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# [Solved] Assignment 219145

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Table1: Temperature, Pressure and Volume DataTemperature of Distilled H2O:Room (or regional) Pressure (atm):Initial Volumeof Air (mL)Final Volume of Air(after reaction) (mL)Volume of O2 Collected(Final Volume – Initial Volume)191.00atm25 mL75 mL.061Table2: Reaction Time DataTime Reaction StartedTime Reaction EndedTotal Reaction Time11:0011:033 minutes 39 secondsPost-Lab Questions1. Calculate the number of moles of O2 produced using the ideal gas law. Then, use this value to calculate the number of moles of hydrogen peroxide you began the experiment with. HINT:Use the balanced equation provided in the lab introduction.· Pv=nRT (· 1.00 (.025L)= n (0.0821 (L x atm)/(mol x K) (292 K)· .025L/ 23.97 = .001 M O22. Calculate the number of moles of hydrogen peroxide you would have if you used 5 mL of a pure hydrogen peroxide solution. HINT: The density of hydrogen peroxide is 1.02 g/mL.
5 mL = .005 L hydrogen peroxide
1.00 (.005L) = n (.0821 Lxatm/ mol x K) (292 K)
.005/ 23.97 = .00021 M Hydrogen peroxide3. Determine the percentage of hydrogen peroxide in your solution.4. Was the calculated percentage of hydrogen peroxide close to the same as the percentage on the label (3%)? Calculate percent error of your value.5. Considering that catalysts are not consumed in a reaction, how do you think increasing the amount of catalyst would affect the reaction rate for the decomposition of hydrogen peroxide?6. What was going on in the graduated cylinder as the H2O was pushed out?7. How would the number of moles (n) of O2 change if your atmospheric pressure was doubled and all other variables stayed the same?

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