Carbon monoxide at a pressure of 1490 mmHg reacts completely with O2 at a pressure of 745 mmHg in a sealed vessel to produce CO2. What is the final pressure in the flask?
2CO2 (g) + O2 (g) -> 2CO2 (g)
I thought it was just a simple addition, but it was wrong.
Just a different perspective. :-)
...............2CO(g) + O2(g) ==> 2CO2(g)
I...............1490.........745.............0
C...............-2x...........-x..............2x
E.................0.............0..............2x
Complete Rxn => 745 + (-x) = 0 => x = 745 mmHg
P(CO2) = 2X = 2(745 mmHg) = 1490 mmHg
First, redo your equation.
................2CO(g) + O2(g) ==> 2CO2(g)
I...............1490.........745.............0
C............-1490........-745............1490
E.................0..............0..............1490
Well, looks like there's some chemical comedy happening in that sealed vessel! Let's break this down with a little joke.
So, we have carbon monoxide (CO) and oxygen (O2) having a reaction party. They're shaking, stirring, and eventually make carbon dioxide (CO2) as a result.
Now, when it comes to pressure, it's like a bad joke – it can't just be added together like that. To figure out the final pressure in the flask, we need to use a different equation. Specifically, we'll be using Dalton's law of partial pressures.
According to Dalton, the total pressure of a mixture of gases is equal to the sum of the partial pressures of each gas. In this case, we have two gases: CO and O2.
Using this law, we'll calculate the final pressure:
Partial pressure of CO2 = Total pressure - Partial pressure of O2
Now, be careful not to laugh too hard!
So, the final pressure in the flask is 1490 mmHg - 745 mmHg = 745 mmHg.
And there you have it, the pressure in the flask is no joke! It's 745 mmHg.
To determine the final pressure in the flask, we need to take into account the stoichiometry of the reaction and the ideal gas law.
First, let's determine the partial pressure of carbon monoxide (CO) before the reaction. The initial pressure of CO is given as 1490 mmHg.
Next, let's consider the balanced equation:
2CO (g) + O2 (g) -> 2CO2 (g)
According to the stoichiometry, we can see that two moles of CO react with one mole of O2 to produce two moles of CO2. Therefore, the reaction consumes CO and O2 in a 2:1 ratio.
Since all the CO reacts, the pressure of CO at the end of the reaction will be zero. This means that the initial pressure of COO2 (pCO2) will be twice the initial pressure of CO, which is 2 * 1490 mmHg = 2980 mmHg.
However, we still need to consider the pressure of the O2 gas. The initial pressure of O2 is given as 745 mmHg. Since we're assuming the reaction goes to completion, all the O2 will be consumed as well.
The final pressure in the flask will be the sum of the partial pressures of the remaining gases, which is the pressure of CO2 (2 * 1490 mmHg) plus the pressure of the remaining O2 (0 mmHg). Therefore, the final pressure in the flask is:
Final pressure = pCO2 + pO2 = 2980 mmHg + 0 mmHg = 2980 mmHg.
To find the final pressure in the flask, we need to apply Dalton's Law of Partial Pressure. According to this law, the total pressure of a mixture of gases is equal to the sum of the partial pressures of each gas.
In this case, we have two gases: carbon monoxide (CO) and oxygen (O2). Initially, the CO is at a pressure of 1490 mmHg, and the O2 is at a pressure of 745 mmHg. We need to determine the final pressure in the flask after the reaction has occurred.
Since the reaction is complete, all of the CO will react with O2, forming carbon dioxide (CO2). The balanced equation given shows that two molecules of CO react with one molecule of O2 to produce two molecules of CO2.
To find the final pressure, we need to consider the moles of each gas present. From the balanced equation, we can see that two moles of CO react with one mole of O2. Therefore, the number of moles of CO remaining after the reaction is half the original amount. The number of moles of O2 remaining is determined by stoichiometry as well.
Now, let's calculate the moles of each gas in the mixture:
Moles of CO = (initial pressure of CO) / (total pressure) * Total moles
= (1490 mmHg) / (1490 mmHg + 745 mmHg) * Total moles
Moles of O2 = (initial pressure of O2) / (total pressure) * Total moles
= (745 mmHg) / (1490 mmHg + 745 mmHg) * Total moles
Since there is complete reaction, the total moles of CO and O2 are the same.
Using these calculated moles, we can determine the final pressure of the mixture.
Final pressure = (moles of CO2) / (Total moles) * (total pressure)
= (2 moles) / (Total moles) * (1490 mmHg + 745 mmHg)
Note: To solve this problem completely, we need to know the moles or any other information related to the amount of each gas present initially. Without this information, we cannot provide the exact value.