Question

NUMBER 3 AND 4 PLEASE! I JUST NEED #3 AND 4

Ozone Formation: NO2(g), is formed by chemical reactions involving N2(g) and O2(g) at the high temperatures Inside internal combustion engines in our cars. In the presence of sunlight, NO2(g) reacts with O2(g) to generate Os(s) as described by the following overall reaction: NO2(g) + O2(g) = NO(g) + O2(g) E = 306.5 l/mol This chemical process occurs in two steps, as illustrated in the potential energy diagram. In the second step of the process, the radical species o formed in the first step reacts with O2. =107 kJ/mol O + NO NO+O, 1. Write the chemical equations for each reaction step and determine the molecularity of each of them. Identify all reaction intermediates. NO, Reaction Path 2. Express the rate law for each reaction step and derive the rate law for the overall process. 3. Estimate how many times faster this reaction is at 30 °C versus 10 °C. 4. Based on the information provided, evaluate whether this process can be expected to occur to a large extent at 25 °C (Hint: Express Ko for the process and calculate its value) Substance srº(J/mol K) O2(g) 161.1 03(8) 238.9 NO(g) 211.2 NO2(g) 240.1

Answer 1

3. Ea = 306.5 kJ/mol

Thus, According to Arrhenius Equation,

ln(k2/k1) = (306500/8.314)*(1/283 - 1/303)

So, k2/k1 = 5423.378

Hence, the Reaction will be 5423.378 times Faster, at 30°C.

4. For the Reaction, ∆H = 306.5 -107 = 199.5 kJ/mol

∆S = 48.9 J/mol-K

Hence, ∆G = 199500 - (298*48.9) = +184.928 kJ/mol

Thus, the Process isn't expected to occur to a large extent, at 25°C, Since ∆G for this process is positive.