1. Calculate the delta H for H2(g)+O2(g)=H2O2(g)

Bond energies, kJ x mol^-1
H-H 432
H-O 459
O-O 207
O=O 494

I did:
deltaH=(2(207kJ/mol))-(432kJ/mol+494kJ/mol)
=-512kJ
is that correct?

2. The theoretical effect of an increase in T can be explained in terms of the collision because it affects

a.The fraction of collisions that are effective and the required activation energy for a reaction.

b.The total number of collisions that occur and the required activation energy for a reaction.

I stumped whether the answer is a or b

1. I didn't calculate it but dHrxn = dHreactants + dHproducts.

2. It depends upon the fraction of the molecules that have the required activation energy as well as that constant "A" in the equation so I would go with a.

I thought the equtaion was

dH=[sum of dH of products]-[sum of dH of reactants]?

Your equation is correct if you are dealing with delta Ho BUT you have bond energies. Delta H from bond energies are calculated a different way.

1. To calculate the delta H for the given reaction, you need to use the bond energies of the bonds broken and formed in the reaction. The reaction can be represented as:

H2(g) + O2(g) → H2O2(g)

The delta H (enthalpy change) is equal to the sum of the bond energies of the bonds broken minus the sum of the bond energies of the bonds formed.

Using the bond energies provided:
Bond energy of H-H is 432 kJ/mol,
Bond energy of H-O is 459 kJ/mol,
Bond energy of O-O is 207 kJ/mol,
Bond energy of O=O is 494 kJ/mol.

In the reaction, one molecule of H-H bond and one molecule of O=O bond are broken, and one molecule of H-O bond is formed.
So, the calculation would be:

ΔH = (2 * bond energy of O-O) - (bond energy of H-H + bond energy of O=O)
= (2 * 207 kJ/mol) - (432 kJ/mol + 494 kJ/mol)
= 414 kJ/mol - 926 kJ/mol
= -512 kJ/mol

Your calculation of -512 kJ/mol for the delta H is correct.

2. The theoretical effect of an increase in temperature (T) can be explained in terms of the collision theory of chemical reactions.
According to the collision theory, for a reaction to occur, reactant molecules must collide with each other with sufficient energy (activation energy) and in the correct orientation. Increasing the temperature increases the average kinetic energy of molecules, making them move faster. This leads to several possibilities:

a. The fraction of collisions that are effective: An increase in temperature increases the kinetic energy of molecules, allowing them to overcome the activation energy barrier more frequently. This means that a larger fraction of collisions between reactant molecules will have sufficient energy to lead to a reaction.

b. The total number of collisions that occur: Increasing the temperature also increases the speed and frequency of molecular collisions. As a result, more collisions will occur per unit time, leading to an increase in the total number of collisions.

c. The required activation energy for a reaction: Although the temperature increase affects the fraction of effective collisions, it does not directly change the required activation energy for a reaction. The activation energy is a characteristic property of a specific reaction and largely depends on the nature and structure of the reactants and products involved.

Therefore, the answer to the question is a. An increase in temperature affects the fraction of collisions that are effective and the required activation energy for a reaction.