This question seems really dumb but I rather be safe than sorry. I read about enthalpy change in my chemistry textbook which defines enthalpy change as the energy exchange between a chemical reaction and the surroundings at constant pressure. My confusion came from when I read this BBC Bitesize page about energy change, can someone explain to me how they're different as they use the same units?
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$\begingroup$ Gibbs free energy and enthalpy generally share the same units, which might be a source of your confusion. Enthalpy is a component of Gibbs free energy. $\endgroup$– kskinnerx16Commented May 28, 2020 at 14:38
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2 Answers
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It seems that you do not understand the difference between the internal energy $\ce{U}$ and the enthalpy $\ce{H}$. I will try to present these fundamental notions in a different way.
The internal energy U is the sum of all the energies stored between the atoms in the chemical bonds. If a chemical reaction happens, the bonds are losing or absorbing energy. Heat is getting in or out of the system, and this heat $\Delta$Q can be measured and correspond to the change in internal energy. $\ce\Delta$U = $\Delta$Q, if the volume is constant. But if the transformation is carried out in the atmosphere, at constant pressure, a part of this energy is converted into work $\Delta$w to repel the atmosphere, because the volume of the system may change. Usually this volume change is small (except if a gas is produced or consumed). And this volume change may sometimes be difficult to know with precision. But it must be taken into account to calculate $\Delta$U. So if this small amount of work is neglected, a correction must be added to the internal energy U, that depends on pressure and volume. Neglecting this volume correction gives a sort of “apparent internal energy”, which is called enthalpy H, with H = U + PV. And then, in all transformations made at ordinary pressure, the following expression is valid : $\Delta$H = $\Delta$Q.
The difference between $\Delta$U and $\Delta$H is similar to the difference between the weight of an object and its weight in the vacuum. After Archimedes, all objects dipped in a fluid lose in weight the amount of fluid (air or water) they displace. When an object is weighed in air, the obtained value is the apparent weight, smaller than the weight in a vacuum. It should be increased by the Archimedes’ principle. Nobody does it. It is about the same for the enthalpy. Enthalpy is a sort of apparent internal energy.
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From 1st law of thermodynamics we get dU=DQ+DW(I wrote 'D' in place of 'd' because U is a total differential and a function of state only but heat or work are no such property of system, they are transient form of energy and so it's 'D') Any way if Pop is the external pressure on system then the work term becomes DW= - Pop.dV dU=DQ-Pop.dV=DQ-PdV(if Pop=pressure of gas=P,i.e. reversible) =>dU=DQ if there is no change in volume of the system,i.e. (DQ)v=dU Whereas, dH=d(U+PV)=dU+pdV+VdP=DQ+vdP =>dH=DQp(if the whole process is isobaric i.e. constant pressure and reversible)
So you get internal energy change of a system as the heat exchange in constant volume process and enthalpy change of a system as the heat exchnge in a constant pressure process
Actually U and H both are measures of some amount of energy. Try to understand their definations. Definition of H is H=U+PV. Don't consider only chemical reactions. It's valid any where.
Actually most of the chemical reactions are done in the lab at constant pressure and so the heat exchange of chemical reactions is more convenient to express in terms of ∆H.
