Question

Consider the reaction 2H₂(g) + O₂(g) → 2H₂O(g).

The following bond energies have been determined experimentally:

• H-H 436 kJ/mol
• O=O 497 kJ/mol
• H-O 464 kJ/mol

. Suppose we start with 4 mol of hydrogen. How many moles of oxygen are needed to "use up" all of the hydrogen? How many moles of water molecules are produced?

n of O₂ = 2 (Correct)

n of H₂O = 4 (Correct)



B. For the number of moles in part A, what is the difference in chemical energy between the reactants and the products for this reaction? (Give the absolute value only.)

|ΔE_chem| = 974 kJ (This is Correct)


C. Does your answer to part B represent a net input of energy, or a net release of energy if the reaction proceeds in the direction indicated? Does the chemical energy increase, or decrease?

There is a net   release  of energy, and the chemical energy   decreases (These are correct as well)

F. If the system was at 300 K before the reaction took place, what is the final temperature after the reaction?

(Part F is the only part I need)

Answer 1