Chemical Reactions

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).

Combination Reactions

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 ) + 3Br₂ (g ) → 2AlBr₃ (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:

S₈ (s ) + 8O₂ (g ) → 8SO₂ (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 ) + O₂ (g ) → 2CO₂ (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 ) + CO₂ (g ) → CaCO₃ (s )

Decomposition Reactions

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 ) + O₂ (g )

Some nonmetal oxides, such as the halogen oxides, also decompose upon heating:

2Cl₂O₅ (g ) → 2Cl₂ (g ) + 5O₂ (g )

Other nonmetal oxides, such as dinitrogen pentoxide, decompose to an element and a compound:

2N₂O₅ (g ) → O₂ (g ) + 4NO₂ (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(NO₃)₂ (s ) → Mg(NO₂)₂ (s ) + O₂ (g )

All other metal nitrates decompose to metal oxides, along with nitrogen dioxide and oxygen:

2Cu(NO₃)₂ (s ) → 2CuO (s ) + 4NO₂ (g ) + O₂ (g )

Salts of the halogen oxoanions decompose to halides and oxygen upon heating:

2KBrO₃ (s ) → 2KBr (s ) + 3O₂ (g )

Carbonates, except for those of the alkali metals, decompose to oxides and carbon dioxide.

CaCO₃ (s ) → CaO (s ) + CO₂ (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.

CaSO₄ · 2H₂O (s ) → CaSO₄ (s ) + 2H₂O (g )

Metal hydroxides are converted to metal oxides by heating:

2Fe(OH)₃ (s ) → Fe₂O₃ (s ) + 3H₂O (g )

Most oxoacids lose water until no hydrogen remains, leaving a nonmetal oxide:

H₂SO₄ (l ) → H₂O (g ) + SO₃ (g )

Oxoanion salts that contain hydrogen ions break down into the corresponding oxoanion salts and oxoacids:

Ca(HSO₄)₂ (s ) → CaSO₄ (s ) + H₂SO₄ (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.

(NH₄)₂Cr₂O₇ (s ) → Cr₂O₃ (s ) + 4H₂O (g ) + N₂ (g )

2NH₄NO₃ (s ) → 2N₂ (g ) + 4H₂O (g ) + O₂ (g )

Ammonium salts, which do not contain an oxidizing agent, lose ammonia gas upon heating:

(NH₄)₂SO₄ (s ) → 2NH₃ (g ) + H₂SO₄ (l )

Single-Displacement Reactions

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 ) + Fe₂O₃ (s ) → Al₂O₃ (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 Ca²⁺ 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 ) + 2H₂O (l ) → Ca(OH)₂ (aq ) + H₂ (g )

Some metals, such as magnesium, do not react with cold water, but react slowly with steam:

Mg (s ) + 2H₂O (g ) → Mg(OH)₂ (aq ) + H₂ (g )

Still less active metals, such as iron, do not react with water at all, but react with acids.

Fe (s ) + 2HCl (aq ) → FeCl₂ (aq ) + H₂ (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 ) + 2AgNO₃ (aq ) → Cu(NO₃)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 FeCl₂ by zinc metal but not by tin.

Double-Displacement Reactions

Aqueous barium chloride reacts with sulfuric acid to form solid barium sulfate and hydrochloric acid:

BaCl₂ (aq ) + H₂SO₄ (aq ) → BaSO₄ (s ) + 2HCl (aq )

Sodium sulfide reacts with hydrochloric acid to form sodium chloride and hydrogen sulfide gas:

Na₂S (aq ) + 2HCl (aq ) → 2NaCl (aq ) + H₂S (g )

Potassium hydroxide reacts with nitric acid to form water and potassium nitrate:

KOH (aq ) + HNO₃ (aq ) → H₂O (l ) + KNO₃ (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

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: Ba²⁺ (aq ), Cl⁻ (aq ), Na⁺ (aq ), and SO₄²⁻ (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.

BaCl₂ (aq ) + Na₂SO₄ (aq ) → BaSO₄ (s ) + 2NaCl (aq )

Gas-Formation Reactions

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 NH₃ (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 ) → ZnCl₂ (aq ) + H₂S (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 (H₂CO₃), an unstable substance that readily decomposes into water and carbon dioxide. Calcium carbonate reacts with hydrochloric acid to form calcium chloride and carbonic acid:

CaCO₃ (s ) + 2HCl (aq ) → CaCl₂ (aq ) + H₂CO₃ (aq )

Carbonic acid decomposes into water and carbon dioxide:

H₂CO₃ (aq ) → H₂O (l ) + CO₂ (g )

The net reaction is:

CaCO₃ (s ) + 2HCl (aq) → CaCl₂ (aq ) + H₂O (l ) + CO₂ (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, H₂O, from hydrogen and hydroxide ions.

HCl (aq ) + NaOH (aq ) → NaCl (aq ) + H₂O (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.

see also Acid-Base Chemistry; Solution Chemistry; Thermodynamics.

James P. Birk

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Internet Resources

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ChemWeb 2000. Available from <http://library.thinkquest.org/19957/>.

Reaction, chemical

When a chemical reaction occurs, at least one product is formed that is different from the substances present before the change occurred. As an example, it is possible to pass an electric current through a sample of water and obtain a mixture of oxygen and hydrogen gases. That change is a chemical reaction because neither oxygen nor hydrogen were present as elements before the change took place.

Any chemical change involves two sets of substances: reactants and products. A reactant is an element or compound present before a chemical change takes place. In the example above, only one reactant was present: water. A product is an element or compound formed as a result of the chemical reaction. In the preceding example, both hydrogen and oxygen are products of the reaction.

Chemical reactions are represented by means of chemical equations. A chemical equation is a symbolic statement that represents the changes that occur during a chemical reaction. The statement consists of the symbols of the elements and the formulas of the products and reactants, along with other symbols that represent certain conditions present in the reaction. For example, the arrow (or yields) sign, *, separates the reactants from the products in a reaction. The chemical equation that represents the electrolysis of water is 2 H₂O → 2 H₂ + O₂.

Types of chemical reactions

Most chemical reactions can be categorized into one of about five general types: synthesis, decomposition, single replacement, double replacement, and oxidation-reduction. A miscellaneous category is also needed for reactions that do not fit into one of these five categories.

Characteristics of each type.

Synthesis: Two substances combine to form one new substance:

In general: A + B → AB

For example:

2 Na + Cl₂ → 2 NaCl or CaO + H₂O → Ca(OH)₂

Decomposition: One substance breaks down to form two new substances:

In general: AB → A + B

For example: 2 H₂O → 2 H₂ + O₂

Single Replacement: An element and a compound react such that the element replaces one other element in the compound:

In general: A + BC → AC + B

For example: Mg + 2 HCl → MgCl₂ + H₂

Double Replacement: Two compounds react with each other in such a way that they exchange partners with each other:

In general: AB + CD → AD + CB

For example:

NaBr + HCl → NaCl + HBr

Oxidation-reduction: One or more elements in the reaction changes its oxidation state during the reaction: In general: A³⁺ → A⁶⁺

For example: Cr³⁺ → Cr⁶⁺

Energy changes and chemical kinetics

Chemical reactions are typically accompanied by energy changes. The equation for the synthesis of ammonia from its elements is N₂ + 3 H₂ → 2 NH₃, but that reaction takes place only under very special conditions—namely at a high temperature and pressure and in the presence of a catalyst. Energy changes that occur during chemical reactions are the subject of a field of science known as thermodynamics.

In addition, chemical reactions are often a good deal more complex than a chemical equation might lead one to believe. For example, one can write the equation for the synthesis of hydrogen iodide from its elements, as follows: H₂ + I₂ → 2 HI. In fact, chemists know that this reaction does not take place in a single step. Instead, it occurs in a series of reactions in which hydrogen and iodine atoms react with each other one at a time. The final equation, H₂ + I₂ → 2 HI, is actually no more than a summary of the net result of all those reactions. The field of chemistry that deals with the details of chemical reactions is known as chemical kinetics.

chemical reaction process by which one or more substances may be transformed into one or more new substances. Energy is released or is absorbed, but no loss in total molecular weight occurs. When, for example, water is decomposed, its molecules, each of which consists of one atom of oxygen and two of hydrogen, are broken down; the hydrogen atoms then combine in pairs to form hydrogen molecules and the oxygen atoms to form oxygen molecules. In a chemical reaction, substances lose their characteristic properties. Water, for example, a liquid which neither burns nor supports combustion, is decomposed to yield flammable hydrogen and combustion-supporting oxygen. In some reactions heat is given off (exothermic reactions), and in others heat is absorbed (endothermic reactions). Furthermore, the new substances formed differ from the original substances in the energy they contain. Chemical reactions are classified according to the kind of change that takes place. When a compound, which consists of two or more elements or groups of elements, is broken down into its constituents, the reaction is called simple decomposition. When two compounds react with one another to form two new compounds, the reaction is called double decomposition. In so-called replacement reactions the place of one of the elements in a compound is taken by another element reacting with the compound. When elements combine to form a compound, the reaction is termed chemical combination. Oxidation and reduction reactions are extremely important. Reversible reactions are those in which the chemical change taking place may be paralleled by another change back to the original substances. The rates at which chemical reactions proceed depend upon various factors, e.g., upon temperature, pressure, and the concentration of the substances involved and, sometimes, upon the use of a chemical called a catalyst . In some chemical reactions, such as that of photographic film, light is an important factor. The changes taking place in a chemical reaction are represented by a chemical equation . An element's activity, i.e., its tendency to enter into compounds, varies from one element to another.

chemical reaction A change in which one or more chemical elements or compounds (the reactants) form new compounds (the products). All reactions are to some extent reversible; i.e. the products can also react to give the original reactants. However, in many cases the extent of this back reaction is negligibly small, and the reaction is regarded as irreversible. See also endergonic reaction; exergonic reaction.

chemical reaction Change or process in which chemical substances convert into other substances. This involves the breaking and formation of chemical bonds. Reaction mechanisms include endothermic (heat imput), exothermic (heat output), replacement, combination, decomposition and oxidation reactions.

chemical reaction see chemical reaction .