I solved part I but I'm having a hard time with part 2: Question 1. A student mixes 5.00 mL of 2.00 x 10‐3 M Fe(NO3)3 with 5.00 mL 2.00 x 10‐3 M KSCN. She finds that in the equilibrium mixture the concentration of FeSCN+2 is 1.40 x 10‐4 M. What is the initial concentration in solution of the Fe+3 and SCN‐? What is the equilibrium constant for the reaction? My solution: number of Fe +3 = M*V = 2*10^-3 M * 0.005 L = 1*10^-5 M number of SCN-= M*V = 2*10^-3 M * 0.005 L = 1*10^-5 M Total volume = 5 +5 = 10 mL = 0.01 L Initial concentration of Fe+3, [Fe+3] = number of moles / total volume =1*10^-5 / 0.01 =1*10^-3 M Initial concentration of SCN-, [SCN-] = number of moles / total volume =1*10^-5 / 0.01 =1*10^-3 M ------------------------------------------------------------------------------- Fe+3 + SCN - --->FeSCN+2 10^-3 10^-3 0 (initial) 10^-3-x 10^-3-x x (at equilibrium) here x = 1.4*10^-4 M At equilibrium: [FeSCN+2] = x = 1.4*10^-4 M [Fe+3] = 10^-3-x = 10^-3-1.4*10^-4 =8.6*10^-4 M [SCN-]= 10^-3-x = 10^-3-1.4*10^-4 =8.6*10^-4 M Kc = [FeSCN+2] / {[Fe+3]*[SCN-]} Kc = (1.4*10^-4) / {(8.6*10^-4 )*(8.6*10^-4 )} =189.3 this is the part where I get stuck: Question 2: Assume that the reaction studied is actually: (photo attached) a) What is the equation for determining the equilibrium constant? b) Using the information from question 1 and assuming [Fe(SCN)2+] = 1.40 x 10-4 M, calculate [Fe3+]e and [SCN-]e. c) Determine the numerical value of K.
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I solved part I but I'm having a hard time with part 2:
Question 1.
A student mixes 5.00 mL of 2.00 x 10‐3 M Fe(NO3)3 with 5.00 mL 2.00 x 10‐3 M KSCN. She finds that in the equilibrium mixture the concentration of FeSCN+2 is 1.40 x 10‐4 M.
What is the initial concentration in solution of the Fe+3 and SCN‐?
What is the equilibrium constant for the reaction?
My solution:
number of Fe +3 = M*V = 2*10^-3 M * 0.005 L = 1*10^-5 M
number of SCN-= M*V = 2*10^-3 M * 0.005 L = 1*10^-5 M
Total volume = 5 +5 = 10 mL = 0.01 L
Initial concentration of Fe+3,
[Fe+3] = number of moles / total volume
=1*10^-5 / 0.01
=1*10^-3 M
Initial concentration of SCN-,
[SCN-] = number of moles / total volume
=1*10^-5 / 0.01
=1*10^-3 M
-------------------------------------------------------------------------------
Fe+3 + SCN - --->FeSCN+2
10^-3 10^-3 0 (initial)
10^-3-x 10^-3-x x (at equilibrium)
here x = 1.4*10^-4 M
At equilibrium:
[FeSCN+2] = x = 1.4*10^-4 M
[Fe+3] = 10^-3-x = 10^-3-1.4*10^-4 =8.6*10^-4 M
[SCN-]= 10^-3-x = 10^-3-1.4*10^-4 =8.6*10^-4 M
Kc = [FeSCN+2] / {[Fe+3]*[SCN-]}
Kc = (1.4*10^-4) / {(8.6*10^-4 )*(8.6*10^-4 )}
=189.3
this is the part where I get stuck:
Question 2:
Assume that the reaction studied is actually:
(photo attached)
a) What is the equation for determining the equilibrium constant?
b) Using the information from question 1 and assuming [Fe(SCN)2+] = 1.40 x 10-4 M, calculate [Fe3+]e and [SCN-]e.
c) Determine the numerical value of K.
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