An experiment was carried out to determine the value of the equilibrium constant Kc for the reaction.
Total moles of Ag+ present = 3.6 x 10-3 moles
Total moles of NH3 present = 6.9 x 10-3 moles
Measured concentration of Ag(NH3)2+ at equilibrium= 3.4* 10-2 M
Total solution volume = 100 mL
(a). Calculate the equilibrium concentration of Ag+ (uncomplexed).
(b). Calculate the equilibrium concentration of NH3 (uncomplexed).
(c). Calculate the value of the equilibrium constant (Kc).
I GOT 1.7*10^7 for question C. IS IT RIGHT?????????
What's the reaction? I'm guessing at
Ag^+ + 2NH3 ==> Ag(NH3)2^+
If that's the reaction my number is 1.75E7 which I would round to 1.8E7 instead of your 1.7E7. Good work.
THANX
uju
How
Well, my mathematical skills may be a bit rusty, but let's take a look at your answer for question C. The equilibrium constant, Kc, is calculated by comparing the concentrations of the products to the concentrations of the reactants, with each concentration raised to the power of its stoichiometric coefficient.
If we have the equation: Ag+ + 2NH3 ⇌ Ag(NH3)2+
We can see that the stoichiometric coefficient of Ag(NH3)2+ is 1, while the stoichiometric coefficient of Ag+ is also 1. The stoichiometric coefficient of NH3 is 2.
To calculate Kc, we use the formula:
Kc = [Ag(NH3)2+]/[Ag+] * [NH3]^2
Now, plug in the values:
[Ag(NH3)2+] = 3.4 * 10^-2 M
[Ag+] = ? (This is what we're trying to find, so let's call it x)
[NH3] = 6.9 * 10^-3 M
So, Kc = (3.4 * 10^-2 M) / (x M) * (6.9 * 10^-3 M)^2
To find x, we need to use the total moles of Ag+ present, which you provided as 3.6 * 10^-3 moles, and the total solution volume of 100 mL.
We can calculate the concentration of Ag+ using the formula:
[Ag+] = moles of Ag+ / total solution volume
[Ag+] = (3.6 * 10^-3 moles) / (100 mL) * (1 L / 1000 mL)
Now, convert the concentration of Ag+ from mol/L to M by dividing by 1000:
[Ag+] = (3.6 * 10^-3 moles) / (100 mL) * (1 L / 1000 mL) * (1000 mol / 1 L)
After doing the math, you should find the equilibrium concentration of Ag+ (uncomplexed). Once you have that, you can plug it into the Kc formula we discussed earlier to calculate Kc.
I hope my explanation didn't make you lose your equilibrium!
To calculate the equilibrium concentrations and the value of the equilibrium constant, we need to use the given information and apply the principles of equilibrium expressions and stoichiometry.
(a) Calculate the equilibrium concentration of Ag+ (uncomplexed):
1. Start with the total moles of Ag+ present: 3.6 x 10^-3 moles.
2. Convert moles to concentration by dividing by the total volume of the solution (100 mL = 0.1 L): (3.6 x 10^-3 moles) / 0.1 L = 0.036 M.
Therefore, the equilibrium concentration of Ag+ (uncomplexed) is 0.036 M.
(b) Calculate the equilibrium concentration of NH3 (uncomplexed):
1. Start with the total moles of NH3 present: 6.9 x 10^-3 moles.
2. Convert moles to concentration by dividing by the total volume of the solution (100 mL = 0.1 L): (6.9 x 10^-3 moles) / 0.1 L = 0.069 M.
Therefore, the equilibrium concentration of NH3 (uncomplexed) is 0.069 M.
(c) Calculate the value of the equilibrium constant (Kc):
The balanced chemical equation for the reaction is:
Ag+(aq) + 2NH3(aq) ⇌ Ag(NH3)2+(aq)
The equilibrium constant expression, Kc, can be determined by taking the ratio of the products' concentration to the reactants' concentration, each raised to the power of their respective stoichiometric coefficients.
Kc = [Ag(NH3)2+] / ([Ag+] * [NH3]^2)
Given that the measured concentration of Ag(NH3)2+ at equilibrium is 3.4 x 10^-2 M, and the equilibrium concentrations of Ag+ and NH3 (uncomplexed) are calculated as 0.036 M and 0.069 M, respectively, we can substitute these values into the equilibrium constant expression to find Kc.
Kc = (3.4 x 10^-2) / (0.036 * (0.069)^2)
Upon calculation, we find that Kc is approximately equal to 1.87 x 10^3 (rounded to three significant figures).
Therefore, the value you obtained for the equilibrium constant, 1.7 x 10^7, is not correct. The correct value is 1.87 x 10^3.