Enthalpy and the Ideal Gas Law: Understanding Constant Pressure Reactions

In summary, the ΔE = q - w formula is used to calculate the change in internal energy for a reaction, with the addition of the ideal gas term pΔV for constant pressure reactions. However, for constant pressure reactions, the change in enthalpy is equal to the heat released, making it a more convenient measure. The work term pΔV may be included for calculations involving a bomb calorimeter, but is not necessary for general calculations. The relationship between ΔH and ΔU is always applicable and can be used to convert between the two values.
  • #1
Zopiclone
2
0
Hi! I'm new to the forums and currently reading about Chemical Thermodynamics. So here's what I know:

ΔE = q - w

So for constant volume reactions, no work is done hence:

ΔE = q

But for constant pressure reactions, heat be may released (for exothermic reactions) and work is done hence:

ΔE = q - pΔV

so ΔE ≠ q because some of the heat does work.

So we define enthalpy H = E + PV, so, ΔH = ΔE + PΔV = q - pΔV + pΔV = q

and hence for constant pressure reactions ΔH = q.

My confusion is that the book says that the work term pΔV = (Δn)RT for these processes, but how can we hold the temperature constant since it is obviously changing as a result of the heat of the reaction?

Any help would be greatly appreciated. Thanks!
 
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  • #2
The formula pΔV = (Δn)RT must be wrong, I would say. Could you tell me which book you have and the equation number? Normally for a constant pressure process the number of moles is constant. That should mean the temperature changes (for an ideal gas). Unless they are thinking of a chemical reaction, of course, but with a constant T??
 
Last edited:
  • #3
Zopiclone said:
Hi! I'm new to the forums and currently reading about Chemical Thermodynamics. So here's what I know:

ΔE = q - w

So for constant volume reactions, no work is done hence:

ΔE = q

But for constant pressure reactions, heat be may released (for exothermic reactions) and work is done hence:

ΔE = q - pΔV

so ΔE ≠ q because some of the heat does work.

So we define enthalpy H = E + PV, so, ΔH = ΔE + PΔV = q - pΔV + pΔV = q

and hence for constant pressure reactions ΔH = q.

My confusion is that the book says that the work term pΔV = (Δn)RT for these processes, but how can we hold the temperature constant since it is obviously changing as a result of the heat of the reaction?

Any help would be greatly appreciated. Thanks!
Everything you said is correct. With regard to your question, the temperature is being held constant at the initial temperature by either adding heat or removing heat from the reaction mixture as necessary (say, by contact with a constant temperature bath held at the original temperature). When they talk about THE "Heat of Reaction," they are referring specifically to holding both the temperature and the pressure constant. The term you are referring to has to be included if the number of moles of products differs from the number of moles of reactants. In such a case, the change in internal energy will differ from the change in enthalpy (even if the temperature and pressure are constant).
 
  • #4
Thank you very much for your help. I think the ideal gas term term is just included as a convenient conversion between change in internal energy and enthalpy change when the heat is measured in a bomb calorimeter (constant volume). What do you think?
 
  • #5
Zopiclone said:
Thank you very much for your help. I think the ideal gas term term is just included as a convenient conversion between change in internal energy and enthalpy change when the heat is measured in a bomb calorimeter (constant volume). What do you think?
No. Although applicable to a bomb calorimeter, the relationship is much more general than just this one application. If you know the ΔH for the reaction, you can always get the ΔU, and vice versa.
 

1. What is enthalpy?

Enthalpy is a thermodynamic property that represents the total internal energy of a system, including both its internal energy and the energy required to create its volume and pressure. It is often denoted as H and has units of joules (J) in the SI system.

2. How is enthalpy related to ideal gas law?

The ideal gas law is an equation that relates the pressure, volume, temperature, and number of moles of an ideal gas. Enthalpy is related to the ideal gas law through the equation H = U + PV, where U is the internal energy, P is the pressure, and V is the volume. This relationship is used to calculate the enthalpy change of a gas during a process.

3. What is the significance of the ideal gas law in thermodynamics?

The ideal gas law is significant in thermodynamics because it provides a simple and accurate way to describe the behavior of gases under various conditions. It is often used in calculations to determine the properties of gases and to understand the relationships between different thermodynamic properties.

4. How does enthalpy change during a physical or chemical process?

Enthalpy can change during a physical or chemical process in three ways: through a change in temperature, a change in pressure, or a change in phase. These changes can be calculated using the ideal gas law and the specific heat capacity of the substance.

5. Can the ideal gas law be applied to all gases?

The ideal gas law is a theoretical model that applies to ideal gases, which are gases that have no volume and do not interact with each other. While it is a good approximation for many real gases under certain conditions, it may not accurately describe the behavior of all gases. In some cases, more complex equations, such as the van der Waals equation, may be necessary to accurately describe the behavior of real gases.

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