a. Exactly 4.000 mol of NH3 is burned in a constant volume (bomb) calorimeter immersed into 100.0kg of water. The temperature of water increases by 3.619 degrees C after the reaction. Calculate the heat released by the reaction. Neglect the heat capacity of the bomb and the calorimeter; assume that all the heat is spent on warming up water. Specific heat capacity for water is Cm= 4.184 J/ (gx°C)

b. Explain why the absolute values of the standard enthalpy change in the reaction and the heat measured by the calorimeter are not equal. Assume that there are no measurement errors of any kind.

My work:

a.
qcal= C cal* dlta T
qcal= 4.184 * 3.619 = 15.142

This answer is wrong.

b. I don't have any idea on this part of the question

Can someone explain how to do this problem. Thanks!

You didn't count the water. The problem tells you to neglect the calorimeter so

q = mass H2O (100,000g) x specific heat H2O x delta T. I don't know what part b is talking about.

Is delta T (3.619-0)?????

yes

Yay! Okay thanks so much! :)

To calculate the heat released by the reaction in a bomb calorimeter, you need to use the formula:

q = Cm * m * ΔT

Where:
q is the heat released by the reaction (in joules)
Cm is the specific heat capacity of water (in joules per gram per degree Celsius)
m is the mass of the water (in grams)
ΔT is the change in temperature of the water (in degrees Celsius)

Given:
Cm = 4.184 J/(g°C)
m = 100,000 g (because 100.0 kg of water)
ΔT = 3.619°C

First, convert the moles of NH3 burned to grams. The molar mass of NH3 is:

NH3: N(14.01 g/mol) + H(1.01 g/mol * 3) = 17.03 g/mol

So, 4.000 mol NH3 * 17.03 g/mol = 68.12 g NH3

Now, use the equation q = Cm * m * ΔT:

q = 4.184 J/(g°C) * 100,000 g * 3.619°C = 1,250,940 J

Therefore, the heat released by the reaction is 1,250,940 joules.

Now, let's move on to part b.

The absolute value of the standard enthalpy change represents the heat change that occurs at standard conditions (usually 298 K and 1 bar pressure) when reactants are converted to products, and is denoted as ΔH°.

On the other hand, the heat measured by the calorimeter represents the actual heat change that occurs during a specific experimental setup, taking into account factors such as heat loss to the surroundings, experimental errors, and other deviations from the ideal conditions.

In real-world scenarios, it is difficult to replicate ideal conditions perfectly, and there will always be some amount of heat loss or gain to the surroundings. This can result in differences between the heat measured by the calorimeter and the absolute value of the standard enthalpy change in the reaction.

Hence, the absolute values of the standard enthalpy change and the heat measured in the calorimeter are not equal because they take into account different conditions and factors present in the experimental setup.