While there are a rather bewildering array of possible reactions, the truth is that most chemical reactions fall into a rather limited number of basic types. This is lucky for the student of chemistry, since recognizing which type involved in a particular chemical reaction greatly simplifies our task, and should enable you to achieve a reasonable level of confidence. While each particular reaction involves specific molecules and specific conditions (e.g. temperature, solvent) and are therefore somewhat different; very often some common rules apply. So rather than memorizing large numbers of seemingly unrelated reactions, we will present a short introduction to the most common reaction types, specifically acid-base (nucleophile/electrophile), reduction-oxidation (do not worry if these names do not mean anything specific to you at the moment, they will).

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

One important point to remember is that, whatever the reaction type, real reactions are systems composed of reactants, products, and the environment in which the reaction occurs. Reactants can behave quite differently in the gaseous phase from how they behave in a aqueous or a non-aqueous system, or at high or low temperature. In the next chapter we will consider whether (thermodynamics) and how fast (kinetics) a particular reaction occurs under specific conditions, and this will, in turn, lead us to consider equilibrium and non-equilibrium systems.

Question to answer:

• What (in your own words) do you mean when you say a chemical reaction.
• How can you tell when a chemical reaction has occurred?
• Give some examples of reactions that you know – or have learned about in previous courses.
• Can you classify or sort the reactions that you have noted? What criteria would you use to classify them?
• What does acidic (or basic) mean to you?
  

Questions to ponder:

• What reactions going on around you right now?

Acid-Base reactions: A guide for beginners

Let us begin with the reaction type that we saw at the end of the last chapter, the reaction of HCl with water, a classic acid-base reaction. To understand how these types of reactions are related, we need to learn how to identify their essential and common components. Our first hurdle is the fact that the terms acid and acidity (and to a lesser extent, bases and basicity) have entered the language of everyday life. Most people have some notion of acids and acidity. Examples of common usage include: acid rain, stomach acid, acid reflux, acid tongue, etc. You might hear someone talk about wine that tastes acidic, by which (we think) they mean sour – and most people would understand that the wine tastes like vinegar (which contains acetic acid). You have also probably heard or even learned about measurements of acidity that involve “pH”, but what is pH exactly? What is an acid, and why would you want to “neutralize” it? Are acidic things “bad”? Do we need to avoid them at all costs and under all circumstances? While the term “base” is less common, you are likely to be familiar with materials that are basic in the chemical sense (not in the normal language sense of being fundamental). Bases are often called alkalis – as in alkaline batteries or alkali metals, they are slippery to the touch and bitter tasting.

The idea of acids and bases is complex and appears on the surface to involve a number of different types of reactions. Not surprisingly, then a number of definitions of acid-base reactions have been developed and used over the years – but not to worry, each definition turns out to be consistent with the others – they represent the evolution of an idea. Subsequent definitions encompass the original ideas about acid-base while widening them; making them more broadly applicable, that is, they cover a larger array of reactions with similar characteristics. We will start with the simplest model of acids and bases – the Arrhenius model because it is the most common introduction to acid-base chemistry (perhaps you will remember being taught this way). However, the Arrhenius model of acids and bases is actually of limited usefulness. Once we have introduced it, we will proceed to more sophisticated, and more generally useful, models. Our model-by-model consideration of acid-base models should help you appreciate how our understanding of acid-base chemistry has become increasingly general, precise and powerful over time. At the same time, it might help to keep this simple rule in mind; that all acid-base reactions begin and end with polarized molecules. As we go through the descriptions of various models for acid-base reactions, see if you can identify the polar groups, and how they interact with each other during the reaction.

Arrhenius acids and bases:

In the Arrhenius model an acid is a compound that, when dissolved in water, dissociates at the molecular level to produce a proton (H+) and a negatively charged ion (an anion). In fact, naked protons (H+) do not roam around in solution – they are always associated with at least one, and more likely multiple water molecules. Generally, chemists use a shorthand for this situation, either referring to the H+ in aqueous solution as a hydronium ion (denoted as H₃O+) or even more simply as H+ , but do not forget, this is a short-hand, H+ stands for aqueous H+ or H+(aq).

An example of an Arrhenius acid reaction is: HCl(g) + H₂O ↔ H₃O+ (aq) + Cl–(aq),
or
more simply (and more true to the original theory): HCl(g) ↔ H+ (aq) + Cl–(aq) or HCl(aq).

But this is really quite a weird way to present the actual situation, since the HCl molecule does not interact with a single water molecule, but rather interacts with water as a solvent. When hydrogen chloride (HCl) gas is dissolved in water, it dissociates into H+ (aq) and Cl– (aq) almost completely. For all intents and purposes there are no HCl molecules in the solution. An aqueous solution of HCl is known as hydrochloric acid, which distinguishes it from the gas hydrogen chloride. This complete dissociation is a characteristic of what are known as strong acids (but not all acids are strong acids!)

An Arrhenius base is a compound that, when dissolved in water, generates hydroxide (–OH) ions. The most common examples of Arrhenius bases are the Group I (alkali metal) hydroxides, such as sodium hydroxide:
NaOH(s) + H2O ↔ Na+(aq) + –OH(aq) or NaOH(aq).

Again, this is a reaction system that involves both NaOH and liquid water. The process of forming a solution of sodium hydroxide is just like that involved in the interaction between sodium chloride (NaCl) and water, in that the ions (Na+ and –OH) separate and are solvated by the water.

As we will see shortly, some acids (and bases) do not ionize completely, that is when they dissolve in water some of the acid molecules remain intact. When this occurs we use double headed arrows (↔) to indicate that the reaction is reversible, and both reactants and products are present in the same reaction mixture. We will have much more to say about which direction (and how far) a reaction proceeds in the next chapter. For now, it is enough to understand that acid-base reactions (and in fact all reactions) are reversible at the molecular level. In the case of these simple Arrhenius acids and bases, however, the reaction proceeds almost exclusively to the right.

An Arrhenius acid-base reaction occurs when a dissolved (aqueous) acid and a dissolved (aqueous) base are mixed together. The product of such a reaction is usually said to be a salt plus water and the reaction is often called a neutralization reaction; the acid “neutralizes” the base, and vice versa. The equation may be written like this:
         HCl(aq) + NaOH(aq) ↔ H₂O(l) + NaCl(aq)

When the reaction is written in this “molecular” form it is quite difficult to see what is actually happening. If we rewrite the equation to show all of the species involved, and assume that the number of HCl and NaOH molecules are equal, we get:
          H+ (aq) + Cl– (aq) + Na+ (aq) + –OH (aq) ↔ H₂O(l) + Na+ (aq) + Cl– (aq)

Na+ (aq) and Cl– (aq) appear on both sides of the equation, that is, they are unchanged and do not react. The only actual reaction that occurs, is the formation of water 
          H+ (aq) + –OH (aq) ↔ H₂O(l)

The formation of water (not the formation of a salt) is the “signature” of an Arrhenius acid - base reaction. A number of common “strong” acids, including hydrochloric acid (HCl), sulfuric acid (H₂SO₄), and nitric acid (HNO₃) will, when reacted with a strong base (which, like strong acids, dissociate completely in water) such as NaOH or KOH, produce a similar reaction - they produce water.

Such acid-base reactions are always exothermic, and we can measure the temperature change and calculate the corresponding enthalpy change (ΔH) for the reaction. It turns out that it does not matter which strong acid or strong base you choose, the enthalpy change is always the same (about 58kJ/mol of H2O produced). This is because the same and only net reaction that takes place in a solution of a strong acid and strong base is:
       H+ (aq) + –OH (aq) ↔ H₂O(l).

One other factor to note is that the overall reaction involves a new bond being formed between the proton (H+) and the oxygen of the hydroxide (–OH). It makes sense that something with a positive charge should be attracted (and bond with) a negatively charged species (although you should recall why the Na+ and Cl– do not combine to form sodium chloride solid in aqueous solution). Whether such a reaction occurs (bond formation or not bond formation) depends on the exact nature of the system, and the enthalpy and entropy changes that are associated with the change.

We will return to this idea later.

Question to answer:

• What would be the reaction if equal amounts if equimolar HNO₃ and KOH were mixed?
• How about equal amounts if equimolar H₂SO₄ and KOH?(What would the products be?)
• How about equal amounts if equimolar H₃PO₄ and KOH?
• How many moles of NaOH would be needed to react fully with one mole of H₃PO₄?

Brønsted-Lowry Acids and Bases.

While the Arrhenius acid-base model is fairly easy to understand, it is very restrictive in terms of the kinds of reactions to which it applies. So rather than continuing down this road, chemists found that they needed to expand their model of acids and bases and how they react. The first of these expansions was the Brønsted-Lowry model. In the Brønsted-Lowry model, an acid is characterized as a proton (H+) donor and a base as a proton acceptor. If we revisit the reactions we looked at earlier, but now in terms of the Brønsted-Lowry acid-base model we see that HCl is the proton donor, it gives away H+ and water is the proton acceptor. In this scheme, HCl is the acid and water is the base.

HCl(g) + H₂O(l)                 ↔     H₃O+ (aq) + Cl–(aq)
Acid    + Base conjugate          acid            + conjugate base

The resulting species are called the conjugate acid (so H₃O+ is the conjugate acid of H₂O), and the conjugate base (Cl– is the conjugate base of HCl). This is because H₃O+ could (and generally does) donate its H+ to another molecule, most often another water molecule, while Cl– can accept an H+.

A major (and important difference) between the Brønsted-Lowry and Arrhenius acid base models is that a Brønsted-Lowry acid must always have an accompanying base to react with - the two are inseparable. A proton donor has to have something to donate the protons to (a base), and in this case it is water. Remember that bond breaking always requires energy, while bond formation always releases energy. A reaction in which the only thing that happens is the breaking of a bond (for example the Cl–H bond in HCl) will always requires the input of energy. Acid-base reactions are typically exothermic; they release energy to the surroundings) and that released energy is associated with the interaction between the H+ and the base. That is the proton does not drop off the acid and then bond with the base, instead the acid–H bond starts to break as the base–H bond starts to form. One way that we can visualize this process is to draw out the Lewis structures of the molecules involved and see how this proton is transferred.

In this representation we use a dotted line to show the growing attraction between the partial positive charge on the H of the H-Cl molecule and the partial negative charge on the oxygen. This interaction results in the destabilization of the H-Cl bond. Because the Cl is more electronegative than the H, the electrons of the original H-Cl bond remain with the Cl (which becomes Cl-), while H+ forms a new bond with a water molecule.

H₂O(l) + H₂O(l) ↔ H₃O+ (aq)         + –OH(aq)
Acid Base              conjugate acid + conjugate base

Again we can visualize this process by drawing out the Lewis structures of the water molecules to see how the proton is able to move from one water molecule to another - never being “alone” - always interacting with the lone pairs on the oxygens.

Question to answer:

• Which do you think is more likely to happen? The reaction H₂O + H₂O →?
• Or the reverse process H₃O+ + –OH → ?   Could they both happen at once?
  What do you think the relative amounts of H₂O, H₃O+ + –OH might be in a pure sample of liquid water? How would you measure the relative amounts?
• Now that you know HCl is an acid and ammonia is a base, can you predict the reaction that occurs between them?
• Is water a necessary component of a Bronsted-Lowry acid base reaction? how about for a Arrhenius acid base reaction?
  

Questions to ponder:

• Which theory is more useful? Why?
• Are both theories “correct”?

How to spot an acid

Moving on from water, can we predict whether a compound would be an acid, a base, or neither? As we have seen, many properties of materials can be predicted by considering their molecular structure.

It is in this context that Lewis structures are particularly useful. Let us first consider the strong acids: hydrochloric acid (HCl(aq)), nitric acid (HNO₃ (aq)), sulfuric acid (H₂SO₄ (aq)), hydrogen bromide (HBr(aq)), hydrogen iodide (HI (aq)). What structural feature do these substances have in common? Well, from their formulae it is clear that they all contain hydrogen, but there are many compounds that contain hydrogen that are not acidic. For example methane (CH₄) and other hydrocarbons are not acidic; they do not donate protons to other molecules. When acids are written in their normal form it can be very difficult to see any similarities, but if we draw out the Lewis structures some commonalities emerge.

One common feature of an acid is that the proton that gets donated (or picked off) is bonded to a highly electronegative atom, most generally an oxygen or a halogen (Cl, Br, I). So, once you know what to look for it is quite easy to spot the potentially acidic sites in a molecule - for example, in the figure above, you could circle the “vulnerable” Hs. This a useful skill (the ability to spot donate-able Hs) that allows you to predict properties of more complex molecules. But what exactly are the reasons why a hydrogen covalently bonded to an oxygen (or a halogen) is potentially acidic (donate-able)?

First, let us consider the O-H bond. Based on our discussion of water molecules, we predict that it will be polarized, with a partial positive charge on the H and as partial negative on the O. In water, the H is (on average) also part of a hydrogen bond (intermolecular interaction) to the oxygen of another water molecule. It turns out that it does not take much energy to break the original O–H bond. Remember the H+ does not just drop off the acid, but forms a bond with the base molecule. In fact strong acid-base reactions are typically exothermic, meaning that the new bond formed between the proton (H+) and the base is stronger than the bond that was broken to release the H+. The released energy raises the temperature of the surroundings. In an aqueous solution of a strong acid, hydrogens ions are moving rapidly and randomly from one oxygen to another; the energy for all this bond breaking coming from the thermal motion of water molecules.

We also have to consider what happens to the oxygen that gets left behind. When the acidic hydrogen is transferred, it leaves behind the electrons that were in the bond, giving that atom more electrons than it started with. The species left behind must be stable even with those extra electrons (the negative charge). In the example below, chloride ion Cl–(aq) is “left” when the proton gets transferred away. We know chloride is stable, and common. It is not surprising to find it as one of the products of the reaction.
HCl(g) + H₂O(l) ↔ H₃O+ (aq)         + Cl–(aq)
Acid        Base       conjugate acid + conjugate base

Now let us recall that the definition of electronegativity is a measure of the ability to attract (and retain) electrons. It therefore makes sense that a negatively-charged electronegative atom, like chlorine or oxygen, will be more stable than a negatively-charged, but less electronegative atom, such as carbon.

Question to answer:

• What other atoms besides chlorine or oxygen would be electronegative enough to stabilize those extra electrons?
• Draw out the reactions of each of the strong acids with water. (HCl(aq)), nitric acid (HNO₃ (aq)), sulfuric acid (H₂SO₄ (aq)), hydrogen bromide (HBr (aq)), hydrogen iodide (HI (aq)). What commonalities are there? What are the differences?
• Draw out the structures of methanol (CH₃OH), acetic acid (CH₃COOH) and methane (CH₄), and write a potential reaction with water. Label the conjugate acid - base pairs. Which reactions do you think are likely to occur? Why?
  

Questions for later:

• What other methods might be available to stabilize the electrons (recall that one model of bonding allows for molecular orbitals that extend over more than two atoms) - we will return to this idea later.

Question to answer:

• Draw out molecular level pictures of a dilute solution of a strong acid, and a weak acid.
• Draw out molecular level pictures of a concentrated solution of a strong acid, and a weak acid.
• What are the similarities and differences between all the representations you have drawn?
• If you need to dilute a concentrated solution of a strong acid with water, considering your knowledge of the energy changes associated with the reaction of an strong acid with water.
• Which makes more sense (from a safety point of view): should you -
  • Add water slowly (dropwise) to the concentrated, strong acid, or
  • Add the concentrated strong acid dropwise to water? Why?

How to spot a base

There is an equally simple method to figure out which compounds are potential bases. Let us take a look at some common bases. The first bases that most people encounter are the metal hydroxides such as NaOH, KOH, Mg(OH)₂ etc. The metal ions, generated when these compounds dissolve in water, typically do not play any role in acid base reactions. The base in these compounds is the hydroxide (–OH). Another common class of bases are molecules, like NH₃, that contain nitrogen. There many kinds of such “nitrogenous bases” some of which play an critical role in biological systems.

Oxygen holds onto its electron pairs a little less tightly, and so is more likely than F to donate a lone pair to a proton, but nitrogen, being even less electronegative has a lone pair that is more available, making most nitrogenous compounds basic.

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:

• Why did we not include CH₄ or Neon in this analysis?
• Do you think compounds with ammonium (NH₄+) are basic? why or why not?
• Can you draw the structure of a basic compound that has not been mentioned in the text yet?
• Draw out the reactions of CH₃NH₂ and CH₃OH with water. Label the conjugate acid and base pairs. Which do you think is most likely to occur? Why
• If you had to design an experiment to figure out whether a compound was an acid or a base (or both) how would you do it?
• What experimental evidence would you accept as evidence that you had an acid or a base?