[jamesrb]

I recently performed a bomb calorimetry experiment using a Parr 6200 bomb calorimeter in an undergraduate P. chem lab and am having some trouble analyzing the data. The lab manual asks to calculate the heat capacity (also called the energy equivalent) for the calorimeter in J/°C using Q_tot/ΔT. The specific heat of combustion of benzoic acid is given to  26.434 kJ/g. So if Q_tot is equal to Q_mm where Q_m is the heat release per gram of benzoic acid (26434 J/g) and m is the mass of benzoic acid (0.994 g), for my experiment I get:

Q_tot = Q_mm = (26434 J/g) × (0.994 g) 26275.396 J

Using C = Q_tot/ΔT:

C= (26275.396 J)/(2.6576 °C) = 9886.8899 J/°C

Where I get confused is because I have "EE" recorded in my notes as 2379.22 cal/K. This experiment was performed several weeks ago and I vaguely remembering someone telling me that this is the energy equivalence (and thus heat capacity) of the calorimeter. Converting that to calories/K I get:

(9886.8899 J/°C)(0.239 cal/1J)(1°C/274.15K)= 10.314 cal/K

Which is no where near 2379.22. I am wondering what the EE value I recorded means. Is that really the heat capacity of the bomb calorimeter? I am leaning towards the first calculated value above to be the true heat capacity. If anyone familiar with bomb calorimeters cares to comment I would appreciate it.

Moving on I need to calculated the heat of combustion for the naphthalene and sucrose samples in kJ/mol. From my understanding we are supposed to use ΔH = ΔU + Δn_gasRT. I am however unclear what value to use for T. Is that supposed to be my temperature change recorded from the calorimeter the standard temperature of 298.15 K? I also eventually have to apply several correction factors to convert this value to standard conditions. I am unsure what the purpose of recording ΔT for each combustion was. Any insight anyone could provide would greatly be appreciated.

[Corribus]

Just for future reference, it's always helpful to know what exactly you're trying to do in your lab. That said, it seems clear a major part of your problem is this:

(9886.8899 J/°C)(0.239 cal/1J)(1°C/274.15K)

That's not an appropriate conversion factor for °C to K. These units are pretty much converted 1:1 (a 1° increase in temperature is the same as a 1 K increase in temperature).  9887 J is approximately 2361 calories. So it doesn't look like you're actually that far off.

For the second part, the change in temperature (usually of a water reservoir) is measured because this is what allows you to determine the amount of heat liberated by the reaction being combusted.

http://en.wikipedia.org/wiki/Calorimeter#Bomb_calorimeters

Oh, you probably realize this now but I'll say it anyway: it helps not to wait weeks before you write up your labs.

[jamesrb]

So using the heat capacity of the bomb (9.887 kJ/mol) and the change in temperature for one run of benzene (2.6576 °C) I get:

Δ_cH=-(9.887 kJ/mol)(2.6576 °C)=-26.28 kJ/mol °C

This is nowhere near the given literature values of -3273 kJ/mol.

[Corribus]

You're still only providing fragments of what you're doing, so it's hard to troubleshoot your work. But it's also very clear you're being sloppy with your units, which probably isn't helping.

Take a look at examples 6 and 7 at this page and see if it doesn't help you understand the proper way to do this kind of calculation.

http://www.science.uwaterloo.ca/~cchieh/cact/c120/calorimetry.html

If it doesn't, please provide a step-by-step account of what you did in your experiment and a full account of your calculations (without omitting steps). This should make it much easier to see what you're doing wrong.