At last we have arrived at the place where many chemistry courses begin – chemical reactions. At this point, it is worth thinking about what a chemical reaction is, exactly, what processes are not chemical reactions, how chemical reactions occur, and how they are characterized (we will return to reaction behaviors in greater detail in chapter 8). First, as seems obvious, chemical reactions are linked to change but not all change involves a chemical reaction. When liquid water boils or freezes it undergoes a change of state (a phase change), but the water molecules remain intact - they are still discrete water molecules, H₂O. In ice, they remain more or less anchored to one another through H-bonds while in liquid and water vapor they they are constantly moving with respect to one another, and the interactions that occur between the molecules are transient. We can write out this transition in symbolic form as :
H₂O (solid) ↔ H₂O (liquid) ↔ H₂O (vapor).

The double arrows mean that the changes are reversible (in this case, reversibility is a function of temperature, which controls whether the interactions between molecules are stable (as in ice) or transient, as in liquid water or basically non-existent as in water vapor. What you notice immediately is that the H₂O remains represent and unchanged in each - that is there are water molecules present in each phase. This help explain away a common misconception, namely what is in the bubbles found in boiling water. Since boiling does not break the bonds in a water molecule, the bubbles are composed of water vapor. In contrast, within liquid water, there is a chemical reaction that is going on – the disassociation of water into –OH and H+, which we will discuss in more detail shortly. However a naked proton (H+ that is a proton as a discrete separate species) does not exist in water, this reaction is more accurately written as:

H₂O + H₂O ↔ H₃O⁺ + ^-OH.

Here we see the signature of a chemical reaction, the molecules on the two sides of the equation are different - covalent bonds are broken (an O-H bond in one water molecule) and formed (a H-O bond in the other). All chemical reactions can be recognized in this way. An interesting feature of the water dissociation reaction, which illustrates another feature of chemical reactions, is that reactions can vary in terms of the extent to which they occur. In liquid water, which has a concentration of about ~55M, very few molecules undergo this reaction - in fact in pure water the concentration of H₃O⁺ is only 10-7 M, that is eight orders of magnitude less than the concentration of water molecules. Another interesting feature of this reaction is that (as indicated by the double arrows ↔) the reaction is going in both directions - that is water reacts with itself to give H₃O⁺ + ^-OH, and at the same time H₃O⁺ + ^-OH are reacting to give water molecules. We say the reaction is at equilibrium, and in this case the position of the equilibrium is mostly on the reactants (H₂O) side.

In contrast, other reactions, will go essentially to completion. For example pure ethanol (CH₃CH₂OH), which is approximately 17.1M will burn in air (oxygen) and the reaction appears to go to completion:

CH₃CH₂OH + 3 O₂ ↔ 2 CO₂ + 3 H₂O

There is very little ethanol left if this reaction occurs in the presence of enough oxygen (however there will be small but detectable traces of ethanol - forensic scientists are able to detect the presence of substances such hydrocarbons at the scene of a fire, even though the amounts are extremely small) In fact,the reaction is essentially irreversible. Another interesting feature of the ethanol burning reaction is that pure ethanol can sit in contact with the atmosphere (which depending upon where you live contains ~20% O₂), and be quite stable. It takes a spark or a little heat before the reaction occurs (for example, vodka - which is about 50% ethanol, will not burst into flames without a little help!) Many reactions need a spark of energy to get them started, but once started release enough energy to keep them going. As we saw with the process of solution, some reactions release energy (are exothermic) and some require energy (are endothermic). It is important to note that this overall energy change is not related to the spark or energy that is required to get some reactions started.

Another feature of reactions is that some are faster than others. For example if we added hydrogen chloride gas to water, a reaction would occur almost instantaneously:

HCl(g) + H₂O(l) ↔ H₃O+(aq) + Cl–(aq)

That is, almost no time would elapse between dissolving the HCl and the reaction occurring. We say the rate of the reaction is very fast. Alternatively, when iron nails are left out in the weather they form rust - a complex mixture of iron oxides and hydroxides. The reaction is slow and can take many years, however in hot climates the reaction goes faster, just as when we cook food the reactions that take place occur at a faster rate than they would at room temperature.

This chapter is an introduction to how molecules come to be “reorganized” during a chemical reaction, and in chapter 8 we will discuss why reactions happen. However, it will help if we get a few ideas about reactions straight now. As we have seen previously, bonded atoms are typically more stable than unbonded atoms. For a reaction to occur, some bonds have to break, and new ones have to form. What would lead to a bond breaking? Why would new bonds be formed? What are the factors that affect whether reactions occur, how much energy is released or absorbed, where they come to equilibrium, and how fast they occur?

For a reaction to occur the reacting molecules have to collide, that is they have to bump into each other to have a chance of reacting at all. An important point to remember is that molecules are not sitting still, they are moving (if they are in liquid or gaseous phase), or they are vibrating if they are in a solid. Remember that the temperature of a system of molecules, is a function of average kinetic energy of those molecules. Normally, it is enough to define the kinetic energy of a molecule as 1/2 mv² , but in fact if we are being completely rigorous this equation applies only to rigid bodies. Molecules are more complex, since they can flex, bend, rotate around bonds and vibrate, and many reactions occur in solution where molecules are constantly in contact with each other - bumping and transferring energy, which may appear as kinetic energy, or vibrational energy. Nevertheless, we can keep things simple for now (as long as we remember what simplifications we are assuming). Recall that while the temperature is proportional to the average kinetic energy of the molecules, this does not mean that all the molecules in the system are moving with the same velocity. There is typically a broad range of velocities of molecules, even if they are all the same type, and since reaction mixtures must have more than one type of molecule in them there is an even broader range of molecular velocities. Since the system has only a single temperature, all types of molecules must have the same average kinetic energy , which means that the more massive molecules are moving more slowly, on average, than the less massive molecules. At the same time, all the molecules are (of course) moving, so they will inevitably collide with one another and, if the system has a rigid boundary, with that boundary. We have previously described the distribution of velocities found in the system in terms of a distribution of velocity (or speed) and the percent or even absolute number of molecules with that speed (the Boltzmann distribution). This figure shows speed distribution of a number of different molecules in a gas at 293.15 K. Such a distribution is characterized by the most common particular speed (the peak of the curve), but many molecules are moving much faster (and so have a much higher kinetic energy) and other molecules that are moving much slower (and so have much less kinetic energy) that the average kinetic energy of the population. This means that when any two molecules collide with one another, the “intensity” of that interaction can vary dramatically. Some collisions will involve relatively little energy, some a lot!

Collisions and chemical reactions:

These collisions may or may not lead to a chemical reaction, so let’s consider, in a general way, what happens during a chemical reaction. To focus our attention, we will consider a specific interaction, namely the reaction of hydrogen and oxygen gases to form water. [2H₂ + O₂ ↔ 2H₂O]. This is, in fact, a very complex reaction, and so we will think about in only a cartoonish (but accurate) way. If we have a closed flask of pure oxygen, and we add some hydrogen (H2) to the flask, the two types of gas molecules will quickly mix (why is that?). Some of the molecules will collide with each other, but the overwhelming majority of these collisions will be unproductive, neither the hydrogen molecule (H2) nor the oxygen molecule (O₂) will be altered, although there will be changes in their respective ‘kinetic energies. However, if we add kinetic energy (from a burning match, and of course the burning match is itself a chemical reaction of another sort), the average kinetic energy of all molecules in the heated region increases, the energy that can be transferred upon collision increases, and therefore the probability that a particular collision will lead to a bond breaking and the probability of the H₂ + O₂ reaction will increase. In addition, because, the stability of the bonds in H₂O are greater than those of H₂ and O₂, the reaction will release energy to the surroundings. This energy can take the the form of kinetic energy (which will increase the temperature) and electromagnetic energy (which will result in the emission of photons of light). In this way the system becomes self-sustaining, that it, it no longer needs the burning match, the energy released as the reaction continues will be enough to keep new molecules reacting. The reaction of H₂ and O₂ is explosive (it rapidly released thermal energy and light), but only after that initial spark has been supplied.

We can plot out the behavior of the reaction, as a function of time, beginning with the addition of the burning match. Now it is worth keeping in mind that the reaction converts H₂ and O₂ into water. That means that the concentrations of H₂ and O₂ in the system must be decreasing as the reaction proceeds, while the concentration of H₂O is increasing. As the reaction proceeds, the probability of productive collisions between H₂ and O₂ molecules decreases. We can think of it this way, the rate that the reaction occurs in the forward (to the right) direction is based on the probability of productive collisions between molecules of H2 and O₂, and this in turn depends upon their relative concentration (one reason hydrogen will not burn in the absence of O₂). As the concentrations of the two molecules decreases, the reaction rate slows down. Now normally, the water molecules produced by burning disperse and the concentration (molecules per unit volume) of H₂O would never grow very large, but if the molecules were contained in a space, then their concentrations would increase, and eventually, the back reaction might begin to occur. At some point, the reaction would reach equilibrium, at which point the rate forward and back reactions would be equal and opposite.

Question to answer:

• What do we mean by rate of reaction? How would you determine a reaction rate?
• What has to happen for something to react?
• How does the concentration of reactants and products influence the rate of a reaction?
• Are chemical reactions possible in the solid phase?
• What factors are required for a reaction to reach a stable (albeit dynamic) equilibrium.
• Why is a burning building unlikely to react equilibrium?
  

Questions to ponder:

• What is required for a reaction to go backwards?