(adsbygoogle = window.adsbygoogle || []).push({}); 1. The problem statement, all variables and given/known data

A quantity of 4.00 3 10 2 mL of 0.600 M HNO 3 is mixed with 4.00 3 10 2 mL of 0.300 M Ba(OH) 2 in a constant-pressure calorimeter of negligible heat capacity. The initial temperature of both solutions is the same at 18.46°C. What is the final temperature of the solution? (Use the result in Example 6.8 for your calculation.)

The result in example 6.8 is the heat of neutralization = -56.2 kJ/mol, I believe.

2. Relevant equations

q = msDelta-T

3. The attempt at a solution

I've asked almost everybody I know who knows anything about chemistry to give this a shot, but no luck. One got x = 120 (no idea where that came from). I've spent hours working on this problem, and have tried various methods, anything I can think of. Two of my closer shots were:

Taking 6.8's 2810J, I tried writing (with C meaning "degrees Celsius"):

2810J = (400mL + 400mL)(4.184 J/g * C)(x - 18.46C)

2810 = 3347.2x - 61789.312

x = 19.2995 (degrees Celsius).

This is incorrect.

I tried taking the number of moles, determining the masses, then combining those, to write; I continue to assume q = 2810J, as have other people in past forum posts I've found via Google. Not entirely sure what else it could be.):

q = (35.684g)(4.184 J/g * C)(x - 18.46C)

2810J = 149.3x - 2711.288

x = 36.98

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# Constant-Pressure Calorimetry question

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