.25×10-4 7 3.07×10-4 11 3.58×10-4 13 4.84×10-4 17 7.49×10-4 23 1.06×10-3 28 1.22×10-3 30 It is known that the variation of the rate constant k with the absolute temperature T is described by the Arrhenius equation: k = A exp [( ?Ea )/(RT)] where Ea is the activation energy, R is the universal gas constant and A is the pre-exponential factor (units of the rate constant). Taking the natural logarithm of both sides affords: ln k = ln A ? Ea RT For a plot of y = ln k versus x = 1/T, calculate the slope of the best straight line using linear regression. Calculate the activation energy Ea.
Description Aspirin can be made in the laboratory by reacting acetic anhydride (C4H6O3) with salicylic acid (C7H6O3) to form aspirin (C9H8O4) and acetic acid (C2H4O2). The balanced equation is: C4H6O3 + C7H6O3 ---àC9H8O4 + C2H4O2 In a laboratory synthesis, a student begins with 5.00 mL of acetic anhydride (density = 1.08 g/mL) and 2.08 g of salicylic acid. Once the reaction is complete, the student collects 2.01 g of aspirin. Determine the limiting reactant, theoretical yield of aspirin, and percent yield for the reaction.
