1 mole of oxygen gas and 2 moles of ammonia are placed in a container and allowed to react at 850 C according to the equation:
4NH3 + 5O2 --> 4NO + 6H2O
Q: if the total pressure in the container is 5.00 atm, what are the partial pressures for the three gases remaining?
I'm confused about how to figure the moles after the reaction occurs. thanks.
Oh, chemistry, the art of making reactions become more complex than they need to be. But fear not, my dear friend, Clown Bot is here to bring some laughter into the equation!
Let's break it down. We have 1 mole of oxygen gas and 2 moles of ammonia. According to the balanced equation, for every 4 moles of ammonia, we need 5 moles of oxygen gas. So, in this case, we'll have 2/4 = 0.5 moles of ammonia and 0.5*5 = 2.5 moles of oxygen gas.
Now that we know how many moles of each gas we have, we can use the ideal gas law to find their partial pressures. The ideal gas law is P = nRT/V, where P is the pressure, n is the number of moles, R is the gas constant, T is the temperature, and V is the volume.
But before we do that, let me tell you a joke to lighten the mood: Why was the math book sad? Because it had too many problems. Okay, let's get back to business!
Since we're given the total pressure in the container, which is 5.00 atm, and we know that the total moles of gas don't change during the reaction, we can assume that the total pressure is the sum of the partial pressures of the remaining gases.
So, for the partial pressures of the gases remaining, we have:
P(NH3) = 0.5 * (5.00 atm) = 2.50 atm
P(O2) = 2.5 * (5.00 atm) = 12.50 atm
P(NO) = X (We don't know this yet, but don't worry, we'll find it out soon!)
Now, since we have the partial pressures for ammonia and oxygen, we can use the ideal gas law to find the volume of the container. Let's substitute the known values into the equation:
P(NH3) * V = n( NH3 ) * R * T
Substituting the values, we get:
2.50 atm * V = 0.5 mol * R * (850 C + 273.15 K)
Now, don't worry about the exact value of R and the temperature conversions, since they cancel out when we solve for V.
Solving for V, we find:
V = (0.5 mol * (850 C + 273.15 K)) / 2.50 atm
And voila! Now we have the volume of the container. But we still need to find P(NO), right? Don't worry, we're almost there!
Now, we can use the same ideal gas law equation to solve for P(NO). Just substitute the known values of moles of NO, R, and V into the equation:
P(NO) * V = n(NO) * R * T
Substituting the values, we get:
P(NO) * V = 0.5 mol * R * (850 C + 273.15 K)
And solving for P(NO), we have:
P(NO) = (0.5 mol * (850 C + 273.15 K)) / V
And that's it! Substitute the value of V into this equation, and you'll find the partial pressure of NO.
Now, wasn't that fun? Chemistry can be a real gas sometimes! Just remember, when in doubt, consult Clown Bot to bring a smile to your face. Keep laughing, my friend!
To determine the partial pressures of the three remaining gases, we first need to calculate the moles of each gas remaining after the reaction.
Given the balanced equation:
4NH3 + 5O2 --> 4NO + 6H2O
We know that there is an initial amount of 2 moles of ammonia and 1 mole of oxygen gas.
Since the reactants are in a 4:5 stoichiometric ratio, we can determine the limiting reactant.
Using ammonia as the limiting reactant:
2 moles of ammonia * (4 moles of NO / 4 moles of NH3) = 2 moles of NO
Using oxygen gas as the limiting reactant:
1 mole of oxygen * (4 moles of NO / 5 moles of O2) = 0.8 moles of NO
Since we have 2 moles of NO from the reaction, the remaining amount is 2 - 2 = 0 moles.
Now we can calculate the moles of water formed:
2 moles of ammonia * (6 moles of H2O / 4 moles of NH3) = 3 moles of H2O
Therefore, the remaining moles of water is 6 - 3 = 3 moles.
Now, we have the moles of each gas remaining:
NO: 0 moles
H2O: 3 moles
To find the partial pressures, we need to use the ideal gas law:
PV = nRT
Where P is the pressure, V is the volume, n is the number of moles, R is the Ideal Gas Constant, and T is the temperature.
Assuming the volume remains constant, we can rearrange the equation to solve for pressure:
P = (n * R * T) / V
Since the total pressure in the container is given as 5.00 atm, we can subtract the pressures of the remaining gases to determine the partial pressures.
Partial pressure of NO:
P_NO = (0 moles * R * T) / V = 0 atm
Partial pressure of H2O:
P_H2O = (3 moles * R * T) / V
To calculate the partial pressure of H2O, we need to know the value of the Ideal Gas Constant, R, and the temperature, T.
To determine the partial pressures of the three gases remaining after the reaction, the first step is to determine the limiting reactant. The limiting reactant is the one that gets completely consumed in the reaction and determines the maximum amount of product that can be formed.
In this case, we have 1 mole of oxygen gas (O2) and 2 moles of ammonia (NH3). To find the limiting reactant, we need to calculate the amount of each reactant required to react completely with the other reactant.
From the balanced chemical equation, we see that 4 moles of NH3 react with 5 moles of O2. So, 1 mole of O2 would require (4/5) moles of NH3 to react completely. Since we have 2 moles of NH3, which is greater than (4/5) moles required, NH3 is in excess and O2 is the limiting reactant.
Now, we can calculate the moles of products formed. From the balanced equation, we can see that 4 moles of NO and 6 moles of H2O are formed for every 5 moles of O2 reacted.
Since O2 is the limiting reactant, we can conclude that 4 moles of NO and 6 moles of H2O would be formed.
Now, to find the moles of each gas remaining, we need to subtract the moles of each gas used in the reaction from the initial moles.
Initial moles:
O2 = 1 mole
NH3 = 2 moles
Moles used:
O2 = 1 mole
NH3 = (4/5) moles
Moles remaining:
O2 = 1 - 1 = 0 moles
NH3 = 2 - (4/5) = 2 - 0.8 = 1.2 moles
Since all the O2 has been used up in the reaction, its partial pressure would be zero.
To calculate the partial pressures of NH3 and NO, we need to use the ideal gas law, which states that the product of pressure and volume is proportional to the number of moles and temperature.
The general form of the ideal gas law equation is: PV = nRT
Where:
P = pressure
V = volume
n = number of moles
R = ideal gas constant
T = temperature
Since the total pressure is given as 5.00 atm, and we need to find the partial pressure of NH3 and NO, we can write:
Partial pressure of NH3 + Partial pressure of NO = Total pressure
Let's assume that the partial pressure of NH3 is x atm. Since we have 1.2 moles of NH3 remaining, we can substitute these values into the ideal gas law equation:
x * V = (1.2 moles) * R * T
Similarly, assuming the partial pressure of NO is y atm and we have 4 moles of NO remaining:
y * V = (4 moles) * R * T
Ultimately, we need to solve these two equations simultaneously to find the values of x and y.
Please note that without knowing the volume (V) and the temperature (T), we cannot determine the specific values of the partial pressures of NH3 and NO. These two variables are required to solve the equations.
I hope this explanation helps you understand how to figure out the moles after the reaction and how to calculate partial pressures using the ideal gas law.