A student placed a mixture of BaO and CaO in a 1.46L flask containing carbon dioxide at 35C and 746mmHg. After the reactions were complete she found that the CO2 pressure had dropped to 2.40 x 10^2mmHg. How many moles of CO2 were consumed in the reactions? Assume volumes of the solids are negligible.
So far, I know to convert to Kelvin, so T = 308K. P = 0.98atm
Initial P = 746 mm Hg.
Final P = 240
delta P = 746-240 = 506 mm Hg.
Change that to atm and use PV = nRT. YOu know P,V, R, and T. Calculate n.
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To find the number of moles of CO2 consumed in the reaction, we can use the Ideal Gas Law equation:
PV = nRT
Where:
P = pressure
V = volume
n = number of moles
R = ideal gas constant
T = temperature
First, we need to convert the given pressures from mmHg to atm. Since 1 atm = 760 mmHg, we can convert the given pressures as follows:
Initial pressure of CO2 (P1) = 746 mmHg / 760 mmHg/atm = 0.98 atm
Final pressure of CO2 (P2) = 2.40 x 10^2 mmHg / 760 mmHg/atm = 0.316 atm
Now, we can rearrange the Ideal Gas Law equation to solve for the number of moles (n):
n = PV / RT
Substituting the known values:
n = (P1V) / (RT)
Since we assume the volume of solids is negligible, the volume (V) is the same before and after the reaction. Therefore, we can omit it from the equation:
n = P1 / (RT)
Now, let's plug in the values:
P1 = 0.98 atm
R = 0.0821 L·atm/(mol·K)
T = 308 K
n = 0.98 atm / (0.0821 L·atm/(mol·K) * 308 K)
Now we can calculate the number of moles of CO2 consumed in the reaction.