Volume 1

General Information • Nomenclature of Oxides • Basic, Acidic, Amphoteric, and Neutral Oxides • Preparation of Oxides • Chemical Properties of Oxides • Exercises
1.14.1. General Information. As we already know, an oxide is a chemical compound comprising oxygen and another element. At room temperature and atmospheric pressure, an oxide can be gaseous (CO2, CO, SO2, NO), liquid (H2O, SO3, Mn2O7) or solid (CaO, Al2O3, ZnO, P2O5).

Oxides are ubiquitous in nature. Water is an oxide. (In principle, H2O could be called hydrogen oxide, but that would sound strange.) Silicon oxide, SiO2, accounts for over 10% of the Earth's crust by mass. Numerous ores, minerals, and rocks are metal oxides, including many precious stones (Figure 1-53). Sapphires and rubies are aluminum oxide (Al2O3). Amethyst is silicon oxide (SiO2). Emerald (beryllium aluminum silicate, Be3Al2Si6O18) is a complex oxide, an oxide that contains at least two elements besides oxygen. All of these oxides are white when pure. The gemstones are colored by trace impurities of other elements, such as titanium (sapphire), iron (amethyst), and chromium (ruby, emerald). Carbon dioxide is present in the atmosphere, although only in small quantities.
Figure 1-53. From left to right: Logan Sapphire, Carmen Lúcia Ruby, 96-carat amethyst, and Hooker Emerald of the National Museum of Natural History (Washington, DC, USA).

1.14.2. Nomenclature of Oxides. Oxides are easy to name. Name the element comprising the oxide, followed by the word "oxide". For instance, MgO is magnesium oxide, Li2O is lithium oxide, and Al2O3 is aluminum oxide. If an element has more than one valence state, its valence in the oxide is specified as a Roman numeral in parentheses after the name of the element. For instance, CuO is copper (II) oxide, whereas Cu2O is copper (I) oxide. Similarly, SO2 is sulfur (IV) oxide and SO3 is sulfur (VI) oxide. Quite often the composition is specified by Greek prefixes that indicate the number of oxygen atoms forming the oxide: "mono" (or "mon") indicates one; "di" indicates two; "tri", "tetra" (or "tetr"), and "penta" (or "pent") are three, four, and five, respectively. In this way, CO, SO2, SO3, OsO4, and P2O5 are called carbon monoxide, sulfur dioxide, sulfur trioxide, osmium tetroxide, and phosphorus pentoxide, respectively.

1.14.3. Basic, Acidic, Amphoteric, and Neutral Oxides. There are four types of oxides: basic, acidic, amphoteric, and neutral.

Basic oxides are metal oxides. They react with acids to give a salt and water (Figure 1-54).
Figure 1-54. Reactions of selected basic oxides with acids.

Acidic oxides are nonmetal oxides. They react with bases to give a salt and water (Figure 1-55).
Figure 1-55. Reactions of selected acidic oxides with bases.
Digression. Some acidic oxides are called anhydrides or acid anhydrides. The term anhydride (from Greek άνυδρος) meaning "waterless" has been used by chemists since the mid-19th century to name acidic oxides that are obtained by the elimination of water from the corresponding acids. For example, SO3 is sometimes called sulfuric anhydride because it is formed upon elimination of H2O from sulfuric acid, H2SO4 (at a very high temperature). Likewise, CO2 is on occasion called carbonic anhydride because it is produced when carbonic acid loses water, H2CO3 = H2O + CO2. Note that the molecules of acids do not contain water molecules per se, but only the H and O atoms which combine to give rise to H2O on thermal decomposition of the acid.
Amphoteric oxides are those that react with both acids and bases to give a salt and water. These oxides are formed by some metalloids and metals that are borderline between typical metals and nonmetals. For example, zinc oxide is amphoteric, as it reacts with acids as a basic oxide and with bases as an acidic oxide (Figure 1-56).
Figure 1-56. Amphoteric zinc oxide reacts with both acids and bases.

The salt produced in the reaction of ZnO with NaOH (Figure 1-56, bottom) is called sodium zincate. Actually, sodium zincate has a composition and structure that differ from those described by the formula Na2ZnO2. For simplicity, however, it is this formula that we will use in our course.

Another example of an amphoteric oxide is aluminum oxide, Al2O3. Those metals that form amphoteric oxides also form amphoteric hydroxides. Just like amphoteric oxides, amphoteric hydroxides, such as Zn(OH)2 or Al(OH)3, react with both acids and bases to give a salt and water. We will learn more about amphoteric hydroxides later in this Volume and in Volumes 2 and 3.

Neutral oxides represent a rather small group of oxides that do not react with either acids or bases. Examples of neutral oxides include carbon monoxide (CO) and nitrous oxide (N2O).

In summary, metals conventionally form basic oxides, whereas acidic oxides are derived from nonmetals. Some metal oxides are amphoteric, such as ZnO and Al2O3. (Interestingly, there are even acidic metal oxides. Those are known only for some transition metals in high-valent states, such as CrO3 and Mn2O7. We will learn about the transition metals in Volume 2.)

1.14.4. Preparation of Oxides. There are many ways to make oxides.

(1). Burning simple substances in air or oxygen (Figure 1-57).
Figure 1-57. Many simple substances such as magnesium and phosphorus react with oxygen to give oxides.

(2). Burning complex substances (Figure 1-58).
Figure 1-58. Some complex substances react with oxygen to give oxides.

(3). Thermal decomposition (dehydration) of some bases (Figure 1-59). Water-insoluble metal hydroxides decompose more readily than water-soluble ones. For example, blue Cu(OH)2 (insoluble in water) decomposes to black CuO at ~200 oC, whereas Ca(OH)2 (soluble in water) starts to decompose only above 500 oC.
Figure 1-59. Many metal hydroxides decompose on heating to give a metal oxide and water.

(4). Thermal decomposition of some acids (Figure 1-60).
Figure 1-60. Some acids decompose to give a nonmetal oxide and water.

(5). Thermal decomposition of some oxygen-containing salts (Figure 1-61).
Figure 1-61. Some salts of oxygen-containing acids decompose on heating to give oxides.

1.14.5. Chemical Properties of Oxides.

(1). Metal and nonmetal oxides react with water to give a metal hydroxide (base) or an acid, respectively (Figure 1-62). These hydration reactions readily occur if the acid or base produced is soluble in water. In contrast, water-insoluble metal hydroxides and acids cannot be easily obtained by hydration of the corresponding oxide. For example, Cu(OH)2 and H2SiO3 that are water-insoluble are not formed on addition of water to CuO and SiO2, respectively.
Figure 1-62. Hydration reactions of oxides readily occur if the resultant acid or base is soluble in water.

(2). Acidic (and amphoteric) oxides react with bases to give a salt and water (Figure 1-63).
Figure 1-63. Acidic oxides react with bases to give a salt and water.

(3). Basic (and amphoteric) oxides react with acids to give a salt and water (Figure 1-64).
Figure 1-64. Basic oxides react with acids to give a salt and water.

(4). A more reactive element can displace a less reactive element from its oxide. The aforementioned reaction of copper (II) oxide with hydrogen and the thermite reaction of iron (III) oxide with Al are typical examples of this type of reaction (Figure 1-65).
Figure 1-65. Displacement reactions of oxides.

A remarkable example of this type of displacement reaction is the burning of magnesium (Mg) in carbon dioxide (Figure 1-66). Watch Video 1-36 to see how intense this combustion is even when very cold (-78 oC) dry ice, which is solid CO2, is used for the reaction.
Figure 1-66. Burning of magnesium in CO2.
Video 1-36. Magnesium burns in CO2. (source).

(5). Basic oxides can react with acidic oxides (Figure 1-67).
Figure 1-67. Some basic oxides can react with certain acidic oxides.

(6). Oxides of some metals such as silver and mercury readily decompose on heating to give the metal and oxygen (Figure 1-68).
Figure 1-68. Thermal decomposition of mercury and silver oxides.

(7). Oxides of elements in a low valent state can react with O2 to produce an oxide of the same element in a higher valent state (Figure 1-69).
Figure 1-69. A low-valent oxide can react with oxygen to give a higher-valent oxide.

1.14.6. Exercises.

1. There are liquid, solid, and gaseous oxides at ambient temperature and pressure. True or false? Answer

2. Name the following oxides: (a) K2O; (b) Fe2O3; (c) FeO; (d) Mn2O7; (e) Cu2O; (f) CuO; (g) Al2O3; (h) CO2; (i) CO; (j) SO3; (k) SO2; (l) MgO; (m) H2O; (n) SiO2; (o) CaO; (p) ZnO. Answer

3. Write chemical formulas for the following oxides: (a) sodium oxide; (b) phosphorus (V) oxide; (c) carbon dioxide; (d) calcium oxide; (e) titanium (IV) oxide; (f) carbon monoxide; (g) silver oxide; (h) nitrogen (IV) oxide; (i) nitrogen (I) oxide; (j) tungsten (VI) oxide. Answer

4. Flammable substances burn in pure oxygen (a) more vigorously than in air; (b) as vigorously as in air; (c) less vigorously than in air. Answer

5. Propose chemical transformations for making slacked lime from limestone and write balanced chemical equations for the reactions. Answer

6. Finish and balance the following chemical equations.

(a) burning of magnesium, Mg + O2 =

(b) thermal decomposition of silver oxide, Ag2O =

(c) thermal decomposition of calcium hydroxide, Ca(OH)2 =

(d) burning of sulfur, S + O2 =

(e) hydration of sulfur trioxide, SO3 + H2O =

(f) reaction of silver oxide with sodium hydroxide, Ag2O + NaOH =

(g) combustion of carbon monoxide, CO + O2 =

(h) combustion of magnesium in carbon dioxide, Mg + CO2 =

(i) hydration of silicon dioxide, SiO2 + H2O =

(j) reduction of copper (II) oxide with hydrogen, CuO + H2 =