C2O4 (aq) + H2O2 (aq) + 2 H+ (aq) → 2 CO2 (g) + 2 H₂O (1) Use the experimental data in the table to determine the rate law for this reaction (all concentrations in mol/L). Exp. [CO2] [H₂O₂] [H+] Rate (in M/s) 123+ 0.390 0.225 0.100 3.34x106 0.780 0.225 0.100 6.69x106 0.390 0.675 0.100 3.01x10-5 4 0.780 0.225 0.200 1.33x10-5

Express the rate of the reaction 2NO(g)+ O2(g)–→ 2NO2(g) a. in terms of ANO2]. rate = Η ΔΝΟ,1/Δt b. in terms of A[O2]. _ A[02]/At rate =

Given the following results: Rate of Run Reaction [1] (M/s) 1.76 0.358 0.476 4.72 0.358 0.843 5.25 0.672 0.476 1 2 3 [H₂O2] Calculate the rate constant(k) for each run. Assuming the Rate Law = k[1]²[H₂O₂] calculate the average rate constant,k, for this reaction. Give your answer to 2 places after the decimal.

Given the following results: Run Rate of Reaction (M/s) [I-] [H2O2] 1 1.99 0.516 0.561 2 3.15 0.516 0.643 3 5.22 0.817 0.561 Calculate the rate constant(k) for each run. Assuming the Rate Law = k[I-]2[H2O2] calculate the average rate constant,k, for this reaction. Give your answer to 2 places after the decimal.

Find the half-life of CO in the presence of Ru, we need to consider the pseudo-first order reaction with the given rate law ____ Min  Rate = (0.845 min^(-1)) [CO].    Step 1:  CO (g) +2 Ru(s) →2 RuCO (s)    Step 2:  2RuCO(s) + 3H2(g) →2 CH4(g) + 2 Ru(s)

It may not surprise you to know that products can show up in the rate law. When the product(s) is(are) a positive order, we describe the rate law as autocatalytic. For example, under conditions of high nitric acid concentrations, this reaction: HNO, + H₂C(OH): → HNO; + HCOOH + H₂O has a rate law: rate = k [H.C(OH).] [HNO₂] Given a rate constant of 0.0255 dm3 mol-1 s-1 and initial concentrations [HNO,] = 5.0 mol dm-3, [H.C(OH).]. = 150 mmol dm-3, [HNO:]. = 8.2 mmol dm-3, how long will it take for the methanediol concentration to fall to 52 mmol dm-3?

Given the following kinetic data for the first order production of H2O2, what is the rate constant (k) for the reaction? Do not put units in your answer. Trial Time (h) [H2O2] (M) ln[H2O2] 1 0 1.000 0.0 2 6.00 0.500 −0.693 3 12.00 0.250 −1.386 4 18.00 0.125 −2.079 5 24.00 0.0625 −2.772

-1 At a certain temperature this reaction follows first-order kinetics with a rate constant of 0.0482 s H,CO3 (aq) → H,O (aq)+ CO, (aq) Suppose a vessel contains H,CO, at a concentration of 0.300M. Calculate how long it takes for the concentration of H,CO, to decrease by 88.0%. You may assume no other reaction is important. Round your answer to 2 significant digits. S x10

Ammonium ion reacts slowly with nitrite ion: NH4+ (aq) + NO2-(aq)-> N2(g) + 2H20 (1) Rate data for this reaction, measured at a certain temperature, are shown in the attached table. Use your rate law, and the rate and concentrations for any of the experiments, to calculate the rate constant (k) for the reaction in /Ms. Type your response: Expt. [NH4+] (M) [NO₂] (M) Rate (M s-1) 1 0.150 0.220 7.00 x 10-6 2 0.600 0.220 2.80 x 10-5 3 0.600 0.110 1.40 x 10-5

5. You studied the chemical reaction, 2NO2(g) → 2NO(g) + O2(g), at 25°C by monitoring the concentration of NO2(g) as a function of time and constructed the following graph. y= 0.5408x + 125.0 R2=0.998 Time (s) What is the rate constant for this reaction at 25°C? Include the proper units. (W/T) [(3)ONI/T