A 2.75 g sample of the hydrocarbon acetone, CH3COCH3, is burned in a bomb calorimeter with 975 mL of water, initially at 23.50 degrees celsius. The bomb is constructed of 285.0 g of nickel metal having a specific heat capacity of Cp = 0.826 J/ g degrees C. The final temperature of the bomb and the water after the combustion process increases to 29.55 degrees celsius. calculate the following:

(a)The heat flow at constant volume, qv for this combustion (in kJ/mol).

i know qv = delta E. and that delta E = q + w. but how do i find the work? i know work = force / distance. (i'm struggling because my book gives an example where the work is already given so it is not helpful at all.)

(b)The energy released per mole of acetone (in kj/mole).

(c)calculation (b) amounts to the heat of combustion per mole, delta h comb/mole for this compound. Balance the reaction for the complete combustion of acetone, find delta n and then find the value of delta H per mole of acetone (kJ/mole).

i don't know how to even start part (b).

thank you!

You know the delta T for the metal and the water.

heat=masswater*cw*deltaT +massnickel*cn*deltatT
b) use that heat, per mole of acetone consumed (convert grams to moles)
c) Use Hess'law.

To solve this problem, let's break it down step by step.

(a) To find the heat flow at constant volume, qv, we need to calculate the heat released during the combustion process.

The heat flow at constant volume (qv) is equal to the change in internal energy (ΔE) of the system.

We can find ΔE by using the equation ΔE = q - w, where q is the heat supplied to the system and w is the work done by the system.

In this case, the combustion of acetone is occurring in a bomb calorimeter, which is a device that allows for constant volume conditions. Hence, the work done (w) is zero, since there is no expansion or compression of gases.

Therefore, ΔE = qv.

Now, let's find the value of qv.

We can use the equation qv = mcΔT, where m is the mass of the water, c is the specific heat capacity of water, and ΔT is the change in temperature.

Given:
- Mass of water (m) = 975 mL = 975 g (since 1 mL of water has a mass of 1 g)
- Specific heat capacity of water (c) = 4.18 J/g°C (this is a commonly used value)
- Initial temperature (Ti) of water = 23.50°C
- Final temperature (Tf) of water = 29.55°C

Substituting these values into the equation, we have:
qv = (975 g) * (4.18 J/g°C) * (29.55°C - 23.50°C)

Now calculate qv in joules.

(b) To find the energy released per mole of acetone, we need to convert the heat flow (qv) calculated in part (a) to energy per mole of acetone.

First, we need to determine the moles of acetone used in the combustion. To do this, we'll need the molar mass of acetone.

The molecular formula of acetone is CH3COCH3, so its molar mass can be calculated as follows:
Molar mass of acetone = (mass of carbon x atomic mass of carbon) + (mass of hydrogen x atomic mass of hydrogen) + (mass of oxygen x atomic mass of oxygen)

Given:
- Mass of carbon (C) = 1 x atomic mass of carbon = 1 x 12.01 g/mol
- Mass of hydrogen (H) = 6 x atomic mass of hydrogen = 6 x 1.01 g/mol
- Mass of oxygen (O) = 1 x atomic mass of oxygen = 1 x 16.00 g/mol

Now calculate the molar mass of acetone.

Once you have the molar mass of acetone, you can calculate the moles of acetone used in the combustion by dividing the mass of acetone (given as 2.75 g) by its molar mass.

Now that you have the moles of acetone, you can divide the heat flow (qv) calculated in part (a) by the moles of acetone to find the energy released per mole of acetone.

(c) Finally, to find the heat of combustion per mole (ΔHcomb/mol) for acetone, you need to balance the combustion reaction for acetone. The balanced equation will give you the stoichiometric coefficients of the reactants and products.

Once you have the balanced equation, you can determine the change in the number of moles (Δn) of acetone from the balanced equation.

The heat of combustion per mole (ΔHcomb/mol) can be obtained by dividing the energy released per mole of acetone (obtained in part (b)) by the change in the number of moles of acetone.

Hope this helps!