Before venturing into enthalpy, let’s review some basic concepts of thermodynamics.
Let’s just follow the definitions of textbook for some concepts.
Energy: the capacity of an object (or substance) to do work (denoted by w) or supply heat (q). In physics and chemistry, work is a force on something, causing the object to move. If nothing moves, there is no work (w = 0).
An object, or a substance contains all kinds of energy which is called the internal energy (U or E). The internal energy is the sum of the system’s kinetic energy and potential energy.
The change in the energy of the system (denoted by ΔE) is equal to the sum of work and heat:
ΔE = q + w
NOTE: An object does not contain heat (q) or work (w). it only contains energy.
Potential energy is the energy due to the composition, position, or condition of an object. Kinetic energy is the energy an object has because of its motion (moving). Take a ball as an instance, when it falls down from a hill, its potential energy is converted to kinetic energy while its energy is still conserved according to the First Law of Thermodynamics (in this case the ball is an isolated system). The ball’s overall energy can be increased by putting work on it (kicking).
NOTE: In chemistry, potential and kinetic energy is the energy of the particles (atoms or molecules).
Chemical Energy and Thermal Energy
Chemical energy is a kind of potential energy – which is the energy stored in the bonds. Many substances contain chemical energy such as wood, food, or petroleum. Thermal energy is the kinetic energy of the atoms or molecules. The faster the atoms of molecules move, the more thermal energy they possess; therefore, they have high temperature, which is the quantitative measure of thermal energy. High thermal energy means high temperature. Heat is the transfer of energy by thermal interactions (motion of atoms and molecules), in other word, heat is the transfer of thermal energy.
Now, with that fundamental knowledge, let’s jump into enthalpy – the abstract concept invented by human beings.
So what is enthalpy?
Enthalpy is a measure of the heat energy (or heat) that is released or absorbed when bonds are broken and formed during a reaction that’s run at constant pressure. Scientists are interested in how much an enthalpy of a system changes rather than its enthalpy itself because to know the total energy or enthalpy of a system is extremely difficult. Therefore, we define DH as a change in enthalpy to measure it.
Energy of a system can be transferred by two mechanisms: work or heat (as the equation ΔE = q + w tells us). In case of heat, the energy is transferred to the motion of atoms and molecules; on the other hand, work is the transferred of energy to the motion of object.
To describe how energy flows, scientists use three designations: the system, the surroundings, and the universe. We can call anything the system and the others will become the surroundings. Mostly, what we are trying to study is called the system. For example, consider reaction of methane (CH4) with oxygen (O2) to produce oxygen (CO2) and H2O (l), the reaction itself is the system and whatever else (the air surrounding it) is the surroundings. When heat flows out of the system, we say the energy of the system decreased and the energy of the surroundings increased. The energy of the system increased when the system gains energy and the energy of the surroundings decreased.
And here, we have the equation for the enthalpy of a system:
H = E + PV
So if we study the change of the enthalpy, the equation will become:
ΔH = ΔE + ΔPV = q + w + ΔPV
Now because we are dealing with most of those chemical reactions that happen at constant pressure such as in a beaker, flask, or even a boiling soup can although the pressure changes a little but is negligible compared to other term in the equation. Therefore, the equation then becomes:
ΔH = q + w + PΔV
And now another assumption is that the work is the one that the pressure does to change the volume, which is known as pressure-volume work. By definition, PΔV and w are always opposite ( w > 0 → ΔV <0 when work is done on the system and the volume decreases and vice versa):
ΔH = q – PΔV + PΔV = q → ΔH = q
So the heat flow (q) and the enthalpy change (ΔH) are equal for a process under constant pressure. The heat is the energy that is transferred out or into the bonds when the chemical reaction occurs.
- When a bond is formed, energy is released. ΔH < 0
- Energy must be put into the system to break a bond. ΔH > 0
In a chemical reaction, energy is put into the system to break the bonds of the reactants. Then energy is released when the bonds are formed to form products. Therefore, the change in enthalpy is equal to the enthalpy of the products minus that of the reactants. (Remember, from the start we have been implying that all the reactions are happening under the constant pressure):
ΔH = Hproducts – Hreactants
If ΔH > 0, we call the reaction endothermic, meaning that the reaction must absorb heat from the surroundings so that it can occur. If ΔH < 0, the reaction is exothermic, the reaction releases heat to the surroundings.