A chemical reaction is a process in which one set of chemical substances (reactants) is converted into another (products). It involves making and breaking chemical bonds and the rearrangement of atoms. Chemical reactions are represented by balanced chemical equations, with chemical formulas symbolizing reactants and products. For specific chemical reactants, two questions may be posed about a possible chemical reaction. First, will a reaction occur? Second, what are the possible products if a reaction occurs? This
entry will focus only on the second question. The most reliable answer is obtained by conducting an experiment—mixing the reactants and then isolating and identifying the products. We can also use periodicity, since elements within the same group in the Periodic Table undergo similar reactions. Finally, we can use rules to help predict the products of reactions, based on the classification of inorganic chemical reactions into four general categories: combination, decomposition, single-displacement, and double-displacement reactions.
Reactions may also be classified according to whether the oxidation number of one or more elements changes. Those reactions in which a change in oxidation number occurs are called oxidation–reduction reactions . One element increases its oxidation number (is oxidized), while the other decreases its oxidation number (is reduced).
In combination reactions, two substances, either elements or compounds, react to produce a single compound. One type of combination reaction involves two elements. Most metals react with most nonmetals to form ionic compounds. The products can be predicted from the charges expected for cations of the metal and anions of the nonmetal. For example, the product of the reaction between aluminum and bromine can be predicted from the following charges: 3+ for aluminum ion and 1− for bromide ion. Since there is a change in the oxidation numbers of the elements, this type of reaction is an oxidation–reduction reaction:
2Al ( s ) + 3Br 2 ( g ) → 2AlBr 3 ( s )
Similarly, a nonmetal may react with a more reactive nonmetal to form a covalent compound. The composition of the product is predicted from the common oxidation numbers of the elements, positive for the less reactive and negative for the more reactive nonmetal (usually located closer to the upper right side of the Periodic Table). For example, sulfur reacts with oxygen gas to form gaseous sulfur dioxide:
S 8 ( s ) + 8O 2 ( g ) → 8SO 2 ( g )
A compound and an element may unite to form another compound if in the original compound, the element with a positive oxidation number has an accessible higher oxidation number. Carbon monoxide, formed by the burning of hydrocarbons under conditions of oxygen deficiency, reacts with oxygen to form carbon dioxide:
2CO ( g ) + O 2 ( g ) → 2CO 2 ( g )
The oxidation number of carbon changes from +2 to +4 so this reaction is an oxidation–reduction reaction.
Two compounds may react to form a new compound. For example, calcium oxide (or lime) reacts with carbon dioxide to form calcium carbonate (limestone):
CaO ( s ) + CO 2 ( g ) → CaCO 3 ( s )
When a compound undergoes a decomposition reaction, usually when heated, it breaks down into its component elements or simpler compounds. The products of a decomposition reaction are determined largely by the identity of the anion in the compound. The ammonium ion also has characteristic decomposition reactions.
A few binary compounds decompose to their constituent elements upon heating. This is an oxidation–reduction reaction since the elements undergo a change in oxidation number. For example, the oxides and halides of noble metals (primarily Au, Pt, and Hg) decompose when heated. When red solid mercury(II) oxide is heated, it decomposes to liquid metallic mercury and oxygen gas:
2HgO ( s ) → 2Hg ( l ) + O 2 ( g )
Some nonmetal oxides, such as the halogen oxides, also decompose upon heating:
2Cl 2 O 5 ( g ) → 2Cl 2 ( g ) + 5O 2 ( g )
Other nonmetal oxides, such as dinitrogen pentoxide, decompose to an element and a compound:
2N 2 O 5 ( g ) → O 2 ( g ) + 4NO 2 ( g )
Many metal salts containing oxoanions decompose upon heating. These salts either give off oxygen gas, forming a metal salt with a different nonmetal anion, or they give off a nonmetal oxide, forming a metal oxide. For example, metal nitrates containing Group 1A or 2A metals or aluminum decompose to metal nitrites and oxygen gas:
Mg(NO 3 ) 2 ( s ) → Mg(NO 2 ) 2 ( s ) + O 2 ( g )
All other metal nitrates decompose to metal oxides, along with nitrogen dioxide and oxygen:
2Cu(NO 3 ) 2 ( s ) → 2CuO ( s ) + 4NO 2 ( g ) + O 2 ( g )
Salts of the halogen oxoanions decompose to halides and oxygen upon heating:
2KBrO 3 ( s ) → 2KBr ( s ) + 3O 2 ( g )
Carbonates, except for those of the alkali metals, decompose to oxides and carbon dioxide.
CaCO 3 ( s ) → CaO ( s ) + CO 2 ( g )
A number of compounds—hydrates, hydroxides, and oxoacids—that contain water or its components lose water when heated. Hydrates, compounds that contain water molecules, lose water to form anhydrous compounds, free of molecular water.
CaSO 4 · 2H 2 O ( s ) → CaSO 4 ( s ) + 2H 2 O ( g )
Metal hydroxides are converted to metal oxides by heating:
2Fe(OH) 3 ( s ) → Fe 2 O 3 ( s ) + 3H 2 O ( g )
Most oxoacids lose water until no hydrogen remains, leaving a nonmetal oxide:
H 2 SO 4 ( l ) → H 2 O ( g ) + SO 3 ( g )
Oxoanion salts that contain hydrogen ions break down into the corresponding oxoanion salts and oxoacids:
Ca(HSO 4 ) 2 ( s ) → CaSO 4 ( s ) + H 2 SO 4 ( l )
Finally, some ammonium salts undergo an oxidation–reduction reaction when heated. Common salts of this type are ammonium dichromate, ammonium permanganate, ammonium nitrate, and ammonium nitrite. When these salts decompose, they give off nitrogen gas and water.
(NH 4 ) 2 Cr 2 O 7 ( s ) → Cr 2 O 3 ( s ) + 4H 2 O ( g ) + N 2 ( g )
2NH 4 NO 3 ( s ) → 2N 2 ( g ) + 4H 2 O ( g ) + O 2 ( g )
Ammonium salts, which do not contain an oxidizing agent, lose ammonia gas upon heating:
(NH 4 ) 2 SO 4 ( s ) → 2NH 3 ( g ) + H 2 SO 4 ( l )
In a single-displacement reaction, a free element displaces another element from a compound to produce a different compound and a different free element. A more active element displaces a less active element from its compounds. These are all oxidation–reduction reactions. An example is the thermite reaction between aluminum and iron(III) oxide:
2Al ( s ) + Fe 2 O 3 ( s ) → Al 2 O 3 ( s ) + 2Fe ( l )
The element displaced from the compound is always the more metallic element—the one nearer the bottom left of the Periodic Table. The displaced element need not always be a metal, however. Consider a common type of single-displacement reaction, the displacement of hydrogen from water or from acids by metals.
The very active metals react with water. For example, calcium reacts with water to form calcium hydroxide and hydrogen gas. Calcium metal has an oxidation number of 0, whereas Ca 2+ in Ca(OH) 2 has an oxidation number of +2, so calcium is oxidized. Hydrogen's oxidation number changes from +1 to 0, so it is reduced.
Ca ( s ) + 2H 2 O ( l ) → Ca(OH) 2 ( aq ) + H 2 ( g )
Some metals, such as magnesium, do not react with cold water, but react slowly with steam:
Mg ( s ) + 2H 2 O ( g ) → Mg(OH) 2 ( aq ) + H 2 ( g )
Still less active metals, such as iron, do not react with water at all, but react with acids.
Fe ( s ) + 2HCl ( aq ) → FeCl 2 ( aq ) + H 2 ( g )
Metals that are even less active, such as copper, generally do not react with acids.
To determine which metals react with water or with acids, we can use an activity series (see Figure 1), a list of metals in order of decreasing activity. Elements at the top of the series react with cold water. Elements above hydrogen in the series react with acids; elements below hydrogen do not react to release hydrogen gas.
The displacement of hydrogen from water or acids is just one type of single-displacement reaction. Other elements can also be displaced from their compounds. For example, copper metal reduces aqueous solutions of ionic silver compounds, such as silver nitrate, to deposit silver metal. The copper is oxidized.
Cu ( s ) + 2AgNO 3 ( aq ) → Cu(NO 3 ) 2 ( aq ) + 2Ag ( s )
The activity series can be used to predict which single-displacement reactions will take place. The elemental metal produced is always lower in the activity series than the displacing element. Thus, iron could be displaced from FeCl 2 by zinc metal but not by tin.
|K||These metals will displace hydrogen gas from water|
|Zn||These metals will displace hydrogen gas from acids|
|Hg||These metals will not displace hydrogen gas from water or acids|
Aqueous barium chloride reacts with sulfuric acid to form solid barium sulfate and hydrochloric acid:
BaCl 2 ( aq ) + H 2 SO 4 ( aq ) → BaSO 4 ( s ) + 2HCl ( aq )
Sodium sulfide reacts with hydrochloric acid to form sodium chloride and hydrogen sulfide gas:
Na 2 S ( aq ) + 2HCl ( aq ) → 2NaCl ( aq ) + H 2 S ( g )
Potassium hydroxide reacts with nitric acid to form water and potassium nitrate:
KOH ( aq ) + HNO 3 ( aq ) → H 2 O ( l ) + KNO 3 ( aq )
These double-displacement reactions have two major features in common. First, two compounds exchange ions or elements to form new compounds. Second, one of the products is either a compound that will separate from the reaction mixture in some way (commonly as a solid or gas) or a stable covalent compound, often water.
Double-displacement reactions can be further classified as precipitation, gas formation, and acid–base neutralization reactions.
Precipitation reactions are those in which the reactants exchange ions to form an insoluble salt—one which does not dissolve in water. Reaction occurs when two ions combine to form an insoluble solid or precipitate. We predict whether such a compound can be formed by consulting solubility rules (see Table 1). If a possible product is insoluble, a precipitation reaction should occur.
A mixture of aqueous solutions of barium chloride and sodium sulfate contains the following ions: Ba 2+ ( aq ), Cl − ( aq ), Na + ( aq ), and SO 4 2− ( aq ). According to solubility rules, most sulfate, sodium, and chloride salts are soluble. However, barium sulfate is insoluble. Since a barium ion and sulfate ion could combine to form insoluble barium sulfate, a reaction occurs.
|SOME SOLUBILITY RULES FOR INORGANIC SALTS IN WATER|
|Na + , K + , NH 4 +||Most salts of sodium, potassium, and ammonium ions are soluble.|
|NO 3 −||All nitrates are soluble.|
|SO 4 2−||Most sulfates are soluble. Exceptions: BaSO 4 , SrSO 4 , PbSO 4 , CaSO 4 , Hg 2 SO 4 , and Ag 2 SO 4 .|
|Cl − , Br − , I − ,||Most chlorides, bromides, and iodides are soluble. Exceptions: AgX, Hg 2 X 2 , PbX 2 , and HgI 2 .|
|Ag +||Silver salts, except AgNO 3 , are insoluble.|
|O 2− , OH −||Oxides and hydroxides are insoluble. Exceptions: NaOH, KOH, NH 4 OH, Ba(OH) 2 , and Ca(OH) 2 (somewhat soluble).|
|S 2−||Sulfides are insoluble. Exceptions: salts of Na + , K + , NH 4 + and the alkaline earth metal ions.|
|CrO 4 2−||Most chromates are insoluble. Exceptions: salts of K + , Na + , NH 4 + , Mg 2+ , Ca 2+ , Al 3+ , and Ni 2+ .|
|CO 3 2− , PO4 3− , SO 3 2− , SiO 3 2−||Most carbonates, phosphates, sulfites, and silicates are insoluble. Exceptions: salts of K + , Na + , and NH 4 + .|
BaCl 2 ( aq ) + Na 2 SO 4 ( aq ) → BaSO 4 ( s ) + 2NaCl ( aq )
A double-displacement reaction should also occur if an insoluble gas is formed. All gases are soluble in water to some extent, but only a few gases [HCl ( g ) and NH 3 ( g )] are highly soluble. All other gases, generally binary covalent compounds, are sufficiently insoluble to provide a driving force if they are formed as a reaction product. For example, many sulfide salts will react with acids to form gaseous hydrogen sulfide:
ZnS ( s ) + 2HCl ( aq ) → ZnCl 2 ( aq ) + H 2 S ( g )
Insoluble gases are often formed by the breakdown of an unstable double-displacement reaction product. For example, carbonates react with acids to form carbonic acid (H 2 CO 3 ), an unstable substance that readily decomposes into water and carbon dioxide. Calcium carbonate reacts with hydrochloric acid to form calcium chloride and carbonic acid:
CaCO 3 ( s ) + 2HCl ( aq ) → CaCl 2 ( aq ) + H 2 CO 3 ( aq )
Carbonic acid decomposes into water and carbon dioxide:
H 2 CO 3 ( aq ) → H 2 O ( l ) + CO 2 ( g )
The net reaction is:
CaCO 3 ( s ) + 2HCl (aq) → CaCl 2 ( aq ) + H 2 O ( l ) + CO 2 ( g )
Sulfites react with acids in a similar manner to release sulfur dioxide.
Acid-Base Neutralization Reactions
A neutralization reaction is a double-displacement reaction of an acid and a base. Acids are compounds that can release hydrogen ions; bases are compounds that can neutralize acids by reacting with hydrogen ions. The most common bases are hydroxide and oxide compounds of the metals. Normally, an acid reacts with a base to form a salt and water. Neutralization reactions occur because of the formation of the very stable covalent water molecule, H 2 O, from hydrogen and hydroxide ions.
HCl ( aq ) + NaOH ( aq ) → NaCl ( aq ) + H 2 O ( l )
Recognizing the pattern of reactants (element or compound, and the number of each) allows us to assign a possible reaction to one of the described classes. Recognizing the class of reaction allows us to predict possible products with some reliability.
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