Write the rate equation for the relative disappearance of dinitrogen pentoxide. Write the rate equations for the relative appearance of nitrogen dioxide and oxygen gas. When writing the equations, remember to think about how many NO2 and O2 molecules are formed from 2 N2O5.

Question

Write the rate equation for the relative disappearance of dinitrogen pentoxide. Write the rate equations for the relative appearance of nitrogen dioxide and oxygen gas. When writing the equations, remember to think about how many NO2 and O2 molecules are formed from 2 N2O5.

Answer 1

To write the rate equation for the relative disappearance of dinitrogen pentoxide (N2O5), we need to consider the stoichiometry of the reaction. The balanced chemical equation for the reaction can be written as:

2 N2O5 → 4 NO2 + O2

In this equation, we can see that two molecules of N2O5 react to form four molecules of nitrogen dioxide (NO2) and one molecule of oxygen gas (O2). The rate equation for the disappearance of N2O5 can be written as:

Rate of disappearance of N2O5 = k[N2O5]

Here, "k" is the rate constant, and [N2O5] represents the concentration of N2O5.

Next, to write the rate equation for the relative appearance of nitrogen dioxide (NO2), we need to consider the stoichiometry once again. From the balanced chemical equation, we can see that four molecules of NO2 are formed from two molecules of N2O5. Therefore, the rate equation for the appearance of NO2 can be written as:

Rate of appearance of NO2 = 2 × k[N2O5]

Here, the factor of 2 accounts for the stoichiometry, indicating that for every two molecules of N2O5 consumed, four molecules of NO2 are formed. The rate constant (k) remains the same as in the disappearance of N2O5 equation.

Finally, for the relative appearance of oxygen gas (O2), we can see that one molecule of O2 is formed from two molecules of N2O5. Therefore, the rate equation for the appearance of O2 can be written as:

Rate of appearance of O2 = 1/2 × k[N2O5]

Here, the 1/2 factor accounts for the stoichiometry, indicating that for every two molecules of N2O5 consumed, one molecule of O2 is formed. The rate constant (k) remains the same as in the previous equations.

It's important to note that the rate constants (k) may have different values for each equation, as they depend on factors such as temperature, catalysts, and reactant concentrations.