Before you can balance a chemical equation, you have to know the formulas for all the reactants and products.
If the names are given for these substances, you have to know how to write formulas from the names (Chap. 6).
If only reactants are given, you have to know how to predict the products from the reactants. This latter topic is the subject of this section.
To simplify the discussion, we will classify simple chemical reactions into five types:
Type 1: combination reactions Type 2: decomposition reactions Type 3: substitution reactions
Type 4: double-substitution reactions Type 5: combustion reactions
More complex oxidation-reduction reactions will be discussed in Chap. 14.
Combination Reactions
A combination reaction is a reaction of two reactants to produce one product. The simplest combination reactions are the reactions of two elements to form a compound. After all, if two elements are treated with each other, they can either react or not. There generally is no other possibility, since neither can decompose. In most reactions like this, there will be a reaction. The main problem is to write the formula of the product correctly and then to balance the equation. In this process, first determine the formula of the product from the rules of chemical combination (Chap. 5). Only after the formulas of the reactants and products have all been written down, balance the equation by adjusting the coefficients.
EXAMPLE 8.7. Complete and balance the following equations:
(a) Na+F2−→ (b) Mg+O2−→ (c) K+ S−→
Ans. The products are determined first (from electron dot structures, if necessary) to be NaF, MgO, and K2S, respectively.
These are placed to the right of the respective arrows, and the equations are then balanced.
(a) 2 Na+ F2−→2 NaF (b) 2 Mg+ O2−→2 MgO (c) 2 K+ S−→ K2S EXAMPLE 8.8. Write a complete, balanced equation for the reaction of each of the following pairs of elements:
(a) aluminum and sulfur (b) aluminum and iodine (c) aluminum and oxygen
Ans. (a) The reactants are Al and S. The Al can lose three electrons [it is in periodic group IIIA (3)], and each sulfur atom can gain two electrons [it is in periodic group VIA (16)]. The ratio of aluminum to sulfur atoms is thus 2 : 3, and the compound which will be formed is Al2S3:
Al+S −→Al2S3 (unbalanced) 2 Al+3 S−→Al2S3
(b) The reactants are Al and I2. (In its elemental form, iodine is stable as diatomic molecules.) The combination of a group IIIA (13) metal and a group VIIA (17) nonmetal produces a salt with a 1 : 3 ratio of atoms: AlI3.
Al+I2 −→AlI3 (unbalanced) 2 Al+3 I2−→2 AlI3
(c) Al+O2 −→Al2O3 (unbalanced)
4 Al+3 O2−→2 Al2O3
It is possible for an element and a compound of that element or for two compounds containing a common element to react by combination. For example,
MgO+SO3 −→MgSO4 PtF2 +F2 −→PtF4
Decomposition Reactions
The second type of simple reaction isdecomposition. This reaction is also easy to recognize. Typically, only one reactant is given. A type of energy, such as heat or electricity, may also be indicated. The reactant usually decomposes to its elements, to an element and a simpler compound, or to two simpler compounds.
Binary compoundsmay yield two elements or an element and a simpler compound.Ternary(three-element) compoundsmay yield an element and a compound or two simpler compounds. These possibilities are shown in Fig. 8-2.
Two elements
An element and a compound
Two compounds Binary compound
Ternary compound
Fig. 8-2. Decomposition possibilities
Acatalystis a substance that speeds up a chemical reaction without undergoing a permanent change in its own composition. Catalysts are often but not always noted above or below the arrow in the chemical equation.
Since a small quantity of catalyst is sufficient to cause a large quantity of reaction, the amount of catalyst need not be specified; it is not balanced as the reactants and products are. In this manner, the equation for a common laboratory preparation of oxygen is written as
2 KClO3 MnO2
−→ 2 KCl+3 O2
EXAMPLE 8.9. Write a complete, balanced equation for the reaction that occurs when (a) Ag2O is heated, (b) H2O is electrolyzed, and (c) CaCO3is heated.
Ans. (a) With only one reactant, what can happen? No simpler compound of Ag and O is evident, and the compound decomposes to its elements. Remember that oxygen occurs in diatomic molecules when it is uncombined:
Ag2O−→Ag+ O2 (unbalanced) 2 Ag2O−→4 Ag+O2
(b) Note that in most of these cases, energy of some type is added to make the compound decompose.
2 H2O
electricity
−−−−→2 H2+ O2
(c) A ternary compound does not yield three elements; this one yields two simpler compounds.
CaCO3−→CaO+CO2
Substitution or Replacement Reactions
Elements have varying abilities to combine. Among the most reactive metals are the alkali metals and the alkaline earth metals. On the opposite end of the scale of reactivities, among the least active metals or the most stable metals are silver and gold, prized for their lack of reactivity.Reactivemeans the opposite ofstable, but means the same asactive.
When a free element reacts with a compound of different elements, the free element will replace one of the elements in the compound if the free element is more reactive than the element it replaces. In general, a free metal will replace the metal in the compound, or a free nonmetal will replace the nonmetal in the compound. A new compound and a new free element are produced. As usual, the formulas of the products are written according to the rules in Chap. 5. The formula of a product does not depend on the formula of the reacting element or compound. For example, consider the reactions of sodium with iron(II) chloride and of fluorine with aluminum oxide:
2 Na+FeCl2−→2 NaCl+Fe 6 F2+2 Al2O3−→4 AlF3+3 O2
Sodium, a metal, replaces iron, another metal. Fluorine, a nonmetal, replaces oxygen, another nonmetal.
(In some high-temperature reactions, a nonmetal can displace a relatively inactive metal from its compounds.) The formulas are written on the basis of the rules of chemical bonding (Chap. 5).
You can easily recognize the possibility of a substitution reaction because you are given a free element and a compound of different elements.
EXAMPLE 8.10. Look only at the reactants in the following equations. Tell which of the reactions represent substitution reactions.
(a) 3 Mg+N2−→Mg3N2
(b) 2 Li+MgO−→Li2O+Mg
(c) 2 KClO3−→2 KCl+3 O2
(d) 2 CrCl2+Cl2−→2 CrCl3
Ans. Reaction (b) only. (a) is a combination, (c) is a decomposition, and (d) is a combination. Note that in (d) elemental chlorine is added to a compound of chlorine.
If the free element is less active than the corresponding element in the compound, no reaction will take place.
A short list of metals and an even shorter list of nonmetals in order of their reactivities are presented in Table 8-1.
The metals in the list range from very active to very stable; the nonmetals listed range from very active to fairly active. A more comprehensive list, a table of standard reduction potentials, is presented in general chemistry textbooks.
Table 8-1 Relative Reactivities of Some Metals and Nonmetals Metals Nonmetals
Most active metals Alkali and alkaline F Most active nonmetal earth metals O
Al Cl
Zn Br
Fe I Less active nonmetals
Pb H Cu Less active metals Ag Au
EXAMPLE 8.11. Complete and balance the following equations. If no reaction occurs, indicate that fact by writing “NR”.
(a) KCl+Fe−→ (b) KF+Cl2−→
Ans. (a) KCl+Fe−→NR (b) KF+Cl2−→NR
In each of these cases, the free element is less active than the corresponding element in the compound, and cannot replace that element from its compound.
In substitution reactions, hydrogen in its compounds with nonmetals often acts as a metal; hence, it is listed among the metals in Table 8-1.
EXAMPLE 8.12. Complete and balance the following equation. If no reaction occurs, indicate that fact by writing “NR”.
Al+HCl−→
Ans. 2 Al+6 HCl−→2 AlCl3+3 H2
Aluminum is more reactive than hydrogen (Table 8-1) and replaces it from its compounds. Note that free hydrogen is in the form H2(Sec. 5.2).
In substitution reactions with acids, metals that can form two different ions in their compounds generally form the one with the lower charge. For example, iron can form Fe2+and Fe3+. In its reaction with HCl, FeCl2
is formed. In contrast, in combination with the free element, the higher-charged ion is often formed if sufficient nonmetal is available.
2 Fe+3 Cl2−→2 FeCl3 See Table 6-4 for the charges on some common metal ions.
Double-Substitution or Double-Replacement Reactions
Double-substitutionordouble-replacement reactions, also calleddouble-decomposition reactionsor me- tathesis reactions, involve two ionic compounds, most often in aqueous solution. In this type of reaction, the cations simply swap anions. The reaction proceeds if a solid or a covalent compound is formed from ions in solution. All gases at room temperature are covalent. Some reactions of ionic solids plus ions in solution also occur. Otherwise, no reaction takes place. For example,
AgNO3+NaCl−→AgCl+NaNO3
HCl(aq)+NaOH−→NaCl+H2O CaCO3(s)+2 HCl−→CaCl2+CO2+H2O
In the first reaction, two ionic compounds in water are mixed. The AgCl formed by the swapping of anions is insoluble, causing the reaction to proceed. The solid AgCl formed from solution is an example of aprecipitate.
In the second reaction, a covalent compound, H2O, is formed from its ions in solution, H+and OH−, causing the reaction to proceed. In the third reaction, a solid reacts with the acid in solution to produce two covalent compounds.
Since it is useful to know what state each reagent is in, we often designate the state in the equation. The designation (s) for solid, (1) for liquid, (g) for gas, or (aq) for aqueous solution may be added to the formula.
Thus, a reaction of silver nitrate with sodium chloride in aqueous solution, yielding solid silver chloride and aqueous sodium nitrate, may be written as
AgNO3(aq)+NaCl(aq)−→AgCl(s)+NaNO3(aq)
Just as with replacement reactions, double-replacement reactions may or may not proceed. They need a driving force. In replacement reactions the driving force is reactivity; here it is insolubility or covalence. In order for you to be able to predict if a double-replacement reaction will proceed, you must know some solubilities of ionic compounds. A short list of solubilities is given in Table 8-2.
EXAMPLE 8.13. Complete and balance the following equation. If no reaction occurs, indicate that fact by writing “NR.”
NaCl+KNO3−→
Ans. NaCl+KNO3−→NR
If a double-substitution reaction had taken place, NaNO3 and KCl would have been produced. However, both of these are soluble and ionic; hence, there is no driving force and therefore no reaction.
Table 8-2 Some Solubility Classes
Soluble Insoluble
Chlorates BaSO4
Acetates Most sulfides
Nitrates Most oxides
Alkali metal salts Most carbonates Ammonium salts Most phosphates
Chlorides, except for. . . . AgCl, PbCl2, Hg2Cl2, CuCl
In double-replacement reactions, the charges on the metal ions (and indeed on nonmetal ions if they do not form covalent compounds) generally remain the same throughout the reaction.
EXAMPLE 8.14. Complete and balance the following equations. If no reaction occurs, indicate that fact by writing “NR”.
(a) FeCl3+AgNO3−→ (b) FeCl2+AgNO3−→
Ans. (a) FeCl3(aq)+3 AgNO3(aq)−→Fe(NO3)3(aq)+3 AgCl(s) (b) FeCl2(aq)+2 AgNO3(aq)−→Fe(NO3)2(aq)+2 AgCl(s)
(a) If you start with Fe3+, you wind up with Fe3+. (b) If you start with Fe2+, you wind up with Fe2+.
NH4OH and H2CO3are unstable. If one of these products were expected as a product of a reaction, either NH3plus H2O or CO2plus H2O would be obtained instead:
NH4OH−→NH3+H2O H2CO3−→CO2+H2O
Combustion Reactions
Reactions of elements and compounds with oxygen are so prevalent that they may be considered a separate type of reaction. Compounds of carbon, hydrogen, oxygen, sulfur, nitrogen, and other elements may be burned.
Of greatest importance, if a reactant contains carbon, then carbon monoxide or carbon dioxide will be produced, depending upon how much oxygen is available. Reactants containing hydrogen always produce water on burning.
SO2and NO are other products of burning in oxygen. (To produce SO3requires a catalyst in a combustion reaction with O2.)
EXAMPLE 8.15. Complete and balance the following equations:
(a) C4H8+O2(limited amount)−→ (b) C4H8+O2(excess amount)−→
Ans. (a) C4H8+4 O2−→4 CO+4 H2O (b) C4H8+6 O2−→4 CO2+4 H2O
If sufficient O2is available (6 mol O2per mole C4H8), CO2is the product. In both cases, H2O is produced.
Acids and Bases
Generally, acids react according to the rules for replacement and double-replacement reactions given above.
They are so important, however, that a special nomenclature has developed for acids and their reactions. Acids were introduced in Sec. 6.4. They may be identified by their formulas that have the H representing hydrogen written first, and by their names that contain the wordacid. An acid will react with a base to form a saltand water. The process is calledneutralization. Neutralization reactions will be used as examples in Sec. 11.3, on titration.
HNO3+NaOH−→NaNO3+H2O
A salt
The driving force for such reactions is the formation of water, a covalent compound.
As pure compounds, acids are covalent. When placed in water, they react with the water to form ions; it is said that theyionize. If they react 100% with the water, they are said to bestrong acids. The seven common strong acids are listed in Table 8-3. All the rest areweak; that is, the rest ionize only a few percent and largely stay in their covalent forms. Both strong and weak acids react 100% with metal hydroxides. All soluble metal hydroxides are ionic in water.
Table 8-3 The Seven Common Strong Acids HCl, HClO3, HClO4, HBr, HI, HNO3, H2SO4(first proton only)
Solved Problems
INTRODUCTION
8.1. How many oxygen atoms are there in each of the following, perhaps part of a balanced chemical equation?
(a) 7 H2O, (b) 4 Ba(NO3)2, (c) 2 CuSO4ã5H2O, and (d) 4 VO(ClO3)2. Ans. (a) 7, (b) 24, (c) 18, and (d) 28.
BALANCING SIMPLE EQUATIONS
8.2. Balance the following equation: C+Cu2O−→heat CO+Cu Ans. C+Cu2O−→CO+2 Cu
8.3. Balance the following equation:
Ca(HCO3)2+H3PO4−→Ca3(PO4)2+H2O+CO2
Ans. ? Ca(HCO3)2+? H3PO4 −→1 Ca3(PO4)2+? H2O+? CO2
3 Ca(HCO3)2+2 H3PO4 −→1 Ca3(PO4)2+6 H2O+6 CO2
2nd 2nd 3rd 3rd
3 Ca(HCO3)2+2 H3PO4 −→Ca3(PO4)2+6 H2O+6 CO2 8.4. Balance the following equation:
NH4Cl+KOH−→NH3+H2O+KCl
Ans. 1 NH4Cl+? KOH −→? NH3+? H2O+? KCl
Balance N and Cl : 1 NH4Cl+? KOH −→1 NH3+? H2O + 1 KCl Balance K : 1 NH4Cl + 1 KOH−→1 NH3+? H2O + 1 KCl Balance H : 1 NH4Cl + 1 KOH−→1 NH3 + 1 H2O + 1 KCl Eliminate 1s : NH4Cl + KOH −→NH3 + H2O + KCl 8.5. Balance the following equation
NH4Cl+NaOH+AgCl−→Ag(NH3)2Cl+NaCl+H2O Ans. ? NH4Cl+? NaOH+? AgCl −→1 Ag(NH3)2Cl+? NaCl+? H2O
2 NH4Cl+2 NaOH+1 AgCl−→1 Ag(NH3)2Cl+2 NaCl+2 H2O
2nd 4th 2nd 3rd 5th
2 NH4Cl+2 NaOH+AgCl −→Ag(NH3)2Cl+2 NaCl+2 H2O 8.6. Balance the following chemical equations:
(a) NCl3+H2O−→HClO+NH3
(b) NaOH + H3PO4 −→Na2HPO4+H2O (c) Al+HCl−→AlCl3+H2
(d) HCl+Mg−→MgCl2+H2
(e) SrCO3+HClO4−→Sr(ClO4)2+CO2+H2O (f) KC2H3O2+HBr−→KBr+HC2H3O2
(g) Ba(OH)2+H3PO4−→BaHPO4+H2O (h) HCl+Na3PO4−→NaCl+NaH2PO4
Ans. (a) NCl3+3 H2O−→3 HClO+NH3
(b) 2 NaOH+H3PO4−→Na2HPO4+2 H2O
(c) 2 Al+6 HCl−→2 AlCl3+3 H2
(d) 2 HCl+Mg−→MgCl2+H2
(e) SrCO3+2 HClO4−→Sr(ClO4)2+CO2+H2O (f) KC2H3O2+HBr−→KBr+HC2H3O2
(g) Ba(OH)2+H3PO4−→BaHPO4+2 H2O (h) 2 HCl+Na3PO4−→2 NaCl+NaH2PO4
8.7. Write balanced equations for each of the following reactions:
(a) Sodium plus oxygen yields sodium peroxide.
(b) Mercury(II) oxide, when heated, yields mercury and oxygen.
(c) Carbon plus oxygen yields carbon monoxide.
(d) Sulfur plus oxygen yields sulfur dioxide.
(e) Propane(C3H8)plus oxygen yields carbon dioxide plus water.
(f) Ethane(C2H6)plus oxygen yields carbon monoxide plus water.
(g) Ethylene(C2H4)plus oxygen yields carbon dioxide plus water.
Ans. (a) 2 Na+O2−→Na2O2 (b) 2 HgO−→heat 2 Hg+O2
(c) 2 C+O2−→2 CO (d) S+O2−→SO2
(e) C3H8+5 O2−→3 CO2+4 H2O (f) 2 C2H6+5 O2−→4 CO+6 H2O (g) C2H4+3 O2−→2 CO2+2 H2O
8.8. Write balanced chemical equations for the following reactions:
(a) Sodium plus fluorine yields sodium fluoride.
(b) Potassium chlorate, when heated, yields potassium chloride plus oxygen.
(c) Zinc plus copper(II) nitrate yields zinc nitrate plus copper.
(d) Magnesium hydrogen carbonate plus heat yields magnesium carbonate plus carbon dioxide plus water.
(e) Magnesium hydrogen carbonate plus hydrobromic acid yields magnesium bromide plus carbon dioxide plus water.
(f) Lead(II) acetate plus sodium chromate yields lead(II) chromate plus sodium acetate.
Ans. (a) 2 Na+F2−→2 NaF (Remember that free fluorine is F2.) (b) 2 KClO3−→2 KCl+3 O2(Remember that free oxygen is O2.) (c) Zn+Cu(NO3)2−→Zn(NO3)2+Cu
(d) Mg(HCO3)2−→MgCO3+H2O+CO2
(e) Mg(HCO3)2+2 HBr−→MgBr2+2 H2O+2 CO2 (f) Pb(C2H3O2)2+Na2CrO4−→PbCrO4+2 NaC2H3O2
8.9. Balance the following equation: Ba(ClO4)2+Na2SO4−→BaSO4+NaClO4
Ans. Ba(ClO4)2+Na2SO4−→BaSO4+2 NaClO4
The oxygen atoms need not be considered individually if the ClO4−and SO42−ions are considered as groups.
(See step 3, Sec. 8.2.)
8.10. Write balanced chemical equations for the following reactions:
(a) Phosphorus pentachloride plus water yields phosphoric acid plus hydrogen chloride.
(b) Sodium hydroxide plus sulfuric acid yields sodium hydrogen sulfate plus water.
(c) Ethane, C2H6, plus oxygen yields carbon dioxide plus water.
(d) Octane, C8H18,plus oxygen yields carbon monoxide plus water.
(e) Copper(II) chloride plus hydrosulfuric acid yields copper(II) sulfide plus hydrochloric acid.
(f) Barium hydroxide plus chloric acid yields barium chlorate plus water.
(g) Copper(II) sulfate plus water yields copper(II) sulfate pentahydrate.
(h) Copper(II) chloride plus sodium iodide yields copper(I) iodide plus iodine plus sodium chloride.
Ans. (a) PCl5+4 H2O−→H3PO4+5 HCl (b) NaOH+H2SO4−→NaHSO4+H2O (c) 2 C2H6+7 O2−→4 CO2+6 H2O (d) 2 C8H18+17 O2−→16 CO+18 H2O (e) CuCl2+H2S−→CuS+2 HCl
(f) Ba(OH)2+2 HClO3−→Ba(ClO3)2+2 H2O (g) CuSO4+5 H2O−→CuSO4ã5 H2O
(h) 2 CuCl2+4 NaI−→2 CuI+I2+4 NaCl
8.11. Balance the following equation: Cr+CrCl3−→CrCl2
Ans. Balance the Cl first, since the Cr appears in two reactants. Here, the Cr happens to be balanced automatically.
Cr+2 CrCl3−→3 CrCl2
8.12. Write balanced chemical equations for the following reactions: (a) Hydrogen chloride is produced by the reaction of hydrogen and chlorine. (b) Hydrogen combines with chlorine to yield hydrogen chloride.
(c) Chlorine reacts with hydrogen to give hydrogen chloride.
Ans. (a), (b), and (c). H2+Cl2−→2 HCl
8.13. Write balanced chemical equations for the following reactions: (a) Hydrogen fluoride is produced by the reaction of hydrochloric acid and sodium fluoride. (b) Hydrochloric acid combines with sodium fluoride to yield hydrogen fluoride. (c) Sodium fluoride reacts with hydrochloric acid to give hydrofluoric acid.
Ans. (a), (b), and (c). HCl + NaF−→NaCl + HF 8.14. Balance the following chemical equations:
(a) Na2CO3+HClO3−→NaClO3+CO2+H2O (b) Ca(HCO3)2+HCl−→CaCl2+CO2+H2O (c) BiCl3+H2O−→BiOCl+HCl
(d) H2S+O2−→H2O+SO2 (e) Cu2S+O2−→Cu+SO2
(f) NH3+O2−→NO+H2O (g) H2O2 −→H2O+O2
Ans. (a) Na2CO3+2 HClO3−→2 NaClO3+CO2+H2O (b) Ca(HCO3)2+2 HCl−→CaCl2+2 CO2+2 H2O (c) BiCl3+H2O−→BiOCl+2 HCl
(d) 2 H2S+3 O2−→2 H2O+2 SO2
(e) Cu2S+O2−→2 Cu+SO2
(f) 4 NH3+5 O2−→4 NO+6 H2O (g) 2 H2O2−→2 H2O+O2
8.15. Balance the following chemical equations:
(a) Pb(NO3)2+KI−→PbI2+KNO3 (b) H2S+CuCl2−→HCl+CuS
(c) Hg2Cl2+NH3−→HgNH2Cl+Hg+NH4Cl
(d) NH3+CuCl2−→Cu(NH3)4Cl (e) As2S3+Na2S−→NaAsS2
Ans. (a) Pb(NO3)2+2 KI−→PbI2+2 KNO3
(b) H2S+CuCl2−→2 HCl+CuS
(c) Hg2Cl2+2 NH3−→HgNH2Cl+Hg+NH4Cl
(d) 4 NH3+CuCl2−→Cu(NH3)4Cl (e) As2S3+Na2S−→2 NaAsS2
8.16. Why is the catalyst not merely placed on both sides of the arrow, since it comes out of the reaction with the same composition as it started with?
Ans. That would imply a certain mole ratio to the other reactants and products, which is not correct.
PREDICTING THE PRODUCTS OF A REACTION
8.17. In the list of reactivities of metals, Table 8-1, are all alkali metals more reactive than all alkaline earth metals, or are all elements of both groups of metals more active than any other metals?
Ans. Both groups of metals are more active than any other metals. Actually, some alkaline earth metals are more active than some alkali metals, and vice versa.
8.18. (a) What type of reaction requires knowledge of reactivities of elements? (b) What type requires knowl- edge of solubility properties of compounds?
Ans. (a) Substitution reaction (b) Double-substitution reaction 8.19. Complete and balance the following equations:
(a) CH4+O2(limited amount) −→ (b) CH4+O2(excess amount) −→
Ans. (a) 2 CH4 + 3 O2(limited amount) −→2 CO + 4 H2O If suffient O2is available, CO2is the product. (b) CH4+2 O2(excess amount) −→CO2+2 H2O
8.20. What type of chemical reaction is represented by each of the following? Complete and balance the equation for each.
(a) Cl2+AlBr3−→
(b) Cl2+K−→
(c) Na+AlCl3−→
(d) FeCl2+AgC2H3O2−→
(e) C3H8+O2(limited)−→
Ans. (a) Substitution 3 Cl2+2 AlBr3−→2 AlCl3+3 Br2
(b) Combination Cl2+2 K−→2 KCl
(c) Substitution 3 Na+AlCl3−→Al+3 NaCl
(d) Double substitution FeCl2+2 AgC2H3O2−→Fe(C2H3O2)2+2 AgCl (e) Combustion 2 C3H8+7 O2−→6 CO+8 H2O
8.21. What type of chemical reaction is represented by each of the following? Complete and balance the equation for each.
(a) Cl2+CrCl2 −→
(b) CO+O2−→
(c) MgCO3−→heat
(d) AlCl3+Cl2−→
(e) C5H12+O2(limited quantity)−→