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Unformatted text preview: Worksheet 18  Equilibrium Balance the following reaction, and use it to answer the following 8 questions:
__ N2 (9) + .52; H2 (9) a A NH3 (9) 1. Starting with 0.500 M N; and 0.800 M H2, the reaction is allowed to
proceed until it reaches equilibrium. In the absence of products, the value
of Q, the reaction quotient, is zero, less than any possible value of K. The [NH,~,],,q 2 0.150 M at 250°C. Determine K for this reaction at this
temperature. a) Start by ﬁlling in the lCE table with the Initial concentrations.
[N2] (M) lei (M) [NHal (M)
Initial 0 500 ﬂ .5’00 0 ._.._.........._._ Change “76 *3X tﬂx Equilibrium r 550 ’ 35 x 5’00 ' 3X a? )L
.42f . 675' . /.S‘0
b) Next, ﬁll in the Changes in the system, in terms of x, which wiil determined by the stoichiometry. Reactants are used
up () and products are formed (+). 0) Next, write equilibrium expressions, in terms of initial
concentrations and the changes that occur (x). d) Solve for x from the information you have. x=r0?5’ [MHJJCJrrl‘Sthi 3C: rot15M
e) Write the equation for K, the equilibrium constant, in terms
of the concentrations of reactants and products. DU”)? [thé’l'ldji 541570375 f) Finally, determine the numerical value of the equilibrium
constant, K at this temperature. K= ,JH’ K: 3 2. Convert K. based on concentration. to KP, based on partial pressures.
The term An means the number of moles of gas in the products minus the number of moles of gas in the reactants. If An = 0. K = Kp. An= g~(3+/):~; l
KP=K(RT)A“= K1,: (,zyg)[‘0320()(250f3?3ﬂ : 1.57 X/O'f 3. Calculate the partial pressures of N2, H2 and NH3 assuming a 10.0 L
volume, at 250°C. a) First calculate the moles of each gas present at equilibrium. moIN2= .425‘_/}2€X/0(0L= 4%‘227/ Ala.  //( 4 A7
moIH2= .§}y;Jg/dtd;z §.}r/nj#p m/cq
molNH3= ./so m; X “3.01., AW mi Nil3 “23"”
b) Now ﬁnd the total pressure, using the ideal gas law. 12. 3am mm): awry/anew me)
P: c) Find the partial pressures of each of the gases using their mol fractions: PN2 = (mol Nglmoitotal)met= 1f. 3 M6.” PH2 = (mol Hzlmol total) x Ptot : 94,; Mam gt ash47v» me : (mol NH3/mol total) x Plot = 4. Calculate Kp using these values. Is it the same as your previous answer,
calculated from K in question 2? “1&5; : (‘5 ‘44) L .. Mg. 2r ﬂ"
[IN 214,6); Sﬂﬂu ﬁlm/67¢
if} MJWMMS / ,. Now, imagine that the volume of the container increases from 10.0 L to
50.0 L. Initially, there will be the same number of moles of each of the
gases that were present at equilibrium. However, they are now in a larger
volume, so their concentrations will change. a) Determine the new concentrations of N2, H2 and NH3. if 1
[N2]: m 5.0i6 L
{H2}: gl/SM 5194:! L
[NH3]= ,05’1'0/‘1 l’m’w’
571.0 L b) Using these new initial concentrations, determine the value of Q,
the reaction quotient g.
Q = product ip = [I a 5) : (€296,
[reactant]: a (I) (i [I 5% 3 0) Compare Q to K. Which way will the reaction proceed, under these new experimental conditions? él§g> I z )1 J7 , .
62 >K W AM 53M“ d) Compare the number of moles of gas on the reactant and product
sides of the equation. nreactant : nproduct : l lncreasin the volume of the system J'fts the equilibrium toward
theEac‘ﬁPl products, which has 1w: smaller number of moles 0 gas. 6. Start the reaction with 1.00 M N2, 1.00 M H; and 1.00 M NHa at the same
temperature, 250° C What is the value of 0 under these initial conditions? b
(I? 6'0) ; /, 0’0 Compare Q and K. Which way will this reaction proceed? /I 07) > .« Z 77 5/ Q > K
Which compound will decrease in concentration? W decﬂ KJHJ’ Complete the ICE table for this set of initial conditions. [N2] (M) [H2] (M) [NHsi (M) Initial 1.00 1.00 1.00
Change ﬂm + g}  a 56 Equilibrium )0?) ﬁt LOW}? A6002 Write the expression for K.
{MD a 2le W
{Mar 11’ ){x'emf‘ 5‘3”); K: 7. The equilibrium constant of this reaction changes with temperature.
At T = 500°C, KP = 3.6 x 10'2 and at 400°C, Kp = 41. 3) Raising the temperature increases dﬁecrease? Kp. b) Raising the temperature drives the “Median. c) Shouid heat be treated as @r’ a reactant? d) Is this an endothermic eaction? ...
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This note was uploaded on 12/05/2011 for the course CHEM 231 taught by Professor Beck during the Spring '11 term at Indiana.
 Spring '11
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