1893. 1 - 1893.2. Collision Theory
| Temperature (K) | $\boldsymbol{k}\left(\mathrm{L} \mathrm{mol}^{-\mathbf{1}} \mathbf{s}^{\mathbf{- 1}}\right)$ |
| :--- | :--- |
| 555 |  |
| 575 | $6.23 \times 10^{-7}$ |
| 645 | $2.42 \times 10^{-6}$ |
| 700 | $1.44 \times 10^{-4}$ |

What is the value of the activation energy (in $\mathrm{kJ} / \mathrm{mol}$ ) for this reaction?
63. The element Co exists in two oxidation states, $\mathrm{Co}(\mathrm{II})$ and $\mathrm{Co}(\mathrm{III})$, and the ions form many complexes. The rate at which one of the complexes of Co (III) was reduced by Fe (II) in water was measured. Determine the activation energy of the reaction from the following data:

| $\boldsymbol{T}(\mathrm{K})$ | $\boldsymbol{k}\left(\mathbf{s}^{-1}\right)$ |  |
| :--- | :--- | :--- |
| 293 |  | 0.054 |
| 298 | 0.100 |  |


1893. 1 - 1893.2. Collision Theory
64. The hydrolysis of the sugar sucrose to the sugars glucose and fructose,
$\mathrm{C}_{12} \mathrm{H}_{22} \mathrm{O}_{11}+\mathrm{H}_{2} \mathrm{O} \longrightarrow \mathrm{C}_{6} \mathrm{H}_{12} \mathrm{O}_{6}+\mathrm{C}_{6} \mathrm{H}_{12} \mathrm{O}_{6}$
follows a first-order rate law for the disappearance of sucrose: rate $=k\left[\mathrm{C}_{12} \mathrm{H}_{22} \mathrm{O}_{11}\right]$ (The products of the reaction, glucose and fructose, have the same molecular formulas but differ in the arrangement of the atoms in their molecules.)
(a) In neutral solution, $k=2.1 \times 10^{-11} \mathrm{~s}^{-1}$ at $27^{\circ} \mathrm{C}$ and $8.5 \times 10^{-11} \mathrm{~s}^{-1}$ at $37^{\circ} \mathrm{C}$. Determine the activation energy, the frequency factor, and the rate constant for this equation at $47^{\circ} \mathrm{C}$ (assuming the kinetics remain consistent with the Arrhenius equation at this temperature).
(b) When a solution of sucrose with an initial concentration of 0.150 M reaches equilibrium, the concentration of sucrose is $1.65 \times 10^{-7} \mathrm{M}$. How long will it take the solution to reach equilibrium at $27^{\circ} \mathrm{C}$ in the absence of a catalyst? Because the concentration of sucrose at equilibrium is so low, assume that the reaction is irreversible.
(c) Why does assuming that the reaction is irreversible simplify the calculation in part (b)?


1893. 1 - 1893.2. Collision Theory
(c) Why does assuming that the reaction is irreversible simplify the calculation in part (b)?
65. Use the PhET Reactions \& Rates interactive simulation (http://openstax.org/l/16PHETreaction) to simulate a system. On the "Single collision" tab of the simulation applet, enable the "Energy view" by clicking the " + " icon. Select the first $A+B C \longrightarrow A B+C$ reaction (A is yellow, B is purple, and C is navy blue). Using the "straight shot" default option, try launching the $A$ atom with varying amounts of energy. What changes when the Total Energy line at launch is below the transition state of the Potential Energy line? Why? What happens when it is above the transition state? Why?
66. Use the PhET Reactions \& Rates interactive simulation (http://openstax.org/l/16PHETreaction) to simulate a system. On the "Single collision" tab of the simulation applet, enable the "Energy view" by clicking the " + " icon. Select the first $A+B C \longrightarrow A B+C$ reaction (A is yellow, $B$ is purple, and $C$ is navy blue). Using the "angled shot" option, try launching the $A$ atom with varying angles, but with more Total energy than the transition state. What happens when the $A$ atom hits the $B C$ molecule from different directions? Why?


1893. 1 - 1893.3. Reaction Mechanisms
67. Why are elementary reactions involving three or more reactants very uncommon?
68. In general, can we predict the effect of doubling the concentration of $A$ on the rate of the overall reaction $A+B \longrightarrow C$ ? Can we predict the effect if the reaction is known to be an elementary reaction?
69. Define these terms:
(a) unimolecular reaction
(b) bimolecular reaction
(c) elementary reaction
(d) overall reaction
70. What is the rate law for the elementary termolecular reaction $A+2 B \longrightarrow$ products? For
$3 A \longrightarrow$ products?
71. Given the following reactions and the corresponding rate laws, in which of the reactions might the elementary reaction and the overall reaction be the same?
(a) $\mathrm{Cl}_{2}+\mathrm{CO} \longrightarrow \mathrm{Cl}_{2} \mathrm{CO}$
rate $=k\left[\mathrm{Cl}_{2}\right]^{3 / 2}[\mathrm{CO}]$
(b) $\mathrm{PCl}_{3}+\mathrm{Cl}_{2} \longrightarrow \mathrm{PCl}_{5}$ rate $=k\left[\mathrm{PCl}_{3}\right]\left[\mathrm{Cl}_{2}\right]$
(c) $2 \mathrm{NO}+\mathrm{H}_{2} \longrightarrow \mathrm{~N}_{2}+\mathrm{H}_{2} \mathrm{O}_{2}$
rate $=k[\mathrm{NO}]\left[\mathrm{H}_{2}\right]$
(d) $2 \mathrm{NO}+\mathrm{O}_{2} \longrightarrow 2 \mathrm{NO}_{2}$
rate $=k[\mathrm{NO}]^{2}\left[\mathrm{O}_{2}\right]$
(e) $\mathrm{NO}+\mathrm{O}_{3} \longrightarrow \mathrm{NO}_{2}+\mathrm{O}_{2}$
rate $=k[\mathrm{NO}]\left[\mathrm{O}_{3}\right]$
72. Write the rate law for each of the following elementary reactions:
(a) $\mathrm{O}_{3} \xrightarrow{\text { sunlight }} \mathrm{O}_{2}+\mathrm{O}$


1893. 1 - 1893.3. Reaction Mechanisms
72. Write the rate law for each of the following elementary reactions:
(a) $\mathrm{O}_{3} \xrightarrow{\text { sunlight }} \mathrm{O}_{2}+\mathrm{O}$
(b) $\mathrm{O}_{3}+\mathrm{Cl} \longrightarrow \mathrm{O}_{2}+\mathrm{ClO}$
(c) $\mathrm{ClO}+\mathrm{O} \longrightarrow \mathrm{Cl}+\mathrm{O}_{2}$
(d) $\mathrm{O}_{3}+\mathrm{NO} \longrightarrow \mathrm{NO}_{2}+\mathrm{O}_{2}$
(e) $\mathrm{NO}_{2}+\mathrm{O} \longrightarrow \mathrm{NO}+\mathrm{O}_{2}$
73. Nitrogen monoxide, NO , reacts with hydrogen, $\mathrm{H}_{2}$, according to the following equation:
$2 \mathrm{NO}+2 \mathrm{H}_{2} \longrightarrow \mathrm{~N}_{2}+2 \mathrm{H}_{2} \mathrm{O}$
What would the rate law be if the mechanism for this reaction were:
$2 \mathrm{NO}+\mathrm{H}_{2} \longrightarrow \mathrm{~N}_{2}+\mathrm{H}_{2} \mathrm{O}_{2}$ (slow)
$\mathrm{H}_{2} \mathrm{O}_{2}+\mathrm{H}_{2} \longrightarrow 2 \mathrm{H}_{2} \mathrm{O}$ (fast)
74. Experiments were conducted to study the rate of the reaction represented by this equation. ${ }^{-3}$
$2 \mathrm{NO}(g)+2 \mathrm{H}_{2}(g) \longrightarrow \mathrm{N}_{2}(g)+2 \mathrm{H}_{2} \mathrm{O}(g)$
Initial concentrations and rates of reaction are given here.


1893. 1 - 1893.3. Reaction Mechanisms
| Experiment | Initial Concentration
 $[\mathrm{NO}]\left(\mathrm{mol} \mathrm{L}^{-1}\right)$ | Initial Concentration,
 $\left[\mathrm{H}_{2}\right]\left(\mathrm{mol} \mathrm{L}^{-1} \mathrm{~min}^{-1}\right)$ | Initial Rate of Formation
 of $\mathrm{N}_{2}\left(\mathrm{~mol} \mathrm{~L}^{-1} \mathrm{~min}^{-1}\right)$ |
| :---: | :---: | :---: | :---: |
| 1 | 0.0060 | 0.0010 | $1.8 \times 10^{-4}$ |
| 2 | 0.0060 | 0.0020 | $3.6 \times 10^{-4}$ |
| 3 | 0.0010 | 0.0060 | $0.30 \times 10^{-4}$ |
| 4 | 0.0020 | 0.0060 | $1.2 \times 10^{-4}$ |

Consider the following questions:
(a) Determine the order for each of the reactants, NO and $\mathrm{H}_{2}$, from the data given and show your reasoning.
(b) Write the overall rate law for the reaction.
(c) Calculate the value of the rate constant, $k$, for the reaction. Include units.
(d) For experiment 2, calculate the concentration of NO remaining when exactly one-half of the original amount of $\mathrm{H}_{2}$ had been consumed.
(e) The following sequence of elementary steps is a proposed mechanism for the reaction.


1893. 1 - 1893.3. Reaction Mechanisms
Step 1: $\mathrm{NO}+\mathrm{NO} \rightleftharpoons \mathrm{N}_{2} \mathrm{O}_{2}$
Step 2: $\mathrm{N}_{2} \mathrm{O}_{2}+\mathrm{H}_{2} \rightleftharpoons \mathrm{H}_{2} \mathrm{O}+\mathrm{N}_{2} \mathrm{O}$
Step 3: $\mathrm{N}_{2} \mathrm{O}+\mathrm{H}_{2} \rightleftharpoons \mathrm{~N}_{2}+\mathrm{H}_{2} \mathrm{O}$
Based on the data presented, which of these is the rate determining step? Show that the mechanism is consistent with the observed rate law for the reaction and the overall stoichiometry of the reaction.
75. The reaction of CO with $\mathrm{Cl}_{2}$ gives phosgene $\left(\mathrm{COCl}_{2}\right)$, a nerve gas that was used in World War I. Use the mechanism shown here to complete the following exercises:
$\mathrm{Cl}_{2}(g) \rightleftharpoons 2 \mathrm{Cl}(g)$ (fast, $\mathrm{k}_{1}$ represents the forward rate constant, $k_{-1}$ the reverse rate constant)
$\mathrm{CO}(g)+\mathrm{Cl}(g) \longrightarrow \mathrm{COCl}(g)$ (slow, $k_{2}$ the rate constant)
$\mathrm{COCl}(g)+\mathrm{Cl}(g) \longrightarrow \mathrm{COCl}_{2}(g)$ (fast, $k_{3}$ the rate constant)
(a) Write the overall reaction.
(b) Identify all intermediates.
(c) Write the rate law for each elementary reaction.
(d) Write the overall rate law expression.


1893. 1 - 1893.4. Catalysis
76. Account for the increase in reaction rate brought about by a catalyst.
77. Compare the functions of homogeneous and heterogeneous catalysts.
78. Consider this scenario and answer the following questions: Chlorine atoms resulting from decomposition of chlorofluoromethanes, such as $\mathrm{CCl}_{2} \mathrm{~F}_{2}$, catalyze the decomposition of ozone in the atmosphere. One simplified mechanism for the decomposition is:
$\mathrm{O}_{3} \xrightarrow{\text { sunlight }} \mathrm{O}_{2}+\mathrm{O}$
$\mathrm{O}_{3}+\mathrm{Cl} \longrightarrow \mathrm{O}_{2}+\mathrm{ClO}$
$\mathrm{ClO}+\mathrm{O} \longrightarrow \mathrm{Cl}+\mathrm{O}_{2}$
(a) Explain why chlorine atoms are catalysts in the gas-phase transformation:
$2 \mathrm{O}_{3} \longrightarrow 3 \mathrm{O}_{2}$
(b) Nitric oxide is also involved in the decomposition of ozone by the mechanism:
$\mathrm{O}_{3} \xrightarrow{\text { sunlight }} \mathrm{O}_{2}+\mathrm{O}$
$\mathrm{O}_{3}+\mathrm{NO} \longrightarrow \mathrm{NO}_{2}+\mathrm{O}_{2}$
$\mathrm{NO}_{2}+\mathrm{O} \longrightarrow \mathrm{NO}+\mathrm{O}_{2}$
Is NO a catalyst for the decomposition? Explain your answer.
79. Water gas is a 1:1 mixture of carbon monoxide and hydrogen gas and is called water gas because it is formed from steam and hot carbon in the following reaction: $\mathrm{H}_{2} \mathrm{O}(g)+\mathrm{C}(s) \rightleftharpoons \mathrm{H}_{2}(g)+\mathrm{CO}(g)$. Methanol, a liquid fuel that could possibly replace gasoline, can be prepared from water gas and hydrogen at high temperature and pressure in the presence of a suitable catalyst. What will happen to the concentrations of $\mathrm{H}_{2}$, CO , and $\mathrm{CH}_{3} \mathrm{OH}$ at equilibrium if more catalyst is added?


1893. 1 - 1893.4. Catalysis
80. Nitrogen and oxygen react at high temperatures. What will happen to the concentrations of $\mathrm{N}_{2}, \mathrm{O}_{2}$, and NO at equilibrium if a catalyst is added?
81. For each of the following pairs of reaction diagrams, identify which of the pair is catalyzed:
(a)


1893. 1 - 1893.4. Catalysis
(a)
(b)

(a)

(b)

(b)
82. For each of the following pairs of reaction diagrams, identify which of the pairs is catalyzed:
(a)

83. For each of the following reaction diagrams, estimate the activation energy $\left(E_{\mathrm{a}}\right)$ of the reaction:
(a)

(b)

84. For each of the following reaction diagrams, estimate the activation energy $\left(E_{\mathrm{a}}\right)$ of the reaction:
(a)

(b)

85. Assuming the diagrams in Exercise 17.83 represent different mechanisms for the same reaction, which of the reactions has the faster rate?
86. Consider the similarities and differences in the two reaction diagrams shown in Exercise 17.84. Do these diagrams represent two different overall reactions, or do they represent the same overall reaction taking place by two different mechanisms? Explain your answer.


1894. CHAPTER 18 <br> Representative Metals, Metalloids, and Nonmetals -
Figure 18.1 Purity is extremely important when preparing silicon wafers. Technicians in a cleanroom prepare silicon without impurities (left). The CEO of VLSI Research, Don Hutcheson, shows off a pure silicon wafer (center). A silicon wafer covered in Pentium chips is an enlarged version of the silicon wafers found in many electronics used today (right). (credit middle: modification of work by "Intel Free Press"/Flickr; credit right: modification of work by Naotake Murayama)


1895. CHAPTER OUTLINE - 1895.1. Periodicity
18.2 Occurrence and Preparation of the Representative Metals
18.3 Structure and General Properties of the Metalloids
18.4 Structure and General Properties of the Nonmetals
18.5 Occurrence, Preparation, and Compounds of Hydrogen
18.6 Occurrence, Preparation, and Properties of Carbonates
18.7 Occurrence, Preparation, and Properties of Nitrogen
18.8 Occurrence, Preparation, and Properties of Phosphorus
18.9 Occurrence, Preparation, and Compounds of Oxygen
18.10 Occurrence, Preparation, and Properties of Sulfur
18.11 Occurrence, Preparation, and Properties of Halogens
18.12 Occurrence, Preparation, and Properties of the Noble Gases

INTRODUCTION The development of the periodic table in the mid-1800s came from observations that there was a periodic relationship between the properties of the elements. Chemists, who have an understanding of the variations of these properties, have been able to use this knowledge to solve a wide variety of technical challenges. For example, silicon and other semiconductors form the backbone of modern electronics because of our ability to fine-tune the electrical properties of these materials. This chapter explores important properties of representative metals, metalloids, and nonmetals in the periodic table.


1896. LEARNING OBJECTIVES -
By the end of this section, you will be able to:

- Classify elements
- Make predictions about the periodicity properties of the representative elements

We begin this section by examining the behaviors of representative metals in relation to their positions in the periodic table. The primary focus of this section will be the application of periodicity to the representative metals.

It is possible to divide elements into groups according to their electron configurations. The representative elements are elements where the $s$ and $p$ orbitals are filling. The transition elements are elements where the $d$ orbitals (groups 3-11 on the periodic table) are filling, and the inner transition metals are the elements where the $f$ orbitals are filling. The $d$ orbitals fill with the elements in group 11; therefore, the elements in group 12 qualify as representative elements because the last electron enters an $s$ orbital. Metals among the representative elements are the representative metals. Metallic character results from an element's ability to lose its outer valence electrons and results in high thermal and electrical conductivity, among other physical and chemical properties. There are 20 nonradioactive representative metals in groups $1,2,3,12,13,14$, and 15 of the periodic table (the elements shaded in yellow in Figure 18.2). The radioactive elements copernicium, flerovium, polonium, and livermorium are also metals but are beyond the scope of this chapter.

In addition to the representative metals, some of the representative elements are metalloids. A metalloid is an element that has properties that are between those of metals and nonmetals; these elements are typically semiconductors.

The remaining representative elements are nonmetals. Unlike metals, which typically form cations and ionic compounds (containing ionic bonds), nonmetals tend to form anions or molecular compounds. In general, the combination of a metal and a nonmetal produces a salt. A salt is an ionic compound consisting of cations and anions.

FIGURE 18.2 The location of the representative metals is shown in the periodic table. Nonmetals are shown in green, metalloids in purple, and the transition metals and inner transition metals in blue.


1896. LEARNING OBJECTIVES -
FIGURE 18.2 The location of the representative metals is shown in the periodic table. Nonmetals are shown in green, metalloids in purple, and the transition metals and inner transition metals in blue.

Most of the representative metals do not occur naturally in an uncombined state because they readily react with water and oxygen in the air. However, it is possible to isolate elemental beryllium, magnesium, zinc, cadmium, mercury, aluminum, tin, and lead from their naturally occurring minerals and use them because they react very slowly with air. Part of the reason why these elements react slowly is that these elements react with air to form a protective coating. The formation of this protective coating is passivation. The coating is a nonreactive film of oxide or some other compound. Elemental magnesium, aluminum, zinc, and tin are important in the fabrication of many familiar items, including wire, cookware, foil, and many household and personal objects. Although beryllium, cadmium, mercury, and lead are readily available, there are limitations in their use because of their toxicity.


1897. Group 1: The Alkali Metals -
The alkali metals lithium, sodium, potassium, rubidium, cesium, and francium constitute group 1 of the periodic table. Although hydrogen is in group 1 (and also in group 17), it is a nonmetal and deserves separate consideration later in this chapter. The name alkali metal is in reference to the fact that these metals and their oxides react with water to form very basic (alkaline) solutions.

The properties of the alkali metals are similar to each other as expected for elements in the same family. The alkali metals have the largest atomic radii and the lowest first ionization energy in their periods. This combination makes it very easy to remove the single electron in the outermost (valence) shell of each. The easy loss of this valence electron means that these metals readily form stable cations with a charge of $1+$. Their
reactivity increases with increasing atomic number due to the ease of losing the lone valence electron (decreasing ionization energy). Since oxidation is so easy, the reverse, reduction, is difficult, which explains why it is hard to isolate the elements. The solid alkali metals are very soft; lithium, shown in Figure 18.3, has the lowest density of any metal $\left(0.5 \mathrm{~g} / \mathrm{cm}^{3}\right)$.

The alkali metals all react vigorously with water to form hydrogen gas and a basic solution of the metal hydroxide. This means they are easier to oxidize than is hydrogen. As an example, the reaction of lithium with water is:

FIGURE 18.3 Lithium floats in paraffin oil because its density is less than the density of paraffin oil.
Alkali metals react directly with all the nonmetals (except the noble gases) to yield binary ionic compounds containing $1+$ metal ions. These metals are so reactive that it is necessary to avoid contact with both moisture and oxygen in the air. Therefore, they are stored in sealed containers under mineral oil, as shown in Figure 18.4 , to prevent contact with air and moisture. The pure metals never exist free (uncombined) in nature due to their high reactivity. In addition, this high reactivity makes it necessary to prepare the metals by electrolysis of alkali metal compounds.


1897. Group 1: The Alkali Metals -
FIGURE 18.4 To prevent contact with air and water, potassium for laboratory use comes as sticks or beads stored under kerosene or mineral oil, or in sealed containers. (credit: http://images-of-elements.com/potassium.php)

Unlike many other metals, the reactivity and softness of the alkali metals make these metals unsuitable for structural applications. However, there are applications where the reactivity of the alkali metals is an advantage. For example, the production of metals such as titanium and zirconium relies, in part, on the ability of sodium to reduce compounds of these metals. The manufacture of many organic compounds, including
certain dyes, drugs, and perfumes, utilizes reduction by lithium or sodium.
Sodium and its compounds impart a bright yellow color to a flame, as seen in Figure 18.5. Passing an electrical discharge through sodium vapor also produces this color. In both cases, this is an example of an emission spectrum as discussed in the chapter on electronic structure. Streetlights sometime employ sodium vapor lights because the sodium vapor penetrates fog better than most other light. This is because the fog does not scatter yellow light as much as it scatters white light. The other alkali metals and their salts also impart color to a flame. Lithium creates a bright, crimson color, whereas the others create a pale, violet color.

FIGURE 18.5 Dipping a wire into a solution of a sodium salt and then heating the wire causes emission of a bright yellow light, characteristic of sodium.


1898. LINK TO LEARNING -
This video (http://openstax.org/l/16alkalih2o) demonstrates the reactions of the alkali metals with water.


1899. Group 2: The Alkaline Earth Metals -
The alkaline earth metals (beryllium, magnesium, calcium, strontium, barium, and radium) constitute group 2 of the periodic table. The name alkaline metal comes from the fact that the oxides of the heavier members of the group react with water to form alkaline solutions. The nuclear charge increases when going from group 1 to group 2. Because of this charge increase, the atoms of the alkaline earth metals are smaller and have higher first ionization energies than the alkali metals within the same period. The higher ionization energy makes the alkaline earth metals less reactive than the alkali metals; however, they are still very reactive elements. Their reactivity increases, as expected, with increasing size and decreasing ionization energy. In chemical reactions, these metals readily lose both valence electrons to form compounds in which they exhibit an oxidation state of $2+$. Due to their high reactivity, it is common to produce the alkaline earth metals, like the alkali metals, by
electrolysis. Even though the ionization energies are low, the two metals with the highest ionization energies (beryllium and magnesium) do form compounds that exhibit some covalent characters. Like the alkali metals, the heavier alkaline earth metals impart color to a flame. As in the case of the alkali metals, this is part of the emission spectrum of these elements. Calcium and strontium produce shades of red, whereas barium produces a green color.

Magnesium is a silver-white metal that is malleable and ductile at high temperatures. Passivation decreases the reactivity of magnesium metal. Upon exposure to air, a tightly adhering layer of magnesium oxycarbonate forms on the surface of the metal and inhibits further reaction. (The carbonate comes from the reaction of carbon dioxide in the atmosphere.) Magnesium is the lightest of the widely used structural metals, which is why most magnesium production is for lightweight alloys.


1899. Group 2: The Alkaline Earth Metals -
Magnesium (shown in Figure 18.6), calcium, strontium, and barium react with water and air. At room temperature, barium shows the most vigorous reaction. The products of the reaction with water are hydrogen and the metal hydroxide. The formation of hydrogen gas indicates that the heavier alkaline earth metals are better reducing agents (more easily oxidized) than is hydrogen. As expected, these metals react with both acids and nonmetals to form ionic compounds. Unlike most salts of the alkali metals, many of the common salts of the alkaline earth metals are insoluble in water because of the high lattice energies of these compounds, containing a divalent metal ion.

FIGURE 18.6 From left to right: $\mathrm{Mg}(s)$, warm water at pH 7 , and the resulting solution with a pH greater than 7 , as indicated by the pink color of the phenolphthalein indicator. (credit: modification of work by Sahar Atwa)

The potent reducing power of hot magnesium is useful in preparing some metals from their oxides. Indeed, magnesium's affinity for oxygen is so great that burning magnesium reacts with carbon dioxide, producing elemental carbon:

$$
2 \mathrm{Mg}(s)+\mathrm{CO}_{2}(g) \longrightarrow 2 \mathrm{MgO}(s)+\mathrm{C}(s)
$$

For this reason, $\mathrm{CO}_{2}$ fire extinguisher will not extinguish a magnesium fire. Additionally, the brilliant white light emitted by burning magnesium makes it useful in flares and fireworks.


1900. Group 12 -
The elements in group 12 are transition elements; however, the last electron added is not a $d$ electron, but an $s$ electron. Since the last electron added is an $s$ electron, these elements qualify as representative metals, or post-transition metals. The group 12 elements behave more like the alkaline earth metals than transition metals. Group 12 contains the four elements zinc, cadmium, mercury, and copernicium. Each of these elements has two electrons in its outer shell $\left(n s^{2}\right)$. When atoms of these metals form cations with a charge of $2+$, where the two outer electrons are lost, they have pseudo-noble gas electron configurations. Mercury is sometimes an exception because it also exhibits an oxidation state of $1+i n$ compounds that contain a diatomic $\mathrm{Hg}_{2}{ }^{2+}$ ion. In their elemental forms and in compounds, cadmium and mercury are both toxic.

Zinc is the most reactive in group 12, and mercury is the least reactive. (This is the reverse of the reactivity
trend of the metals of groups 1 and 2, in which reactivity increases down a group. The increase in reactivity with increasing atomic number only occurs for the metals in groups 1 and 2.) The decreasing reactivity is due to the formation of ions with a pseudo-noble gas configuration and to other factors that are beyond the scope of this discussion. The chemical behaviors of zinc and cadmium are quite similar to each other but differ from that of mercury.

Zinc and cadmium have lower reduction potentials than hydrogen, and, like the alkali metals and alkaline earth metals, they will produce hydrogen gas when they react with acids. The reaction of zinc with hydrochloric acid, shown in Figure 18.7, is:

FIGURE 18.7 Zinc is an active metal. It dissolves in hydrochloric acid, forming a solution of colorless $\mathrm{Zn}^{2+}$ ions, $\mathrm{Cl}^{-}$ ions, and hydrogen gas.


1900. Group 12 -
Zinc is a silvery metal that quickly tarnishes to a blue-gray appearance. This change in color is due to an adherent coating of a basic carbonate, $\mathrm{Zn}_{2}(\mathrm{OH})_{2} \mathrm{CO}_{3}$, which passivates the metal to inhibit further corrosion. Dry cell and alkaline batteries contain a zinc anode. Brass ( Cu and Zn ) and some bronze ( Cu , Sn , and sometimes Zn ) are important zinc alloys. About half of zinc production serves to protect iron and other metals from corrosion. This protection may take the form of a sacrificial anode (also known as a galvanic anode, which is a means of providing cathodic protection for various metals) or as a thin coating on the protected metal. Galvanized steel is steel with a protective coating of zinc.


1902. Sacrificial Anodes -
A sacrificial anode, or galvanic anode, is a means of providing cathodic protection of various metals. Cathodic protection refers to the prevention of corrosion by converting the corroding metal into a cathode. As a cathode, the metal resists corrosion, which is an oxidation process. Corrosion occurs at the sacrificial anode instead of at the cathode.

The construction of such a system begins with the attachment of a more active metal (more negative reduction potential) to the metal needing protection. Attachment may be direct or via a wire. To complete the circuit, a salt bridge is necessary. This salt bridge is often seawater or ground water. Once the circuit is complete, oxidation (corrosion) occurs at the anode and not the cathode.

The commonly used sacrificial anodes are magnesium, aluminum, and zinc. Magnesium has the most negative reduction potential of the three and serves best when the salt bridge is less efficient due to a low electrolyte concentration such as in freshwater. Zinc and aluminum work better in saltwater than does magnesium. Aluminum is lighter than zinc and has a higher capacity; however, an oxide coating may passivate the aluminum. In special cases, other materials are useful. For example, iron will protect copper.

Mercury is very different from zinc and cadmium. Mercury is the only metal that is liquid at $25^{\circ} \mathrm{C}$. Many metals dissolve in mercury, forming solutions called amalgams (see the feature on Amalgams), which are alloys of mercury with one or more other metals. Mercury, shown in Figure 18.8, is a nonreactive element that is more difficult to oxidize than hydrogen. Thus, it does not displace hydrogen from acids; however, it will react with strong oxidizing acids, such as nitric acid:

$$
\begin{gathered}
\mathrm{Hg}(l)+\mathrm{HCl}(a q) \longrightarrow \text { no reaction } \\
3 \mathrm{Hg}(l)+8 \mathrm{HNO}_{3}(a q) \longrightarrow 3 \mathrm{Hg}\left(\mathrm{NO}_{3}\right)_{2}(a q)+4 \mathrm{H}_{2} \mathrm{O}(l)+2 \mathrm{NO}(g)
\end{gathered}
$$


1902. Sacrificial Anodes -
The clear NO initially formed quickly undergoes further oxidation to the reddish brown $\mathrm{NO}_{2}$.

FIGURE 18.8 From left to right: $\mathrm{Hg}\left(\mathrm{l}, \mathrm{Hg}+\right.$ concentrated $\mathrm{HCl}, \mathrm{Hg}+$ concentrated $\mathrm{HNO}_{3}$. (credit: Sahar Atwa)
Most mercury compounds decompose when heated. Most mercury compounds contain mercury with a 2+oxidation state. When there is a large excess of mercury, it is possible to form compounds containing the $\mathrm{Hg}_{2}{ }^{2+}$ ion. All mercury compounds are toxic, and it is necessary to exercise great care in their synthesis.


1904. Amalgams -
An amalgam is an alloy of mercury with one or more other metals. This is similar to considering steel to be an alloy of iron with other metals. Most metals will form an amalgam with mercury, with the main exceptions being iron, platinum, tungsten, and tantalum.

Due to toxicity issues with mercury, there has been a significant decrease in the use of amalgams. Historically, amalgams were important in electrolytic cells and in the extraction of gold. Amalgams of the alkali metals still find use because they are strong reducing agents and easier to handle than the pure alkali metals.

Prospectors had a problem when they found finely divided gold. They learned that adding mercury to their pans collected the gold into the mercury to form an amalgam for easier collection. Unfortunately, losses of small amounts of mercury over the years left many streams in California polluted with mercury.

Dentists use amalgams containing silver and other metals to fill cavities. There are several reasons to use an amalgam including low cost, ease of manipulation, and longevity compared to alternate materials. Dental amalgams are approximately $50 \%$ mercury by weight, which, in recent years, has become a concern due to the toxicity of mercury.

After reviewing the best available data, the Food and Drug Administration (FDA) considers amalgam-based fillings to be safe for adults and children over six years of age. Even with multiple fillings, the mercury levels in the patients remain far below the lowest levels associated with harm. Clinical studies have found no link between dental amalgams and health problems. Health issues may not be the same in cases of
children under six or pregnant women. The FDA conclusions are in line with the opinions of the Environmental Protection Agency (EPA) and Centers for Disease Control (CDC). The only health consideration noted is that some people are allergic to the amalgam or one of its components.


1905. Group 13 -
Group 13 contains the metalloid boron and the metals aluminum, gallium, indium, and thallium. The lightest element, boron, is semiconducting, and its binary compounds tend to be covalent and not ionic. The remaining elements of the group are metals, but their oxides and hydroxides change characters. The oxides and hydroxides of aluminum and gallium exhibit both acidic and basic behaviors. A substance, such as these two, that will react with both acids and bases is amphoteric. This characteristic illustrates the combination of nonmetallic and metallic behaviors of these two elements. Indium and thallium oxides and hydroxides exhibit only basic behavior, in accordance with the clearly metallic character of these two elements. The melting point of gallium is unusually low (about $30^{\circ} \mathrm{C}$ ) and will melt in your hand.

Aluminum is amphoteric because it will react with both acids and bases. A typical reaction with an acid is:

$$
2 \mathrm{Al}(s)+6 \mathrm{HCl}(a q) \longrightarrow 2 \mathrm{AlCl}_{3}(a q)+3 \mathrm{H}_{2}(g)
$$

The products of the reaction of aluminum with a base depend upon the reaction conditions, with the following being one possibility:

$$
2 \mathrm{Al}(s)+2 \mathrm{NaOH}(a q)+6 \mathrm{H}_{2} \mathrm{O}(l) \longrightarrow 2 \mathrm{Na}\left\mathrm{Al}(\mathrm{OH})_{4}\right+3 \mathrm{H}_{2}(g)
$$


1905. Group 13 -
With both acids and bases, the reaction with aluminum generates hydrogen gas.
The group 13 elements have a valence shell electron configuration of $n s^{2} n p^{1}$. Aluminum normally uses all of its valence electrons when it reacts, giving compounds in which it has an oxidation state of $3+$. Although many of these compounds are covalent, others, such as $\mathrm{AlF}_{3}$ and $\mathrm{Al}_{2}\left(\mathrm{SO}_{4}\right)_{3}$, are ionic. Aqueous solutions of aluminum salts contain the cation $\left[\mathrm{Al}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]^{3+}$, abbreviated as $\mathrm{Al}^{3+}(a q)$. Gallium, indium, and thallium also form ionic compounds containing $\mathrm{M}^{3+}$ ions. These three elements exhibit not only the expected oxidation state of $3+$ from the three valence electrons but also an oxidation state (in this case, $1+$ ) that is two below the expected value. This phenomenon, the inert pair effect, refers to the formation of a stable ion with an oxidation state two lower than expected for the group. The pair of electrons is the valence $s$ orbital for those elements. In general, the inert pair effect is important for the lower $p$-block elements. In an aqueous solution, the $\mathrm{Tl}^{+}(a q)$ ion is more stable than is $\mathrm{Tl}^{3+}(a q)$. In general, these metals will react with air and water to form $3+\mathrm{ions}$; however, thallium reacts to give thallium(I) derivatives. The metals of group 13 all react directly with nonmetals such as sulfur, phosphorus, and the halogens, forming binary compounds.


1905. Group 13 -
The metals of group 13 ( $\mathrm{Al}, \mathrm{Ga}, \mathrm{In}$, and Tl ) are all reactive. However, passivation occurs as a tough, hard, thin film of the metal oxide forms upon exposure to air. Disruption of this film may counter the passivation, allowing the metal to react. One way to disrupt the film is to expose the passivated metal to mercury. Some of the metal dissolves in the mercury to form an amalgam, which sheds the protective oxide layer to expose the metal to further reaction. The formation of an amalgam allows the metal to react with air and water.


1906. LINK TO LEARNING -
Although easily oxidized, the passivation of aluminum makes it very useful as a strong, lightweight building material. Because of the formation of an amalgam, mercury is corrosive to structural materials made of aluminum. This video (http://openstax.org/l/16aluminumhg) demonstrates how the integrity of an aluminum beam can be destroyed by the addition of a small amount of elemental mercury.

The most important uses of aluminum are in the construction and transportation industries, and in the manufacture of aluminum cans and aluminum foil. These uses depend on the lightness, toughness, and strength of the metal, as well as its resistance to corrosion. Because aluminum is an excellent conductor of heat and resists corrosion, it is useful in the manufacture of cooking utensils.

Aluminum is a very good reducing agent and may replace other reducing agents in the isolation of certain metals from their oxides. Although more expensive than reduction by carbon, aluminum is important in the isolation of $\mathrm{Mo}, \mathrm{W}$, and Cr from their oxides.


1907. Group 14 -
The metallic members of group 14 are tin, lead, and flerovium. Carbon is a typical nonmetal. The remaining elements of the group, silicon and germanium, are examples of semimetals or metalloids. Tin and lead form the stable divalent cations, $\mathrm{Sn}^{2+}$ and $\mathrm{Pb}^{2+}$, with oxidation states two below the group oxidation state of $4+$. The stability of this oxidation state is a consequence of the inert pair effect. Tin and lead also form covalent compounds with a formal $4+$-oxidation state. For example, $\mathrm{SnCl}_{4}$ and $\mathrm{PbCl}_{4}$ are low-boiling covalent liquids.

FIGURE 18.9 (a) Tin(II) chloride is an ionic solid; (b) tin(IV) chloride is a covalent liquid.
Tin reacts readily with nonmetals and acids to form tin(II) compounds (indicating that it is more easily oxidized than hydrogen) and with nonmetals to form either tin(II) or tin(IV) compounds (shown in Figure 18.9), depending on the stoichiometry and reaction conditions. Lead is less reactive. It is only slightly easier to oxidize than hydrogen, and oxidation normally requires a hot concentrated acid.


1907. Group 14 -
Many of these elements exist as allotropes. Allotropes are two or more forms of the same element in the same physical state with different chemical and physical properties. There are two common allotropes of tin. These allotropes are grey (brittle) tin and white tin. As with other allotropes, the difference between these forms of tin is in the arrangement of the atoms. White tin is stable above $13.2^{\circ} \mathrm{C}$ and is malleable like other metals. At low temperatures, gray tin is the more stable form. Gray tin is brittle and tends to break down to a powder. Consequently, articles made of tin will disintegrate in cold weather, particularly if the cold spell is lengthy. The change progresses slowly from the spot of origin, and the gray tin that is first formed catalyzes further change. In a way, this effect is similar to the spread of an infection in a plant or animal body, leading people to call this process tin disease or tin pest.

The principal use of tin is in the coating of steel to form tin plate-sheet iron, which constitutes the tin in tin cans. Important tin alloys are bronze ( Cu and Sn ) and solder ( Sn and Pb ). Lead is important in the lead storage batteries in automobiles.


1908. Group 15 -
Bismuth, the heaviest member of group 15, is a less reactive metal than the other representative metals. It readily gives up three of its five valence electrons to active nonmetals to form the tri-positive ion, $\mathrm{Bi}^{3+}$. It forms compounds with the group oxidation state of $5+$ only when treated with strong oxidizing agents. The stability of the $3+$-oxidation state is another example of the inert pair effect.


1909. LEARNING OBJECTIVES -
By the end of this section, you will be able to:

- Identify natural sources of representative metals
- Describe electrolytic and chemical reduction processes used to prepare these elements from natural sources

Because of their reactivity, we do not find most representative metals as free elements in nature. However, compounds that contain ions of most representative metals are abundant. In this section, we will consider the two common techniques used to isolate the metals from these compounds-electrolysis and chemical reduction.

These metals primarily occur in minerals, with lithium found in silicate or phosphate minerals, and sodium and potassium found in salt deposits from evaporation of ancient seas and in silicates. The alkaline earth metals occur as silicates and, with the exception of beryllium, as carbonates and sulfates. Beryllium occurs as the mineral beryl, $\mathrm{Be}_{3} \mathrm{Al}_{2} \mathrm{Si}_{6} \mathrm{O}_{18}$, which, with certain impurities, may be either the gemstone emerald or aquamarine. Magnesium is in seawater and, along with the heavier alkaline earth metals, occurs as silicates, carbonates, and sulfates. Aluminum occurs abundantly in many types of clay and in bauxite, an impure aluminum oxide hydroxide. The principle tin ore is the oxide cassiterite, $\mathrm{SnO}_{2}$, and the principle lead and thallium ores are the sulfides or the products of weathering of the sulfides. The remaining representative metals occur as impurities in zinc or aluminum ores.


1910. Electrolysis -
Ions of metals in of groups 1 and 2, along with aluminum, are very difficult to reduce; therefore, it is necessary to prepare these elements by electrolysis, an important process discussed in the chapter on electrochemistry. Briefly, electrolysis involves using electrical energy to drive unfavorable chemical reactions to completion; it is useful in the isolation of reactive metals in their pure forms. Sodium, aluminum, and magnesium are typical examples.

The Preparation of Sodium
The most important method for the production of sodium is the electrolysis of molten sodium chloride; the set-up is a Downs cell, shown in Figure 18.10. The reaction involved in this process is:

$$
2 \mathrm{NaCl}(l) \xrightarrow[600^{\circ} \mathrm{C}]{\text { electrolysis }} 2 \mathrm{Na}(l)+\mathrm{Cl}_{2}(g)
$$

The electrolysis cell contains molten sodium chloride (melting point $801^{\circ} \mathrm{C}$ ), to which calcium chloride has been added to lower the melting point to $600^{\circ} \mathrm{C}$ (a colligative effect). The passage of a direct current through the cell causes the sodium ions to migrate to the negatively charged cathode and pick up electrons, reducing the ions to sodium metal. Chloride ions migrate to the positively charged anode, lose electrons, and undergo oxidation to chlorine gas. The overall cell reaction comes from adding the following reactions:

$$
\begin{aligned}
& \text { at the cathode: } 2 \mathrm{Na}^{+}+2 \mathrm{e}^{-} \longrightarrow 2 \mathrm{Na}(l) \\
& \text { at the anode: } 2 \mathrm{Cl}^{-} \longrightarrow \mathrm{Cl}_{2}(g)+2 \mathrm{e}^{-} \\
& \text {overall change: } 2 \mathrm{Na}^{+}+2 \mathrm{Cl}^{-} \longrightarrow 2 \mathrm{Na}(l)+\mathrm{Cl}_{2}(g)
\end{aligned}
$$


1910. Electrolysis -
Separation of the molten sodium and chlorine prevents recombination. The liquid sodium, which is less dense than molten sodium chloride, floats to the surface and flows into a collector. The gaseous chlorine goes to storage tanks. Chlorine is also a valuable product.

FIGURE 18.10 Pure sodium metal is isolated by electrolysis of molten sodium chloride using a Downs cell. It is not possible to isolate sodium by electrolysis of aqueous solutions of sodium salts because hydrogen ions are more easily reduced than are sodium ions; as a result, hydrogen gas forms at the cathode instead of the desired sodium metal. The high temperature required to melt NaCl means that liquid sodium metal forms.

The Preparation of Aluminum
The preparation of aluminum utilizes a process invented in 1886 by Charles M. Hall, who began to work on the problem while a student at Oberlin College in Ohio. Paul L. T. Héroult discovered the process independently a month or two later in France. In honor to the two inventors, this electrolysis cell is known as the Hall-Héroult cell. The Hall-Héroult cell is an electrolysis cell for the production of aluminum. Figure 18.11 illustrates the Hall-Héroult cell.

The production of aluminum begins with the purification of bauxite, the most common source of aluminum. The reaction of bauxite, $\mathrm{AlO}(\mathrm{OH})$, with hot sodium hydroxide forms soluble sodium aluminate, while clay and other impurities remain undissolved:

$$
\mathrm{AlO}(\mathrm{OH})(s)+\mathrm{NaOH}(a q)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{Na}\left\mathrm{Al}(\mathrm{OH})_{4}\right
$$

After the removal of the impurities by filtration, the addition of acid to the aluminate leads to the reprecipitation of aluminum hydroxide:


1910. Electrolysis -
After the removal of the impurities by filtration, the addition of acid to the aluminate leads to the reprecipitation of aluminum hydroxide:

$$
\mathrm{Na}\left\mathrm{Al}(\mathrm{OH})_{4}\right+\mathrm{H}_{3} \mathrm{O}^{+}(a q) \longrightarrow \mathrm{Al}(\mathrm{OH})_{3}(s)+\mathrm{Na}^{+}(a q)+2 \mathrm{H}_{2} \mathrm{O}(l)
$$

The next step is to remove the precipitated aluminum hydroxide by filtration. Heating the hydroxide produces aluminum oxide, $\mathrm{Al}_{2} \mathrm{O}_{3}$, which dissolves in a molten mixture of cryolite, $\mathrm{Na}_{3} \mathrm{AlF}_{6}$, and calcium fluoride, $\mathrm{CaF}_{2}$. Electrolysis of this solution takes place in a cell like that shown in Figure 18.11. Reduction of aluminum ions to the metal occurs at the cathode, while oxygen, carbon monoxide, and carbon dioxide form at the anode.

FIGURE 18.11 An electrolytic cell is used for the production of aluminum. The electrolysis of a solution of cryolite and calcium fluoride results in aluminum metal at the cathode, and oxygen, carbon monoxide, and carbon dioxide at the anode.

The Preparation of Magnesium
Magnesium is the other metal that is isolated in large quantities by electrolysis. Seawater, which contains approximately $0.5 \%$ magnesium chloride, serves as the major source of magnesium. Addition of calcium hydroxide to seawater precipitates magnesium hydroxide. The addition of hydrochloric acid to magnesium hydroxide, followed by evaporation of the resultant aqueous solution, leaves pure magnesium chloride. The electrolysis of molten magnesium chloride forms liquid magnesium and chlorine gas:


1910. Electrolysis -
$$
\begin{gathered}
\mathrm{MgCl}_{2}(a q)+\mathrm{Ca}(\mathrm{OH})_{2}(a q) \longrightarrow \mathrm{Mg}(\mathrm{OH})_{2}(s)+\mathrm{CaCl}_{2}(a q) \\
\mathrm{Mg}(\mathrm{OH})_{2}(s)+2 \mathrm{HCl}(a q) \longrightarrow \mathrm{MgCl}_{2}(a q)+2 \mathrm{H}_{2} \mathrm{O}(l) \\
\mathrm{MgCl}_{2}(l) \longrightarrow \mathrm{Mg}(l)+\mathrm{Cl}_{2}(g)
\end{gathered}
$$

Some production facilities have moved away from electrolysis completely. In the next section, we will see how the Pidgeon process leads to the chemical reduction of magnesium.


1911. Chemical Reduction -
It is possible to isolate many of the representative metals by chemical reduction using other elements as reducing agents. In general, chemical reduction is much less expensive than electrolysis, and for this reason, chemical reduction is the method of choice for the isolation of these elements. For example, it is possible to produce potassium, rubidium, and cesium by chemical reduction, as it is possible to reduce the molten chlorides of these metals with sodium metal. This may be surprising given that these metals are more reactive than sodium; however, the metals formed are more volatile than sodium and can be distilled for collection. The removal of the metal vapor leads to a shift in the equilibrium to produce more metal (see how reactions can be driven in the discussions of Le Châtelier's principle in the chapter on fundamental equilibrium concepts).

The production of magnesium, zinc, and tin provide additional examples of chemical reduction.

The Preparation of Magnesium
The Pidgeon process involves the reaction of magnesium oxide with elemental silicon at high temperatures to form pure magnesium:

$$
\mathrm{Si}(s)+2 \mathrm{MgO}(s) \xrightarrow{\Delta} \mathrm{SiO}_{2}(s)+2 \mathrm{Mg}(g)
$$

Although this reaction is unfavorable in terms of thermodynamics, the removal of the magnesium vapor produced takes advantage of Le Châtelier's principle to continue the forward progress of the reaction. Over $75 \%$ of the world's production of magnesium, primarily in China, comes from this process.

The Preparation of Zinc
Zinc ores usually contain zinc sulfide, zinc oxide, or zinc carbonate. After separation of these compounds from the ores, heating in air converts the ore to zinc oxide by one of the following reactions:

$$
\begin{gathered}
2 \mathrm{ZnS}(s)+3 \mathrm{O}_{2}(g) \xrightarrow{\Delta} 2 \mathrm{ZnO}(s)+2 \mathrm{SO}_{2}(g) \\
\mathrm{ZnCO}_{3}(s) \xrightarrow{\Delta} \mathrm{ZnO}(s)+\mathrm{CO}_{2}(g)
\end{gathered}
$$

Carbon, in the form of coal, reduces the zinc oxide to form zinc vapor:


1911. Chemical Reduction -
Carbon, in the form of coal, reduces the zinc oxide to form zinc vapor:

$$
\mathrm{ZnO}(s)+\mathrm{C}(s) \longrightarrow \mathrm{Zn}(g)+\mathrm{CO}(g)
$$

The zinc can be distilled (boiling point $907^{\circ} \mathrm{C}$ ) and condensed. This zinc contains impurities of cadmium ( 767 $\left.{ }^{\circ} \mathrm{C}\right)$, iron $\left(2862^{\circ} \mathrm{C}\right)$, lead $\left(1750^{\circ} \mathrm{C}\right)$, and arsenic $\left(613^{\circ} \mathrm{C}\right)$. Careful redistillation produces pure zinc. Arsenic and cadmium are distilled from the zinc because they have lower boiling points. At higher temperatures, the zinc is distilled from the other impurities, mainly lead and iron.


1912. The Preparation of Tin -
The ready reduction of tin(IV) oxide by the hot coals of a campfire accounts for the knowledge of tin in the ancient world. In the modern process, the roasting of tin ores containing $\mathrm{SnO}_{2}$ removes contaminants such as arsenic and sulfur as volatile oxides. Treatment of the remaining material with hydrochloric acid removes the oxides of other metals. Heating the purified ore with carbon at temperature above $1000{ }^{\circ} \mathrm{C}$ produces tin:

$$
\mathrm{SnO}_{2}(s)+2 \mathrm{C}(s) \xrightarrow{\Delta} \mathrm{Sn}(s)+2 \mathrm{CO}(g)
$$

The molten tin collects at the bottom of the furnace and is drawn off and cast into blocks.


1912. The Preparation of Tin - 1912.1. Structure and General Properties of the Metalloids
LEARNING OBJECTIVES
By the end of this section, you will be able to:

- Describe the general preparation, properties, and uses of the metalloids
- Describe the preparation, properties, and compounds of boron and silicon

A series of six elements called the metalloids separate the metals from the nonmetals in the periodic table. The metalloids are boron, silicon, germanium, arsenic, antimony, and tellurium. These elements look metallic; however, they do not conduct electricity as well as metals so they are semiconductors. They are semiconductors because their electrons are more tightly bound to their nuclei than are those of metallic conductors. Their chemical behavior falls between that of metals and nonmetals. For example, the pure metalloids form covalent crystals like the nonmetals, but like the metals, they generally do not form monatomic anions. This intermediate behavior is in part due to their intermediate electronegativity values. In this section, we will briefly discuss the chemical behavior of metalloids and deal with two of these elements-boron and silicon-in more detail.

The metalloid boron exhibits many similarities to its neighbor carbon and its diagonal neighbor silicon. All three elements form covalent compounds. However, boron has one distinct difference in that its $2 s^{2} 2 p^{1}$ outer electron structure gives it one less valence electron than it has valence orbitals. Although boron exhibits an oxidation state of $3+$ in most of its stable compounds, this electron deficiency provides boron with the ability to form other, sometimes fractional, oxidation states, which occur, for example, in the boron hydrides.


1912. The Preparation of Tin - 1912.1. Structure and General Properties of the Metalloids
Silicon has the valence shell electron configuration $3 s^{2} 3 p^{2}$, and it commonly forms tetrahedral structures in which it is $s p^{3}$ hybridized with a formal oxidation state of $4+$. The major differences between the chemistry of carbon and silicon result from the relative strength of the carbon-carbon bond, carbon's ability to form stable bonds to itself, and the presence of the empty $3 d$ valence-shell orbitals in silicon. Silicon's empty $d$ orbitals and boron's empty $p$ orbital enable tetrahedral silicon compounds and trigonal planar boron compounds to act as Lewis acids. Carbon, on the other hand, has no available valence shell orbitals; tetrahedral carbon compounds cannot act as Lewis acids. Germanium is very similar to silicon in its chemical behavior.

Arsenic and antimony generally form compounds in which an oxidation state of $3+$ or $5+$ is exhibited; however, arsenic can form arsenides with an oxidation state of $3-$. These elements tarnish only slightly in dry air but readily oxidize when warmed.

Tellurium combines directly with most elements. The most stable tellurium compounds are the tellurides-salts of $\mathrm{Te}^{2-}$ formed with active metals and lanthanides-and compounds with oxygen, fluorine, and chlorine, in which tellurium normally exhibits an oxidation state $2+$ or $4+$. Although tellurium(VI) compounds are known (for example, $\mathrm{TeF}_{6}$ ), there is a marked resistance to oxidation to this maximum group oxidation state.


1913. Structures of the Metalloids -
Covalent bonding is the key to the crystal structures of the metalloids. In this regard, these elements resemble nonmetals in their behavior.

Elemental silicon, germanium, arsenic, antimony, and tellurium are lustrous, metallic-looking solids. Silicon and germanium crystallize with a diamond structure. Each atom within the crystal has covalent bonds to four neighboring atoms at the corners of a regular tetrahedron. Single crystals of silicon and germanium are giant, three-dimensional molecules. There are several allotropes of arsenic with the most stable being layer like and containing puckered sheets of arsenic atoms. Each arsenic atom forms covalent bonds to three other atoms within the sheet. The crystal structure of antimony is similar to that of arsenic, both shown in Figure 18.12. The structures of arsenic and antimony are similar to the structure of graphite, covered later in this chapter. Tellurium forms crystals that contain infinite spiral chains of tellurium atoms. Each atom in the chain bonds to two other atoms.


1914. LINK TO LEARNING -
Explore a cubic diamond (http://openstax.org/l/16crystal) crystal structure.

FIGURE 18.12 (a) Arsenic and (b) antimony have a layered structure similar to that of (c) graphite, except that the layers are puckered rather than planar. (d) Elemental tellurium forms spiral chains.

Pure crystalline boron is transparent. The crystals consist of icosahedra, as shown in Figure 18.13, with a boron atom at each corner. In the most common form of boron, the icosahedra pack together in a manner similar to the cubic closest packing of spheres. All boron-boron bonds within each icosahedron are identical and are approximately 176 pm in length. In the different forms of boron, there are different arrangements and connections between the icosahedra.

FIGURE 18.13 An icosahedron is a symmetrical, solid shape with 20 faces, each of which is an equilateral triangle. The faces meet at 12 corners.

The name silicon is derived from the Latin word for flint, silex. The metalloid silicon readily forms compounds containing Si-O-Si bonds, which are of prime importance in the mineral world. This bonding capability is in contrast to the nonmetal carbon, whose ability to form carbon-carbon bonds gives it prime importance in the plant and animal worlds.

Occurrence, Preparation, and Compounds of Boron and Silicon
Boron constitutes less than $0.001 \%$ by weight of the earth's crust. In nature, it only occurs in compounds with oxygen. Boron is widely distributed in volcanic regions as boric acid, $\mathrm{B}(\mathrm{OH})_{3}$, and in dry lake regions, including the desert areas of California, as borates and salts of boron oxyacids, such as borax, $\mathrm{Na}_{2} \mathrm{~B}_{4} \mathrm{O}_{7} \cdot 10 \mathrm{H}_{2} \mathrm{O}$.


1914. LINK TO LEARNING -
Elemental boron is chemically inert at room temperature, reacting with only fluorine and oxygen to form boron trifluoride, $\mathrm{BF}_{3}$, and boric oxide, $\mathrm{B}_{2} \mathrm{O}_{3}$, respectively. At higher temperatures, boron reacts with all nonmetals, except tellurium and the noble gases, and with nearly all metals; it oxidizes to $\mathrm{B}_{2} \mathrm{O}_{3}$ when heated with concentrated nitric or sulfuric acid. Boron does not react with nonoxidizing acids. Many boron compounds react readily with water to give boric acid, $\mathrm{B}(\mathrm{OH})_{3}$ (sometimes written as $\mathrm{H}_{3} \mathrm{BO}_{3}$ ).

Reduction of boric oxide with magnesium powder forms boron ( $95-98.5 \%$ pure) as a brown, amorphous powder:

$$
\mathrm{B}_{2} \mathrm{O}_{3}(s)+3 \mathrm{Mg}(s) \longrightarrow 2 \mathrm{~B}(s)+3 \mathrm{MgO}(s)
$$

An amorphous substance is a material that appears to be a solid, but does not have a long-range order like a true solid. Treatment with hydrochloric acid removes the magnesium oxide. Further purification of the boron begins with conversion of the impure boron into boron trichloride. The next step is to heat a mixture of boron trichloride and hydrogen:

$$
2 \mathrm{BCl}_{3}(g)+3 \mathrm{H}_{2}(g) \xrightarrow{1500^{\circ} \mathrm{C}} 2 \mathrm{~B}(s)+6 \mathrm{HCl}(g) \quad \Delta H^{\circ}=253.7 \mathrm{~kJ}
$$


1914. LINK TO LEARNING -
Silicon makes up nearly one-fourth of the mass of the earth's crust-second in abundance only to oxygen. The crust is composed almost entirely of minerals in which the silicon atoms are at the center of the silicon-oxygen tetrahedron, which connect in a variety of ways to produce, among other things, chains, layers, and threedimensional frameworks. These minerals constitute the bulk of most common rocks, soil, and clays. In addition, materials such as bricks, ceramics, and glasses contain silicon compounds.

It is possible to produce silicon by the high-temperature reduction of silicon dioxide with strong reducing agents, such as carbon and magnesium:

$$
\begin{aligned}
& \mathrm{SiO}_{2}(s)+2 \mathrm{C}(s) \xrightarrow{\Delta} \mathrm{Si}(s)+2 \mathrm{CO}(g) \\
& \mathrm{SiO}_{2}(s)+2 \mathrm{Mg}(s) \xrightarrow{\Delta} \mathrm{Si}(s)+2 \mathrm{MgO}(s)
\end{aligned}
$$

Extremely pure silicon is necessary for the manufacture of semiconductor electronic devices. This process begins with the conversion of impure silicon into silicon tetrahalides, or silane ( $\mathrm{SiH}_{4}$ ), followed by
decomposition at high temperatures. Zone refining, illustrated in Figure 18.14, completes the purification. In this method, a rod of silicon is heated at one end by a heat source that produces a thin cross-section of molten silicon. Slowly lowering the rod through the heat source moves the molten zone from one end of the rod to other. As this thin, molten region moves, impurities in the silicon dissolve in the liquid silicon and move with the molten region. Ultimately, the impurities move to one end of the rod, which is then cut off.

FIGURE 18.14 A zone-refining apparatus used to purify silicon.
This highly purified silicon, containing no more than one part impurity per million parts of silicon, is the most important element in the computer industry. Pure silicon is necessary in semiconductor electronic devices such as transistors, computer chips, and solar cells.


1914. LINK TO LEARNING -
Like some metals, passivation of silicon occurs due the formation of a very thin film of oxide (primarily silicon dioxide, $\mathrm{SiO}_{2}$ ). Silicon dioxide is soluble in hot aqueous base; thus, strong bases destroy the passivation. Removal of the passivation layer allows the base to dissolve the silicon, forming hydrogen gas and silicate anions. For example:

$$
\mathrm{Si}(s)+4 \mathrm{OH}^{-}(a q) \longrightarrow \mathrm{SiO}_{4}{ }^{4-}(a q)+2 \mathrm{H}_{2}(g)
$$

Silicon reacts with halogens at high temperatures, forming volatile tetrahalides, such as $\operatorname{SiF}_{4}$.
Unlike carbon, silicon does not readily form double or triple bonds. Silicon compounds of the general formula $\mathrm{SiX}_{4}$, where X is a highly electronegative group, can act as Lewis acids to form six-coordinate silicon. For example, silicon tetrafluoride, $\mathrm{SiF}_{4}$, reacts with sodium fluoride to yield $\mathrm{Na}_{2}\left[\mathrm{SiF}_{6}\right]$, which contains the octahedral $\left[\mathrm{SiF}_{6}\right]^{2-}$ ion in which silicon is $s p^{3} d^{2}$ hybridized:

$$
2 \mathrm{NaF}(s)+\mathrm{SiF}_{4}(g) \longrightarrow \mathrm{Na}_{2} \mathrm{SiF}_{6}(s)
$$

Antimony reacts readily with stoichiometric amounts of fluorine, chlorine, bromine, or iodine, yielding trihalides or, with excess fluorine or chlorine, forming the pentahalides $\mathrm{SbF}_{5}$ and $\mathrm{SbCl}_{5}$. Depending on the stoichiometry, it forms antimony(III) sulfide, $\mathrm{Sb}_{2} \mathrm{~S}_{3}$, or antimony(V) sulfide when heated with sulfur. As expected, the metallic nature of the element is greater than that of arsenic, which lies immediately above it in group 15.


1915. Boron and Silicon Halides -
Boron trihalides $-\mathrm{BF}_{3}, \mathrm{BCl}_{3}, \mathrm{BBr}_{3}$, and $\mathrm{BI}_{3}-$ can be prepared by the direct reaction of the elements. These nonpolar molecules contain boron with $s p^{2}$ hybridization and a trigonal planar molecular geometry. The fluoride and chloride compounds are colorless gasses, the bromide is a liquid, and the iodide is a white crystalline solid.

Except for boron trifluoride, the boron trihalides readily hydrolyze in water to form boric acid and the corresponding hydrohalic acid. Boron trichloride reacts according to the equation:

$$
\mathrm{BCl}_{3}(g)+3 \mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{B}(\mathrm{OH})_{3}(a q)+3 \mathrm{HCl}(a q)
$$

Boron trifluoride reacts with hydrofluoric acid, to yield a solution of fluoroboric acid, $\mathrm{HBF}_{4}$ :

$$
\mathrm{BF}_{3}(a q)+\mathrm{HF}(a q)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{H}_{3} \mathrm{O}^{+}(a q)+\mathrm{BF}_{4}^{-}(a q)
$$

In this reaction, the $\mathrm{BF}_{3}$ molecule acts as the Lewis acid (electron pair acceptor) and accepts a pair of electrons from a fluoride ion:

All the tetrahalides of silicon, $\mathrm{SiX}_{4}$, have been prepared. Silicon tetrachloride can be prepared by direct chlorination at elevated temperatures or by heating silicon dioxide with chlorine and carbon:

$$
\mathrm{SiO}_{2}(s)+2 \mathrm{C}(s)+2 \mathrm{Cl}_{2}(g) \xrightarrow{\Delta} \mathrm{SiCl}_{4}(g)+2 \mathrm{CO}(g)
$$

Silicon tetrachloride is a covalent tetrahedral molecule, which is a nonpolar, low-boiling $\left(57^{\circ} \mathrm{C}\right)$, colorless liquid.


1915. Boron and Silicon Halides -
Silicon tetrachloride is a covalent tetrahedral molecule, which is a nonpolar, low-boiling $\left(57^{\circ} \mathrm{C}\right)$, colorless liquid.

It is possible to prepare silicon tetrafluoride by the reaction of silicon dioxide with hydrofluoric acid:

$$
\mathrm{SiO}_{2}(s)+4 \mathrm{HF}(g) \longrightarrow \mathrm{SiF}_{4}(g)+2 \mathrm{H}_{2} \mathrm{O}(l) \quad \Delta H^{\circ}=-191.2 \mathrm{~kJ}
$$

Hydrofluoric acid is the only common acid that will react with silicon dioxide or silicates. This reaction occurs because the silicon-fluorine bond is the only bond that silicon forms that is stronger than the silicon-oxygen bond. For this reason, it is possible to store all common acids, other than hydrofluoric acid, in glass containers.

Except for silicon tetrafluoride, silicon halides are extremely sensitive to water. Upon exposure to water, $\mathrm{SiCl}_{4}$ reacts rapidly with hydroxide groups, replacing all four chlorine atoms to produce unstable orthosilicic acid, $\mathrm{Si}(\mathrm{OH})_{4}$ or $\mathrm{H}_{4} \mathrm{SiO}_{4}$, which slowly decomposes into $\mathrm{SiO}_{2}$.

Boron and Silicon Oxides and Derivatives
Boron burns at $700^{\circ} \mathrm{C}$ in oxygen, forming boric oxide, $\mathrm{B}_{2} \mathrm{O}_{3}$. Boric oxide is necessary for the production of heat-resistant borosilicate glass, like that shown in Figure 18.15 and certain optical glasses. Boric oxide dissolves in hot water to form boric acid, $\mathrm{B}(\mathrm{OH})_{3}$ :

$$
\mathrm{B}_{2} \mathrm{O}_{3}(s)+3 \mathrm{H}_{2} \mathrm{O}(l) \longrightarrow 2 \mathrm{~B}(\mathrm{OH})_{3}(a q)
$$


1915. Boron and Silicon Halides -
FIGURE 18.15 Laboratory glassware, such as Pyrex and Kimax, is made of borosilicate glass because it does not break when heated. The inclusion of borates in the glass helps to mediate the effects of thermal expansion and contraction. This reduces the likelihood of thermal shock, which causes silicate glass to crack upon rapid heating or cooling. (credit: "Tweenk"/Wikimedia Commons)

The boron atom in $\mathrm{B}(\mathrm{OH})_{3}$ is $s p^{2}$ hybridized and is located at the center of an equilateral triangle with oxygen atoms at the corners. In solid $\mathrm{B}(\mathrm{OH})_{3}$, hydrogen bonding holds these triangular units together. Boric acid, shown in Figure 18.16, is a very weak acid that does not act as a proton donor but rather as a Lewis acid, accepting an unshared pair of electrons from the Lewis base $\mathrm{OH}^{-}$:

$$
\mathrm{B}(\mathrm{OH})_{3}(a q)+2 \mathrm{H}_{2} \mathrm{O}(l) \rightleftharpoons \mathrm{B}(\mathrm{OH})_{4}^{-}(a q)+\mathrm{H}_{3} \mathrm{O}^{+}(a q) \quad K_{\mathrm{a}}=5.8 \times 10^{-10}
$$

FIGURE 18.16 Boric acid has a planar structure with three -OH groups spread out equally at $120^{\circ}$ angles from each other.

Heating boric acid to $100^{\circ} \mathrm{C}$ causes molecules of water to split out between pairs of adjacent -OH groups to form metaboric acid, $\mathrm{HBO}_{2}$. At about $150^{\circ} \mathrm{C}$, additional B-O-B linkages form, connecting the $\mathrm{BO}_{3}$ groups together with shared oxygen atoms to form tetraboric acid, $\mathrm{H}_{2} \mathrm{~B}_{4} \mathrm{O}_{7}$. Complete water loss, at still higher temperatures, results in boric oxide.


1915. Boron and Silicon Halides -
Borates are salts of the oxyacids of boron. Borates result from the reactions of a base with an oxyacid or from the fusion of boric acid or boric oxide with a metal oxide or hydroxide. Borate anions range from the simple trigonal planar $\mathrm{BO}_{3}{ }^{3-}$ ion to complex species containing chains and rings of three- and four-coordinated boron atoms. The structures of the anions found in $\mathrm{CaB}_{2} \mathrm{O}_{4}, \mathrm{~K}\left[\mathrm{~B}_{5} \mathrm{O}_{6}(\mathrm{OH})_{4}\right] \cdot 2 \mathrm{H}_{2} \mathrm{O}$ (commonly written $\mathrm{KB}_{5} \mathrm{O}_{8} \cdot 4 \mathrm{H}_{2} \mathrm{O}$ ) and $\mathrm{Na}_{2}\left[\mathrm{~B}_{4} \mathrm{O}_{5}(\mathrm{OH})_{4}\right] \cdot 8 \mathrm{H}_{2} \mathrm{O}$ (commonly written $\mathrm{Na}_{2} \mathrm{~B}_{4} \mathrm{O}_{7} \cdot 10 \mathrm{H}_{2} \mathrm{O}$ ) are shown in Figure 18.17. Commercially, the most important borate is borax, $\mathrm{Na}_{2}\left[\mathrm{~B}_{4} \mathrm{O}_{5}(\mathrm{OH})_{4}\right] \cdot 8 \mathrm{H}_{2} \mathrm{O}$, which is an important component of some laundry detergents. Most of the supply of borax comes directly from dry lakes, such as Searles Lake in California, or is prepared from kernite, $\mathrm{Na}_{2} \mathrm{~B}_{4} \mathrm{O}_{7} \cdot 4 \mathrm{H}_{2} \mathrm{O}$.

(a)

(b)

(c)


1915. Boron and Silicon Halides -
(a)

(b)

(c)

FIGURE 18.17 The borate anions are (a) $\mathrm{CaB}_{2} \mathrm{O}_{4}$, (b) $\mathrm{KB}_{5} \mathrm{O}_{8} \cdot 4 \mathrm{H}_{2} \mathrm{O}$, and (c) $\mathrm{Na}_{2} \mathrm{~B}_{4} \mathrm{O}_{7} \cdot 10 \mathrm{H}_{2} \mathrm{O}$. The anion in $\mathrm{CaB}_{2} \mathrm{O}_{4}$ is an "infinite" chain.

Silicon dioxide, silica, occurs in both crystalline and amorphous forms. The usual crystalline form of silicon dioxide is quartz, a hard, brittle, clear, colorless solid. It is useful in many ways-for architectural decorations, semiprecious jewels, and frequency control in radio transmitters. Silica takes many crystalline forms, or polymorphs, in nature. Trace amounts of $\mathrm{Fe}^{3+}$ in quartz give amethyst its characteristic purple color. The term quartz is also used for articles such as tubing and lenses that are manufactured from amorphous silica. Opal is a naturally occurring form of amorphous silica.


1915. Boron and Silicon Halides -
The contrast in structure and physical properties between silicon dioxide and carbon dioxide is interesting, as illustrated in Figure 18.18. Solid carbon dioxide (dry ice) contains single $\mathrm{CO}_{2}$ molecules with each of the two oxygen atoms attached to the carbon atom by double bonds. Very weak intermolecular forces hold the molecules together in the crystal. The volatility of dry ice reflect these weak forces between molecules. In contrast, silicon dioxide is a covalent network solid. In silicon dioxide, each silicon atom links to four oxygen atoms by single bonds directed toward the corners of a regular tetrahedron, and $\mathrm{SiO}_{4}$ tetrahedra share oxygen atoms. This arrangement gives a three dimensional, continuous, silicon-oxygen network. A quartz crystal is a macromolecule of silicon dioxide. The difference between these two compounds is the ability of the group 14 elements to form strong $\pi$ bonds. Second-period elements, such as carbon, form very strong $\pi$ bonds, which is why carbon dioxide forms small molecules with strong double bonds. Elements below the second period, such as silicon, do not form $\pi$ bonds as readily as second-period elements, and when they do form, the $\pi$ bonds are weaker than those formed by second-period elements. For this reason, silicon dioxide does not contain $\pi$ bonds but only $\sigma$ bonds.

FIGURE 18.18 Because carbon tends to form double and triple bonds and silicon does not, (a) carbon dioxide is a discrete molecule with two $\mathrm{C}=\mathrm{O}$ double bonds and (b) silicon dioxide is an infinite network of oxygen atoms bridging between silicon atoms with each silicon atom possessing four $\mathrm{Si}-\mathrm{O}$ single bonds. (credit a photo: modification of work by Erica Gerdes; credit b photo: modification of work by Didier Descouens)


1915. Boron and Silicon Halides -
At $1600^{\circ} \mathrm{C}$, quartz melts to yield a viscous liquid. When the liquid cools, it does not crystallize readily but usually supercools and forms a glass, also called silica. The $\mathrm{SiO}_{4}$ tetrahedra in glassy silica have a random arrangement characteristic of supercooled liquids, and the glass has some very useful properties. Silica is highly transparent to both visible and ultraviolet light. For this reason, it is important in the manufacture of lamps that give radiation rich in ultraviolet light and in certain optical instruments that operate with ultraviolet light. The coefficient of expansion of silica glass is very low; therefore, rapid temperature changes do not cause it to fracture. CorningWare and other ceramic cookware contain amorphous silica.

Silicates are salts containing anions composed of silicon and oxygen. In nearly all silicates, $s p^{3}$-hybridized silicon atoms occur at the centers of tetrahedra with oxygen at the corners. There is a variation in the silicon-to-oxygen ratio that occurs because silicon-oxygen tetrahedra may exist as discrete, independent units or may share oxygen atoms at corners in a variety of ways. In addition, the presence of a variety of cations gives rise to the large number of silicate minerals.

Many ceramics are composed of silicates. By including small amounts of other compounds, it is possible to modify the physical properties of the silicate materials to produce ceramics with useful characteristics.


1916. LEARNING OBJECTIVES -
By the end of this section, you will be able to:

- Describe structure and properties of nonmetals

The nonmetals are elements located in the upper right portion of the periodic table. Their properties and behavior are quite different from those of metals on the left side. Under normal conditions, more than half of the nonmetals are gases, one is a liquid, and the rest include some of the softest and hardest of solids. The nonmetals exhibit a rich variety of chemical behaviors. They include the most reactive and least reactive of elements, and they form many different ionic and covalent compounds. This section presents an overview of the properties and chemical behaviors of the nonmetals, as well as the chemistry of specific elements. Many of these nonmetals are important in biological systems.

In many cases, trends in electronegativity enable us to predict the type of bonding and the physical states in compounds involving the nonmetals. We know that electronegativity decreases as we move down a given group and increases as we move from left to right across a period. The nonmetals have higher electronegativities than do metals, and compounds formed between metals and nonmetals are generally ionic in nature because of the large differences in electronegativity between them. The metals form cations, the nonmetals form anions, and the resulting compounds are solids under normal conditions. On the other hand, compounds formed between two or more nonmetals have small differences in electronegativity between the atoms, and covalent bonding-sharing of electrons-results. These substances tend to be molecular in nature and are gases, liquids, or volatile solids at room temperature and pressure.

In normal chemical processes, nonmetals do not form monatomic positive ions (cations) because their ionization energies are too high. All monatomic nonmetal ions are anions; examples include the chloride ion, $\mathrm{Cl}^{-}$, the nitride ion, $\mathrm{N}^{3-}$, and the selenide ion, $\mathrm{Se}^{2-}$.


1916. LEARNING OBJECTIVES -
The common oxidation states that the nonmetals exhibit in their ionic and covalent compounds are shown in Figure 18.19. Remember that an element exhibits a positive oxidation state when combined with a more electronegative element and that it exhibits a negative oxidation state when combined with a less electronegative element.

FIGURE 18.19 Nonmetals exhibit these common oxidation states in ionic and covalent compounds.
The first member of each nonmetal group exhibits different behaviors, in many respects, from the other group members. The reasons for this include smaller size, greater ionization energy, and (most important) the fact that the first member of each group has only four valence orbitals (one $2 s$ and three $2 p$ ) available for bonding, whereas other group members have empty $d$ orbitals in their valence shells, making possible five, six, or even more bonds around the central atom. For example, nitrogen forms only $\mathrm{NF}_{3}$, whereas phosphorus forms both $\mathrm{PF}_{3}$ and $\mathrm{PF}_{5}$.

Another difference between the first group member and subsequent members is the greater ability of the first member to form $\pi$ bonds. This is primarily a function of the smaller size of the first member of each group, which allows better overlap of atomic orbitals. Nonmetals, other than the first member of each group, rarely form $\pi$ bonds to nonmetals that are the first member of a group. For example, sulfur-oxygen $\pi$ bonds are well known, whereas sulfur does not normally form stable $\pi$ bonds to itself.

The variety of oxidation states displayed by most of the nonmetals means that many of their chemical reactions involve changes in oxidation state through oxidation-reduction reactions. There are five general aspects of the oxidation-reduction chemistry:

1. Nonmetals oxidize most metals. The oxidation state of the metal becomes positive as it undergoes oxidation and that of the nonmetal becomes negative as it undergoes reduction. For example:

$$
\begin{array}{cccc}
4 \mathrm{Fe}(s)+ & 3 \mathrm{O}_{2}(g) & \longrightarrow \mathrm{Fe}_{2} \mathrm{O}_{3}(s) \\
0 & 0 & +3-2
\end{array}
$$


1916. LEARNING OBJECTIVES -
2. With the exception of nitrogen and carbon, which are poor oxidizing agents, a more electronegative nonmetal oxidizes a less electronegative nonmetal or the anion of the nonmetal:

$$
\begin{array}{cccc}
\begin{array}{llll}
\mathrm{S}(s)+ & \mathrm{O}_{2}(g) & \longrightarrow & 2 \mathrm{SO}_{2}(s) \\
0 & 0 & & +4-2 \\
& & & \\
\mathrm{Cl}_{2}(g)+ & 2 \mathrm{I}^{-}(a q) & \longrightarrow & \mathrm{I}_{2}(s)+2 \mathrm{Cl}^{-}(a q) \\
0 & & 0
\end{array}
\end{array}
$$

3. Fluorine and oxygen are the strongest oxidizing agents within their respective groups; each oxidizes all the elements that lie below it in the group. Within any period, the strongest oxidizing agent is in group 17. A nonmetal often oxidizes an element that lies to its left in the same period. For example:

$$
\begin{array}{llll}
2 \operatorname{As}(s)+ & 3 \mathrm{Br}_{2}(l) \longrightarrow & 2 \mathrm{AsBr}_{3}(s) \\
0 & +3-1
\end{array}
$$

4. The stronger a nonmetal is as an oxidizing agent, the more difficult it is to oxidize the anion formed by the nonmetal. This means that the most stable negative ions are formed by elements at the top of the group or in group 17 of the period.
5. Fluorine and oxygen are the strongest oxidizing elements known. Fluorine does not form compounds in which it exhibits positive oxidation states; oxygen exhibits a positive oxidation state only when combined with fluorine. For example:

$$
\begin{array}{lll}
2 \mathrm{~F}_{2}(g)+2 \mathrm{OH}^{-}(a q) \longrightarrow & \mathrm{OF}_{2}(g)+2 \mathrm{~F}^{-}(a q)+\mathrm{H}_{2} \mathrm{O}(l) \\
0 & +2 \quad-1
\end{array}
$$


1916. LEARNING OBJECTIVES -
With the exception of most of the noble gases, all nonmetals form compounds with oxygen, yielding covalent oxides. Most of these oxides are acidic, that is, they react with water to form oxyacids. Recall from the acid-base chapter that an oxyacid is an acid consisting of hydrogen, oxygen, and some other element. Notable exceptions are carbon monoxide, CO , nitrous oxide, $\mathrm{N}_{2} \mathrm{O}$, and nitric oxide, NO . There are three characteristics of these acidic oxides:

1. Oxides such as $\mathrm{SO}_{2}$ and $\mathrm{N}_{2} \mathrm{O}_{5}$, in which the nonmetal exhibits one of its common oxidation states, are acid anhydrides and react with water to form acids with no change in oxidation state. The product is an oxyacid. For example:

$$
\begin{gathered}
\mathrm{SO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{H}_{2} \mathrm{SO}_{3}(a q) \\
\mathrm{N}_{2} \mathrm{O}_{5}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow 2 \mathrm{HNO}_{3}(a q)
\end{gathered}
$$

2. Those oxides such as $\mathrm{NO}_{2}$ and $\mathrm{ClO}_{2}$, in which the nonmetal does not exhibit one of its common oxidation states, also react with water. In these reactions, the nonmetal is both oxidized and reduced. For example:

$$
\begin{array}{ll}
3 \mathrm{NO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow & 2 \mathrm{HNO}_{3}(a q)+ \\
+4 & +5
\end{array}
$$

Reactions in which the same element is both oxidized and reduced are called disproportionation reactions.
3. The acid strength increases as the electronegativity of the central atom increases. To learn more, see the discussion in the chapter on acid-base chemistry.


1916. LEARNING OBJECTIVES -
The binary hydrogen compounds of the nonmetals also exhibit an acidic behavior in water, although only HCl , HBr , and HI are strong acids. The acid strength of the nonmetal hydrogen compounds increases from left to right across a period and down a group. For example, ammonia, $\mathrm{NH}_{3}$, is a weaker acid than is water, $\mathrm{H}_{2} \mathrm{O}$, which is weaker than is hydrogen fluoride, HF. Water, $\mathrm{H}_{2} \mathrm{O}$, is also a weaker acid than is hydrogen sulfide, $\mathrm{H}_{2} \mathrm{~S}$, which is weaker than is hydrogen selenide, $\mathrm{H}_{2} \mathrm{Se}$. Weaker acidic character implies greater basic character.


1917. Structures of the Nonmetals -
The structures of the nonmetals differ dramatically from those of metals. Metals crystallize in closely packed arrays that do not contain molecules or covalent bonds. Nonmetal structures contain covalent bonds, and many nonmetals consist of individual molecules. The electrons in nonmetals are localized in covalent bonds,
whereas in a metal, there is delocalization of the electrons throughout the solid.
The noble gases are all monatomic, whereas the other nonmetal gases-hydrogen, nitrogen, oxygen, fluorine, and chlorine-normally exist as the diatomic molecules $\mathrm{H}_{2}, \mathrm{~N}_{2}, \mathrm{O}_{2}, \mathrm{~F}_{2}$, and $\mathrm{Cl}_{2}$. The other halogens are also diatomic; $\mathrm{Br}_{2}$ is a liquid and $\mathrm{I}_{2}$ exists as a solid under normal conditions. The changes in state as one moves down the halogen family offer excellent examples of the increasing strength of intermolecular London forces with increasing molecular mass and increasing polarizability.

Oxygen has two allotropes: $\mathrm{O}_{2}$, dioxygen, and $\mathrm{O}_{3}$, ozone. Phosphorus has three common allotropes, commonly referred to by their colors: white, red, and black. Sulfur has several allotropes. There are also many carbon allotropes. Most people know of diamond, graphite, and charcoal, but fewer people know of the recent discovery of fullerenes, carbon nanotubes, and graphene.

Descriptions of the physical properties of three nonmetals that are characteristic of molecular solids follow.


1918. Carbon -
Carbon occurs in the uncombined (elemental) state in many forms, such as diamond, graphite, charcoal, coke, carbon black, graphene, and fullerene.

Diamond, shown in Figure 18.20, is a very hard crystalline material that is colorless and transparent when pure. Each atom forms four single bonds to four other atoms at the corners of a tetrahedron ( $s p^{3}$ hybridization); this makes the diamond a giant molecule. Carbon-carbon single bonds are very strong, and, because they extend throughout the crystal to form a three-dimensional network, the crystals are very hard and have high melting points $\left(\sim 4400^{\circ} \mathrm{C}\right)$.

FIGURE 18.20 (a) Diamond and (b) graphite are two forms of carbon. (c) In the crystal structure of diamond, the covalent bonds form three-dimensional tetrahedrons. (d) In the crystal structure of graphite, each planar layer is composed of six-membered rings. (credit a: modification of work by "Fancy Diamonds"/Flickr; credit b: modification of work from http://images-of-elements.com/carbon.php)

Graphite, also shown in Figure 18.20, is a soft, slippery, grayish-black solid that conducts electricity. These properties relate to its structure, which consists of layers of carbon atoms, with each atom surrounded by three other carbon atoms in a trigonal planar arrangement. Each carbon atom in graphite forms three $\sigma$ bonds, one to each of its nearest neighbors, by means of $s p^{2}$-hybrid orbitals. The unhybridized $p$ orbital on each carbon atom will overlap unhybridized orbitals on adjacent carbon atoms in the same layer to form $\pi$ bonds. Many resonance forms are necessary to describe the electronic structure of a graphite layer; Figure 18.21 illustrates two of these forms.

(a)

(b)

FIGURE 18.21 (a) Carbon atoms in graphite have unhybridized $p$ orbitals. Each $p$ orbital is perpendicular to the plane of carbon atoms. (b) These are two of the many resonance forms of graphite necessary to describe its electronic structure as a resonance hybrid.


1918. Carbon -
Atoms within a graphite layer are bonded together tightly by the $\sigma$ and $\pi$ bonds; however, the forces between layers are weak. London dispersion forces hold the layers together. To learn more, see the discussion of these weak forces in the chapter on liquids and solids. The weak forces between layers give graphite the soft, flaky character that makes it useful as the so-called "lead" in pencils and the slippery character that makes it useful as a lubricant. The loosely held electrons in the resonating $\pi$ bonds can move throughout the solid and are responsible for the electrical conductivity of graphite.

Other forms of elemental carbon include carbon black, charcoal, and coke. Carbon black is an amorphous form of carbon prepared by the incomplete combustion of natural gas, $\mathrm{CH}_{4}$. It is possible to produce charcoal and coke by heating wood and coal, respectively, at high temperatures in the absence of air.

Recently, new forms of elemental carbon molecules have been identified in the soot generated by a smoky flame and in the vapor produced when graphite is heated to very high temperatures in a vacuum or in helium. One of these new forms, first isolated by Professor Richard Smalley and coworkers at Rice University, consists of icosahedral (soccer-ball-shaped) molecules that contain 60 carbon atoms, $\mathrm{C}_{60}$. This is buckminsterfullerene (often called bucky balls) after the architect Buckminster Fuller, who designed domed structures, which have a similar appearance (Figure 18.22).

FIGURE 18.22 The molecular structure of $\mathrm{C}_{60}$, buckminsterfullerene, is icosahedral.


1920. Nanotubes and Graphene -
Graphene and carbon nanotubes are two recently discovered allotropes of carbon. Both of the forms bear some relationship to graphite. Graphene is a single layer of graphite (one atom thick), as illustrated in Figure 18.23, whereas carbon nanotubes roll the layer into a small tube, as illustrated in Figure 18.23.

FIGURE 18.23 (a) Graphene and (b) carbon nanotubes are both allotropes of carbon.
Graphene is a very strong, lightweight, and efficient conductor of heat and electricity discovered in 2003. As in graphite, the carbon atoms form a layer of six-membered rings with $s p^{2}$-hybridized carbon atoms at the corners. Resonance stabilizes the system and leads to its conductivity. Unlike graphite, there is no stacking of the layers to give a three-dimensional structure. Andre Geim and Kostya Novoselov at the University of Manchester won the 2010 Nobel Prize in Physics for their pioneering work characterizing graphene.

The simplest procedure for preparing graphene is to use a piece of adhesive tape to remove a single layer of graphene from the surface of a piece of graphite. This method works because there are only weak London dispersion forces between the layers in graphite. Alternative methods are to deposit a single layer of carbon atoms on the surface of some other material (ruthenium, iridium, or copper) or to synthesize it at the surface of silicon carbide via the sublimation of silicon.

There currently are no commercial applications of graphene. However, its unusual properties, such as high electron mobility and thermal conductivity, should make it suitable for the manufacture of many advanced electronic devices and for thermal management applications.

Carbon nanotubes are carbon allotropes, which have a cylindrical structure. Like graphite and graphene, nanotubes consist of rings of $s p^{2}$-hybridized carbon atoms. Unlike graphite and graphene, which occur in layers, the layers wrap into a tube and bond together to produce a stable structure. The walls of the tube may be one atom or multiple atoms thick.

Carbon nanotubes are extremely strong materials that are harder than diamond. Depending upon the shape of the nanotube, it may be a conductor or semiconductor. For some applications, the conducting form is preferable, whereas other applications utilize the semiconducting form.


1920. Nanotubes and Graphene -
The basis for the synthesis of carbon nanotubes is the generation of carbon atoms in a vacuum. It is possible to produce carbon atoms by an electrical discharge through graphite, vaporization of graphite
with a laser, and the decomposition of a carbon compound.
The strength of carbon nanotubes will eventually lead to some of their most exciting applications, as a thread produced from several nanotubes will support enormous weight. However, the current applications only employ bulk nanotubes. The addition of nanotubes to polymers improves the mechanical, thermal, and electrical properties of the bulk material. There are currently nanotubes in some bicycle parts, skis, baseball bats, fishing rods, and surfboards.


1921. Phosphorus -
The name phosphorus comes from the Greek words meaning light bringing. When phosphorus was first isolated, scientists noted that it glowed in the dark and burned when exposed to air. Phosphorus is the only member of its group that does not occur in the uncombined state in nature; it exists in many allotropic forms. We will consider two of those forms: white phosphorus and red phosphorus.

White phosphorus is a white, waxy solid that melts at $44.2^{\circ} \mathrm{C}$ and boils at $280^{\circ} \mathrm{C}$. It is insoluble in water (in which it is stored-see Figure 18.24), is very soluble in carbon disulfide, and bursts into flame in air. As a solid, as a liquid, as a gas, and in solution, white phosphorus exists as $\mathrm{P}_{4}$ molecules with four phosphorus atoms at the corners of a regular tetrahedron, as illustrated in Figure 18.24. Each phosphorus atom covalently bonds to the other three atoms in the molecule by single covalent bonds. White phosphorus is the most reactive allotrope and is very toxic.

FIGURE 18.24 (a) Because white phosphorus bursts into flame in air, it is stored in water. (b) The structure of white phosphorus consists of $\mathrm{P}_{4}$ molecules arranged in a tetrahedron. (c) Red phosphorus is much less reactive than is white phosphorus. (d) The structure of red phosphorus consists of networks of $\mathrm{P}_{4}$ tetrahedra joined by P-P single bonds. (credit a: modification of work from http://images-of-elements.com/phosphorus.php)


1921. Phosphorus -
Heating white phosphorus to $270-300^{\circ} \mathrm{C}$ in the absence of air yields red phosphorus. Red phosphorus (shown in Figure 18.24 ) is denser, has a higher melting point $\left(\sim 600^{\circ} \mathrm{C}\right)$, is much less reactive, is essentially nontoxic, and is easier and safer to handle than is white phosphorus. Its structure is highly polymeric and appears to contain three-dimensional networks of $\mathrm{P}_{4}$ tetrahedra joined by P-P single bonds. Red phosphorus is insoluble in solvents that dissolve white phosphorus. When red phosphorus is heated, $\mathrm{P}_{4}$ molecules sublime from the solid.


1922. Sulfur -
The allotropy of sulfur is far greater and more complex than that of any other element. Sulfur is the brimstone referred to in the Bible and other places, and references to sulfur occur throughout recorded history-right up to the relatively recent discovery that it is a component of the atmospheres of Venus and of Io, a moon of Jupiter. The most common and most stable allotrope of sulfur is yellow, rhombic sulfur, so named because of the shape of its crystals. Rhombic sulfur is the form to which all other allotropes revert at room temperature. Crystals of rhombic sulfur melt at $113^{\circ} \mathrm{C}$. Cooling this liquid gives long needles of monoclinic sulfur. This form is stable from $96^{\circ} \mathrm{C}$ to the melting point, $119^{\circ} \mathrm{C}$. At room temperature, it gradually reverts to the rhombic form.

Both rhombic sulfur and monoclinic sulfur contain $\mathrm{S}_{8}$ molecules in which atoms form eight-membered, puckered rings that resemble crowns, as illustrated in Figure 18.25. Each sulfur atom is bonded to each of its
two neighbors in the ring by covalent S-S single bonds.

FIGURE 18.25 These four sulfur allotropes show eight-membered, puckered rings. Each sulfur atom bonds to each of its two neighbors in the ring by covalent $\mathrm{S}-\mathrm{S}$ single bonds. Here are (a) individual $\mathrm{S}_{8}$ rings, (b) $\mathrm{S}_{8}$ chains formed when the rings open, (c) longer chains formed by adding sulfur atoms to $\mathrm{S}_{8}$ chains, and (d) part of the very long sulfur chains formed at higher temperatures.


1922. Sulfur -
When rhombic sulfur melts, the straw-colored liquid is quite mobile; its viscosity is low because $\mathrm{S}_{8}$ molecules are essentially spherical and offer relatively little resistance as they move past each other. As the temperature rises, S-S bonds in the rings break, and polymeric chains of sulfur atoms result. These chains combine end to end, forming still longer chains that tangle with one another. The liquid gradually darkens in color and becomes so viscous that finally (at about $230^{\circ} \mathrm{C}$ ) it does not pour easily. The dangling atoms at the ends of the chains of sulfur atoms are responsible for the dark red color because their electronic structure differs from those of sulfur atoms that have bonds to two adjacent sulfur atoms. This causes them to absorb light differently and results in a different visible color. Cooling the liquid rapidly produces a rubberlike amorphous mass, called plastic sulfur.

Sulfur boils at $445{ }^{\circ} \mathrm{C}$ and forms a vapor consisting of $\mathrm{S}_{2}, \mathrm{~S}_{6}$, and $\mathrm{S}_{8}$ molecules; at about $1000^{\circ} \mathrm{C}$, the vapor density corresponds to the formula $\mathrm{S}_{2}$, which is a paramagnetic molecule like $\mathrm{O}_{2}$ with a similar electronic structure and a weak sulfur-sulfur double bond.


1922. Sulfur -
As seen in this discussion, an important feature of the structural behavior of the nonmetals is that the elements usually occur with eight electrons in their valence shells. If necessary, the elements form enough covalent bonds to supplement the electrons already present to possess an octet. For example, members of group 15 have five valence electrons and require only three additional electrons to fill their valence shells. These elements form three covalent bonds in their free state: triple bonds in the $\mathrm{N}_{2}$ molecule or single bonds to three different atoms in arsenic and phosphorus. The elements of group 16 require only two additional electrons. Oxygen forms a double bond in the $\mathrm{O}_{2}$ molecule, and sulfur, selenium, and tellurium form two single bonds in various rings and chains. The halogens form diatomic molecules in which each atom is involved in only one bond. This provides the electron required necessary to complete the octet on the halogen atom. The noble gases do not form covalent bonds to other noble gas atoms because they already have a filled outer shell.


1922. Sulfur - 1922.1. Occurrence, Preparation, and Compounds of Hydrogen
LEARNING OBJECTIVES
By the end of this section, you will be able to:

- Describe the properties, preparation, and compounds of hydrogen

Hydrogen is the most abundant element in the universe. The sun and other stars are composed largely of hydrogen. Astronomers estimate that $90 \%$ of the atoms in the universe are hydrogen atoms. Hydrogen is a component of more compounds than any other element. Water is the most abundant compound of hydrogen found on earth. Hydrogen is an important part of petroleum, many minerals, cellulose and starch, sugar, fats, oils, alcohols, acids, and thousands of other substances.

At ordinary temperatures, hydrogen is a colorless, odorless, tasteless, and nonpoisonous gas consisting of the diatomic molecule $\mathrm{H}_{2}$. Hydrogen is composed of three isotopes, and unlike other elements, these isotopes have different names and chemical symbols: protium, ${ }^{1} \mathrm{H}$, deuterium, ${ }^{2} \mathrm{H}$ (or " D "), and tritium ${ }^{3} \mathrm{H}$ (or " T "). In a naturally occurring sample of hydrogen, there is one atom of deuterium for every 7000 H atoms and one atom of radioactive tritium for every $10^{18} \mathrm{H}$ atoms. The chemical properties of the different isotopes are very similar because they have identical electron structures, but they differ in some physical properties because of their differing atomic masses. Elemental deuterium and tritium have lower vapor pressure than ordinary hydrogen. Consequently, when liquid hydrogen evaporates, the heavier isotopes are concentrated in the last portions to evaporate. Electrolysis of heavy water, $\mathrm{D}_{2} \mathrm{O}$, yields deuterium. Most tritium originates from nuclear reactions.


1923. Preparation of Hydrogen -
Elemental hydrogen must be prepared from compounds by breaking chemical bonds. The most common methods of preparing hydrogen follow.

From Steam and Carbon or Hydrocarbons
Water is the cheapest and most abundant source of hydrogen. Passing steam over coke (an impure form of elemental carbon) at $1000^{\circ} \mathrm{C}$ produces a mixture of carbon monoxide and hydrogen known as water gas:

$$
\mathrm{C}(\mathrm{~s})+\mathrm{H}_{2} \mathrm{O}(\mathrm{~g}) \xrightarrow{1000^{\circ} \mathrm{C}} \mathrm{CO}(g)+\mathrm{H}_{2}(g)
$$

water gas
Water gas is as an industrial fuel. It is possible to produce additional hydrogen by mixing the water gas with steam in the presence of a catalyst to convert the CO to $\mathrm{CO}_{2}$. This reaction is the water gas shift reaction.

It is also possible to prepare a mixture of hydrogen and carbon monoxide by passing hydrocarbons from natural gas or petroleum and steam over a nickel-based catalyst. Propane is an example of a hydrocarbon reactant:

$$
\mathrm{C}_{3} \mathrm{H}_{8}(\mathrm{~g})+3 \mathrm{H}_{2} \mathrm{O}(\mathrm{~g}) \xrightarrow[\text { catalyst }]{90{ }^{\circ} \mathrm{C}} 3 \mathrm{CO}(\mathrm{~g})+7 \mathrm{H}_{2}(\mathrm{~g})
$$

Electrolysis
Hydrogen forms when direct current electricity passes through water containing an electrolyte such as $\mathrm{H}_{2} \mathrm{SO}_{4}$, as illustrated in Figure 18.26. Bubbles of hydrogen form at the cathode, and oxygen evolves at the anode. The
net reaction is:

$$
2 \mathrm{H}_{2} \mathrm{O}(l)+\text { electrical energy } \longrightarrow 2 \mathrm{H}_{2}(g)+\mathrm{O}_{2}(g)
$$


1923. Preparation of Hydrogen -
$$
2 \mathrm{H}_{2} \mathrm{O}(l)+\text { electrical energy } \longrightarrow 2 \mathrm{H}_{2}(g)+\mathrm{O}_{2}(g)
$$

FIGURE 18.26 The electrolysis of water produces hydrogen and oxygen. Because there are twice as many hydrogen atoms as oxygen atoms and both elements are diatomic, there is twice the volume of hydrogen produced at the cathode as there is oxygen produced at the anode.

Reaction of Metals with Acids
This is the most convenient laboratory method of producing hydrogen. Metals with lower reduction potentials reduce the hydrogen ion in dilute acids to produce hydrogen gas and metal salts. For example, as shown in Figure 18.27, iron in dilute hydrochloric acid produces hydrogen gas and iron(II) chloride:

$$
\mathrm{Fe}(s)+2 \mathrm{H}_{3} \mathrm{O}^{+}(a q)+2 \mathrm{Cl}^{-}(a q) \longrightarrow \mathrm{Fe}^{2+}(a q)+2 \mathrm{Cl}^{-}(a q)+\mathrm{H}_{2}(g)+2 \mathrm{H}_{2} \mathrm{O}(l)
$$

FIGURE 18.27 The reaction of iron with an acid produces hydrogen. Here, iron reacts with hydrochloric acid. (credit: Mark Ott)


1924. Reaction of Ionic Metal Hydrides with Water -
It is possible to produce hydrogen from the reaction of hydrides of the active metals, which contain the very strongly basic $\mathrm{H}^{-}$anion, with water:

$$
\mathrm{CaH}_{2}(s)+2 \mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{Ca}^{2+}(a q)+2 \mathrm{OH}^{-}(a q)+2 \mathrm{H}_{2}(g)
$$

Metal hydrides are expensive but convenient sources of hydrogen, especially where space and weight are important factors. They are important in the inflation of life jackets, life rafts, and military balloons.