Naked Science Forum
Non Life Sciences => Chemistry => Topic started by: Kryptid on 03/06/2011 20:44:11
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I'm aware of the principle behind chemosynthesis; reacting certain chemicals together can result in an exothermic reaction which also generates carbohydrates as a by-product. That's why they don't need energy from the Sun in order to produce sugars.
There are different kinds of reactions used in chemosynthesis. However, I'm having difficulty understanding one such equation in particular:
6CO2 (Carbon Dioxide) + 6H2O (Water) + 3H2S (Hydrogen Sulfide) → C6H12O6 (Glucose) + 3H2SO4 (Sulfuric Acid)
Here's what is confusing me; I calculated the energy released by the reaction based on heat of formation data, and the result I got was it is endothermic. If my calculations were right, this reaction should consume energy, not liberate it.
Here are the enthalpies of formation for the selected chemicals (in kilojoules per mole):
Carbon Dioxide (g) = -393.509
Water (l) = -285.83
Hydrogen Sulfide (g) = - 20.63
Glucose (s) = -1,271
Sulfuric Acid (l) = -814
And changing the equation to include these numbers:
6(-393.509 kJ/mol) + 6(-285.83 kJ/mol) + 3(-20.63 kJ/mol) → 1(-1,271 kJ/mol) + 3(-814 kJ/mol)
= (-2,361.054 kJ/mol) + (-1,714.98 kJ/mol) + (-61.89 kJ/mol) → (-1,271 kJ/mol) + (-2,442)
= (-4,137.924 kJ/mol) → (3,713 kJ/mol)
= +424.924 kJ/mol
I did notice one possible mistake I made; I used the heats of formation of the chemicals in their standard states (glucose as a solid, and carbon dioxide and hydrogen sulfide as gases). In reality, all of the chemicals would be in water solution. Would that make enough of a difference for the reaction to be exothermic? Should I collect enthalpy of solution data and do a recalculation?
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Whether a reaction happens spontaneously (all by itself)is a function of its change in Gibbs' free energy, not enthalpy. While enthalpy is a part of Gibbs' free energy, so is the change in entropy. There is another consideration on spontaneity that is not addressed when calculating energy from standard states: the concentration ("activity") of each of the species involved in the reaction. This is a sort of hidden aspect to "standard state or standard conditions". This would mean that for your calculation to represent the actual value found for energy changes in the real system, a cell wold have to be in a gas whose pressure of CO2 gas and H2S gas would each need to be about 1 atmosphere. I hope this helps.
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Interesting. I've heard about Gibb's Free Energy, but I don't know much about it. I suppose you do have a point, since supposedly a mixture of salt and ice will spontaneously drop in temperature (which would be a positive change in enthalpy). Is Gibb's Free Energy responsible for that too?
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The Gibbs Free Energy includes both the enthalpy (or "heat of reaction") and an entropy-based term:
ΔG = ΔH - T.ΔS where the symbols are G -- gibbs free energy; H -- enthalpy; T -- kelvin temperature; S -- entropy, and Δ signifies "change in..."
For a reaction to be spontaneous ΔG would need to be negative. The substances not being in "standard state" is an issue, but in this case it would make only a minute difference.
For the reaction you cite, it should come as no surprise that an input of energy is required. As you correctly point out the reaction is endothermic -- absorbs heat energy.
The entropy term is also unfavourable for spontaneous reaction because you have two gas reactants (very disordered, high entropy) on the reactant side, and only condensed phases (much more ordered with lower entropy) on the product side.
The bottom line is that you cannot get from low energy substances (water, carbon dioxide, hydrogen sulfide) to high energy substances (glucose) without an input of energy from somewhere. If not the sun, then ...??
Unless we are talking about synthesizing glucose from high energy substances, or a hidden source of energy input, this type of chemosynthesis is science fiction.
(An afterthought ... atmospheric oxygen is a high energy substance. If this reaction were coupled with conversion of large amounts of H2S to H2SO4 via reaction with atmospheric oxygen, that could possibly provide the required energy)
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The bottom line is that you cannot get from low energy substances (water, carbon dioxide, hydrogen sulfide) to high energy substances (glucose) without an input of energy from somewhere. If not the sun, then ...??
Is the thermal energy released by oceanic thermal vents sufficient, or do the animals that exist there simply tap into the chemicals released by the vent?
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OK Geezer, in your latest post you have provided a context that I had not understood.
In the ocean vents, there would be three possible sources of energy to drive the sort of reaction you describe
(1) Steep thermal gradients -- large temperature variations over a small distance or a short time
(2) Steep concentration gradients -- streams of high or low mineral content which could be exploited in a similar way.
(3) Probably traces of high energy chemicals in the mix issuing from the vent. Or water incompatible ones.
The reaction between sulfuric acid and water releases a lot of energy for example.
Microorganisms could possibly develop that could exploit any or all of these.
Atmospheric oxygen would certainly not be a participant at those levels.
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Thank you Damocles!