Reaction: 5C + 2SO2 -> CS2 + 4CO
1) How many moles of CS2 are there for when 2.7 moles of C reacts?
2) How many moles of carbon are needed to react with 5.44 moles of SO2?
3) How many moles of carbon monoxide form at the same time that .246 moles of CS2 forms?
4) How many moles of SO2 are required to make 118 moles of CS2?
1.) 5 moles of C = 1 mole of CS2
Therefore, 2.7 moles of C = 2.7/5 = 0.54 moles of CS2 (Answer)
2.) 2 moles of SO2 ...
This solution is comprised of the step by step instruction required in order to solve the different questions associated with the reaction given in this problem. Reasoning is provided so that an understanding of how the answers were obtained can be acquired.
Lab : Gases, Pressure and Stoichiometry
Add 0.5g of sodium bicarbonate, NaHCO3 into a flask. Record the initial pressure and temperature in the flask. Add 10 mL of water to the flask to dissolve the NaHCO3 then record the pressure as the initial pressure for the experiment. Record the gas volume as the initial volume. Add 1 mL of 1M HCl, allow the temperature in the flask to return to room temperature, and then record the pressure in the flask. Continue to add 1 mL increments of HCl solution. (observe whether a reaction has occurred).
These are the initial conditions of the closed flask with just the sodium carbonate inside. Assume that the sodium bicarbonate powder takes up a negligible volume in the flask.
(a) Initial pressure (atm): 1.072 atm
(b) Initial gas volume (mL): 139.72 mL
(c) Initial temperature (C): 21C
These are the final conditions of the closed flask:
(a) Final pressure (atm): 2.405 atm
(b) Final gas volume (mL): 131.92mL
(c) Final temperature (C): 21 C
1. Use the initial conditions and the ideal gas law to calculate the number of moles of air molecules (N2, O2, etc.) trapped in the flask. The gas constant, R = 0.082057 (L*atm)/(K*mole), so make sure to convert your temperature value to Kelvin.
2. The final conditions represent the air and CO2 from the reaction held in a smaller volume than at the start. Calculate the total moles of air and CO2 molecules from the ideal gas law, using the final conditions of pressure, volume and temperature. Record the total moles here.
3. Subtract the moles of air (from the initial conditions) from the total moles of gas at the end to find the moles of CO2 produced by the reaction.Record the moles of CO2 gas produced by the reaction:
4. You now have enough information to balance the equation: NaHCO3(s) + HCl(aq) --> CO2(g) + H2O(l) + NaCl(aq)
Calculate the following:
(a) Moles of NaHCO3 added to the flask (MW=84.007 g/mol):
(b) Moles of CO2 produced
(c) The ratio of (moles of CO2) divided by (moles of NaHCO3)
(d) What is the nearest whole number ratio for your experimentally-determined ratio? For example, if the experimental ratio is 1.75, then you would use a whole number ratio of 2 to 1. If the experimental ratio is 3.2, then you would use a whole number ratio of 3 to 1.
(e) Use the whole number ratio to balance the equation for the reaction. If your whole number ration is 2 to 1, then place a 2 next to the CO2 and a 1 next to the NaHCO3. Continue from there to balance the number of atoms on each side of the equation. Write the balanced equation here:
(f) Calculate the experimental error from the difference between the experimental ratio and the whole number ratio as follows:
% error = |real ratio - whole number ratio| / (whole number ratio) * 100%
5. (a) Could you boil off the liquid at the end of the experiment and weigh the NaCl left in the flask in order to balance the equation?
(b) Air usually contains water vapor. The air trapped in the flask was dry air with no water vapor. How would the presence of water vapor affect your calculations?
(c) Carbon dioxide is slightly soluble in water and its solubility increases with increasing pressure. How would the fact that some of the carbon dioxide produced by the reaction is dissolved in the aqueous solution affect your calculations?