General Chemistry II Lab Manual, 2017 Revision 79 Questions: 1) 5.0 moles of NH3

Question

General Chemistry II Lab Manual, 2017 Revision 79 Questions: 1) 5.0 moles of NH3 and 4.0 moles of H2 are introduced into a 4.00 L flask. At equilibrium, 1.5 moles of NHs gas remain. Calculate Ke for this reaction. 3H2(g) + N2(g) 2NHs (8) 2) The Kc value for the following equilibrium at 500°C is 49. 112 (g) + 12(g) 2H1(g) If 2.00 moles of H2 and 2.00 moles of I2 are introduced into a 2.00 L flask at 500°C, how many moles of HI are present at equilibrium? What would happen to the values you calculated for Kc in the experiment if the Beer's law constant (k) value you used was high (e.g. you used 6000 M when the actual constant had a value of 5,000 M)? Explain your answer. 3)

Explanation / Answer

1)
initially,
[NH3] = moles of NH3 / volume in L
= 5.0 moles / 4.00 L
= 1.25 M
[H2] = moles of H2 / volume in L
= 4.0 moles / 4.00 L
= 1.00 M

At equilibrium:
[NH3] = moles of NH3 / volume in L
= 1.5 moles / 4.00 L
= 0.375 M

                    [H2]                [N2]                [NH3]             

initial             1.0                 0                   1.25              

change              +3x                 +1x                 -2x               

equilibrium         1.0+3x              +1x                 1.25-2x           

Given at equilibrium,
[NH3] = 0.375
1.25-2x = 0.375
x = 0.4375

Equilibrium constant expression is
Kc = [NH3]^2/[H2]^3*[N2]
Kc = (1.25-2x)^2/(1.0+3x)^3*(x)
Kc = (1.25-2*0.4375)^2/(1.0+3*0.4375)^3*(0.4375)
Kc = (0.375)^2/(2.3125)^3*(0.4375)
Kc = 2.6*10^-2

AnsweR: 2.6*10^-2

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