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Problem #1: When 0.0322 mol of NO and 1.70 g of bromine are placed in a 1.00 L reaction vessel and sealed, the mixture reacts and the following equilibrium is established:
2NO(g) + Br2(g) ⇌ 2NOBr(g)
At 25.0 °C the equilibrium of nitrosyl bromide is 0.438 atm. What is the Kp?
1) Determine pressure exerted by initial amount of NO:
P = (nRT)/V
P = [ (0.0322 mol) (0.08206 L atm mol¯1 K¯1) (298 K) ] / 1.00 L
P = 0.787415 atm (keeping a few guard digits)
2) Determine pressure exerted by initial amount of Br2:
P = (nRT)/V
P = (gRT)/(molar mass times V)
P = [ (1.70 g) (0.08206 L atm mol¯1 K¯1) (298 K) ] / [ (159.808 g mol¯1) (1.00 L) ]
P = 0.260135 atm
3) Determine equilibrium pressure of NO:
From the balanced equation, the NO : NOBr ratio is 1:1. Therefore, 0.438 atm of NO reacted.
0.787415 atm - 0.438 atm = 0.349415 atm
4) Determine equilibrium pressure of Br2:
From the balanced equation, the Br2 : NOBr ratio is 1:2. Therefore, 0.438/2 = 0.219 atm of Br2 reacted.
0.260135 atm - 0.219 atm = 0.041135 atm
5) Write the Kp expression and solve it:
Kp = PNOBr2 / [ ( PNO2) (PBr2) ]
Kp = (0.438)2 / [ (0.349415)2 (0.041135) ]
Kp = 38.2 (to 3 sig fig.)
Problem #2: A 1.00 L vessel contains at equilibrium 0.300 mol of N2, 0.400 mol H2, and 0.100 mol NH3. If the temp is maintained constant, how many moles of H2 must be introduced into the vessel in order to double the equilibrium concentration of NH3?
1) Solve for the Kc first:
Kc = [NH3]2 / ( [N2] [H2]3 )
x = (0.100)2 /[ (0.300) (0.400)3 ]
x = 0.5208333
2) Create a new equilibrium expression with [NH3] = 0.200. Use the stoichiometry of the equation to determine the [N2] when the [NH3] is doubled. Set x = [H2]. Remember the cube:
[N2] = 0.250 since 0.050 moles of N2 are required to make the added 0.100 mole of NH3
0.52083 = (0.200)2 /[ (0.250) (x)3 ]
x = 0.675 M
3) This is the [H2] present in the new equilibrium conditions (remember, we are determining how much H2 got added, which is why it is larger than 0.400). Figure out how much had to have been used (from the balanced equation). Add that to the equilibrium amount. This is the total initial [H2]:
the molar ratio between NH3 and H2 is 2:3; therefore, 0.15 mole of H2 got used up in making the added 0.100 mol of NH3 in the new equilibrium
0.675 + 0.150 = 0.825 M (this is the initial amount of H2 in the equilibrium that eventually produced 0.200 mol of NH3)
4) Subtract 0.400 to get the amount of H2 that was added:
0.825 - 0.400 = 0.425 mol of H2 added to push the [NH3] from 0.100 to 0.200 in the new equilibrium
Comment on the above solution: I posted the above answer on sci.chem many years ago. I got this response (which has an error in it):
After solving for Keq at the initial conditions, I set the ammonia concentration to 0.200 M, the nitrogen concentration to 0.250 M, and the hydrogen concentration to (0.400 + x) M. After plugging the new values into the expression for Keq and solving for x, my calculations yielded a value of 0.274746132241 moles of hydrogen needed. (aren't TI-85's wonderful). This is of course too many sig figs but when plugged into the problem it does agree with the original Keq to 10 places. My answer would then be that 0.275 moles of hydrogen are needed in order to double the concentration of the ammonia under the conditions given. (My assumption of 0.250 moles for nitrogen is based upon the balanced equation and the idea that it takes 0.050 moles of nitrogen to make 0.100 moles of ammonia.)
Someone else then posted this
And it takes 0.15 moles of H2 to make the extra 0.1 moles of ammonia. John [N.B. that's me, the ChemTeam!] got it right. The number you came up with is the change in H2 concentration observed after the extra hydrogen is added. However, 0.15 mol of the H2 added goes into making the extra 0.1 mol of ammonia. The final H2 concentration should be 0.4 + x - 0.15, or 0.25 + x.
Problem #3: Nitric oxide and bromine at initial pressures of 98.4 and 41.3 torr, respectively, were allowed to react at 300. K. At equilibrium the total pressure was 110.5 torr. The reaction is as follows.
2 NO(g) + Br2(g) ⇌ 2 NOBr(g)
Determine the Kp.
1) Set up an ICEbox:
PNO PBr2 PNOBr Initial 98.4 41.3 0 Change - 2x -x +2x Equilibrium 98.4 - 2x 41.3 - x +2x
2) Determine the value of x:
At equilibrium, the sum of the partial pressures is equal to 110.5 torr.
(98.4 - 2x) + (41.3 - x) + 2x = 110.5
x = 29.2 torr
3) At equilibrium, the partial pressures will be:
NO: 40.0 torr
Br2: 12.1 torr
NOBr: 58.4 torr
4) Determine the Kp:
Kp = (PNOBr)2 / [ (PNO)2 (PBr2) ]
Kp = (58.4)2 / [ (40.0)2 (12.1) ]
Kp = 0.176 torr¯1
5) Additional problem: What would Kp be if the pressures had been given in atmospheres?
0.176 torr¯1 x 760 torr atm¯1 = 137 atm¯1
6) Additional problem (much harder!): What would be the partial pressures of all the species if NO and Br2, both at an initial pressure of 0.300 atm, were allowed to come to equilibrium at this temperature? Sorry, no solution will be provided.
Problem #4: Consider the following equilibrium:
2 CH3OH(g) ⇌ CH3OCH3(g) + H2O(g)
If Kp is 13.54, what is the ratio between PCH3OH and PCH3OCH3?
1) Let the three partial pressures be:
PCH3OH = x
PCH3OCH3 = y
PH2O = y
The key assumption is that PCH3OCH3 = PH2O. This comes from assuming all the two products came from the methanol reacting. If we do not assume this, then we have no sure knowledge about the partial pressures of the two products and we cannot solve the problem.
2) Insert values into the Kp expression and solve:
Kp = [ (PCH3OCH3) (PH2O) ] / (PCH3OH)2
13.54 = [(y) (y)] / (x)2
3.68 = y/x
3) However, we wish the ratio PCH3OH : PCH3OCH3
x/y = 0.272
Problem #5: Consider the reaction:
H2 + I2 ⇌ 2HIwhose Keq = 54.8. If an equimolar mixture of the reactants gives the concentration of the product to be 0.500 M at equilibrium, determine the concentration of hydrogen.
1) The equilibrum expression is:
Keq = [HI]2 / ([H2] [I2])
2) Substituting into this, we have:
54.8 = (0.500)2 / [(x) (x)]
Comment: we know the equilibrium concentrations for H2 and I2 are equal because of the following two reasons: (1) they started out equimolar (that is, in equal amounts) and (2) they were used up in a 1:1 ratio (that is, an equal rate of consumption for both reactants) to make HI.
3) Seeing the the right-hand side is a perfect square, we take the square root and proceed to the answer:
7.403 = 0.500 / x
7.403x = 0.500
x = 0.0675 M (to three sig fig)
Problem #6: When NaF is added slowly to a solution that is 0.025 M Ba2+ and 0.025 M Ca2+ what will the concentration of calcium be when the barium just begins to precipitate? Ksp (BaF2) = 1.0 x 10¯7; Ksp (CaF2) = 1.7 x 10¯10.
1) What is [F¯] when BaF2 just begins to precipitate?
1.0 x 10¯7 = (0.025) (x)2
x = 0.0020 M
2) What is [Ca2+] when [F¯] = 0.0020 M?
1.7 x 10¯10 = (x) (0.0020)2
x = 4.25 x 10¯5 M
Problem #7: The solubility product constant for Cu(IO3)2 is 1.44 x 10-7. What volume of 0.0520 M S2O32- would be required to titrate a 20.00 mL sample of saturated solution of Cu(IO3)2
1) Determine moles of iodate in 20.00 mL:
1.44 x 10¯-7 = (x) (2x)2
x = 0.003302 M
[IO3¯] = 0.006604 M
0.006604 mol/L times 0.02000 L = 1.3208 x 10-4 mol
2) Determine iodate : thiosulfate ratio:
IO3¯(aq) + 6H+(aq) + 6S2O32-(aq) ----> I¯(aq) + 3S4O62-(aq) + 3H2O(aq)
The iodate : thiosulfate ratio is 1 : 6
3) Determine volume of thiosulfate required:
0.0520 mol/L = 1.3208 x 10-4 mol / x
x = 0.00254 L = 2.54 mL (if the ratio were 1:1)
since ratio is 1:6, we do this:
2.54 x 6 = 15.24 mL (this is the answer)
4) An alternate approach (from Yahoo Answers):
1 mole of IO3¯ reacts with 6 moles of S2O32-
Assume 1.00 L of solution present. Therefore:
moles of S2O32- = 6.604 x 10-4 mol x 6 = 0.0396 mol
number of moles = molarity x volume
volume = 0.0396 / 0.0520 = 0.762 L = 762 mL
However, we only titrated 20.00 mL, so:
762 x 0.02 = 15.24 mL
Problem #8: Bromine chloride, BrCl, a reddish covalent gas with properties similar to those of Cl2, may eventually replace Cl2 as a water disinfectant. One (1.00) mole of chlorine and one (1.00) mole of bromine are enclosed in a 8.05 L flask and allowed to reach equilibrium at a certain temperature.
Cl2(g) + Br2(g) ⇌ 2 BrCl(g)
Kc = 11.57 x 10¯2 at the given temperature. What mass of Cl2 is present at equilibrium?
1) Determine molarities of Cl2 and Br2:
[Cl2] = [Br2] = 1.00 / 8.05 = 0.1242236 M (I'll carry several guard digits)
2) Set up an ICEbox:
[Cl2] [Br2] [BrCl] Initial 0.1242236 0.1242236 0 Change - x -x +2x Equilibrium 0.1242236 - x 0.1242236 - x +2x
3) Determine x:
Kc = [BrCl]2 / ([Cl2] [Br2]
11.57 x 10¯2 = (2x)2 / [(0.1242236 - x) (0.1242236 - x)]
The right-hand side is a perfect square.
0.340147 = 2x / (0.1242236 - x)
0.0422543 - 0.340147x = 2x
0.0422543 = 2.340147x
x = 0.01805626 M
4) Determine [Cl2] at equilibrum:
0.1242236 minus 0.01805626 = 0.10616734 M
5) Determine mass of Cl2 present at equilibrium:
MV = g / molar mass
(0.10616734 mol/L) (8.05 L) = x / 70.906 g/mol
x = 60.6 g (to three sig figs)
6) We can determine if this is the correct answer by attempting to recover the Kc:
Kc = (0.03611252)2 / [(0.10617) (0.10617)
The computation is left to the reader.
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