Temperature resulting from the formation of water

[vjk2]

Basically, I need confirmation on the heat produced when four mole of H2O is produced from Hydrogen and oxygen gas.

I'm getting something ridiculous, basically saying that the temperature goes up to ~4000 degrees C.

What I'm doing is using Hess's law to find the energies of reaction. H2 and O2 gas result in 0 KJ/mol. H2O is -285.8 KJ/mol.

4 mols x 285.8 KJ/mol = 1143.2 KJ

I'm taking this and plugging it into the equation

q=mC<>T (change in Temperature)

so,

1143.2 = (18 g/mol H2O * 4 mol H2O = 72 g) * 4.18 j/(g*c) * <>T

<>T = 3797 degrees C

I know that when water forms from oxygen and hydrogen, the result is explosive, but thousands degrees C seems way too high. However...my logic seems right. Thoughts?

[windwitch]

Do not underestimate the amount of energy available in the bonds of water. If you conduct an experiment, you will be able to see that water is able to absorb and produce a large amount of heat in comparison to many other substances.

From your equations alone, it seems that there are no mistakes. HOWEVER, from your equation, it seems to assume that water rises 4000 degrees in temperature without undergoing a change in liquid to gaseous state. Remember that when you are finding a standard enthalpy of formation, you assume that the product is in it's standard state. The standard state for water is liquid. Does it make sense that water, starting from (0 or 25 degrees celsius, I am forgetful of the standard temperature) standard temperature to go to 4000 degrees without having any energy lost in a state change?

[vjk2]

Yes, I've deduced that the phase change does take place.

The issue then is how to calculate it in. What I'm thinking is to

1. take the total energy released by the formation of water.

2. subtract from that figure the amount of energy required to heat water to 100 degrees C by using the equation q = (18*4)(4.18)(75)

3. take the remainder of energy and instead of using the 4.18 heat capacity figure for liquid water, use the heat capacity for vapor water.

Only problem is finding the heat capacity of vapor water.

[Dadface]

Yes, I've deduced that the phase change does take place.

The issue then is how to calculate it in. What I'm thinking is to

1. take the total energy released by the formation of water.

2. subtract from that figure the amount of energy required to heat water to 100 degrees C by using the equation q = (18*4)(4.18)(75)

3. take the remainder of energy and instead of using the 4.18 heat capacity figure for liquid water, use the heat capacity for vapor water.

Only problem is finding the heat capacity of vapor water.

Dont forget the specific latent heat of vapourisation that is the energy needed to change unit mass of water to unit mass of steam at its boiling point.

[Borek]

Also to be precise in this type of calculations you should take into account fact that specific heat is not temperature independent. As long as delta T is in the range of several or even small tens of degress that's usually not a large problem, but when we are talking about temperature changes in the range hundreds or thousands degress, that has to be taken into consideration.

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