In the Classroom
Systems of Chemical Equations as Reasonable Reaction Mechanisms Sergey V. Dorozhkin† Research Institute of Fertilizers and Insectofungicides, Kudrinskaja sq. 1-155, 123242 Moscow D-242, Russia;
[email protected] From my teaching experience, there is a great difference between an ability to write chemical equations and an ability to understand chemical transformations (reaction routes) at the molecular, atomic, or ionic level. For this reason, in lessons on classical inorganic chemistry, I always start with a brief lecture that can be summarized as follows. I discuss six general principles of chemical reactions: 1. Any reaction occurs as a result of collisions among particles (molecules, atoms, ions, or radicals); 2. The probability of simultaneous collisions of more than three particles is equal to zero; 3. Collisions among similarly charged particles or ions are impossible owing to charge repulsion; 4. Any chemical reaction occurs via formation of some intermediates; 5. In most cases, the intermediates can be described as chemical compounds; 6. In most cases, formation of the intermediates can be written using chemical equations.
After the above introduction, I start writing examples of chemical reactions according to the following logic. First, the net equation of a chemical transformation is written. Second, this equation is balanced. To do so, both the classical balancing techniques (reviewed in refs 1 and 2) and the algebraic method (described in refs 3 and 4) can be used. According to the classification introduced by Johnstone (5), these two steps belong to the symbolic level of chemistry teaching. As shown later by Lee (6 ), the symbolic level is not enough for students to understand the chemical reactions and reaction routes: the particle method, describing atoms as beads and molecules as groups of beads, should be used. However, from my experience, systems of chemical equations are also able to provide reasonable suggestions on the reaction routes. Here are some examples from inorganic chemistry. Reasonable Reaction Mechanisms among Simple Inorganic Compounds
A Solid–Liquid Heterogeneous Interaction Chemical equations of interaction between solid hydroxyapatite (chemical formula Ca5(PO4)3OH) and an acid (e.g., HCl) are usually written as follows: Ca5(PO4)3OH + 4HCl = 2CaCl2 + 3CaHPO4 + H2O (1) 2Ca5(PO4)3OH + 14HCl = 7CaCl2 + 3Ca(H2PO4)2 + H2O (2) Ca5(PO4)3OH + 10HCl = 5CaCl2 + 3H3PO4 + H2O (3) † Temporary address until April 2002: Festkörperchemie, Lehrstuhl für Anorganische Chemie I, Fakultät für Chemie, RuhrUniversität Bochum, D-44780 Bochum, Germany.
Initially I explain to my students that the products of this chemical interaction and, therefore, the correct equation depend on amounts of the reagents mixed. In the case where hydroxyapatite is a limiting reagent (there is an excess of acid in the system) the chemical interaction follows eq 3. In the opposite case (an excess of hydroxyapatite), it follows eq 1. Finally, the chemical interaction follows eq 2 when a mixture contains seven moles of HCl per mole of hydroxyapatite. What can be said about chemical eqs 1–3? First, I remind the students that none of them gives a mechanism but merely expresses the net reaction. Second, I point out that single molecules of solid hydroxyapatite are mentioned in left-hand side of these equations. I mention that the presence of single molecules, atoms, ions, and radicals is reasonable for chemical interactions among gases and more or less reasonable for chemical reactions in solutions (especially in the cases when a solvating effect is omitted for simplicity). However, in most cases it is impossible to separate single molecules within the crystal lattice of solids. Third, according to eqs 1–3, four, seven, or even ten molecules of HCl simultaneously interact with each “single molecule” of hydroxyapatite. The latter contradicts the second general principle from the list above. In other words, the molecularity of these reactions is too high. Finally, I remind them of principles 4–6 and start creating a reasonable system of chemical equations. First, I assume that “single molecules” of hydroxyapatite in the solid state are equal to the unit cells: each unit cell is a “single molecule”. According to the second principle, molecules of HCl should interact one after another. For example (the hydration effect is everywhere omitted for simplicity): Ca5(PO4)3OH + HCl = Ca5(PO4)3Cl + H2O
(4)
Initially, HCl (an acid) interacts with hydroxide (a base). This results in replacement of hydroxide by chloride and transformation of “single molecules” of hydroxyapatite into those of chlorapatite. Chlorapatite is a first reasonable intermediate of this interaction. Further, I suggest that water-soluble calcium chloride might be dissolved from the surface: 2Ca5(PO4)3Cl = 3Ca3(PO4)2 + CaCl2
(5)
This results in transformation of chlorapatite into calcium phosphate tribasic (a second intermediate). Then, the latter interacts with the acid: Ca3(PO4)2 + 2HCl = 2CaHPO4 + CaCl2
(6)
Here two molecules of HCl (one after another) interact with calcium phosphate tribasic causing its transformation into calcium hydrophosphate dibasic (a third intermediate) with simultaneous detachment of one molecule of calcium chloride. Equation 6 is the final point of this chemical interaction in the case of an excess of hydroxyapatite (eq 1). However, in the
JChemEd.chem.wisc.edu • Vol. 78 No. 7 July 2001 • Journal of Chemical Education
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In the Classroom
case of an excess of HCl (eq 3), the chemical transformation goes further: 2CaHPO4 + 2HCl = Ca(H2PO4)2 + CaCl2
(7)
Calcium hydrophosphate monobasic (a fourth intermediate) has been obtained as a result. Finally, the last intermediate interacts with two molecules of HCl: Ca(H2PO4)2 + 2HCl = 2H3PO4 + CaCl2
(8)
Thus a system of chemical equations (4–8) has been created. It provides an opportunity to show a reasonable reaction route or mechanism. For better comparison with the particle method described by Lee (6 ), a schematic drawing of this system is shown in Figure 1. To conclude, it is necessary to draw the students’ attention to the fact that, in spite of the differences in net equations 1–3, the initial stages of the reaction route (eqs 4–6) are always the same. This clarifies why the products of a chemical transformation might depend on the amounts of reagents mixed.
A Homogeneous Reaction in Aqueous Solution Before starting explanations on reaction mechanisms in aqueous solutions, I always remind the students that when the influence of the solvating effect is omitted for simplicity, it is possible to operate with single atoms, ions, and molecules rather easily. Later I write examples of chemical equations and begin to construct reasonable reaction mechanisms. For example, the chemical equation for the hydrolysis of SiF62᎑ ion in basic aqueous solutions looks as follows: SiF62᎑ + 4OH᎑ = 6F ᎑ + SiO2 + 2H2O
(9)
Metal fluoride (e.g. NaF, KF) and SiO2 gel are obtained as a result. In this case, I explain to my students that the direct reaction (eq 9) is impossible owing to the second and third principles. Therefore, eq 9 is the net reaction only. To provide chemical interaction between SiF62᎑ and OH᎑, I ask the students to propose a reasonable way for these ions to overcome this mutual charge repulsion. I also add that the simplest way to do so is to remove charge from one of these ions. Usually, the most discerning students suggest dissociation of SiF62᎑ ions: SiF62᎑ = SiF5᎑ + F ᎑ ᎑
SiF5 = SiF4 + F
᎑
(10) (11)
Both reactions are reasonable, because the equilibrium SiF4 + 2HF
H2SiF6
(12)
was found in aqueous medium (7). Moreover, water solutions of pure H2SiF6 were found to consist of a mixture of H2SiF6, HSiF5, SiF4⭈2H2O, SiF3OH⭈2H2O, and HF (8). After their formation, uncharged molecules of SiF4 can interact with one hydroxide: SiF4 + OH᎑ = SiF4OH᎑
(13)
After that, one fluoride leaves: SiF4OH᎑ = SiF3OH + F ᎑ 918
(14)
Figure 1. Diagrammatic representation of chemical interaction between hydroxyapatite and HCl. Intermediates are shown in italic and ions participating in chemical reactions are shown in boldface. “Single molecules” of hydroxyapatite and all the intermediates are shadowed in gray. Hydration effect is omitted everywhere.
Journal of Chemical Education • Vol. 78 No. 7 July 2001 • JChemEd.chem.wisc.edu
In the Classroom 2−
F Si
F F
eq 10
F
−
F
F F
F
Si
F
Further chemical equations can be proposed in a similar way:
F eq 11
Si
F
eq 13
F
SiF3OH + OH ᎑ = SiF3(OH)2᎑ SiF3(OH)2᎑
F
F F
(15)
᎑
᎑
= SiF2(OH)2 + F = SiF2O + H2O + F (16)
SiF2O + OH ᎑ = SiF2OOH ᎑ −
F F
F
Si
F eq 14
F OH
Si
F
F OH
eq 15
F
F
Si
F
eq 17
SiFOOH + OH = SiFO(OH)2 SiFO(OH)2᎑
−
eq 18
F
O
OH
O
OH
−
OH
Si
Si
F
eq 19
F
HO
O
Si
OH eq 20
F
O
O eq 20
Si
HO
(18)
᎑
OH
OH F
eq 16
F
(17)
SiF2OOH ᎑ = SiFOOH + F ᎑ eq 16
F
Si
OH
OH Si
F
−
O
Si
(gel)
O
Figure 2. Diagrammatic representation of hydrolysis of hexafluorosilicate anion in basic aqueous solutions. Hydration effect is omitted.
᎑
᎑
(19) ᎑
= SiO(OH)2 + F = SiO2 + H2O + F (20)
Finally, a single-stage chemical equation (eq 9) has been transformed into the system of ten equations (10, 11, 13–20). As in the previous example, I also draw the students’ attention to the case of excess of SiF 62᎑ in the system: the chemical transformation might finish, say, in eq 16. However, it has no influence on the initial stages of the reaction route. Moreover, after addition of the extra amount of a basic compound, the chemical transformation proceeds until eq 20 is reached. A diagrammatic representation of this example is shown in Figure 2.
A Gas–Liquid Redox Interaction Redox reactions in aqueous solutions seem to be the most popular in teaching the classical inorganic chemistry. However, systems of chemical equations describing reasonable reaction mechanisms can be also created for the redox transformations. For example, let us consider the following interaction: 3Cl2 + 6KOH(T > 60 °C) = KClO3 + 5KCl + 3H2O (21) The same equation written in the ionic form would be
Cl
Cl
+
−
Cl
O
H
+
Cl
Cl
+
Cl
Cl
Cl
O
H
H
O
−
O
Cl
H
Cl
O−
O−
O
H
O
Cl
O
Cl
Cl
Cl
O
−
H
+
Cl
O−
+
H
O
+
Cl −
+
Cl −
H
O
(23)
O
(24)
H
(25)
(26)
3Cl2 + 6OH ᎑ = ClO3᎑ + 5Cl ᎑ + 3H2O
(22)
As described above, I begin by explaining to the students that both “normal” chemical equation (eq 21) and that written in the ionic form (eq 22) describe the net reaction only and do not imply any reaction route or mechanism. Later, based on the six general principles from above, I start creating a reasonable mechanism. First, chlorine interacts with hydroxide: Cl2 + OH ᎑ = HClO + Cl ᎑
(23)
Later, there is a simple acid–base interaction: H
O−
+
H
O
H
+
−
O
Cl
O
(27)
HClO + OH ᎑ = ClO᎑ + H2O
(24)
Further studies are quite similar: Cl −
O
Cl
Cl Cl
O
O
Cl
O
+
Cl −
(28)
Cl2 + ClO᎑ = Cl2O + Cl ᎑
(25)
᎑
Cl2O + OH = HClO2 + Cl ᎑
H
Cl O
᎑
HClO2 + OH = ClO2 + H2O
Cl O
᎑
O
O
−
H
Cl
O
+
Cl −
(29)
᎑
ClO2 + Cl2 = Cl2O2 + Cl ᎑
᎑
Cl2O2 + OH = HClO3 + Cl
O
᎑
H
O−
+
Cl
H
O
O
O
Cl O−
O
+
H
O
H
(30)
Figure 3. Oxidation/reduction of chlorine in hot alkaline solutions. Hydration effect is omitted. These equations correspond to equations having the same number in the text.
(27) (28)
᎑
᎑
HClO3 + OH = ClO3 + H2O O
(26)
(29) (30)
Here Cl2O and Cl2O2 are not “usual” chlorine oxides with oxidation values of chlorine +1 and +2 respectively,; rather, they are chloranhydrides of acids HClO2 and HClO3, respectively. In this case, their structures should be Cl–Cl=O and Cl–ClO2 with oxidation values of chlorine atoms (from left to right) either 0, +2 and 0, +4 or ᎑1, +3 and ᎑1, +5, respectively.
JChemEd.chem.wisc.edu • Vol. 78 No. 7 July 2001 • Journal of Chemical Education
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In the Classroom
Moreover, the logic of chemical equations provides an opportunity to “prolong” further oxidation of chlorates (ClO3᎑) until perchlorates (ClO4᎑ ): ClO3᎑ + Cl2 = Cl2O3 + Cl ᎑
(31)
Cl2O3 + OH ᎑ = HClO4 + Cl ᎑
(32)
᎑
᎑
HClO4 + OH = ClO4 + H2O
F
F Cl O
O
O
−
ClO3᎑ + F2 = FClO3 + F ᎑
(34)
FClO3 + OH ᎑ = HClO4 + F ᎑
(35)
Diagrammatic representation of chemical equations 23– 30 and 33–35 is shown in Figures 3 and 4, respectively. Final Remarks I would like to note that none of the systems of chemical equations discussed above has been experimentally proven completely. Therefore, they might be incorrect. However, according to ref 10, the same can be said about the reaction route of any chemical transformation.
−
O
H
H
O
−
Cl
O
+
F−
(34)
O
F
(33)
In this case, Cl2O3 is not “usual” chlorine oxide either. It is the chloranhydride of HClO4 acid. Its structure should be Cl–ClO3, with oxidation values of chlorine atoms (from left to right) either 0, +6 or ᎑1, +7. One should especially notice that formation of perchlorate (ClO4᎑) anions by passing gaseous chlorine through a hot alkaline solution is impossible: only chlorate (ClO3᎑) anions are obtained in this way. Other methods such as electrochemical oxidation in solutions or solidstate conversion of chlorates are used to produce perchlorates (9). Therefore, chemical equations 31–33 appear to be incorrect for the experimental conditions of reaction 21 because they describe an impossible chemical transformation. This is a good example of some limitations of the systems of chemical equations. However, with minor changes (e.g., by replacing gaseous chlorine with fluorine in eq 31), oxidation of chlorates to perchlorates is possible in a similar way:
920
F O
+
+
O
Cl
O
H
O
O
Cl
O
O
H
O
O−
O
Cl
O
O
O
Cl
O
+
F−
O
+
H
(35)
O
H
(33)
O
Figure 4. Oxidation of chlorate anions in the presence of fluorine. Hydration effect is omitted. These equations correspond to equations having the same number in the text.
Literature Cited 1. 2. 3. 4. 5. 6. 7.
Kolb, D. J. Chem. Educ. 1978, 55, 326–331. Kolb, D. J. Chem. Educ. 1979, 56, 181–184. Tóth, Z. J. Chem. Educ. 1997, 74, 744. Olson, J. A. J. Chem. Educ. 1997, 74, 538–542. Johnstone, A. H. J. Comput. Assisted Learning 1991, 7, 75–83. Lee, K. W. L. J. Chem. Educ. 1999, 76, 1008–1012. Rochow, E. G. The Chemistry of Silicon; Pergamon Texts in Inorganic Chemistry 9; Pergamon: Oxford, 1975. 8. Lenfesty, F. A.; Farr, T. D.; Brosheer, J. C. Ind. Eng. Chem. 1952, 44, 1448–1457. 9. Downs, A. J.; Adams, C. J. The Chemistry of Chlorine, Bromine, Iodine and Astatine; Oxford University Press: Oxford, 1975. 10. Alexander, A. J.; Zare, R. N. J. Chem. Educ. 1998, 75, 1105–1118.
Journal of Chemical Education • Vol. 78 No. 7 July 2001 • JChemEd.chem.wisc.edu
