Calculating deltaH°f for O(g) - 142.0 kJ/mol

[NickP717]

Homework Statement

The bond energy of O2(g) is 142.0 kJ/mol. Calculate deltaH°f for O(g).

Homework Equations

The Attempt at a Solution

no clue

 

[lucky_star]

Hi Nick,
You have an Oxygen molecule and it asks you for the delta Hf of Oxygen atom ONLY. So, i would suggest you to halve the bond energy of O2.

Any other suggestions or am I right?

 

[Blueshift5]

I would like to provide a response to this statement by explaining the concept of standard enthalpy of formation (deltaH°f). The standard enthalpy of formation is the change in enthalpy when one mole of a compound is formed from its constituent elements in their standard states. In this case, we are looking to calculate the deltaH°f for O(g), which is the standard enthalpy of formation for gaseous oxygen.

To calculate this value, we need to consider the reaction in which oxygen gas is formed from its constituent elements, which are O2 and O. The balanced equation for this reaction is:

1/2 O2(g) + 1/2 O2(g) --> O(g)

We can see that the reaction involves the formation of one mole of gaseous oxygen from half a mole of O2 and half a mole of O. Therefore, the deltaH°f for O(g) can be calculated as:

deltaH°f = (1/2*142.0 kJ/mol) + (1/2*0 kJ/mol) - (1/2*142.0 kJ/mol) = 71.0 kJ/mol

This means that the standard enthalpy of formation for gaseous oxygen is 71.0 kJ/mol. This value is half of the bond energy of O2(g) because the reaction only involves the formation of one mole of oxygen, while the bond energy of O2(g) represents the energy required to break two moles of O-O bonds.

In summary, by considering the balanced reaction and using the concept of standard enthalpy of formation, we can calculate the deltaH°f for O(g) as 71.0 kJ/mol. It is important to note that this value may vary slightly depending on the source of the bond energy for O2(g), but the general concept remains the same.