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(a) The decomposition of N2O5(g) is a first order reaction with a rate constant of 5 x 10–4 sec–1 at 45°C.
i.e., 2N2O5(g) = 4NO2(g) + O(g)
If initial concentration of N2O5 is 0.25, calculate its concentration after two minutes. Also calculate half life for the decomposition of N2O5(g).
(b) For an elementary reaction: 2A + B → 3C The rate of appearance of C at time ‘t’ is 1.3 x 10–4 mol l–1 s–1. Calculate at this time:
(i) Rate of reaction (ii) Rate of disappearance of A.


Rate constant K = 5 x 10–4 sec. Initial concentration [A]0 = 0.25 M Final concentration [A]t =? Time taken by the reaction, t = 2 min.
For a first order reaction, rate constant (K) is given by
                        K = 2.303tlog A0At


5×10-4 = 2.303tlog0.25At5×10-4×22.303 = log 0.25 - logAt           0.0004  = log 0.25 - logAt           0.0004 = 0.3979 - log At0.0004 - 0.3979 = -logAt              -0.3975 = -logAt                         At = antilog (0.3975)                          At = 2.6


(b) 2A+B   C

(i) The rate of appearance of C at time t

                              = 1.3 × 10-4 mol L-1 s-1

-ddt[C] = 12ddtA = +12ddtB

(ii) Rate of disappearance of A

                  = 2 × 1.3 × 10-4 mol L-1 S-1=2.6 × 10-4 molL-1 S-1.

                    




































































































































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A first order reaction is 50%. Complete in 30 minutes at 27° and in 10 minutes at 47°C. Calculate the reaction rate constant at 27°C and the energy of activation of the reaction in kJ/mol.

by using half life equation, we get

K = 0.693t1/2K = 0.69330 min = 0.0231 at 27°C or 300 KK = 0.69310 min = 0.0693 at 47°C or 320 K

From the Arrhenius equation

logk2k1 = Ea2.3031T1-T2    Ea = 2.303×R×T1×T2T2-T1log k2k1= 2.303 × 8.314 J K-1 mol-1×300×320 K320 K - 300 K              log 0.06935-10.02315-1 = 43.848 kJ mol-1.
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The rate constant increases 50% when temperature is increased from 298 K to 308 K. THe value of ΔH0 is 15 kJ mol. Calculate the activation energies Eaf and Ear.

The rate constant at 298 K = k.
When the temperature is raised from 298 k to 308 k.
Increase in rate constant

                      = k×50100 = 0.5 k
By using the equation

        Ink2/k1 = [Eaf/R] 1T1-1T2

We have,

     In(1.5 k/k) = [Eaf/R] 1298k-1308k                     = [Eaf/R] 10298 × 208log (2.303 × 1.5) = Eaf8.314×10298×308log 3.4545 = Eaf8.314×10298×3080.4460 = Eaf8.314×10298×308Eaf = 0.4460×8.314×298×30810     =       34.03 kJ/mol

                rH° = 15 kJ mol-1

but            rH° = Eaf-Ear

              Eaf = (34.03-15) kJ mol       = 19.03 kJ mol-1

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The rate law for the gas phase reaction of chlorform with chlorine.
CHCl3(g) + Cl2(g) → CCl 4(g) + HCl(g)
is given by rate = k[CHCl3] [Cl2]1/2. How would the rate of reaction vary when (a) the concentration of CHCl3 is doubled (b) the concentration of Cl2 is doubled. What is the effect of each of these two changes on the rate constant.



The rate law for the gas phase reaction of chlorform with chlorine.
CHCl3(g) + Cl2(g) → CCl 4(g) + HCl(g)
 the rate is given by 

 rate1 =k[CHCl3] [Cl2]1/2 .....1

if the concentration of CHCl3 is doubled
then 

rate2 =k[CHCl3]2 [Cl2]1/2 .....2
divide 1by 2 we get

rate1 =k[CHCl3] [Cl2]1/2 .rate2 =k[CHCl3]2 [Cl2]1/2  rate become doublerate1rate2= 122rate1= rate2thus rate 1 become double

if the concentration of Cl2  is doubled
then .
rate3 =k[CHCl3] [Cl2]2/2........3

divide 1 by 3 
we get 

rate1 =k[CHCl3] [Cl2]1/2 .rate3 =k[CHCl3] [Cl2] rate become doublerate1rate3= 122rate1= rate2thus rate 1 become double

thus the rate reaction is double in both case.


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The energy of activation for a reaction is 100 kJ mol–1. The presence of a catalyst lowers the enregy of activation by 75%. What will be the effect on rate of reaction at 20% C, other things being equal.

According to the Arrhenius equation 

K = Ae–Ea/R
In absence of catalyst k1 = Ae–100/RT
In presence of catalyst k2 = Ae–25/RT

           k1k2 = e-100/RTe-25/RT=e-75/RT 

or       2.303 log10k1k2 = (75/RT)

or         2.303 log10 k1k2 = 758.314×10-3×298

                  k1k2 = 2.35 × 1013


Since Rate k[ ]
n at any temperature for a reaction n and concentration of reactants are same and temperature changes.

r2r1 = k2k1 = 2.35 × 1013





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