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)
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)
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 rea
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