Heat and Chemical Reactions
Recall that Dalton’s atomic model properly stated that in chemical reactions atoms rearrange as the reaction proceeds. In more modern terms this means chemical bonds are broken and formed, broken to split a compound into smaller parts and formed as smaller parts stick together into a bigger compound. In terms of energy these changes are accompanied by a certain amount of heat gain or loss.
Imagine a chunk of pure sodium lying around. In the pure elemental state there is no forming or breaking of bonds, but chemistry knowledge says that exposed sodium will not stay in elemental form for long. Any water it is exposed to, even the humidity in the air, will cause a reaction that gives off a very noticeable amount of heat. At this point a determination for the change in energy can be determined by determining the change in heat and work.
What if the amount of heat for the reaction was desired before the reaction occurred? Determining this would be like determining the chemical potential energy for the sodium. This chemical potential energy is expressed in terms of heat, and is called enthalpy (H).
Chemists have calculated and tabulated enthalpy of reactions. The same problem occurs with enthalpy as with heat - it is impossible to calculate an absolute value. The way chemists got around this problem is to say that the elements in their pure, elemental form have an enthalpy of zero, and then all the other measurements can be made relative to that value. Below is an example list of some enthalpies of formation (the change in enthalpy when the compound is formed from the pure elements).
Why is the value for H2, F2, and Cl2 zero? Remember the results are compared to the elemental forms, which is diatomic for seven elements.
Most of the enthalpies are negative, which means the system loses energy as the products are formed. The study of the change in heat (and energy in general) during a reaction is called thermodynamics.
Thermodynamics has placed reactions into two broad categories: endothermic and exothermic.
Endothermic reactions require energy as the reaction progresses, and this is most often accomplished by absorbing heat. (Think endo = enter and thermo = heat.) The surroundings must provide this energy to the system, and if your hand was the surroundings it would feel cool as the heat leaves your hand. Often chemists denote that heat is needed by writing it as reactant, and the same could be done with any form of energy. Sometimes heat is written over the arrow of an equation to mean heat is added as the reaction goes that direction.
An endothermic reaction has a positive value for the enthalpy (established by "committee"), and so only one of the products listed in the table above would be labeled as an endothermic reaction. Common endothermic reactions include melting, evaporating, vaporizing (boiling), and decomposition. Some chemicals are stored in amber colored bottles because light energy is enough to get the chemical to decompose.
Exothermic reactions are literally the reverse of endothermic reactions, in that they give off energy (usually in the form of heat) as the reaction progresses. (Think exo = exit.) Exothermic reactions would be warm to the touch (provided it is a safe enough to touch). Heat can be written as a product, and sometimes written with an arrow pointing away from the reaction to indicate heat is escaping.
Exothermic reactions have a negative value for enthalpy, as energy is given off. Common exothermic reactions are freezing, condensing, most synthesis reactions and of course, combustions.
Often chemists use a visual representation of the enthalpy of a reaction to get a general idea about the change in enthalpy that occurs as a reaction progresses. Enthalpy is the dependent variable and time is the independent variable (the one the experimenter controls).
A diagram like this is helpful as the enthalpy could be likened to the potential energy of a ball traveling the same curve. Overall the ball would travel downhill and the ball’s potential energy is less than it had to begin with (-Ep). The same is said about the enthalpy of the reaction. The products have less enthalpy than the reactants, so the overall ∆H has a negative value (energy is given off). This reaction would be called an exothermic reaction.
It may be surprising to see that the reactants have to go “uphill” or gain energy before they can form products, even in an exothermic reaction. This amount of energy is called the Energy of Activation (Ea). Most reactions, even exothermic ones, require some amount of energy to get started, just like it takes a match to get a log in a fireplace to burn. The actual amount of energy needed depends on the reaction and if necessary for calculations would have to be specified.
Once the reactants reach the top of the “hill” the products can form spontaneously. Because the “downhill” side of the diagram is bigger than the “uphill” side, the overall reaction has the value of negative enthalpy (exothermic) as would be measured from the ∆H from the graph.
What would the diagram look like for an endothermic reaction? Water can be separated into hydrogen and oxygen, but it takes energy (usually in the form of electricity).
Now the reaction is endothermic because the products end up with more energy then the reactants started with. Notice the value for ∆H is not the entire height of the “hill”, but is only the difference between the products and the reactants. The energy of activation is the larger amount and will never be completely the same as the ∆H. Some of the energy of activation is returned as the products roll “downhill”.
In both exothermic and endothermic reactions the ∆H = ΣHproducts - ΣHreactants. If ΣHproducts - ΣHreactants < 0, then the negative sign indicates it is an exothermic reaction. If ΣHproducts - ΣHreactants > 0, then the positive sign indicates it is an endothermic reaction. (Just like in math class, the symbol Σ (capital sigma) means the sum of.)
For the purpose of comparison, chemists often determine a value for enthalpy under specific conditions. The standard temperature and pressure for the enthalpy of substances is 25 °C (298 K), 1 atmosphere of pressure, and a concentration of 1 M for solutions.
Enthalpy is a state function, which means the value is only dependent on the initial and final enthalpies. There is no need to worry about how the enthalpy changes between those states, which is very convenient as the curves shown in the graphs shown above are idealized. For those reactions for which chemists have successfully completely figured out exactly what happens as the reactants form products there are often many peaks to the graph, not just one big one. In addition, most reactions happen at a rate that is too fast to determine what exactly is happening, so technically for these reactions there would be a known curve. As enthalpy is a state function none of this matters, and the enthalpy calculations are determined as easily as finding the initial and final enthalpies and subtracting.