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3 years ago

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Understanding the high-temperature behavior of nitrogen oxides is essential for controlling pollution generated in automobile en

gines. The decomposition of nitric oxide (no) to n2 and o2 is second order with a rate constant of 0.0796 m−1⋅s−1 at 737∘c and 0.0815 m−1⋅s−1 at 947∘c. You may want to reference (page) section 14.5 while completing this problem. Part a calculate the activation energy for the reaction. Express the activation energy in kilojoules per mole to three significant digits.

1 answer:

disa [49]3 years ago

6 0

Hey there!

Arrhenius equation :

k = A exp ( -Ea/ RT )

ln ( K2 / K1 ) = ( -Ea/ R ) * ( 1 / T2 - 1 / T1 )

K1 = 0.0796 m⁻¹s⁻¹

K2 = 0.0815 m⁻¹s⁻¹

T1 = 737 ºC = 737 + 273.15 => 1010.15 K

T2 = 947ºC = 947 + 273.15 => 1220.15 K

R = 8.314 J / mol*K

ln ( 0.0815 / 0.0796 ) = ( -Ea / 8.314 ) * ( 1 / 1220.15 - 1 / 1010.15 )

Activation energy Ea = 1151 J / mol ≈ 1.15 Kj/mol

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$4CN^ - \ \ \ + \ \ \ NiS \ \ \ \ \ \ \ \underset{}{\stackrel{}{\rightleftharpoons}} \ \ \ \ \ Ni(CN)^{2-}_4 \ \ \ \ \ \ \ \ \ + \ \ \ \ \ \ \ \ \ S^{2-}$

initial [ I] 0.091 0 0

Change [C] -4x +x + x

Equilibrium [E ] 0.091 - 4x x x

Here is x is the amount in term of concentration that is lost by $CN^-$ and gained by $Ni(CN)_4 ^{2-}$ and $S^{2-}$

Gnerally the equilibrium constant for this reaction is mathematically represented as

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