Energy changes
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Hi,
I have been going through Nuffield Topic 5 and have a question on calculating q for a solution (page 101). Can you tell me if my thinking to follow is correct: "In an endothermic reaction, the change in temperature value ('delta' T) will be negative which means q will be negative. When you use q to calculate energy per mol ('delta' H) for an endothermic reaction, it has a positive sign."
I get really confused between q of the surroundings and q of the reaction so if you have any hints on how to remember the difference, it would be really helpful.
Many thanks
Michele
Corrie writes ....
You are not alone in having problems with signs in this sort of calculation. Signs are often best left till the end, using the rise or fall in temperature in the experiment to add a sign to your answer, and then call it 'delta'H. Regard 'q' as just the amount of energy involved in your experiment, either lost or gained by the system, so it does not need a sign.
First work out the amount of energy, q (in J), involved in your experiment, using the 'm x C x temperature change' formula - without any signs.
Then scale this amount of energy up to 1 mole of the reactant or product you are interested in, e.g. 1 mole of copper sulphate, and convert it to kJ - still no sign!
Now, finally, turn your answer into a 'delta'H value by adding a sign, using the temperature change to tell you whether it was an exo- or endothermic reaction. If the temperature went up, then it was an exothermic reaction, and the sign of 'delta'H should be -, and vice versa if the temperature went down.
Hope this helps.
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updated: 20 December 2006
