All chemical reactions are accompanied by energy changes. Under most circumstances (particularly when the pressure and volume are kept constant) these changes can be ascribed to changes in enthalpy ΔH. For example combustion reactions (redox reactions involving oxygen) are a major source of energy for most organisms. In warm blooded organisms, the energy released through such reactions is used, in part, to maintain a set body temperature. Within organisms, combustion reactions occur as highly controlled steps (which is why you do not burst into flames), through the process known as respiration. As we will see later, this energy is harnessed by releasing it during the course of a cascade of coupled reactions, rather than all at once.

7.1 Reactions
7.2 Types (Acid-Base)
7.3 Lewis Acid-Base
7.4 Nucleo- & Electrophiles
7.5 Oxidation-Reduction
7.6 Energetics

Not all biological forms of respiration use molecular oxygen, there are other molecules that serve to accept electrons - this process is known as anaerobic (air-free) respiration. That said, all known organisms use the molecule adenosine triphosphate (ATP) as a convenient place to store energy; ATP is synthesized from adenosine diphosphate (ADP) and inorganic phosphate. As two separate species ADP and inorganic phosphate are more stable than ATP, and the energy captured from the environment to make ATP can be released again (with the formation of ADP and inorganic phosphate).
ADP + Pi + energy ↔ ATP + H₂O

ATP synthesis is another example of a electrophile-nucleophile interaction. For all these type of reactions we can ask the same question, where (ultimately) does the energy come from? That is: when an exothermic reaction occurs and energy is transferred from the system to the surroundings (resulting in a temperature increase in the surroundings and a negative enthalpy change – ΔH), what is exactly the source of the energy that is transferred? (Of course you already know the answer, it has to be the energy released when a bond is formed!)

The defining trait of a chemical reaction is a change in the chemical identity of the reactants, new types of molecules are produced. For this to be true, at least some of the bonds in the starting material must be broken, and new bonds must be formed in the products (otherwise no reaction would have occurred). So to analyze energy changes in chemical reactions we look at which bonds get broken and which are formed, and compare their energies. As we will discuss later, the process is not quite so simple, since the pathway by which reaction occurs may involve higher energy intermediates. It is the pathway of a reaction that determines its rate (how fast it occurs), while the difference between products and reactions determined the extent to which the reaction will occur. However, to a good approximation, the approach we are going to look at to estimate energy changes during a reaction gives us a reasonable answer.

As we have already seen, bond formation releases energy, bond breaking requires energy. Tables of bond dissociation (breaking) energies are found in most chemistry books and can be easily retrieved on the internet. One caveat is that these measurements are typically taken in the gas phase, and refer to a process where the bond is broken homolytically (that is each atom in the original bond ends up with one electron and the species formed are known as radicals. The bond dissociation energy for hydrogen is the energy required to drive the process: H–H(g) -> 2H•

where the dot represents an unpaired electron and ΔH = + 436 kJ/mol.

Note that tables of bond energies record the energy required to break the bond. Since enthalpy is a state function, meaning that its value does not depend on the path taken for the change to occur, we also know the enthalpy for the reverse process. That is: when a hydrogen molecule forms from two hydrogen atoms the process is exothermic 2H • → H–H(g) ΔH = – 436 kJ/mol.

Since we have tables of bond energy values for most common bond types, one way to figure out energy changes (or at least enthalpy changes) for a particular reaction is to analyze the reaction in terms of which bonds must be broken and which bonds are formed. The broken bonds will contribute to a positive term to the total reaction energy change while bond formation will contribute a negative term. For example: let us take a closer look at the combustion of methane.
CH₄(g)+ 2O₂(g) ↔ CO₂(g) + 2H₂O(g)

In the course of this reaction, four C-H bonds will be broken [4 x C–H (436kJ/mol)] and 2 x O=O (498 kJ/mol), while the new bonds formed are 2 x C=O (803 kJ/mol) and 4 x O–H (460 kJ/mol). If you do the math you find that the bond energies of all the bonds broken is 2740 kJ, while the sum of the bond energies formed is –3330 kJ. That is the bonds in the products are 706 kJ more stable than the bonds in the reactants. This may be easier to see if we plot the enthalpy vs reaction progress, it becomes more obvious that the products are lower in energy (more stable).

There are large numbers of us expelling CO₂ into the atmosphere (from burning fossil fuels – and breathing!), at a higher rate than it it is currently being removed (through various types of sequestration, chemical reactions and photosynthesis ). You have almost certainly heard of the greenhouse effect, caused by the build up of CO₂. One of the reasons why CO₂ is difficult to get rid of is that it is stable because it has strong bonds. [Given the notoriety of CO₂ in terms of climate change, we will come back to this topic in (yet to be written) chapter 10).]

7.1 Reactions
7.2 Types (Acid-Base)
7.3 Lewis Acid-Base
7.4 Nucleo- & Electrophiles
7.5 Oxidation-Reduction
7.6 Energetics

Question to answer:

• Many biology texts refer to ideas such as energy is released when the high energy bonds in ATP are broken. In light of what you know is this a reasonable statement? What do they really mean?
• Why do you think the enthalpy change for most Bronsted-Lowry acid base reactions is independent of the nature of the acid or base (hint - what is the reaction that is actually occuring?)
• Using tables of bond dissociation energies, calculate the energy change for the reaction of CH₂=CH₂ + HCl → CH₃₂CH₂Cl. What steps do you have to take to complete this calculation? Make a list.
• If you look up the enthalpy change for this reaction (ΔH°) you will find it is not exactly what you calculated, why do you think that is? (Hint, this reaction typically takes place in a solvent - what role might the solvent play in the reaction?)