6.6
Enthalpy
Chemists ordinarily use a property known as enthalpy (H) to describe the thermodynamics of chemical and physical processes. Enthalpy is defined as the…
Some chemical reactions give off tremendous heat and perform work on the surroundings, such as the combustion of rocket fuel causing a space shuttle to lift off from the ground.
The sum of heat “q,” and work “w,” is the change in internal energy, “ΔE,” as given by the first law of thermodynamics.
For chemical reactions involving gases that occur at atmospheric pressure, the work done is the mechanical work associated with volume changes — either expansion or contraction. Work is therefore equal to the negative value of the pressure times the change in volume.
Substituting for w in the first law of thermodynamics and rearranging the terms in the equation shows that q = ΔE + PΔV, giving the expression for heat flow under constant pressure.
In other chemical reactions, such as the burning of wood to cook food, it is more relevant to quantify the heat given off to facilitate cooking, than to measure the amount of expansion work done on the surrounding.
Since internal energy accounts for both heat and work, ΔE is not used for constant pressure conditions. To exclusively discuss the energy flow in the form of heat, a new thermodynamic function – enthalpy – is defined.
Enthalpy, H, equals the sum of internal energy, E, and pressure-volume work, P-V. Because energy, pressure, and volume are state functions, enthalpy is also a state function.
Absolute enthalpy values for specific substances cannot be measured. Only the change in enthalpy can be determined. Enthalpy change, ΔH, equals the change in internal energy, ΔE + PΔV.
Recalling that the change in energy is the sum of heat and pressure-volume work the equations can be combined to show that under constant pressure conditions ΔH equals the heat, q, gained, or lost by the system.
If the system loses energy to the surroundings in the form of heat — as in the burning of wood — the temperature of the surroundings rises. This is described by a negative sign convention for q. Consequently, ΔH becomes negative, and the process is described as exothermic.
On the contrary, if the system gains energy from the surroundings in the form of heat — such as the reaction occurring in a chemical cold pack — the temperature of the surroundings falls. The heat, in this case, is described by a positive sign convention. This makes ΔH positive, and the process is called endothermic.
Q1: Why is enthalpy more useful than internal energy for measuring heat in chemical reactions?
Enthalpy is more practical because most chemical reactions occur at constant atmospheric pressure, where enthalpy change equals heat flow. Internal energy accounts for both heat and work, making it less convenient for reactions where only heat matters. At constant pressure, ΔH = q, simplifying calculations for everyday chemistry experiments like burning wood or operating a Bunsen burner.
Q2: What is the mathematical relationship between enthalpy and internal energy?
Enthalpy (H) equals internal energy (E) plus the product of pressure and volume (PV). This relationship is expressed as H = E + PV. Since enthalpy incorporates pressure-volume work into internal energy, it provides a complete measure of energy changes at constant pressure without requiring separate work calculations.
Q3: How do exothermic and endothermic reactions differ in terms of enthalpy change?
Exothermic reactions release heat to surroundings, producing negative ΔH values. Endothermic reactions absorb heat from surroundings, producing positive ΔH values. Examples include burning wood (exothermic, ΔH < 0) and chemical cold packs (endothermic, ΔH > 0). The sign of ΔH directly indicates energy flow direction.
Q4: Why can't absolute enthalpy values be measured directly?
Enthalpy is a state function, meaning only changes in enthalpy can be determined, not absolute values for specific substances. Chemists measure ΔH—the difference between final and initial enthalpy states—rather than total enthalpy. This approach is sufficient for tracking energy changes in chemical and physical processes.
Q5: What does pressure-volume work represent in the enthalpy equation?
Pressure-volume work (PΔV) represents mechanical work from gas expansion or contraction during reactions. In the enthalpy equation H = E + PV, the PV term accounts for this expansion work. The sign of ΔV and work are always opposite: expansion (positive ΔV) produces negative work, while contraction (negative ΔV) produces positive work.
Q6: How does enthalpy change when a chemical equation is reversed?
Reversing a chemical equation changes the arithmetic sign of its ΔH. A reaction that is endothermic in one direction becomes exothermic in the opposite direction. This reflects the fundamental principle that energy released in a forward reaction must be absorbed when the reaction proceeds backward.
Q7: When do ΔE and ΔH values differ significantly?
ΔE and ΔH differ significantly when volume changes are substantial, such as during evaporation of water. When volume change is small, like melting ice, ΔE and ΔH are nearly identical. The difference arises because ΔH includes pressure-volume work, which becomes negligible for processes with minimal volume changes.