Reactions : Types of Reactions

At first glance, the many reactions that compounds and elements can undergo are overwhelming. How can such a variety be remembered or characterized? Fortunately, many of the important chemical reactions belong to one of four main types: (a) ion-combination reactions, also called metathetical or precipitation reactions, (b) proton-transfer reactions, also called Lowry-Bronsted acid-base reactions, (c) electron-sharing reactions, also known as Lewis acid-base reactions, and (d) electron-transfer reactions, more commonly known as oxidation-reduction reactions.

Ion-combination Reactions

Although both ionic and covalent compounds can undergo the latter three types of reactions, ion-combination reactions are usually confined to ionic compounds. This is also the simplest type of reaction. A simple ion-combination reaction takes place when a solution of silver nitrate is mixed with a solution of sodium chloride. Both solutions are colorless and clear, but when one is added to the other, a white precipitate forms. When subjected to chemical analysis, the white material can be shown to be silver chloride. Moreover, the solution that remains, which is called the filtrate, is a solution of sodium nitrate. The equation for the reaction therefore is:

NaCl + AgNO₃  AgCl + NaNO₃

This reaction comes to equilibrium, which is why we use double arrows.

Actually, it is more informative to write this equation in several other ways. First, we can show how the compounds exist in aqueous solution.

Na⁺(aq) + Cl^-(aq) + Ag⁺(aq) + NO₃^-(aq)  AgCl(s) + Na⁺(aq) + NO₃^-(aq)

Because both reactants are ionic, they enter the water solution as ions. The silver chloride is also ionic, but it is denoted with a subscript (s) to indicate that it exists primarily as a solid, even though some silver ions and chloride ions remain in solution when the reaction is complete.

We can also convert this ionic equation to a net ionic equation, which eliminates ions that really do not participate in the reaction. A look at the equation above shows that the sodium ions and nitrate ions appear in exactly the same form on both sides of the equation. These ions can be cancelled in the same way that x in the following equation can be cancelled:

10 + x = x + y

These ions are sometimes referred to as spectator ions because, in an anthropomorphic sense, they just "look on" while the silver ion and chloride ion participate in the reaction. After we cancel the sodium and nitrate ions, we obtain the net ionic equation:

Ag⁺(aq) + Cl^-(aq)  AgCl(s)

This reaction has a high extent with an equilibrium constant of 10¹⁰. It is the kind of reaction that is highly desirable for the quantitative analysis of chloride ion. For example, suppose that a qualitiy control chemist is analyzing batches of the herbicide 2,4-D. This compound, 2,4-dichlorophenoxyacetic acid, contains two chlorine atoms per molecule. If the compound is treated with sodium hydroxide under the proper conditions, the hydroxide ions will displace the chlorines as chloride ions. If silver nitrate is then added to the solution, the chloride will undergo an ion combination reation with the silver to form insoluble silver chloride. Because the extent of this reaction is very high, the analyst can assume that all of the chloride in the 2,4-D has been converted to silver chloride. She can then weigh the precipitate, determine the number of moles of silver chloride, divide that number by two to get the number of moles of 2,4-D, and finally convert moles to grams. If a percentage is desired, the grams of 2,4-D must be divided by the original sample weight. This entire procedure is called a gravimetric analysis because the final step involves weighing the precipitate. Currently there are more sophisticated spectrometric ways to analyze materials, but much quality control work still relies on classical methods such as this because of the ease of obtaining the result and the low cost of the reagents.

The analysis could also be carried out by volumetric analysis. A buret, shown in Figure 59, is filled with a solution of known concentration of silver nitrate while the sample to be analyzed is dissolved in water and placed in an Erlenmeyer flask. An indicator, which is an organic compound that will change color when all of the chloride ion has reacted, is also added to the flask. The silver nitrate solution is then added slowly to the flask. As the silver ions react with the chloride ions, a white precipitate of silver chloride is formed. When sufficient silver has been added to precipitate the chloride ion, the next drop of silver nitrate (excess AgNO₃) reacts with the indicator and the color of the solution changes. The analytical chemist can then use the number of milliliters of titrant to determine how many moles of silver ion were added. This is the same as the number of moles of chloride ion present in the original sample. Consequently, the percent chloride can be calculated.

Figure 59. Some volumetric glassware.

Proton-transfer Reactions (Lowry-Bronsted Acid-Base Reactions)

Acids, bases, and salts were among the earliest known categories of matter. The alchemists, for example, knew that acids and bases "killed" each other and that the result was a salt. Lavoisier, one of the most famous chemists of the 18th century, believed that oxygen (which had been discovered by Priestly) was what gave acids their acidic properties and, in fact, gave oxygen its name (oxygen means "acid former"). One of Lavoisier's students, Davy, discovered that hydrochloric acid contained no oxygen, and eventually convinced the chemical community that hydrogen, not oxygen, was the element that conferred acidic properties. It is interesting that Lavoisier continued to defend his oxygen theory for many years.

In the 19th century, Arrhenius concluded that acids produce hydrogen ions in water and bases produce hydroxide ions in water. Because this conclusion, although correct, limits the definition of acids and bases to the solvent water, a number of other definitions have been proposed that make the concept of acids and bases more general. One of the most important definitions of acids and bases was proposed simultaneously by Lowry and Bronsted: an acid is a species that donates a proton (hydrogen ion), while a base is a species that accepts a proton. This definition encompasses reactions carried out in water but also permits the use of other solvents. The Lowry-Bronsted definition makes it clear that acid-base reactions are proton transfer reactions. Consider, for example, the reaction of acetic acid with water (acetic acid is the main ingredient in vinegar):

CH₃CO₂H + H₂O  CH₃CO₂^- + H₃O⁺

The hydrogen attached to the very electronegative oxygen is donated to the water, leaving the acetate ion (CH₃CO₂^-) and protonated water--the hydronium ion (H₃O⁺). Notice that there are two other hydrogens in the reactants: the three attached to the carbon and the two attached to the oxygen of water. Any of these hydrogens could be donated, but generally the hydrogen ion comes from the site that has the least amount of electron density-- in this case, the COOH group. Hence, it is fair to conclude that the acid must have a hydrogen attached to an electonegative site. The base must have a site of high electron density--a pair of electrons.

Now consider the reverse reaction; that is, the reaction of hydronium ion with the acetate ion:

CH₃CO₂^- + H₃O⁺  CH₃CO₂H + H₂O

The hydronium ion contains three hydrogens attached to an electronegative oxygen and can therefore function as an acid. The acetate ion has a number of non-bonded electron pairs, as shown in the following electron-dot formula:

Thus, a proton transfer can also occur in the reverse reaction. Indeed, this reaction is more favorable than the forward reaction. As a result, the reaction comes to an equilibrium that lies on the side of the reactants. To put it another way, there are more transfers of protons from the hydronium ion to the acetate ion than there are transfers from acetic acid to water.

Another way to summarize the extent of the reaction is to say that the equilibrium constant is low--actually about 10^-5. In a 0.1 M solution, only about 1% of the acetic acid molecules have dissociated at any given moment. It is for this reason that acetic acid is said to be a weak acid. Hydrochloric acid, on the other hand, is a strong acid; about 99% of the HCl molecules are dissociated into chloride ions and hydronium ions at any given moment. The equation below indicates that the reaction of HCl with water comes to equilibrium, but the equilibrium constant is high--about 10⁵.

HCl + H₂O  H₃O⁺ + Cl^-

Most common acids are weak acids and do not produce much hydronium ion when they react with water. If they react with a stronger base, such as ammonia, the extent of the reaction is greater. The common strong acids are nitric acid, sulfuric acid, perchloric acid, hydrochloric acid, hydrobromic acid, and hydroiodic acid.

Because hydronium ion is produced by acids when they react with water, a special symbol (pH) is used to designate the concentration of this species. pH is the negative logarithim of the hydronium ion concentration:

pH = - log [H₃O⁺]

In ordinary water, the water molecules react with themselves to form some hydronium ion and an equal amount of hydroxide ion according to the equation:

H₂O + H₂O  H₃O⁺ + OH^-

The equilibrium constant for this reaction is low (10^-14), and the concentration of hydronium ion is also low--1 x 10^-7. The negative logarithim of 1 x 10^-7 is 7.0 and therefore the pH of a neutral solution is 7.0. If the solution contains excess hydronium ions, it has a pH below 7.0 and is said to be acidic. If the solution contains excess hydroxide ions, the pH is above 7.0, and the solution is said to be basic.

When a weak acid is mixed with a substantial amount of its conjugate base (the anion that results when a proton is removed from the acid), a buffer solution is established. For example, if 0.1 mole of acetic acid and 0.1 mole of sodium acetate are added to a liter of water, the resulting solution will have a pH of 5.0. If a small amount of acid or base are added to this solution, the pH will remain at about 5.0. Clearly, something in the solution negates the usual effect of acid or base; this solution maintains a constant pH and is said to be a buffer solution. To demonstrate how dramatic this effect is, consider the addition of 0.01 moles of HCl (a strong acid) to one liter of water. The HCl reacts with the water and produces 0.01 mole of hydronium ion (remember that the reaction of a strong acid with water has a high extent of reaction; that is, essentially all of the HCl is converted to hydronium ion and chloride ion). The pH of the water has therefore changed from 7.0 to 2.0. If 0.01 mole of HCl is added to our acetic acid/sodium acetate buffer, the pH will not change by more than several tenths of a pH unit.

How does the solution maintain a constant pH? The HCl reacts with the sodium acetate in the buffer solution and forms acetic acid as shown below:

HCl + CH₃CO₂^-  CH₃CO₂H + Cl^-

The buffer solution already contains acetic acid, and consequently this reaction only changes the relative amounts of acetic acid and sodium acetate by a small amount. It is this ratio that determines the pH of the solution.

When HCl is added to pure water, the HCl reacts with water to form hydronium ions. Clearly, it is the presence of the acetate ion that prevents the HCl from reacting directly with water molecules. When hydroxide ions are added to our buffer solution, they react with the acetic acid and convert a small amount of acetic acid to acetate ion:

OH^- + CH₃CO₂H  CH₃CO₂^- + H₂O

Again, the ratio of acetic acid to sodium acetate is very little affected.

Buffer solutions can also be prepared from a weak base and its conjugate acid such as NH₃ with NH₄Cl or CO₃^2- with HCO₃^-. Buffer solutions containing the carbonate ion, the hydrogen carbonate ion, carbonic acid and dissolved carbon dioxide are extremely important. These four species are in equilibrium as shown below:

CO₂ + H₂O  H₂CO₃
H₂CO₃ + H₂O  HCO₃^- + H₃O⁺
HCO₃^- + H₂O  CO₃^2- + H₃O⁺

Notice that carbon dioxide, produced in the respiratory process of animals, reacts with water to form carbonic acid, which in turn behaves as a normal weak acid and dissociates into the hydrogen carbonate ion, which, in turn, dissociates to carbonate ion.

This system is responsible for maintaining the pH of blood at 7.35 - 7.45. A change from this value by more than a couple of tenths of a unit is sufficient to cause death. This system is also responsible for the constant pH of the oceans. Carbonate sediments, coral reefs, and even the shells of mollusks (all consist primarily of calcium carbonate) react with acidic pollutants to produce hydrogen carbonate and carbonic acid.

CO₃^2- + H₃O⁺  HCO₃^- + H₂O
HCO₃^- + H₃O⁺  H₂CO₃ + H₂O
H₂CO₃  CO₂ + H₂O

Added base reacts with the carbonic acid to form hydrogen carbonate ion and carbonate ion.

Electron-sharing Reaction (Lewis Acid-base Reactions)

Electron-sharing reactions are better known as Lewis acid-base reactions. This definition of acids and bases was proposed by Lewis in the 1920s. It is more general than the Lowry-Bronsted definition and is extremely important in understanding the myriad of reactions embodied in organic chemistry. According to Lewis, an acid is a species that can accept a pair of electrons and a base is a species that can donate a pair of electrons. The species that is formed when the acid and base react is an adduct in which the electrons provided by the base are shared between both the acid and base. The reaction of a proton with the fluoride ion provides a simple example.

H⁺ + F^- → H-F

The hydrogen ion (proton) has no electrons and is therefore eager to accept electron density from the base, which has three lone pairs of electrons. The adduct allows the pair to be shared between the hydrogen and fluorine. This example should also demonstrate that Lowry-Bronsted acid-base reactions are a subset of Lewis reactions. The reaction of protons with fluoride ions is also a proton-transfer (Lowry-Bronsted) reaction. All Lewis bases have at least one lone pair of electrons and are therefore also Lowry-Bronsted bases. Lewis acids, on the other hand, do not necessarily contain a hydrogen. Other species that can function as Lewis acids include cations, such as Be²⁺, Ag⁺, and Ni²⁺, neutral molecules that can accept a pair of electrons, such as BCl₃, and organic molecules that can rearrange their electron density to accommodate a pair of electrons. For example, the famous Grignard reaction used to make alcohols involves the attack of the base CH₃^- on the carbonyl (C=O) group of ketones or aldehydes.

Electron-transfer Reactions (Oxidation-reduction Reactions)

The last category of reactions is electron-transfer or oxidation-reduction reactions. Oxidation is defined as the loss of electrons, while reduction is the gain of electrons. In the equation below,

Zn + Cu²⁺  Zn²⁺ + Cu

zinc gives up electrons that are acquired by copper ions. The zinc is therefore oxidized and the copper ions are reduced. Each process can be written as a half-reaction:

Zn  Zn²⁺ + 2 e^- (oxidation)
Cu²⁺ + 2 e^-  Cu (reduction)

Because the copper ions are essential in order for the oxidation of the zinc to occur, the copper ions are collectively called the oxidizing agent. Likewise, zinc metal must give up its electrons in order to cause the reduction of the copper ions. Therefore, zinc is the reducing agent in this reaction.

It is sometimes difficult to determine whether a particular reaction is an electron-transfer reaction. In these cases, it is helpful to assign oxidation numbers to all of the elements in the equation. Oxidation numbers are assigned by following a somewhat arbitrary set of rules, but the method gives an approximate idea of the excess or deficiency of charge on each element. Positive oxidation numbers indicate a deficiency of charge, and negative oxidation numbers indicate excess charge. Oxidation is accompanied by an increase in the oxidation number for a particular element. To illustrate, let us consider the reaction of potassium dichromate with potassium chloride to form hypochlorous acid and chromium(III) chloride. The unbalanced equation is given below:

K₂Cr₂O₇ + KCl  HClO + CrCl₃ + KCl

+1  +6  -1		+1  -1		+1  +1  -2		+3  -1		+1  -1
K2Cr2O7	+	KCl		HClO	+	CrCl3	+	KCl

If we now examine each element, we can see that the oxidation number of potassium on both sides of the equation is the same. The same is true for oxygen. However, chromium has a +6 oxidation number on the left and a +3 on the right side. This decrease from +6 to +3 indicates that the chromium is being reduced. In fact, each chromium atom is being fed three electrons during the reaction. Of course, if we have a reduction, we must have an oxidation; we cannot have one without the other. It is chlorine that is being oxidized: the oxidation number on the left is -1, and on the right it is +1 in the hypochlorous acid. This increase in oxidation number represents an oxidation, or loss of electrons.

We also notice that we can get rid of some of the ions, such as the potassium and chloride ions on the right. These are spectator ions. We must be careful as we cancel species to make sure that we know how each compound exists in aqueous solution--the HClO is molecular and does not exist as ions, so we must write it as HClO, whereas K₂Cr₂O₇ can be written in its ionic form--K⁺ Cr₂O₇^2-. The net ionic equation then is

Cr₂O₇^2- + Cl^-  Cr³⁺ + HOCl

In order to balance the equation, it is necessary to add water and hydrogen ions, which finally yields the equation

11 H⁺ + Cr₂O₇^2- + 3 Cl^-  2 Cr³⁺ + 3 HOCl + 4 H₂O

The fact that electrons are transferred from one species to another leads us to suspect that we can separate the oxidation from the reduction and force the electrons to migrate from one process to another through a wire. This is indeed what occurs in a battery or voltaic cell. For example, the "cell," shown in Figure 60, transfers electrons through the wire from the zinc strip to the beaker containing the copper ions. A salt bridge is necessary in order to allow ions to migrate from one side to another. This migration balances the charge that builds up in each half-cell. This particular cell provides a voltage of 1.1 V.

Figure 60. A voltaic cell.
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The voltage that is created in a cell depends on the nature of the species and their concentrations. In a lead storage cell (a car battery), there are six cells that together generate 12 V. The reaction that occurs is

Pb + PbO₂ + 4 H⁺ + 2 SO₄^2-  2 PbSO₄ + 2 H₂O

As current is drawn from the battery, lead and lead oxide are converted into lead sulfate, and sulfuric acid is consumed. The battery can be recharged by supplying electricity (from the alternator), which forces the reverse reaction

2 PbSO₄ + 2 H₂O  Pb + PbO₂ + 4 H⁺ + 2 SO₄^2-

to occur and the original state of the battery is regenerated.

In fact, many reactions that do not occur (the standard free energy change is positive) can be forced to proceed by literally jamming electrons down the throats of the reactants. This process is called electrolysis. The commercial preparation of many of the elements utilizes electrolysis. Sodium metal is prepared by sending a current through a molten mass of sodium chloride. As shown in Figure 61, the sodium ions migrate to the negative electrode and there receive electrons. Thus, the sodium ions are reduced to the metal.

Na⁺ + e^-  Na

At the other electrode, chloride ions (which migrated there because of their attraction to the positive electrode) are converted into diatomic chlorine gas.

Figure 61. An electrolytic cell.

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