HCONH2(g) <-> NH3(g) + CO2(g) Kc= 4.84 at 400 K2(g)
If 0.186 mol of HCONH2(g) dissciates in 2.16L at 400K, what will be the total pressure at equilibrium?
............HCONH2 ==> NH3 + CO2
initial.....0.186 mols..0.....0
change.......-x.........x......x
equil.......0.186-x.....x......x
Substitute the above ICE chart into the Kc expression for the reaction and solve for x. That will give you moles NH3 and moles CO2. Add the moles together and substitute into PV = nRT to solve for pressure.
When I am solving for X my formula comes out to be
4.84 = X^2/ (.186-X)
Can I assume that x << .05 and eliminate it from the denominator and just solve for the equation 4.84 = X^2/.186?
No, you may not make that assumption; however, I may have made an error. Let me think about this awhile. I'll post something different and erase what is there now or post a note that the original response stands.
Okay thanks so much.
OK. I made an error and the problem can't be solved that way. First, we convert moles to molarity.
(HCONH2) = 0.186/2.16L = 0.08611M
............HCONH2 ==> NH3 + CO2
initial.....0.08611M....0.....0
change.......-xM.........xM......xM
equil......0.08611-x M.....xM....xM
Substitute the above ICE chart into the Kc expression for the reaction and solve for x. That will give you M NH3, M CO2 and M HCONH2. Convert M to moles for each (M x L = moles), add the moles together and substitute into PV = nRT to solve for pressure. I will leave the original post there for the time being so you can make comparisons if you wish; however, I'll erase it before I go to bed tonight. I don't like to leave incorrect posts on the board.
Ok thanks so much.
So when I solve for X I got
4.48(0.086 - x) = x^2
0.385 - 4.48x - x^2 = 0
Now i'd substitute that into the quadratic formula? Or is there an easier way to solve for X?
That K is 4.84 and not 4.48 so you need to redo that part. When you solve that quadratic that gives you x (and I don't know of an easier way to do it.) Most calculators now have that built in to solve those things. I programmed mine to do that. That gives you x, use that to determine (HCONH2), (NH3), and (CO2), convert concns to moles, add the moles to find the total, then use PV = nRT to solve for P.
Ok this helps so much. Thanks a lot I appreciate it.
To find the total pressure at equilibrium, we need to first determine the equilibrium concentrations of all the species involved in the reaction.
We are given that 0.186 mol of HCONH2(g) dissociates, so the equilibrium concentration of HCONH2 (denoted as [HCONH2]) is given by:
[HCONH2] = initial concentration of HCONH2 - moles dissociated/volume
= (0.186 mol) / (2.16 L)
= 0.0861 M
Since the stoichiometric coefficient of HCONH2 is 1, the equilibrium concentration of NH3 will also be 0.0861 M.
For CO2, the stoichiometric coefficient is also 1, but since it is not involved in the dissociation of HCONH2, its initial concentration remains unchanged.
Now, we can calculate the total pressure at equilibrium using the ideal gas law, which states:
PV = nRT
Since we have the equilibrium concentrations in Molarity, we need to convert them to partial pressures in atmospheres before plugging them into the equation.
Partial pressure (P) can be calculated using the equation:
P = concentration (M) * (R * T) / (0.0821 L * atm / mol * K)
Here, R is the ideal gas constant (0.0821 L * atm / mol * K), and T is the temperature in Kelvin (400 K).
For HCONH2 and NH3, the partial pressure will be the same since they have the same concentration (0.0861 M). For CO2, we need to use its initial concentration, as it does not change during the reaction.
Let's calculate the partial pressures:
Partial pressure of HCONH2 = (0.0861 M) * (0.0821 L * atm / mol * K) * (400 K) / (0.0821 L * atm / mol * K)
= 3.37 atm (rounded to two decimal places)
Partial pressure of NH3 = (0.0861 M) * (0.0821 L * atm / mol * K) * (400 K) / (0.0821 L * atm / mol * K)
= 3.37 atm (rounded to two decimal places)
Partial pressure of CO2 = (initial concentration of CO2) * (0.0821 L * atm / mol * K) * (400 K) / (0.0821 L * atm / mol * K)
= (initial concentration of CO2) * 400 K
Now, since CO2 is not involved in the dissociation of HCONH2, its initial concentration remains unchanged. Therefore, we do not have enough information to determine the total pressure at equilibrium as we do not know the initial concentration of CO2.
In conclusion, without the information about the initial concentration of CO2, we cannot calculate the total pressure at equilibrium in this scenario.