Question

Cander The reaction 2Hy(6) Op19))

The following bond energes to be determined experimentally:
H-H 436 kJ/mol 0=0 497 kJ/mol H-O 464 kJ/mol
A. Suppose we start with 4 mal of hydrogen. How many moles of oxygen are needed to use up all of the hydrogen? How many moles of water molecules are produced?
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.)
C. Does your answer to part 8 represent a net out of energy, ore net release of energy if the reaction proceeds in the direction indicated? Does the chemical energy increase, or decrease?

Answer 1

The reaction requires 2 moles of oxygen to use up all 4 moles of hydrogen. 2 moles of water molecules are produced. The difference in chemical energy between the reactants and the products is 468.5 kJ/mol, representing a net release of energy.

Step-by-step explanation:

To determine the number of moles of oxygen needed to use up all of the 4 moles of hydrogen, we can use the balanced chemical equation for the reaction:

2H₂ + O₂ → 2H₂O

From the equation, we see that 2 moles of hydrogen react with 1 mole of oxygen to produce 2 moles of water. Therefore, the number of moles of oxygen needed is 4/2 = 2 moles. Similarly, the number of moles of water molecules produced is also 2 moles.

To calculate the difference in chemical energy between the reactants and the products, we need to subtract the sum of the energies required to break the bonds on the reactants side from the sum of the energies released to form the bonds on the products side. From the given information, the energy required to break the bonds is 1840 kJ/mol, while the energy released to form the bonds is 872.8 kJ/mol + 498.7 kJ/mol = 1371.5 kJ/mol. Therefore, the difference in chemical energy is 1840 kJ/mol - 1371.5 kJ/mol = 468.5 kJ/mol. Since the question asks for the absolute value, the difference in chemical energy is 468.5 kJ/mol.

The answer to part B indicates that the reaction releases more energy than it consumes, as the difference in chemical energy is positive. This means that the reaction is exothermic, and the chemical energy decreases during the reaction.

Answer 2

In this reaction, 2 moles of O₂ are needed to use up 4 moles of H₂ and 4 moles of H₂O molecules are produced. The difference in chemical energy between the reactants and the products is 470 kJ/mol. The reaction proceeds in the indicated direction and releases energy, causing a decrease in chemical energy.

Step-by-step explanation:

A. To use up all 4 moles of hydrogen, we need an equal number of moles of oxygen. Since the balanced equation for the reaction is:
2H₂ + O₂ → 2H₂O, we can see that for every 2 moles of H₂, we need 1 mole of O₂. Therefore, with 4 moles of H₂, we would need 2 moles of O₂.
To determine the number of moles of water molecules produced, we can use the balanced equation. For every 2 moles of H₂, we produce 2 moles of H₂O. So with 4 moles of H₂, we would produce 4 moles of H₂O.

B. In part A, we determined that we need 2 moles of O₂ to use up 4 moles of H₂. The difference in chemical energy between the reactants and the products can be calculated by subtracting the sum of the energies required to break the bonds on the reactants side from the sum of the energies released to form the bonds on the products side. The sum of the energies required to break the bonds on the reactant side is 1840 kJ/mol, while the sum of the energies released to form the bonds on the product side is 1370 kJ/mol. Therefore, the difference in chemical energy is 1840 kJ/mol - 1370 kJ/mol = 470 kJ/mol.

C. The positive value of the difference in chemical energy (470 kJ/mol) indicates a net release of energy in the reaction. This means that the reaction proceeds in the direction indicated and releases energy. The chemical energy decreases in the reaction since the energy is released.