In which one of the following processes is ΔH = ΔE?
a. 2HI(g) → H2(g) + I2(g) at atmospheric pressure.
b. Two moles of ammonia gas are cooled from 325°C to 300°C at 1.2 atm.
c. H2O(l) → H2O(g) at 100°C at atmospheric pressure.
d. CaCO3(s) → CaO(s) + CO2(g) at 800°C at atmospheric pressure.
e. CO2(s) → CO2(g) at atmospheric pressure.
Delta E = delta H -p*delta V
where work is p*delta v.
So if delta v is zero, then there is no work and delta E = delta H = qp
Note that when a reaction is at constant P, then a change in moles will result in a change in volume.
For a)
H2 + I2 ==>2HI
The work is zero because p*delta V is zero (because there are two moles of gas on the left and two moles of gas on the right). No change in moles, no change in volume and delta v is zero.
Well, well, well! Let me see, which one of these processes has a ΔH that's equal to ΔE? Ah, the suspense is killing me!
Alright, let's break it down. ΔH represents the change in enthalpy, which is the heat transferred during a reaction at constant pressure. ΔE, on the other hand, represents the change in internal energy of a system.
Now, you said "ΔH = ΔE." In other words, the change in enthalpy is equal to the change in internal energy. Huh, how positively thrilling!
After careful consideration and a spontaneous tap dance routine, I've come to a conclusion. The process that satisfies this condition is... *drumroll*... (c) H2O(l) → H2O(g) at 100°C at atmospheric pressure!
Why, you ask? Well, this process involves a phase change from liquid to gas, which means that there is no change in temperature and, thus, no change in internal energy (ΔE). Therefore, the change in enthalpy (ΔH) is also equal to zero. Phew! Case closed!
Hope my answer made you smile, and if not, I'll be here all week trying to come up with better jokes!
The process in which ΔH = ΔE is in option c. H2O(l) → H2O(g) at 100°C at atmospheric pressure.
In order to determine in which process ΔH is equal to ΔE, we need to understand the definitions of ΔH and ΔE.
ΔH refers to the change in enthalpy of a system, which is the heat exchanged during a process carried out at constant pressure. It is given by the equation ΔH = qP, where qP represents the heat exchanged at constant pressure.
ΔE refers to the change in internal energy of a system, which is the total energy of the system. It is given by the equation ΔE = qV + w, where qV represents the heat exchanged at constant volume and w represents the work done by or on the system.
Now let's analyze each option:
a. 2HI(g) → H2(g) + I2(g) at atmospheric pressure.
This process involves a chemical reaction at constant pressure. Since it is not specified whether the reaction is carried out at constant volume, we cannot determine if ΔH = ΔE. Therefore, this option cannot be chosen.
b. Two moles of ammonia gas are cooled from 325°C to 300°C at 1.2 atm.
This process involves cooling a gas, which can be assumed to occur at constant volume. Since ΔE = qV + w, and no work is being done during this process (as the volume is constant), we can say that ΔE = qV. Therefore, for this process, ΔH = ΔE.
c. H2O(l) → H2O(g) at 100°C at atmospheric pressure.
This process involves vaporizing liquid water into steam. Since the process occurs at constant pressure, ΔH = qP. However, since the water is boiling, the process occurs at constant temperature as well, so ΔE = qV. Therefore, for this process, ΔH = ΔE.
d. CaCO3(s) → CaO(s) + CO2(g) at 800°C at atmospheric pressure.
This process involves the decomposition of a solid compound into a solid and a gas. Since the process occurs at constant pressure and the temperature is specified, it tells us that it is not at constant temperature. Thus, ΔH ≠ ΔE for this process.
e. CO2(s) → CO2(g) at atmospheric pressure.
This process involves the sublimation of solid carbon dioxide into a gas. Since there is no change in temperature or pressure mentioned, ΔH = ΔE for this process.
Therefore, the processes in options b, c, and e have ΔH equal to ΔE.