CLASSES OF CHEMICAL REACTIONS

CHEMISTRY 1a, FALL, 2001

Introduction. This document was drafted with the intent of providing a systematic approach to the study of elementary inorganic reactions. We have considered four classes of reactions in General Chemistry:

Suppose that a pair of reagents is mixed. Our task is to predict the chemical outcome. In most cases, the ensuing reaction will represent more than one class of the four fundamental reactions. Theory might unambiguously to a single prediction which could be confirmed by an experimental observation. More often, theory will yield a small set of alternate explanations. Experimental observations are then required to select the best explanation from the finite set of answers provided by theory. Hence, the marriage of theory and experiment provides far more information than either theory or experiment alone.

Strategy. The strategy of predicting the likely outcome(s) from the combination of chemical reagents can be broken down into the following steps:

An Elementary Example. As an elementary example, let us apply the approach to the chemistry resulting from the mixing of aqueous solutions of acetic acid and sodium hydroxide.

Step 1) The covalently bonded acetic acid is soluble in water. It is a weak acid and hence other than water, it is the predominant species in its aqueous solution. The amounts of hydronium and acetate anion are small and need not be considered here. Ionic sodium hydroxide when dissolved yields sodium cations and hydroxide anions. Hence, our attention in the next step is drawn to three species: undissociated acetic acid, the sodium cation, and the hydroxide anion.

Step 2) First remove the sodium cation from consideration. It is a spectator ion. Acetic acid is a weak acid. The hydroxide anion is a non-special anion (its salts are generally insoluble unless the cation is special), a base, and a ligand (bases are also ligands since both require the ability to share electrons).

Step 3) There is a single match. The acid property of acetic acid matches the basic property of the hydroxide anion. The reaction is an acid-base reaction.

Step 4) The net ionic equation is CH₃COOH(aq) + OH^-(aq) ® CH3COO^-(aq) + H₂O(l). Note that the reaction is not written as H+(aq) + OH^-(aq) ® H₂O(l) as acetic acid is a weak acid and it, not hydronium ion, is the predominant acidic species in the solution.

Step 5) In this case, we have exhausted the possibilities. The products of the reaction will not react further with hydroxide anion. Other types of reactions will not occur. Although the hydroxide anion is a ligand, it will not participate in a complexation reaction as no transition metal is present. Hydroxide will not be precipitated as an insoluble ionic salt as no non-special cation is present.

Example of a Complicated Reaction. The above simple example was chosen to illustrate the application of the method. The method will now be applied to a set of illustrative but non-trivial examples that combine two or more of the classes of reactions. The first is the reaction of manganese(II) sulfide, with cold, dilute nitric acid.

Step 1) Neither the manganese(II) cation nor the sulfide anion are special. Hence the salt, MnS, is insoluble. However, it follows from the Second Law of Thermodynamics that no substance is completely insoluble. When a small amount of the salt does dissolve, it produces the Mn(II) and S^-2 ions. Nitric acid is a strong acid and dissociates nearly 100% in water to yield hydronium and nitrate ions. Four species-nitrate, hydronium, manganese(II), and sulfide-will be the focus of attention in the next step.

Step 2) The solution is cold. Under these conditions, the nitrate anion is inert so remove it from further consideration. The manganese(II) cation is a transition metal species and a non-special cation. The sulfide anion is the conjugate base of a weak acid, HS^-. Sulfur and oxygen are in the same family of the periodic table so one might conclude from the properties of the water (H₂O) system that HS^- is a fairly weak acid. Hence the sulfide anion would be a moderately strong base and also a ligand. The sulfide anion is electron rich and could function as a reducing agent. The sulfide anion is a non-special anion. Finally, the hydronium ion has one principal chemical identity; it is an acid.

Step 3) We'll begin our analysis with the hydronium ion since its spectrum of chemistry is narrow. As an acid, it matches with the base sulfide ion. An acid-base reaction is a very likely outcome. A new precipitation reaction is not an option; the only match of a non-special cation with a non-special anion, is between manganese(II) and sulfide, species produced when MnS dissolved. Reversing the process adds nothing new. The formation of a complex between the transition metal species Mn(II) and the sulfide anion is possible. However, the reaction between sulfide and hydronium ion will remove sulfide from the system. Furthermore, normal complexation calls for 6 ligands per transition-metal cation. However, the sole source of sulfur is the salt MnS which upon dissolution produces only one sulfide anion per manganese(II) cation. Complexation is not favored by this stoichiometry.

Step 4) We can divide the overall reaction into two steps:

Dissolution: MnS(s) ® Mn⁺²(aq) + S^-2(aq)

Acid-base: H⁺(aq) + S^-2(aq) ® HS^-(aq).

Step 5) We are not quite done. The HS- anion is also a base; it is the conjugate base of H₂S and will reaction with a second hydronium ion. The product, H₂S, is uncharged and relatively non-polar. Its solubility in water is limited. It a moderate amount is produced, some will saturate the solution and remainder will come out of solution. H₂S is a small molecule and in its pure, unsolvated form will be expected to be gaseous. Hence, the equation for the second step will be

H⁺(aq) + HS^-(aq) ® H₂S(g)

If we put all the pieces together, we end up with the net result

MnS(s) + 2 H⁺(aq) ® Mn⁺²(aq) + H₂S(g)

In other words, the nitric acid effects the dissolving of the insoluble manganese(II) sulfide by a cascade of event. As soon as a bit of the salt dissolves, the hydronium ions remove the sulfide ions from the solution. More salt dissolves and the sulfide produces is also removed. The process continues until the salt has completely dissolved.

Variations of a Theme. An overview of the net reaction is instructive. Manganese(II) sulfide is insoluble in part because of the strong electrostatic forces in the ionic crystal and the substance dissolves to a very limited extent. This process of dissolution can be nudged to the right with an increase in the concentration of Mn(II) in solution but removing the other product, sulfide. That is function of the hydronium ion. Some sulfides such as silver(I) sulfide are so insoluble that the ability of hydronium ion to reduce the concentration of sulfide ion is not sufficient to effect the dissolving of the sulfide. In such cases, a more potent reagent is required. Hot, concentrated nitric will do the trick. Under the conditions of elevated temperature, the potential of the nitrate anion as a very strong oxidizing agent is released. The half reaction for its reduction is

3 e^- + 4 H⁺(aq) + NO₃^-(aq) ® NO(g) + 2 H₂O(l).

Note the stoichiometric requirement for hydronium ions. The oxidizing agent nitrate anion can not effect an oxidation alone. A similar situation exists for the moderate oxidizing agent hydrogen peroxide. The half reaction in this case is

2 e^- + 2 H⁺(aq) + H₂O₂(aq) ® 2 H₂O(l).

Recall that acidified hydrogen peroxide is one of the reagents in the Five Anion Experiment. The above half reaction shows that it normally functions as an oxidizing agent.

Let us return to the dissolution of silver(I) sulfide in hot, concentrated nitric acid. The nitrate anion oxidizes the sulfide to elemental sulfur.

S^-2(aq) ® S(s) + 2 e^-

The sequence of elementary reactions in the dissolution of silver(I) sulfide is therefore

Ag₂S(s) ® 2 Ag⁺(aq) + S^-2(aq)

3 e^- + 4 H⁺(aq) + NO₃^-(aq) ® NO(g) + 2 H₂O(l)

S^-2(aq) ® S(s) + 2 e^-

If these three steps are combined so that electronic charge is conserved, one obtains for the net reaction

3 Ag₂S(s) + 8 H⁺(aq) + 2 NO₃^-(aq) ® 6 Ag⁺(aq) + 3 S(s) + 2 NO(g) + 4 H₂O(l).

One should perform this chemical operation in a hood. NO is toxic. Furthermore, it reacts with oxygen to form NO₂, an orange- brown toxic substance that is responsible for the color of Los Angeles smog.

The analysis performed to understanding the reaction of manganese(II) sulfide with dilute, cold aqueous nitric acid suggests another means of dissolving the ionic salt. The sulfide anion was identified as a ligand. A complexation reaction was ruled out on stoichiometric grounds. The salt does not provide enough sulfide. However, one can favor the complexation reaction by replacing the nitric acid with an excess of a soluble sulfide salt such as sodium sulfide. In the presence of an excess of sulfide, complexation of the manganese(II) cation would now be possible. As a first and quite reasonable one might consider a coordination number of six. That is, one manganese(II) cation would combine with six sulfide ligands to form the complex MnS₆^-10, a species with six metal-ligand bonds. It turns out that the actual coordination number is 4, the second most common coordination number and the actual overall reaction is

MnS(s) + 3 S^-2(aq) ® MnS₄^-6(aq).

There is no simple explanation for this turn of events although the high negative charge on the complex with a coordination number of 6 is surely a factor. However, it should be noted that the coordination number of 4 is often observed with sulfide and sulfide-like species as ligands. The most important illustration of this phenomenon is the class of proteins containing metals where the amino acid cysteine serves as a ligand. Typically, 4 cysteines coordinate with a transition metal ion such as iron(II) or zinc(II).

An Example of Kitchen Chemistry. In this section, we'll apply the strategy to the design of a chemical solution to the accumulation of solid calcium carbonate in the shower. Water in the Pomona Valley is derived from two sources. The first is ground water that comes from the San Gabriel Mountains. A dominant rock type in the mountains is granite and the water is relatively soft. That is, the calcium ion concentration is not particularly high. There is a problem with the water. The area once had extensive groves of citrus trees that were heavily fertilized with ammonia. The ammonia not utilized by the trees washed into the ground water where it was converted by bacteria to nitrate and nitrite. Infants are born with a special form of hemoglobin called fetal hemoglobin that is replaced with adult hemoglobin as the infant grows. Unfortunately, nitrate reduced the oxygen-binding capacity of fetal hemoglobin. Consequently, nitrate-rich water from the local wells must be diluted with nitrate-lean water from the Colorado River. The Colorado River flows through sedimentary formations and a dominant rock is limestone that is principally the mineral calcite, i.e. calcium carbonate. Colorado River water therefore ends up becoming close to a saturated solution of calcium carbonate; it is hard. In the home, the calcium cations often end up as solid deposits of calcium carbonate in the plumbing. A prominent example in my bathroom is the nozzle in the shower. It is so encrusted with the solid deposit that its effectiveness is reduced. Our task is to discover a means of removing the deposit.

To this end, consider the properties of the species produced when a small amount of calcium carbonate does dissolve. The calcium cation does not possess a particularly rich chemistry. It is not a special cation so many of its salts would be expected to be insoluble. This aspect of its chemistry frustrates the development of an effective strategy. It is not a transition-metal cation and complexation would not be a prominent aspect of its chemistry. However, all metal cations can be induced to form a complex if the ligand is particularly effective. An unusually effective ligand is EDTA^-4, an octopus-like ligand that has 6 binding sites forms 6 metal-ligand bonds per molecule. Complexation might therefore be a possible remedy.

The carbonate anion is a base and consequently also a ligand. An acid could react with the carbonate anion to form monohydrogen carbonate. If the acid is sufficiently strong, it could also transfer a proton to the monohydrogen carbonate anion to form carbonic acid, H₂CO₃. Carbonic acid has a characteristic that is desirable for our purposes. It is unstable and decomposes to form gaseous carbon dioxide and water.

H₂CO₃(aq) ® CO₂(g) + H₂O(l)

The decomposition removes carbonate from the system and the reaction is nudged to the right. This decomposition reaction is the origin of the fizz in Alka Selzer, a mixture of an acid such as citric acid and sodium monohydrogen carbonate. It should also be noted that sulfurous acid, H₂SO₃, also is unstable with respect to decomposition to sulfur dioxide and water.

The carbonate chemistry suggests a possible solution. Remove the nozzle and place it in a ceramic pan filled with vinegar, an aqueous solution of acetic acid. The vinegar could be heated to accelerate the process and enhance the escape of the carbon dioxide. Small, molecules such as carbon dioxide are less soluble in cold water than hot water. The net reaction for our clean up would be

2 CH₃COOH(aq) + CaCO₃(s) ® Ca⁺²(aq) + CO₂(g) + 2 CH₃COO^-(aq) + H₂O(l).

I would recommend not using an aluminum sauce pan for this purpose. Aluminum is a very reactive metal and is a potent reducing agent. Under normal conditions, it is inert as the top layer of aluminum atoms in a piece of aluminum is oxidized by the oxygen in the air to form a protective oxide layer. Neither aluminum(III) nor oxide are special ions. Furthermore, they are highly charged. Consequently, aluminum oxide is very insoluble. However, under acidic condition, the oxide layer will dissolve, exposing a fresh surface of aluminum for further attack. The basis for the enhanced solubility of the oxide is the reaction of the strongly basic oxide with two moles of acid to form water. The first proton transfer yields hydroxide which itself is basic and hence effects a second proton transfer. The net ionic equation for this deleterious reaction is

Al₂O₃(s) + 6 CH₃COOH(aq) ® 2 Al⁺³(aq) + 3 H₂O(l) + 6 CH₃COO^-(aq)

One might also using a EDTA to remove the deposit. EDTA is normally sold as the salt Na₂H₂EDTA· H₂O. In order to produce the required potent ligand, EDTA^-4 , two protons must be removed. Household ammonia, a weak base, should suffice.

H₂EDTA^-2(aq) + 2 NH₃(aq) ® EDTA^-4(aq) + 2 NH₄⁺(aq)

The net reaction for the complexation would be

CaCO3(s) + H₂EDTA^-2(aq) + 2 NH₃(aq) ® EDTA^-4(aq) + 2 NH₄⁺(aq) + Ca⁺²(aq) + CO₃^-2(aq)

A double-pronged attack will not work in this case. That is, the addition of the EDTA salt will reduce the concentration of the acetic acid. It is basic and will react with the acid. Furthermore, the acidic conditions will yield the formation of H₄EDTA, which is not a ligand, rather than the desired EDTA^-4.

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7 October 2001, WES