1960. Properties of the Halogens -
The nonmetal chlorine is more electronegative than any other element except fluorine, oxygen, and nitrogen. In general, very electronegative elements are good oxidizing agents; therefore, we would expect elemental chlorine to oxidize all of the other elements except for these three (and the nonreactive noble gases). Its oxidizing property, in fact, is responsible for its principal use. For example, phosphorus(V) chloride, an important intermediate in the preparation of insecticides and chemical weapons, is manufactured by oxidizing the phosphorus with chlorine:

$$
\mathrm{P}_{4}(s)+10 \mathrm{Cl}_{2}(g) \longrightarrow 4 \mathrm{PCl}_{5}(l)
$$

A great deal of chlorine is also used to oxidize, and thus to destroy, organic or biological materials in water purification and in bleaching.

The chemical properties of bromine are similar to those of chlorine, although bromine is the weaker oxidizing agent and its reactivity is less than that of chlorine.

Iodine is the least reactive of the halogens. It is the weakest oxidizing agent, and the iodide ion is the most easily oxidized halide ion. Iodine reacts with metals, but heating is often required. It does not oxidize other halide ions.

Compared with the other halogens, iodine reacts only slightly with water. Traces of iodine in water react with a mixture of starch and iodide ion, forming a deep blue color. This reaction is a very sensitive test for the presence of iodine in water.


1961. Halides of the Representative Metals -
Thousands of salts of the representative metals have been prepared. The binary halides are an important subclass of salts. A salt is an ionic compound composed of cations and anions, other than hydroxide or oxide ions. In general, it is possible to prepare these salts from the metals or from oxides, hydroxides, or carbonates. We will illustrate the general types of reactions for preparing salts through reactions used to prepare binary halides.

The binary compounds of a metal with the halogens are the halides. Most binary halides are ionic. However, mercury, the elements of group 13 with oxidation states of $3+$, tin(IV), and lead(IV) form covalent binary halides.

The direct reaction of a metal and a halogen produce the halide of the metal. Examples of these oxidationreduction reactions include:

$$
\begin{aligned}
\mathrm{Cd}(s)+\mathrm{Cl}_{2}(g) & \longrightarrow \mathrm{CdCl}_{2}(s) \\
2 \mathrm{Ga}(l)+3 \mathrm{Br}_{2}(l) & \longrightarrow 2 \mathrm{GaBr}_{3}(s)
\end{aligned}
$$


1962. LINK TO LEARNING -
Reactions of the alkali metals with elemental halogens are very exothermic and often quite violent. Under controlled conditions, they provide exciting demonstrations for budding students of chemistry. You can view the initial heating (http://openstax.org/l/16sodium) of the sodium that removes the coating of sodium hydroxide, sodium peroxide, and residual mineral oil to expose the reactive surface. The reaction with chlorine gas then proceeds very nicely.

If a metal can exhibit two oxidation states, it may be necessary to control the stoichiometry in order to obtain the halide with the lower oxidation state. For example, preparation of tin(II) chloride requires a 1:1 ratio of Sn to $\mathrm{Cl}_{2}$, whereas preparation of $\operatorname{tin}(\mathrm{IV})$ chloride requires a 1:2 ratio:

$$
\begin{gathered}
\mathrm{Sn}(s)+\mathrm{Cl}_{2}(g) \longrightarrow \mathrm{SnCl}_{2}(s) \\
\mathrm{Sn}(s)+2 \mathrm{Cl}_{2}(g) \longrightarrow \mathrm{SnCl}_{4}(l)
\end{gathered}
$$

The active representative metals-those that are easier to oxidize than hydrogen-react with gaseous hydrogen halides to produce metal halides and hydrogen. The reaction of zinc with hydrogen fluoride is:

$$
\mathrm{Zn}(s)+2 \mathrm{HF}(g) \longrightarrow \mathrm{ZnF}_{2}(s)+\mathrm{H}_{2}(g)
$$

The active representative metals also react with solutions of hydrogen halides to form hydrogen and solutions of the corresponding halides. Examples of such reactions include:

$$
\begin{aligned}
\mathrm{Cd}(s)+2 \mathrm{HBr}(a q) & \longrightarrow \mathrm{CdBr}_{2}(a q)+\mathrm{H}_{2}(g) \\
\mathrm{Sn}(s)+2 \mathrm{HI}(a q) & \longrightarrow \mathrm{SnI}_{2}(a q)+\mathrm{H}_{2}(g)
\end{aligned}
$$


1962. LINK TO LEARNING -
Hydroxides, carbonates, and some oxides react with solutions of the hydrogen halides to form solutions of halide salts. It is possible to prepare additional salts by the reaction of these hydroxides, carbonates, and oxides with aqueous solution of other acids:

$$
\begin{gathered}
\mathrm{CaCo}_{3}(s)+2 \mathrm{HCl}(a q) \longrightarrow \mathrm{CaCl}_{2}(a q)+\mathrm{CO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(l) \\
\mathrm{TlOH}(a q)+\mathrm{HF}(a q) \longrightarrow \mathrm{TlF}(a q)+\mathrm{H}_{2} \mathrm{O}(l)
\end{gathered}
$$

A few halides and many of the other salts of the representative metals are insoluble. It is possible to prepare these soluble salts by metathesis reactions that occur when solutions of soluble salts are mixed (see Figure 18.61). Metathesis reactions are examined in the chapter on the stoichiometry of chemical reactions.

FIGURE 18.61 Solid $\mathrm{HgI}_{2}$ forms when solutions of KI and $\mathrm{Hg}\left(\mathrm{NO}_{3}\right)_{2}$ are mixed. (credit: Sahar Atwa)
Several halides occur in large quantities in nature. The ocean and underground brines contain many halides. For example, magnesium chloride in the ocean is the source of magnesium ions used in the production of magnesium. Large underground deposits of sodium chloride, like the salt mine shown in Figure 18.62, occur in many parts of the world. These deposits serve as the source of sodium and chlorine in almost all other compounds containing these elements. The chlor-alkali process is one example.

FIGURE 18.62 Underground deposits of sodium chloride are found throughout the world and are often mined. This is a tunnel in the Kłodawa salt mine in Poland. (credit: Jarek Zok)


1963. Interhalogens -
Compounds formed from two or more different halogens are interhalogens. Interhalogen molecules consist of one atom of the heavier halogen bonded by single bonds to an odd number of atoms of the lighter halogen. The structures of $\mathrm{IF}_{3}, \mathrm{IF}_{5}$, and $\mathrm{IF}_{7}$ are illustrated in Figure 18.63. Formulas for other interhalogens, each of which comes from the reaction of the respective halogens, are in Table 18.3.

FIGURE 18.63 The structure of $\mathrm{IF}_{3}$ is T -shaped (left), $\mathrm{IF}_{5}$ is square pyramidal (center), and $\mathrm{IF}_{7}$ is pentagonal bipyramidal (right).

Note from Table 18.3 that fluorine is able to oxidize iodine to its maximum oxidation state, $7+$, whereas bromine and chlorine, which are more difficult to oxidize, achieve only the 5+-oxidation state. A 7+-oxidation state is the limit for the halogens. Because smaller halogens are grouped about a larger one, the maximum number of smaller atoms possible increases as the radius of the larger atom increases. Many of these compounds are unstable, and most are extremely reactive. The interhalogens react like their component halides; halogen fluorides, for example, are stronger oxidizing agents than are halogen chlorides.

The ionic polyhalides of the alkali metals, such as $\mathrm{KI}_{3}, \mathrm{KICl}_{2}, \mathrm{KICl}_{4}, \mathrm{CsIBr}_{2}$, and $\mathrm{CsBrCl}_{2}$, which contain an anion composed of at least three halogen atoms, are closely related to the interhalogens. As seen previously, the formation of the polyhalide anion $\mathrm{I}_{3}{ }^{-}$is responsible for the solubility of iodine in aqueous solutions
containing an iodide ion.

Interhalogens


1963. Interhalogens -
Interhalogens

| YX | $\mathrm{YX}_{3}$ | $\mathrm{YX}_{5}$ | $\mathrm{YX}_{7}$ |
| :--- | :--- | :--- | :--- |
| $\mathrm{ClF}(g)$ | $\mathrm{ClF}_{3}(g)$ | $\mathrm{ClF}_{5}(g)$ |  |
| $\operatorname{BrF}(g)$ | $\mathrm{BrF}_{3}(I)$ | $\mathrm{BrF}_{5}(I)$ |  |
| $\operatorname{BrCl}(g)$ |  |  |  |
| $\mathrm{IF}(s)$ | $\mathrm{IF}_{3}(s)$ | $\mathrm{IF}_{5}(I)$ | $\mathrm{IF}_{7}(g)$ |
| $\mathrm{ICl}(I)$ | $\mathrm{ICl}_{3}(s)$ |  |  |
| $\mathrm{IBr}(s)$ |  |  |  |

TABLE 18.3


1964. Applications -
The fluoride ion and fluorine compounds have many important uses. Compounds of carbon, hydrogen, and fluorine are replacing Freons (compounds of carbon, chlorine, and fluorine) as refrigerants. Teflon is a polymer composed of $-\mathrm{CF}_{2} \mathrm{CF}_{2}-$ units. Fluoride ion is added to water supplies and to some toothpastes as $\mathrm{SnF}_{2}$ or NaF to fight tooth decay. Fluoride partially converts teeth from $\mathrm{Ca}_{5}\left(\mathrm{PO}_{4}\right)_{3}(\mathrm{OH})$ into $\mathrm{Ca}_{5}\left(\mathrm{PO}_{4}\right)_{3} \mathrm{~F}$.

Chlorine is important to bleach wood pulp and cotton cloth. The chlorine reacts with water to form hypochlorous acid, which oxidizes colored substances to colorless ones. Large quantities of chlorine are important in chlorinating hydrocarbons (replacing hydrogen with chlorine) to produce compounds such as tetrachloride $\left(\mathrm{CCl}_{4}\right)$, chloroform $\left(\mathrm{CHCl}_{3}\right)$, and ethyl chloride $\left(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{Cl}\right)$, and in the production of polyvinyl chloride (PVC) and other polymers. Chlorine is also important to kill the bacteria in community water supplies.

Bromine is important in the production of certain dyes, and sodium and potassium bromides are used as sedatives. At one time, light-sensitive silver bromide was a component of photographic film.

Iodine in alcohol solution with potassium iodide is an antiseptic (tincture of iodine). Iodide salts are essential for the proper functioning of the thyroid gland; an iodine deficiency may lead to the development of a goiter. Iodized table salt contains $0.023 \%$ potassium iodide. Silver iodide is useful in the seeding of clouds to induce rain; it was important in the production of photographic film and iodoform, $\mathrm{CHI}_{3}$, is an antiseptic.


1965. LEARNING OBJECTIVES -
By the end of this section, you will be able to:

- Describe the properties, preparation, and uses of the noble gases

The elements in group 18 are the noble gases (helium, neon, argon, krypton, xenon, and radon). They earned the name "noble" because they were assumed to be nonreactive since they have filled valence shells. In 1962, Dr. Neil Bartlett at the University of British Columbia proved this assumption to be false.

These elements are present in the atmosphere in small amounts. Some natural gas contains $1-2 \%$ helium by mass. Helium is isolated from natural gas by liquefying the condensable components, leaving only helium as a gas. The United States possesses most of the world's commercial supply of this element in its helium-bearing gas fields. Argon, neon, krypton, and xenon come from the fractional distillation of liquid air. Radon comes from other radioactive elements. More recently, it was observed that this radioactive gas is present in very
small amounts in soils and minerals. Its accumulation in well-insulated, tightly sealed buildings, however, constitutes a health hazard, primarily lung cancer.

The boiling points and melting points of the noble gases are extremely low relative to those of other substances of comparable atomic or molecular masses. This is because only weak London dispersion forces are present, and these forces can hold the atoms together only when molecular motion is very slight, as it is at very low temperatures. Helium is the only substance known that does not solidify on cooling at normal pressure. It remains liquid close to absolute zero $(0.001 \mathrm{~K})$ at ordinary pressures, but it solidifies under elevated pressure.

Helium is used for filling balloons and lighter-than-air craft because it does not burn, making it safer to use than hydrogen. Helium at high pressures is not a narcotic like nitrogen. Thus, mixtures of oxygen and helium are important for divers working under high pressures. Using a helium-oxygen mixture avoids the disoriented mental state known as nitrogen narcosis, the so-called rapture of the deep. Helium is important as an inert atmosphere for the melting and welding of easily oxidizable metals and for many chemical processes that are sensitive to air.


1965. LEARNING OBJECTIVES -
Liquid helium (boiling point, 4.2 K ) is an important coolant to reach the low temperatures necessary for cryogenic research, and it is essential for achieving the low temperatures necessary to produce superconduction in traditional superconducting materials used in powerful magnets and other devices. This cooling ability is necessary for the magnets used for magnetic resonance imaging, a common medical diagnostic procedure. The other common coolant is liquid nitrogen (boiling point, 77 K ), which is significantly cheaper.

Neon is a component of neon lamps and signs. Passing an electric spark through a tube containing neon at low pressure generates the familiar red glow of neon. It is possible to change the color of the light by mixing argon or mercury vapor with the neon or by utilizing glass tubes of a special color.

Argon was useful in the manufacture of gas-filled electric light bulbs, where its lower heat conductivity and chemical inertness made it preferable to nitrogen for inhibiting the vaporization of the tungsten filament and prolonging the life of the bulb. Fluorescent tubes commonly contain a mixture of argon and mercury vapor. Argon is the third most abundant gas in dry air.

Krypton-xenon flash tubes are used to take high-speed photographs. An electric discharge through such a tube gives a very intense light that lasts only $\frac{1}{50,000}$ of a second. Krypton forms a difluoride, $\mathrm{KrF}_{2}$, which is thermally unstable at room temperature.

Stable compounds of xenon form when xenon reacts with fluorine. Xenon difluoride, $\mathrm{XeF}_{2}$, forms after heating an excess of xenon gas with fluorine gas and then cooling. The material forms colorless crystals, which are stable at room temperature in a dry atmosphere. Xenon tetrafluoride, $\mathrm{XeF}_{4}$, and xenon hexafluoride, $\mathrm{XeF}_{6}$, are prepared in an analogous manner, with a stoichiometric amount of fluorine and an excess of fluorine, respectively. Compounds with oxygen are prepared by replacing fluorine atoms in the xenon fluorides with oxygen.


1965. LEARNING OBJECTIVES -
When $\mathrm{XeF}_{6}$ reacts with water, a solution of $\mathrm{XeO}_{3}$ results and the xenon remains in the $6+$-oxidation state:

$$
\mathrm{XeF}_{6}(s)+3 \mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{XeO}_{3}(a q)+6 \mathrm{HF}(a q)
$$

Dry, solid xenon trioxide, $\mathrm{XeO}_{3}$, is extremely explosive-it will spontaneously detonate. Both $\mathrm{XeF}_{6}$ and $\mathrm{XeO}_{3}$ disproportionate in basic solution, producing xenon, oxygen, and salts of the perxenate ion, $\mathrm{XeO}_{6}{ }^{4-}$, in which xenon reaches its maximum oxidation sate of 8+.

Radon apparently forms $\mathrm{RnF}_{2}$-evidence of this compound comes from radiochemical tracer techniques.
Unstable compounds of argon form at low temperatures, but stable compounds of helium and neon are not known.


1966. Key Terms -
acid anhydride compound that reacts with water to form an acid or acidic solution
alkaline earth metal any of the metals (beryllium, magnesium, calcium, strontium, barium, and radium) occupying group 2 of the periodic table; they are reactive, divalent metals that form basic oxides
allotropes two or more forms of the same element, in the same physical state, with different chemical structures
amorphous solid material such as a glass that does not have a regular repeating component to its three-dimensional structure; a solid but not a crystal
base anhydride metal oxide that behaves as a base towards acids
bicarbonate anion salt of the hydrogen carbonate ion, $\mathrm{HCO}_{3}{ }^{-}$
bismuth heaviest member of group 15; a less reactive metal than other representative metals
borate compound containing boron-oxygen bonds, typically with clusters or chains as a part of the chemical structure
carbonate salt of the anion $\mathrm{CO}_{3}{ }^{2-}$; often formed by the reaction of carbon dioxide with bases
chemical reduction method of preparing a representative metal using a reducing agent
chlor-alkali process electrolysis process for the synthesis of chlorine and sodium hydroxide
disproportionation reaction chemical reaction where a single reactant is simultaneously reduced and oxidized; it is both the reducing agent and the oxidizing agent
Downs cell electrochemical cell used for the commercial preparation of metallic sodium (and chlorine) from molten sodium chloride
Frasch process important in the mining of free sulfur from enormous underground deposits
Haber process main industrial process used to produce ammonia from nitrogen and hydrogen; involves the use of an iron catalyst and elevated temperatures and pressures
halide compound containing an anion of a group 17 element in the 1-oxidation state (fluoride, $\mathrm{F}^{-}$; chloride, $\mathrm{Cl}^{-}$; bromide, $\mathrm{Br}^{-}$; and iodide, $\mathrm{I}^{-}$)
Hall-Héroult cell electrolysis apparatus used to isolate pure aluminum metal from a solution of alumina in molten cryolite


1966. Key Terms -
Hall-Héroult cell electrolysis apparatus used to isolate pure aluminum metal from a solution of alumina in molten cryolite
hydrogen carbonate salt of carbonic acid, $\mathrm{H}_{2} \mathrm{CO}_{3}$ (containing the anion $\mathrm{HCO}_{3}{ }^{-}$) in which one hydrogen atom has been replaced; an acid carbonate; also known as bicarbonate ion
hydrogen halide binary compound formed between hydrogen and the halogens: $\mathrm{HF}, \mathrm{HCl}$, HBr , and HI
hydrogen sulfate $\mathrm{HSO}_{4}{ }^{-}$ion
hydrogen sulfite $\mathrm{HSO}_{3}{ }^{-}$ion
hydrogenation addition of hydrogen $\left(\mathrm{H}_{2}\right)$ to reduce a compound
hydroxide compound of a metal with the hydroxide ion $\mathrm{OH}^{-}$or the group -OH
interhalogen compound formed from two or more different halogens
metal (representative) atoms of the metallic elements of groups $1,2,12,13,14,15$, and 16 , which form ionic compounds by losing electrons from their outer $s$ or $p$ orbitals
metalloid element that has properties that are between those of metals and nonmetals; these elements are typically semiconductors
nitrate $\mathrm{NO}_{3}{ }^{-}$ion; salt of nitric acid
nitrogen fixation formation of nitrogen compounds from molecular nitrogen
Ostwald process industrial process used to convert ammonia into nitric acid
oxide binary compound of oxygen with another element or group, typically containing $\mathrm{O}^{2-}$ ions or the group -O - or $=0$
ozone allotrope of oxygen; $\mathrm{O}_{3}$
passivation metals with a protective nonreactive film of oxide or other compound that creates a barrier for chemical reactions; physical or chemical removal of the passivating film allows the metals to demonstrate their expected chemical reactivity
peroxide molecule containing two oxygen atoms bonded together or as the anion, $\mathrm{O}_{2}{ }^{2-}$
photosynthesis process whereby light energy promotes the reaction of water and carbon dioxide to form carbohydrates and oxygen; this allows photosynthetic organisms to store energy


1966. Key Terms -
photosynthesis process whereby light energy promotes the reaction of water and carbon dioxide to form carbohydrates and oxygen; this allows photosynthetic organisms to store energy
Pidgeon process chemical reduction process used to produce magnesium through the thermal reaction of magnesium oxide with silicon
polymorph variation in crystalline structure that results in different physical properties for the resulting compound
representative element element where the $s$ and $p$ orbitals are filling
representative metal metal among the representative elements
silicate compound containing silicon-oxygen bonds, with silicate tetrahedra connected in rings, sheets, or three-dimensional networks,
depending on the other elements involved in the formation of the compounds
sulfate $\mathrm{SO}_{4}{ }^{2-}$ ion


1967. Summary - 1967.1. Periodicity
This section focuses on the periodicity of the representative elements. These are the elements where the electrons are entering the $s$ and $p$ orbitals. The representative elements occur in groups 1, 2, and 12-18. These elements are representative metals, metalloids, and nonmetals. The alkali metals (group 1) are very reactive, readily form ions with a charge of $1+$ to form ionic compounds that are usually soluble in water, and react vigorously with water to form hydrogen gas and a basic solution of the metal hydroxide. The outermost electrons of the alkaline earth metals (group 2) are more difficult to remove than the outer electron of the alkali metals, leading to the group 2 metals being less reactive than those in group 1. These elements easily form compounds in which the metals exhibit an oxidation state of $2+$. Zinc, cadmium, and mercury (group 12) commonly exhibit the group oxidation state of 2+ (although mercury also exhibits an oxidation state of $1+$ in compounds that contain $\mathrm{Hg}_{2}{ }^{2+}$ ). Aluminum, gallium, indium, and thallium (group 13) are easier to oxidize than is hydrogen. Aluminum, gallium, and indium occur with an oxidation state 3+ (however, thallium also commonly occurs as the $\mathrm{Tl}^{+}$ion). Tin and lead form stable divalent cations and covalent compounds in which the metals exhibit the 4+oxidation state.


1967. Summary - 1967.2. Occurrence and Preparation of the Representative Metals
Because of their chemical reactivity, it is necessary to produce the representative metals in their pure forms by reduction from naturally occurring compounds. Electrolysis is important in the production of sodium, potassium, and aluminum. Chemical reduction is the primary method for the isolation of magnesium, zinc, and tin. Similar procedures are important for the other representative metals.


1967. Summary - 1967.3. Structure and General Properties of the Metalloids
The elements boron, silicon, germanium, arsenic, antimony, and tellurium separate the metals from the nonmetals in the periodic table. These elements, called metalloids or sometimes semimetals, exhibit
sulfite $\mathrm{SO}_{3}{ }^{2-}$ ion
superoxide oxide containing the anion $\mathrm{O}_{2}{ }^{-}$
properties characteristic of both metals and nonmetals. The structures of these elements are similar in many ways to those of nonmetals, but the elements are electrical semiconductors.


1967. Summary - 1967.4. Structure and General Properties of the Nonmetals
Nonmetals have structures that are very different from those of the metals, primarily because they have greater electronegativity and electrons that are more tightly bound to individual atoms. Most nonmetal oxides are acid anhydrides, meaning that they react with water to form acidic solutions. Molecular structures are common for most of the nonmetals, and several have multiple allotropes with varying physical properties.


1967. Summary - 1967.5. Occurrence, Preparation, and Compounds of Hydrogen
Hydrogen is the most abundant element in the universe and its chemistry is truly unique. Although it has some chemical reactivity that is similar to that of the alkali metals, hydrogen has many of the same chemical properties of a nonmetal with a relatively low electronegativity. It forms ionic hydrides with active metals, covalent compounds in which it has an oxidation state of 1- with less electronegative elements, and covalent compounds in which it has an oxidation state of $1+$ with more electronegative nonmetals. It reacts explosively with oxygen, fluorine, and chlorine, less readily with bromine, and much less readily with iodine, sulfur, and nitrogen. Hydrogen reduces the oxides of metals with lower reduction potentials than chromium to form the metal and water. The hydrogen halides are all acidic when dissolved in water.


1967. Summary - 1967.6. Occurrence, Preparation, and Properties of Carbonates
The usual method for the preparation of the carbonates of the alkali and alkaline earth metals is by reaction of an oxide or hydroxide with carbon dioxide. Other carbonates form by precipitation. Metal carbonates or hydrogen carbonates such as limestone $\left(\mathrm{CaCO}_{3}\right)$, the antacid Tums $\left(\mathrm{CaCO}_{3}\right)$, and baking soda $\left(\mathrm{NaHCO}_{3}\right)$ are common examples. Carbonates and hydrogen carbonates decompose in
the presence of acids and most decompose on heating.


1968. Properties of Nitrogen -
Nitrogen exhibits oxidation states ranging from 3to $5+$. Because of the stability of the $\mathrm{N} \equiv \mathrm{N}$ triple bond, it requires a great deal of energy to make compounds from molecular nitrogen. Active metals such as the alkali metals and alkaline earth metals can reduce nitrogen to form metal nitrides. Nitrogen oxides and nitrogen hydrides are also important substances.


1968. Properties of Nitrogen - 1968.1. Occurrence, Preparation, and Properties of Phosphorus
Phosphorus (group 15) commonly exhibits oxidation states of 3 - with active metals and of $3+$ and $5+$ with more electronegative nonmetals. The halogens and oxygen will oxidize phosphorus. The oxides are phosphorus(V) oxide, $\mathrm{P}_{4} \mathrm{O}_{10}$, and phosphorus(III) oxide, $\mathrm{P}_{4} \mathrm{O}_{6}$. The two common methods for preparing orthophosphoric acid, $\mathrm{H}_{3} \mathrm{PO}_{4}$, are either the reaction of a phosphate with sulfuric acid or the reaction of water with phosphorus(V) oxide. Orthophosphoric acid is a triprotic acid that forms three types of salts.


1968. Properties of Nitrogen - 1968.2. Occurrence, Preparation, and Compounds of Oxygen
Oxygen is one of the most reactive elements. This reactivity, coupled with its abundance, makes the chemistry of oxygen very rich and well understood.

Compounds of the representative metals with oxygen exist in three categories (1) oxides, (2) peroxides and superoxides, and (3) hydroxides. Heating the corresponding hydroxides, nitrates, or carbonates is the most common method for producing oxides. Heating the metal or metal oxide in oxygen may lead to the formation of peroxides and superoxides. The soluble oxides dissolve in water to form solutions of hydroxides. Most metals oxides are base anhydrides and react with acids. The hydroxides of the representative metals react with acids in acid-base reactions to form salts and water.

The hydroxides have many commercial uses.
All nonmetals except fluorine form multiple oxides. Nearly all of the nonmetal oxides are acid anhydrides. The acidity of oxyacids requires that the hydrogen atoms bond to the oxygen atoms in the molecule rather than to the other nonmetal atom. Generally, the strength of the oxyacid increases with the number of oxygen atoms bonded to the nonmetal atom and not to a hydrogen.


1968. Properties of Nitrogen - 1968.3. Occurrence, Preparation, and Properties of Sulfur
Sulfur (group 16) reacts with almost all metals and readily forms the sulfide ion, $\mathrm{S}^{2-}$, in which it has as oxidation state of $2-$. Sulfur reacts with most nonmetals.


1968. Properties of Nitrogen - 1968.4. Occurrence, Preparation, and Properties of Halogens
The halogens form halides with less electronegative elements. Halides of the metals vary from ionic to covalent; halides of nonmetals are covalent. Interhalogens form by the combination of two or more different halogens.

All of the representative metals react directly with elemental halogens or with solutions of the hydrohalic acids (HF, $\mathrm{HCl}, \mathrm{HBr}$, and HI ) to produce representative metal halides. Other laboratory preparations involve the addition of aqueous hydrohalic acids to compounds that contain such basic anions, such as hydroxides, oxides, or carbonates.


1968. Properties of Nitrogen - 1968.5. Occurrence, Preparation, and Properties of the Noble Gases
The most significant property of the noble gases (group 18) is their inactivity. They occur in low concentrations in the atmosphere. They find uses as inert atmospheres, neon signs, and as coolants. The three heaviest noble gases react with fluorine to form fluorides. The xenon fluorides are the best characterized as the starting materials for a few other noble gas compounds.


1969. Exercises - 1969.1. Periodicity
1. How do alkali metals differ from alkaline earth metals in atomic structure and general properties?
2. Why does the reactivity of the alkali metals decrease from cesium to lithium?
3. Predict the formulas for the nine compounds that may form when each species in column 1 of the table reacts with each species in column 2.

1 2

| Na | I |
| :---: | :---: |
| Sr | Se |
| Al | O |


1969. Exercises - 1969.1. Periodicity
4. Predict the best choice in each of the following. You may wish to review the chapter on electronic structure for relevant examples.
(a) the most metallic of the elements $\mathrm{Al}, \mathrm{Be}$, and Ba
(b) the most covalent of the compounds $\mathrm{NaCl}, \mathrm{CaCl}_{2}$, and $\mathrm{BeCl}_{2}$
(c) the lowest first ionization energy among the elements $\mathrm{Rb}, \mathrm{K}$, and Li
(d) the smallest among $\mathrm{Al}, \mathrm{Al}^{+}$, and $\mathrm{Al}^{3+}$
(e) the largest among $\mathrm{Cs}^{+}, \mathrm{Ba}^{2+}$, and Xe
5. Sodium chloride and strontium chloride are both white solids. How could you distinguish one from the other?
6. The reaction of quicklime, CaO , with water produces slaked lime, $\mathrm{Ca}(\mathrm{OH})_{2}$, which is widely used in the construction industry to make mortar and plaster. The reaction of quicklime and water is highly exothermic:
$\mathrm{CaO}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{Ca}(\mathrm{OH})_{2}(s) \quad \Delta H=-350 \mathrm{~kJ} \mathrm{~mol}^{-1}$
(a) What is the enthalpy of reaction per gram of quicklime that reacts?
(b) How much heat, in kilojoules, is associated with the production of 1 ton of slaked lime?
7. Write a balanced equation for the reaction of elemental strontium with each of the following:
(a) oxygen
(b) hydrogen bromide
(c) hydrogen
(d) phosphorus
(e) water
8. How many moles of ionic species are present in 1.0 L of a solution marked 1.0 M mercury(I) nitrate?


1969. Exercises - 1969.1. Periodicity
(c) hydrogen
(d) phosphorus
(e) water
8. How many moles of ionic species are present in 1.0 L of a solution marked 1.0 M mercury(I) nitrate?
9. What is the mass of fish, in kilograms, that one would have to consume to obtain a fatal dose of mercury, if the fish contains 30 parts per million of mercury by weight? (Assume that all the mercury from the fish ends up as mercury(II) chloride in the body and that a fatal dose is 0.20 g of $\mathrm{HgCl}_{2}$.) How many pounds of fish is this?
10. The elements sodium, aluminum, and chlorine are in the same period.
(a) Which has the greatest electronegativity?
(b) Which of the atoms is smallest?
(c) Write the Lewis structure for the simplest covalent compound that can form between aluminum and chlorine.
(d) Will the oxide of each element be acidic, basic, or amphoteric?
11. Does metallic tin react with HCl ?
12. What is tin pest, also known as tin disease?
13. Compare the nature of the bonds in $\mathrm{PbCl}_{2}$ to that of the bonds in $\mathrm{PbCl}_{4}$.
14. Is the reaction of rubidium with water more or less vigorous than that of sodium? How does the rate of reaction of magnesium compare?


1969. Exercises - 1969.2. Occurrence and Preparation of the Representative Metals
15. Write an equation for the reduction of cesium chloride by elemental calcium at high temperature.
16. Why is it necessary to keep the chlorine and sodium, resulting from the electrolysis of sodium chloride, separate during the production of sodium metal?
17. Give balanced equations for the overall reaction in the electrolysis of molten lithium chloride and for the reactions occurring at the electrodes. You may wish to review the chapter on electrochemistry for relevant examples.
18. The electrolysis of molten sodium chloride or of aqueous sodium chloride produces chlorine. Calculate the mass of chlorine produced from 3.00 kg sodium chloride in each case. You may wish to review the chapter on electrochemistry for relevant examples.
19. What mass, in grams, of hydrogen gas forms during the complete reaction of 10.01 g of calcium with water?
20. How many grams of oxygen gas are necessary to react completely with $3.01 \times 10^{21}$ atoms of magnesium to yield magnesium oxide?
21. Magnesium is an active metal; it burns in the form of powder, ribbons, and filaments to provide flashes of brilliant light. Why is it possible to use magnesium in construction?
22. Why is it possible for an active metal like aluminum to be useful as a structural metal?
23. Describe the production of metallic aluminum by electrolytic reduction.
24. What is the common ore of tin and how is tin separated from it?
25. A chemist dissolves a 1.497 -g sample of a type of metal (an alloy of $\mathrm{Sn}, \mathrm{Pb}, \mathrm{Sb}$, and Cu ) in nitric acid, and metastannic acid, $\mathrm{H}_{2} \mathrm{SnO}_{3}$, is precipitated. She heats the precipitate to drive off the water, which leaves 0.4909 g of $\operatorname{tin}(\mathrm{IV})$ oxide. What was the percentage of tin in the original sample?
26. Consider the production of 100 kg of sodium metal using a current of $50,000 \mathrm{~A}$, assuming a $100 \%$ yield.
(a) How long will it take to produce the 100 kg of sodium metal?


1969. Exercises - 1969.2. Occurrence and Preparation of the Representative Metals
(a) How long will it take to produce the 100 kg of sodium metal?
(b) What volume of chlorine at $25^{\circ} \mathrm{C}$ and 1.00 atm forms?
27. What mass of magnesium forms when $100,000 \mathrm{~A}$ is passed through a $\mathrm{MgCl}_{2}$ melt for 1.00 h if the yield of magnesium is $85 \%$ of the theoretical yield?


1969. Exercises - 1969.3. Structure and General Properties of the Metalloids
28. Give the hybridization of the metalloid and the molecular geometry for each of the following compounds or ions. You may wish to review the chapters on chemical bonding and advanced covalent bonding for relevant examples.
(a) $\mathrm{GeH}_{4}$
(b) $\mathrm{SbF}_{3}$
(c) $\mathrm{Te}(\mathrm{OH})_{6}$
(d) $\mathrm{H}_{2} \mathrm{Te}$
(e) $\mathrm{GeF}_{2}$
(f) $\mathrm{TeCl}_{4}$
(g) $\mathrm{SiF}_{6}{ }^{2-}$
(h) $\mathrm{SbCl}_{5}$
(i) $\mathrm{TeF}_{6}$
29. Write a Lewis structure for each of the following molecules or ions. You may wish to review the chapter on chemical bonding.
(a) $\mathrm{H}_{3} \mathrm{BPH}_{3}$
(b) $\mathrm{BF}_{4}{ }^{-}$
(c) $\mathrm{BBr}_{3}$
(d) $\mathrm{B}\left(\mathrm{CH}_{3}\right)_{3}$
(e) $\mathrm{B}(\mathrm{OH})_{3}$
30. Describe the hybridization of boron and the molecular structure about the boron in each of the following:
(a) $\mathrm{H}_{3} \mathrm{BPH}_{3}$
(b) $\mathrm{BF}_{4}^{-}$
(c) $\mathrm{BBr}_{3}$
(d) $\mathrm{B}\left(\mathrm{CH}_{3}\right)_{3}$
(e) $\mathrm{B}(\mathrm{OH})_{3}$
31. Using only the periodic table, write the complete electron configuration for silicon, including any empty orbitals in the valence shell. You may wish to review the chapter on electronic structure.
32. Write a Lewis structure for each of the following molecules and ions:
(a) $\left(\mathrm{CH}_{3}\right)_{3} \mathrm{SiH}$
(b) $\mathrm{SiO}_{4}{ }^{4-}$
(c) $\mathrm{Si}_{2} \mathrm{H}_{6}$
(d) $\mathrm{Si}(\mathrm{OH})_{4}$


1969. Exercises - 1969.3. Structure and General Properties of the Metalloids
(b) $\mathrm{SiO}_{4}{ }^{4-}$
(c) $\mathrm{Si}_{2} \mathrm{H}_{6}$
(d) $\mathrm{Si}(\mathrm{OH})_{4}$
(e) $\mathrm{SiF}_{6}{ }^{2-}$
33. Describe the hybridization of silicon and the molecular structure of the following molecules and ions:
(a) $\left(\mathrm{CH}_{3}\right)_{3} \mathrm{SiH}$
(b) $\mathrm{SiO}_{4}{ }^{4-}$
(c) $\mathrm{Si}_{2} \mathrm{H}_{6}$
(d) $\mathrm{Si}(\mathrm{OH})_{4}$
(e) $\mathrm{SiF}_{6}{ }^{2-}$
34. Describe the hybridization and the bonding of a silicon atom in elemental silicon.
35. Classify each of the following molecules as polar or nonpolar. You may wish to review the chapter on chemical bonding.
(a) $\mathrm{SiH}_{4}$
(b) $\mathrm{Si}_{2} \mathrm{H}_{6}$
(c) $\mathrm{SiCl}_{3} \mathrm{H}$
(d) $\mathrm{SiF}_{4}$
(e) $\mathrm{SiCl}_{2} \mathrm{~F}_{2}$
36. Silicon reacts with sulfur at elevated temperatures. If 0.0923 g of silicon reacts with sulfur to give 0.3030 g of silicon sulfide, determine the empirical formula of silicon sulfide.
37. Name each of the following compounds:
(a) $\mathrm{TeO}_{2}$
(b) $\mathrm{Sb}_{2} \mathrm{~S}_{3}$
(c) $\mathrm{GeF}_{4}$
(d) $\mathrm{SiH}_{4}$
(e) $\mathrm{GeH}_{4}$
38. Write a balanced equation for the reaction of elemental boron with each of the following (most of these reactions require high temperature):
(a) $\mathrm{F}_{2}$
(b) $\mathrm{O}_{2}$
(c) S
(d) Se
(e) $\mathrm{Br}_{2}$
39. Why is boron limited to a maximum coordination number of four in its compounds?


1969. Exercises - 1969.3. Structure and General Properties of the Metalloids
(b) $\mathrm{O}_{2}$
(c) S
(d) Se
(e) $\mathrm{Br}_{2}$
39. Why is boron limited to a maximum coordination number of four in its compounds?
40. Write a formula for each of the following compounds:
(a) silicon dioxide
(b) silicon tetraiodide
(c) silane
(d) silicon carbide
(e) magnesium silicide
41. From the data given in Appendix G, determine the standard enthalpy change and the standard free energy change for each of the following reactions:
(a) $\mathrm{BF}_{3}(g)+3 \mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{B}(\mathrm{OH})_{3}(s)+3 \mathrm{HF}(g)$
(b) $\mathrm{BCl}_{3}(g)+3 \mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{B}(\mathrm{OH})_{3}(s)+3 \mathrm{HCl}(g)$
(c) $\mathrm{B}_{2} \mathrm{H}_{6}(g)+6 \mathrm{H}_{2} \mathrm{O}(l) \longrightarrow 2 \mathrm{~B}(\mathrm{OH})_{3}(s)+6 \mathrm{H}_{2}(g)$
42. A hydride of silicon prepared by the reaction of $\mathrm{Mg}_{2} \mathrm{Si}$ with acid exerted a pressure of 306 torr at $26^{\circ} \mathrm{C}$ in a bulb with a volume of 57.0 mL . If the mass of the hydride was 0.0861 g , what is its molecular mass? What is the molecular formula for the hydride?
43. Suppose you discovered a diamond completely encased in a silicate rock. How would you chemically free the diamond without harming it?


1969. Exercises - 1969.4. Structure and General Properties of the Nonmetals
44. Carbon forms a number of allotropes, two of which are graphite and diamond. Silicon has a diamond structure. Why is there no allotrope of silicon with a graphite structure?
45. Nitrogen in the atmosphere exists as very stable diatomic molecules. Why does phosphorus form less stable $\mathrm{P}_{4}$ molecules instead of $\mathrm{P}_{2}$ molecules?
46. Write balanced chemical equations for the reaction of the following acid anhydrides with water:
(a) $\mathrm{SO}_{3}$
(b) $\mathrm{N}_{2} \mathrm{O}_{3}$
(c) $\mathrm{Cl}_{2} \mathrm{O}_{7}$
(d) $\mathrm{P}_{4} \mathrm{O}_{10}$
(e) $\mathrm{NO}_{2}$
47. Determine the oxidation number of each element in each of the following compounds:
(a) HCN
(b) $\mathrm{OF}_{2}$
(c) $\mathrm{AsCl}_{3}$
48. Determine the oxidation state of sulfur in each of the following:
(a) $\mathrm{SO}_{3}$
(b) $\mathrm{SO}_{2}$
(c) $\mathrm{SO}_{3}{ }^{2-}$
49. Arrange the following in order of increasing electronegativity: $\mathrm{F} ; \mathrm{Cl} ; \mathrm{O} ;$ and S .
50. Why does white phosphorus consist of tetrahedral $P_{4}$ molecules while nitrogen consists of diatomic $N_{2}$ molecules?


1969. Exercises - 1969.5. Occurrence, Preparation, and Compounds of Hydrogen
51. Why does hydrogen not exhibit an oxidation state of 1 - when bonded to nonmetals?
52. The reaction of calcium hydride, $\mathrm{CaH}_{2}$, with water can be characterized as a Lewis acid-base reaction: $\mathrm{CaH}_{2}(s)+2 \mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{Ca}(\mathrm{OH})_{2}(a q)+2 \mathrm{H}_{2}(g)$
Identify the Lewis acid and the Lewis base among the reactants. The reaction is also an oxidationreduction reaction. Identify the oxidizing agent, the reducing agent, and the changes in oxidation number that occur in the reaction.
53. In drawing Lewis structures, we learn that a hydrogen atom forms only one bond in a covalent compound. Why?
54. What mass of $\mathrm{CaH}_{2}$ is necessary to react with water to provide enough hydrogen gas to fill a balloon at 20 ${ }^{\circ} \mathrm{C}$ and 0.8 atm pressure with a volume of 4.5 L ? The balanced equation is:
$\mathrm{CaH}_{2}(s)+2 \mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{Ca}(\mathrm{OH})_{2}(a q)+2 \mathrm{H}_{2}(g)$
55. What mass of hydrogen gas results from the reaction of 8.5 g of KH with water?
$\mathrm{KH}+\mathrm{H}_{2} \mathrm{O} \longrightarrow \mathrm{KOH}+\mathrm{H}_{2}$


1969. Exercises - 1969.6. Occurrence, Preparation, and Properties of Carbonates
56. Carbon forms the $\mathrm{CO}_{3}{ }^{2-}$ ion, yet silicon does not form an analogous $\mathrm{SiO}_{3}{ }^{2-}$ ion. Why?
57. Complete and balance the following chemical equations:
(a) hardening of plaster containing slaked lime
$\mathrm{Ca}(\mathrm{OH})_{2}+\mathrm{CO}_{2} \longrightarrow$
(b) removal of sulfur dioxide from the flue gas of power plants
$\mathrm{CaO}+\mathrm{SO}_{2} \longrightarrow$
(c) the reaction of baking powder that produces carbon dioxide gas and causes bread to rise
$\mathrm{NaHCO}_{3}+\mathrm{NaH}_{2} \mathrm{PO}_{4} \longrightarrow$
58. Heating a sample of $\mathrm{Na}_{2} \mathrm{CO}_{3} \cdot \mathrm{xH}_{2} \mathrm{O}$ weighing 4.640 g until the removal of the water of hydration leaves 1.720 g of anhydrous $\mathrm{Na}_{2} \mathrm{CO}_{3}$. What is the formula of the hydrated compound?


1969. Exercises - 1969.7. Occurrence, Preparation, and Properties of Nitrogen
59. Write the Lewis structures for each of the following:
(a) $\mathrm{NH}^{2-}$
(b) $\mathrm{N}_{2} \mathrm{~F}_{4}$
(c) $\mathrm{NH}_{2}{ }^{-}$
(d) $\mathrm{NF}_{3}$
(e) $\mathrm{N}_{3}{ }^{-}$
60. For each of the following, indicate the hybridization of the nitrogen atom (for $\mathrm{N}_{3}{ }^{-}$, the central nitrogen).
(a) $\mathrm{N}_{2} \mathrm{~F}_{4}$
(b) $\mathrm{NH}_{2}{ }^{-}$
(c) $\mathrm{NF}_{3}$
(d) $\mathrm{N}_{3}{ }^{-}$
61. Explain how ammonia can function both as a Brønsted base and as a Lewis base.
62. Determine the oxidation state of nitrogen in each of the following. You may wish to review the chapter on chemical bonding for relevant examples.
(a) $\mathrm{NCl}_{3}$
(b) ClNO
(c) $\mathrm{N}_{2} \mathrm{O}_{5}$
(d) $\mathrm{N}_{2} \mathrm{O}_{3}$
(e) $\mathrm{NO}_{2}{ }^{-}$
(f) $\mathrm{N}_{2} \mathrm{O}_{4}$
(g) $\mathrm{N}_{2} \mathrm{O}$
(h) $\mathrm{NO}_{3}{ }^{-}$
(i) $\mathrm{HNO}_{2}$
(j) $\mathrm{HNO}_{3}$
63. For each of the following, draw the Lewis structure, predict the ONO bond angle, and give the hybridization of the nitrogen. You may wish to review the chapters on chemical bonding and advanced theories of covalent bonding for relevant examples.
(a) $\mathrm{NO}_{2}$
(b) $\mathrm{NO}_{2}{ }^{-}$
(c) $\mathrm{NO}_{2}{ }^{+}$
64. How many grams of gaseous ammonia will the reaction of 3.0 g hydrogen gas and 3.0 g of nitrogen gas produce?


1969. Exercises - 1969.7. Occurrence, Preparation, and Properties of Nitrogen
(c) $\mathrm{NO}_{2}{ }^{+}$
64. How many grams of gaseous ammonia will the reaction of 3.0 g hydrogen gas and 3.0 g of nitrogen gas produce?
65. Although $\mathrm{PF}_{5}$ and $\mathrm{AsF}_{5}$ are stable, nitrogen does not form $\mathrm{NF}_{5}$ molecules. Explain this difference among members of the same group.
66. The equivalence point for the titration of a $25.00-\mathrm{mL}$ sample of CsOH solution with $0.1062 \mathrm{M} \mathrm{HNO}_{3}$ is at 35.27 mL . What is the concentration of the CsOH solution?


1969. Exercises - 1969.8. Occurrence, Preparation, and Properties of Phosphorus
67. Write the Lewis structure for each of the following. You may wish to review the chapter on chemical bonding and molecular geometry.
(a) $\mathrm{PH}_{3}$
(b) $\mathrm{PH}_{4}{ }^{+}$
(c) $\mathrm{P}_{2} \mathrm{H}_{4}$
(d) $\mathrm{PO}_{4}{ }^{3-}$
(e) $\mathrm{PF}_{5}$
68. Describe the molecular structure of each of the following molecules or ions listed. You may wish to review the chapter on chemical bonding and molecular geometry.
(a) $\mathrm{PH}_{3}$
(b) $\mathrm{PH}_{4}{ }^{+}$
(c) $\mathrm{P}_{2} \mathrm{H}_{4}$
(d) $\mathrm{PO}_{4}{ }^{3-}$
69. Complete and balance each of the following chemical equations. (In some cases, there may be more than one correct answer.)
(a) $\mathrm{P}_{4}+\mathrm{Al} \longrightarrow$
(b) $\mathrm{P}_{4}+\mathrm{Na} \longrightarrow$
(c) $\mathrm{P}_{4}+\mathrm{F}_{2} \longrightarrow$
(d) $\mathrm{P}_{4}+\mathrm{Cl}_{2} \longrightarrow$
(e) $\mathrm{P}_{4}+\mathrm{O}_{2} \longrightarrow$
(f) $\mathrm{P}_{4} \mathrm{O}_{6}+\mathrm{O}_{2} \longrightarrow$
70. Describe the hybridization of phosphorus in each of the following compounds: $\mathrm{P}_{4} \mathrm{O}_{10}, \mathrm{P}_{4} \mathrm{O}_{6}, \mathrm{PH}_{4} \mathrm{I}$ (an ionic compound), $\mathrm{PBr}_{3}, \mathrm{H}_{3} \mathrm{PO}_{4}, \mathrm{H}_{3} \mathrm{PO}_{3}, \mathrm{PH}_{3}$, and $\mathrm{P}_{2} \mathrm{H}_{4}$. You may wish to review the chapter on advanced theories of covalent bonding.


1969. Exercises - 1969.8. Occurrence, Preparation, and Properties of Phosphorus
71. What volume of 0.200 M NaOH is necessary to neutralize the solution produced by dissolving 2.00 g of $\mathrm{PCl}_{3}$ is an excess of water? Note that when $\mathrm{H}_{3} \mathrm{PO}_{3}$ is titrated under these conditions, only one proton of the acid molecule reacts.
72. How much $\mathrm{POCl}_{3}$ can form from 25.0 g of $\mathrm{PCl}_{5}$ and the appropriate amount of $\mathrm{H}_{2} \mathrm{O}$ ?
73. How many tons of $\mathrm{Ca}_{3}\left(\mathrm{PO}_{4}\right)_{2}$ are necessary to prepare 5.0 tons of phosphorus if the yield is $90 \%$ ?
74. Write equations showing the stepwise ionization of phosphorous acid.
75. Draw the Lewis structures and describe the geometry for the following:
(a) $\mathrm{PF}_{4}{ }^{+}$
(b) $\mathrm{PF}_{5}$
(c) $\mathrm{PF}_{6}{ }^{-}$
(d) $\mathrm{POF}_{3}$
76. Why does phosphorous acid form only two series of salts, even though the molecule contains three hydrogen atoms?
77. Assign an oxidation state to phosphorus in each of the following:
(a) $\mathrm{NaH}_{2} \mathrm{PO}_{3}$
(b) $\mathrm{PF}_{5}$
(c) $\mathrm{P}_{4} \mathrm{O}_{6}$
(d) $\mathrm{K}_{3} \mathrm{PO}_{4}$
(e) $\mathrm{Na}_{3} \mathrm{P}$
(f) $\mathrm{Na}_{4} \mathrm{P}_{2} \mathrm{O}_{7}$
78. Phosphoric acid, one of the acids used in some cola drinks, is produced by the reaction of phosphorus(V) oxide, an acidic oxide, with water. Phosphorus(V) oxide is prepared by the combustion of phosphorus.
(a) Write the empirical formula of phosphorus(V) oxide.
(b) What is the molecular formula of phosphorus(V) oxide if the molar mass is about 280.


1969. Exercises - 1969.8. Occurrence, Preparation, and Properties of Phosphorus
(a) Write the empirical formula of phosphorus(V) oxide.
(b) What is the molecular formula of phosphorus(V) oxide if the molar mass is about 280.
(c) Write balanced equations for the production of phosphorus(V) oxide and phosphoric acid.
(d) Determine the mass of phosphorus required to make $1.00 \times 10^{4} \mathrm{~kg}$ of phosphoric acid, assuming a yield of $98.85 \%$.


1969. Exercises - 1969.9. Occurrence, Preparation, and Compounds of Oxygen
79. Predict the product of burning francium in air.
80. Using equations, describe the reaction of water with potassium and with potassium oxide.
81. Write balanced chemical equations for the following reactions:
(a) zinc metal heated in a stream of oxygen gas
(b) zinc carbonate heated until loss of mass stops
(c) zinc carbonate added to a solution of acetic acid, $\mathrm{CH}_{3} \mathrm{CO}_{2} \mathrm{H}$
(d) zinc added to a solution of hydrobromic acid
82. Write balanced chemical equations for the following reactions:
(a) cadmium burned in air
(b) elemental cadmium added to a solution of hydrochloric acid
(c) cadmium hydroxide added to a solution of acetic acid, $\mathrm{CH}_{3} \mathrm{CO}_{2} \mathrm{H}$
83. Illustrate the amphoteric nature of aluminum hydroxide by citing suitable equations.
84. Write balanced chemical equations for the following reactions:
(a) metallic aluminum burned in air
(b) elemental aluminum heated in an atmosphere of chlorine
(c) aluminum heated in hydrogen bromide gas
(d) aluminum hydroxide added to a solution of nitric acid
85. Write balanced chemical equations for the following reactions:
(a) sodium oxide added to water
(b) cesium carbonate added to an excess of an aqueous solution of HF
(c) aluminum oxide added to an aqueous solution of $\mathrm{HClO}_{4}$
(d) a solution of sodium carbonate added to solution of barium nitrate
(e) titanium metal produced from the reaction of titanium tetrachloride with elemental sodium
86. What volume of $0.250 \mathrm{MH}_{2} \mathrm{SO}_{4}$ solution is required to neutralize a solution that contains 5.00 g of $\mathrm{CaCO}_{3}$ ?
87. Which is the stronger acid, $\mathrm{HClO}_{4}$ or $\mathrm{HBrO}_{4}$ ? Why?
88. Write a balanced chemical equation for the reaction of an excess of oxygen with each of the following. Remember that oxygen is a strong oxidizing agent and tends to oxidize an element to its maximum oxidation state.
(a) Mg
(b) Rb
(c) Ga


1969. Exercises - 1969.9. Occurrence, Preparation, and Compounds of Oxygen
(a) Mg
(b) Rb
(c) Ga
(d) $\mathrm{C}_{2} \mathrm{H}_{2}$
(e) CO
89. Which is the stronger acid, $\mathrm{H}_{2} \mathrm{SO}_{4}$ or $\mathrm{H}_{2} \mathrm{SeO}_{4}$ ? Why? You may wish to review the chapter on acid-base equilibria.


1969. Exercises - 1969.10. Occurrence, Preparation, and Properties of Sulfur
90. Explain why hydrogen sulfide is a gas at room temperature, whereas water, which has a lower molecular mass, is a liquid.
91. Give the hybridization and oxidation state for sulfur in $\mathrm{SO}_{2}$, in $\mathrm{SO}_{3}$, and in $\mathrm{H}_{2} \mathrm{SO}_{4}$.
92. Which is the stronger acid, $\mathrm{NaHSO}_{3}$ or $\mathrm{NaHSO}_{4}$ ?
93. Determine the oxidation state of sulfur in $\mathrm{SF}_{6}, \mathrm{SO}_{2} \mathrm{~F}_{2}$, and KHS.
94. Which is a stronger acid, sulfurous acid or sulfuric acid? Why?
95. Oxygen forms double bonds in $\mathrm{O}_{2}$, but sulfur forms single bonds in $\mathrm{S}_{8}$. Why?
96. Give the Lewis structure of each of the following:
(a) $\mathrm{SF}_{4}$
(b) $\mathrm{K}_{2} \mathrm{SO}_{4}$
(c) $\mathrm{SO}_{2} \mathrm{Cl}_{2}$
(d) $\mathrm{H}_{2} \mathrm{SO}_{3}$
(e) $\mathrm{SO}_{3}$
97. Write two balanced chemical equations in which sulfuric acid acts as an oxidizing agent.
98. Explain why sulfuric acid, $\mathrm{H}_{2} \mathrm{SO}_{4}$, which is a covalent molecule, dissolves in water and produces a solution that contains ions.
99. How many grams of Epsom salts $\left(\mathrm{MgSO}_{4} \cdot 7 \mathrm{H}_{2} \mathrm{O}\right)$ will form from 5.0 kg of magnesium?


1969. Exercises - 1969.11. Occurrence, Preparation, and Properties of Halogens
100. What does it mean to say that mercury(II) halides are weak electrolytes?
101. Why is $\mathrm{SnCl}_{4}$ not classified as a salt?
102. The following reactions are all similar to those of the industrial chemicals. Complete and balance the equations for these reactions:
(a) reaction of a weak base and a strong acid
$\mathrm{NH}_{3}+\mathrm{HClO}_{4} \longrightarrow$
(b) preparation of a soluble silver salt for silver plating
$\mathrm{Ag}_{2} \mathrm{CO}_{3}+\mathrm{HNO}_{3} \longrightarrow$
(c) preparation of strontium hydroxide by electrolysis of a solution of strontium chloride
$\mathrm{SrCl}_{2}(a q)+\mathrm{H}_{2} \mathrm{O}(l) \xrightarrow{\text { electrolysis }}$
103. Which is the stronger acid, $\mathrm{HClO}_{3}$ or $\mathrm{HBrO}_{3}$ ? Why?
104. What is the hybridization of iodine in $\mathrm{IF}_{3}$ and $\mathrm{IF}_{5}$ ?
105. Predict the molecular geometries and draw Lewis structures for each of the following. You may wish to review the chapter on chemical bonding and molecular geometry.
(a) $\mathrm{IF}_{5}$
(b) $\mathrm{I}_{3}{ }^{-}$
(c) $\mathrm{PCl}_{5}$
(d) $\mathrm{SeF}_{4}$
(e) $\mathrm{ClF}_{3}$
106. Which halogen has the highest ionization energy? Is this what you would predict based on what you have learned about periodic properties?
107. Name each of the following compounds:
(a) $\mathrm{BrF}_{3}$
(b) $\mathrm{NaBrO}_{3}$
(c) $\mathrm{PBr}_{5}$
(d) $\mathrm{NaClO}_{4}$
(e) KClO
108. Explain why, at room temperature, fluorine and chlorine are gases, bromine is a liquid, and iodine is a solid.
109. What is the oxidation state of the halogen in each of the following?
(a) $\mathrm{H}_{5} \mathrm{IO}_{6}$


1969. Exercises - 1969.11. Occurrence, Preparation, and Properties of Halogens
109. What is the oxidation state of the halogen in each of the following?
(a) $\mathrm{H}_{5} \mathrm{IO}_{6}$
(b) $\mathrm{IO}_{4}^{-}$
(c) $\mathrm{ClO}_{2}$
(d) $\mathrm{ICl}_{3}$
(e) $\mathrm{F}_{2}$
110. Physiological saline concentration-that is, the sodium chloride concentration in our bodies-is approximately 0.16 M . A saline solution for contact lenses is prepared to match the physiological concentration. If you purchase 25 mL of contact lens saline solution, how many grams of sodium chloride have you bought?


1969. Exercises - 1969.12. Occurrence, Preparation, and Properties of the Noble Gases
111. Give the hybridization of xenon in each of the following. You may wish to review the chapter on the advanced theories of covalent bonding.
(a) $\mathrm{XeF}_{2}$
(b) $\mathrm{XeF}_{4}$
(c) $\mathrm{XeO}_{3}$
(d) $\mathrm{XeO}_{4}$
(e) $\mathrm{XeOF}_{4}$
112. What is the molecular structure of each of the following molecules? You may wish to review the chapter on chemical bonding and molecular geometry.
(a) $\mathrm{XeF}_{2}$
(b) $\mathrm{XeF}_{4}$
(c) $\mathrm{XeO}_{3}$
(d) $\mathrm{XeO}_{4}$
(e) $\mathrm{XeOF}_{4}$
113. Indicate whether each of the following molecules is polar or nonpolar. You may wish to review the chapter on chemical bonding and molecular geometry.
(a) $\mathrm{XeF}_{2}$
(b) $\mathrm{XeF}_{4}$
(c) $\mathrm{XeO}_{3}$
(d) $\mathrm{XeO}_{4}$
(e) $\mathrm{XeOF}_{4}$
114. What is the oxidation state of the noble gas in each of the following? You may wish to review the chapter on chemical bonding and molecular geometry.
(a) $\mathrm{XeO}_{2} \mathrm{~F}_{2}$
(b) $\mathrm{KrF}_{2}$
(c) $\mathrm{XeF}_{3}{ }^{+}$
(d) $\mathrm{XeO}_{6}{ }^{4-}$
(e) $\mathrm{XeO}_{3}$
115. A mixture of xenon and fluorine was heated. A sample of the white solid that formed reacted with hydrogen to yield 81 mL of xenon (at STP) and hydrogen fluoride, which was collected in water, giving a solution of hydrofluoric acid. The hydrofluoric acid solution was titrated, and 68.43 mL of 0.3172 M sodium hydroxide was required to reach the equivalence point. Determine the empirical formula for the white solid and write balanced chemical equations for the reactions involving xenon.


1969. Exercises - 1969.12. Occurrence, Preparation, and Properties of the Noble Gases
116. Basic solutions of $\mathrm{Na}_{4} \mathrm{XeO}_{6}$ are powerful oxidants. What mass of $\mathrm{Mn}\left(\mathrm{NO}_{3}\right)_{2} \cdot 6 \mathrm{H}_{2} \mathrm{O}$ reacts with 125.0 mL of a 0.1717 M basic solution of $\mathrm{Na}_{4} \mathrm{XeO}_{6}$ that contains an excess of sodium hydroxide if the products include Xe and solution of sodium permanganate?


1970. CHAPTER 19 Transition Metals and Coordination Chemistry -
Figure 19.1 Transition metals often form vibrantly colored complexes. The minerals malachite (green), azurite (blue), and proustite (red) are some examples. (credit left: modification of work by James St. John; credit middle: modification of work by Stephanie Clifford; credit right: modification of work by Terry Wallace)


1971. CHAPTER OUTLINE - 1971.1. Occurrence, Preparation, and Properties of Transition Metals and Their Compounds 19.2 Coordination Chemistry of Transition Metals 19.3 Spectroscopic and Magnetic Properties of Coordination Compounds
INTRODUCTION We have daily contact with many transition metals. Iron occurs everywhere-from the rings in your spiral notebook and the cutlery in your kitchen to automobiles, ships, buildings, and in the hemoglobin in your blood. Titanium is useful in the manufacture of lightweight, durable products such as bicycle frames, artificial hips, and jewelry. Chromium is useful as a protective plating on plumbing fixtures and automotive detailing.

In addition to being used in their pure elemental forms, many compounds containing transition metals have numerous other applications. Silver nitrate is used to create mirrors, zirconium silicate provides friction in automotive brakes, and many important cancer-fighting agents, like the drug cisplatin and related species, are platinum compounds.

The variety of properties exhibited by transition metals is due to their complex valence shells. Unlike most main group metals where one oxidation state is normally observed, the valence shell structure of transition metals means that they usually occur in several different stable oxidation states. In addition, electron transitions in these elements can correspond with absorption of photons in the visible electromagnetic spectrum, leading to colored compounds. Because of these behaviors, transition metals exhibit a rich and fascinating chemistry.


1971. CHAPTER OUTLINE - 1971.2. Occurrence, Preparation, and Properties of Transition Metals and Their Compounds
Transition metals are defined as those elements that have (or readily form) partially filled $d$ orbitals. As shown in Figure 19.2, the $\boldsymbol{d}$-block elements in groups 3-11 are transition elements. The $\boldsymbol{f}$-block elements, also called inner transition metals (the lanthanides and actinides), also meet this criterion because the $d$ orbital is partially occupied before the $f$ orbitals. The $d$ orbitals fill with the copper family (group 11); for this reason, the
next family (group 12) are technically not transition elements. However, the group 12 elements do display some of the same chemical properties and are commonly included in discussions of transition metals. Some chemists do treat the group 12 elements as transition metals.

FIGURE 19.2 The transition metals are located in groups 3-11 of the periodic table. The inner transition metals are in the two rows below the body of the table.

The $d$-block elements are divided into the first transition series (the elements Sc through Cu ), the second transition series (the elements Y through Ag), and the third transition series (the element La and the elements Hf through Au). Actinium, Ac, is the first member of the fourth transition series, which also includes Rf through Rg.

The $f$-block elements are the elements Ce through Lu, which constitute the lanthanide series (or lanthanoid series), and the elements Th through Lr, which constitute the actinide series (or actinoid series). Because lanthanum behaves very much like the lanthanide elements, it is considered a lanthanide element, even though its electron configuration makes it the first member of the third transition series. Similarly, the behavior of actinium means it is part of the actinide series, although its electron configuration makes it the first member of the fourth transition series.


1973. Valence Electrons in Transition Metals -
Review how to write electron configurations, covered in the chapter on electronic structure and periodic properties of elements. Recall that for the transition and inner transition metals, it is necessary to remove the $s$
electrons before the $d$ or $f$ electrons. Then, for each ion, give the electron configuration:
(a) cerium(III)
(b) lead(II)
(c) $\mathrm{Ti}^{2+}$
(d) $\mathrm{Am}^{3+}$
(e) $\mathrm{Pd}^{2+}$

For the examples that are transition metals, determine to which series they belong.


1974. Solution -
For ions, the $s$-valence electrons are lost prior to the $d$ or $f$ electrons.
(a) $\mathrm{Ce}^{3+}[\mathrm{Xe}] 4 f^{1} ; \mathrm{Ce}^{3+}$ is an inner transition element in the lanthanide series.
(b) $\mathrm{Pb}^{2+}[\mathrm{Xe}] 6 s^{2} 5 d^{10} 4 f^{14}$; the electrons are lost from the $p$ orbital. This is a main group element.
(c) titanium(II) $[\mathrm{Ar}] 3 d^{2}$; first transition series
(d) americium(III) $[\mathrm{Rn}] 5 f^{6}$; actinide
(e) palladium(II) $[\mathrm{Kr}] 4 d^{8}$; second transition series


1975. Check Your Learning -
Give an example of an ion from the first transition series with no $d$ electrons.


1976. Answer: -
$\mathrm{V}^{5+}$ is one possibility. Other examples include $\mathrm{Sc}^{3+}, \mathrm{Ti}^{4+}, \mathrm{Cr}^{6+}$, and $\mathrm{Mn}^{7+}$.


1978. Uses of Lanthanides in Devices -
Lanthanides (elements 57-71) are fairly abundant in the earth's crust, despite their historic characterization as rare earth elements. Thulium, the rarest naturally occurring lanthanoid, is more common in the earth's crust than silver ( $4.5 \times 10^{-5} \%$ versus $0.79 \times 10^{-5} \%$ by mass). There are 17 rare earth elements, consisting of the 15 lanthanoids plus scandium and yttrium. They are called rare because they were once difficult to extract economically, so it was rare to have a pure sample; due to similar chemical properties, it is difficult to separate any one lanthanide from the others. However, newer separation methods, such as ion exchange resins similar to those found in home water softeners, make the separation of these elements easier and more economical. Most ores that contain these elements have low concentrations of all the rare earth elements mixed together.

The commercial applications of lanthanides are growing rapidly. For example, europium is important in flat screen displays found in computer monitors, cell phones, and televisions. Neodymium is useful in laptop hard drives and in the processes that convert crude oil into gasoline (Figure 19.3). Holmium is found in dental and medical equipment. In addition, many alternative energy technologies rely heavily on lanthanoids. Neodymium and dysprosium are key components of hybrid vehicle engines and the magnets used in wind turbines.

FIGURE 19.3 (a) Europium is used in display screens for televisions, computer monitors, and cell phones. (b) Neodymium magnets are commonly found in computer hard drives. (credit b: modification of work by "KUERT Datenrettung"/Flickr)

As the demand for lanthanide materials has increased faster than supply, prices have also increased. In 2008 , dysprosium cost $\$ 110 / \mathrm{kg}$; by 2014 , the price had increased to $\$ 470 / \mathrm{kg}$. Increasing the supply of lanthanoid elements is one of the most significant challenges facing the industries that rely on the optical and magnetic properties of these materials.


1978. Uses of Lanthanides in Devices -
The transition elements have many properties in common with other metals. They are almost all hard, highmelting solids that conduct heat and electricity well. They readily form alloys and lose electrons to form stable cations. In addition, transition metals form a wide variety of stable coordination compounds, in which the central metal atom or ion acts as a Lewis acid and accepts one or more pairs of electrons. Many different molecules and ions can donate lone pairs to the metal center, serving as Lewis bases. In this chapter, we shall focus primarily on the chemical behavior of the elements of the first transition series.


1979. Properties of the Transition Elements -
Transition metals demonstrate a wide range of chemical behaviors. As can be seen from their reduction potentials (see Appendix H), some transition metals are strong reducing agents, whereas others have very low reactivity. For example, the lanthanides all form stable 3+ aqueous cations. The driving force for such oxidations is similar to that of alkaline earth metals such as Be or Mg , forming $\mathrm{Be}^{2+}$ and $\mathrm{Mg}^{2+}$. On the other hand, materials like platinum and gold have much higher reduction potentials. Their ability to resist oxidation makes them useful materials for constructing circuits and jewelry.

Ions of the lighter $d$-block elements, such as $\mathrm{Cr}^{3+}, \mathrm{Fe}^{3+}$, and $\mathrm{Co}^{2+}$, form colorful hydrated ions that are stable in water. However, ions in the period just below these $\left(\mathrm{Mo}^{3+}, \mathrm{Ru}^{3+}\right.$, and $\mathrm{Ir}^{2+}$ ) are unstable and react readily with oxygen from the air. The majority of simple, water-stable ions formed by the heavier $d$-block elements are oxyanions such as $\mathrm{MoO}_{4}{ }^{2-}$ and $\mathrm{ReO}_{4}{ }^{-}$.

Ruthenium, osmium, rhodium, iridium, palladium, and platinum are the platinum metals. With difficulty, they form simple cations that are stable in water, and, unlike the earlier elements in the second and third transition series, they do not form stable oxyanions.
Both the $d$ - and $f$-block elements react with nonmetals to form binary compounds; heating is often required. These elements react with halogens to form a variety of halides ranging in oxidation state from $1+$ to $6+$. On heating, oxygen reacts with all of the transition elements except palladium, platinum, silver, and gold. The oxides of these latter metals can be formed using other reactants, but they decompose upon heating. The $f$-block elements, the elements of group 3, and the elements of the first transition series except copper react with aqueous solutions of acids, forming hydrogen gas and solutions of the corresponding salts.


1979. Properties of the Transition Elements -
Transition metals can form compounds with a wide range of oxidation states. Some of the observed oxidation states of the elements of the first transition series are shown in Figure 19.4. As we move from left to right across the first transition series, we see that the number of common oxidation states increases at first to a maximum towards the middle of the table, then decreases. The values in the table are typical values; there are other known values, and it is possible to synthesize new additions. For example, in 2014, researchers were successful in synthesizing a new oxidation state of iridium (9+).

| ${ }^{21} \mathbf{S c}$ | ${ }^{22} \mathrm{Ti}$ | ${ }^{23} \mathrm{~V}$ | ${ }^{24} \mathrm{Cr}$ | ${ }^{25} \mathrm{Mn}$ | ${ }^{26} \mathrm{Fe}$ | ${ }^{27} \mathrm{Co}$ | ${ }^{28} \mathrm{Ni}$ | ${ }^{29} \mathrm{Cu}$ | ${ }^{30} \mathrm{Zn}$ |
| :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: | :---: |
|  |  |  |  |  |  |  |  | $1+$ |  |
|  |  | $2+$ | ${ }^{2+}$ | $2+$ | ${ }^{2+}$ | $2+$ | ${ }^{2+}$ | $2+$ | 2+ |
| $3+$ | $3+$ | $3+$ | $3+$ | $3+$ | $3+$ | $3+$ | $3+$ | $3+$ |  |
|  | $4+$ | $4+$ | $4+$ | $4+$ |  |  |  |  |  |
|  |  | $5+$ |  |  |  |  |  |  |  |
|  |  |  | 6+ | ${ }^{6+}$ | $6+$ |  |  |  |  |
|  |  |  |  | $7+$ |  |  |  |  |  |


1979. Properties of the Transition Elements -
FIGURE 19.4 Transition metals of the first transition series can form compounds with varying oxidation states.
For the elements scandium through manganese (the first half of the first transition series), the highest oxidation state corresponds to the loss of all of the electrons in both the $s$ and $d$ orbitals of their valence shells. The titanium(IV) ion, for example, is formed when the titanium atom loses its two $3 d$ and two $4 s$ electrons. These highest oxidation states are the most stable forms of scandium, titanium, and vanadium. However, it is not possible to continue to remove all of the valence electrons from metals as we continue through the series. Iron is known to form oxidation states from $2+$ to $6+$, with iron(II) and iron(III) being the most common. Most of the elements of the first transition series form ions with a charge of $2+$ or $3+$ that are stable in water, although those of the early members of the series can be readily oxidized by air.

The elements of the second and third transition series generally are more stable in higher oxidation states than are the elements of the first series. In general, the atomic radius increases down a group, which leads to the ions of the second and third series being larger than are those in the first series. Removing electrons from orbitals that are located farther from the nucleus is easier than removing electrons close to the nucleus. For example, molybdenum and tungsten, members of group 6, are limited mostly to an oxidation state of 6+ in aqueous solution. Chromium, the lightest member of the group, forms stable $\mathrm{Cr}^{3+}$ ions in water and, in the absence of air, less stable $\mathrm{Cr}^{2+}$ ions. The sulfide with the highest oxidation state for chromium is $\mathrm{Cr}_{2} \mathrm{~S}_{3}$, which contains the $\mathrm{Cr}^{3+}$ ion. Molybdenum and tungsten form sulfides in which the metals exhibit oxidation states of $4+$ and 6+.


1981. Activity of the Transition Metals -
Which is the strongest oxidizing agent in acidic solution: dichromate ion, which contains chromium(VI), permanganate ion, which contains manganese(VII), or titanium dioxide, which contains titanium(IV)?


1982. Solution -
First, we need to look up the reduction half reactions (in Appendix L) for each oxide in the specified oxidation state:

$$
\begin{array}{cc}
\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}+14 \mathrm{H}^{+}+6 \mathrm{e}^{-} \longrightarrow 2 \mathrm{Cr}^{3+}+7 \mathrm{H}_{2} \mathrm{O} & +1.33 \mathrm{~V} \\
\mathrm{MnO}_{4}^{-}+8 \mathrm{H}^{+}+5 \mathrm{e}^{-} \longrightarrow \mathrm{Mn}^{2+}+\mathrm{H}_{2} \mathrm{O} & +1.51 \mathrm{~V} \\
\mathrm{TiO}_{2}+4 \mathrm{H}^{+}+2 \mathrm{e}^{-} \longrightarrow \mathrm{Ti}^{2+}+2 \mathrm{H}_{2} \mathrm{O} & -0.50 \mathrm{~V}
\end{array}
$$

A larger reduction potential means that it is easier to reduce the reactant. Permanganate, with the largest reduction potential, is the strongest oxidizer under these conditions. Dichromate is next, followed by titanium dioxide as the weakest oxidizing agent (the hardest to reduce) of this set.


1983. Check Your Learning -
Predict what reaction (if any) will occur between HCl and $\mathrm{Co}(s)$, and between HBr and $\mathrm{Pt}(s)$. You will need to use
the standard reduction potentials from Appendix L.


1984. Answer: -
$\mathrm{Co}(s)+2 \mathrm{HCl} \longrightarrow \mathrm{H}_{2}+\mathrm{CoCl}_{2}(a q) ;$ no reaction because $\mathrm{Pt}(s)$ will not be oxidized by $\mathrm{H}^{+}$


1985. Preparation of the Transition Elements -
Ancient civilizations knew about iron, copper, silver, and gold. The time periods in human history known as the Bronze Age and Iron Age mark the advancements in which societies learned to isolate certain metals and use them to make tools and goods. Naturally occurring ores of copper, silver, and gold can contain high concentrations of these metals in elemental form (Figure 19.5). Iron, on the other hand, occurs on earth almost exclusively in oxidized forms, such as rust $\left(\mathrm{Fe}_{2} \mathrm{O}_{3}\right)$. The earliest known iron implements were made from iron meteorites. Surviving iron artifacts dating from approximately 4000 to 2500 BC are rare, but all known examples contain specific alloys of iron and nickel that occur only in extraterrestrial objects, not on earth. It took thousands of years of technological advances before civilizations developed iron smelting, the ability to extract a pure element from its naturally occurring ores and for iron tools to become common.

(a)

(b)

(c)

FIGURE 19.5 Transition metals occur in nature in various forms. Examples include (a) a nugget of copper, (b) a deposit of gold, and (c) an ore containing oxidized iron. (credit a: modification of work by http://images-of-elements.com/copper-2.jpg; credit c: modification of work by http://images-of-elements.com/iron-ore.jpg)

Generally, the transition elements are extracted from minerals found in a variety of ores. However, the ease of their recovery varies widely, depending on the concentration of the element in the ore, the identity of the other elements present, and the difficulty of reducing the element to the free metal.

In general, it is not difficult to reduce ions of the $d$-block elements to the free element. Carbon is a sufficiently strong reducing agent in most cases. However, like the ions of the more active main group metals, ions of the $f$-block elements must be isolated by electrolysis or by reduction with an active metal such as calcium.

We shall discuss the processes used for the isolation of iron, copper, and silver because these three processes illustrate the principal means of isolating most of the $d$-block metals. In general, each of these processes involves three principal steps: preliminary treatment, smelting, and refining.


1985. Preparation of the Transition Elements -
1. Preliminary treatment. In general, there is an initial treatment of the ores to make them suitable for the extraction of the metals. This usually involves crushing or grinding the ore, concentrating the metalbearing components, and sometimes treating these substances chemically to convert them into compounds that are easier to reduce to the metal.
2. Smelting. The next step is the extraction of the metal in the molten state, a process called smelting, which includes reduction of the metallic compound to the metal. Impurities may be removed by the addition of a compound that forms a slag-a substance with a low melting point that can be readily separated from the molten metal.
3. Refining. The final step in the recovery of a metal is refining the metal. Low boiling metals such as zinc and mercury can be refined by distillation. When fused on an inclined table, low melting metals like tin flow away from higher-melting impurities. Electrolysis is another common method for refining metals.


1986. Isolation of Iron -
The early application of iron to the manufacture of tools and weapons was possible because of the wide distribution of iron ores and the ease with which iron compounds in the ores could be reduced by carbon. For a long time, charcoal was the form of carbon used in the reduction process. The production and use of iron became much more widespread about 1620, when coke was introduced as the reducing agent. Coke is a form of carbon formed by heating coal in the absence of air to remove impurities.

The first step in the metallurgy of iron is usually roasting the ore (heating the ore in air) to remove water, decomposing carbonates into oxides, and converting sulfides into oxides. The oxides are then reduced in a blast furnace that is $80-100$ feet high and about 25 feet in diameter (Figure 19.6) in which the roasted ore, coke, and limestone (impure $\mathrm{CaCO}_{3}$ ) are introduced continuously into the top. Molten iron and slag are withdrawn at the bottom. The entire stock in a furnace may weigh several hundred tons.

FIGURE 19.6 Within a blast furnace, different reactions occur in different temperature zones. Carbon monoxide is generated in the hotter bottom regions and rises upward to reduce the iron oxides to pure iron through a series of reactions that take place in the upper regions.

Near the bottom of a furnace are nozzles through which preheated air is blown into the furnace. As soon as the air enters, the coke in the region of the nozzles is oxidized to carbon dioxide with the liberation of a great deal of heat. The hot carbon dioxide passes upward through the overlying layer of white-hot coke, where it is reduced to carbon monoxide:

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

The carbon monoxide serves as the reducing agent in the upper regions of the furnace. The individual reactions are indicated in Figure 19.6.

The iron oxides are reduced in the upper region of the furnace. In the middle region, limestone (calcium carbonate) decomposes, and the resulting calcium oxide combines with silica and silicates in the ore to form
slag. The slag is mostly calcium silicate and contains most of the commercially unimportant components of the ore:


1986. Isolation of Iron -
$$
\mathrm{CaO}(s)+\mathrm{SiO}_{2}(s) \longrightarrow \mathrm{CaSiO}_{3}(l)
$$

Just below the middle of the furnace, the temperature is high enough to melt both the iron and the slag. They collect in layers at the bottom of the furnace; the less dense slag floats on the iron and protects it from oxidation. Several times a day, the slag and molten iron are withdrawn from the furnace. The iron is transferred to casting machines or to a steelmaking plant (Figure 19.7).

FIGURE 19.7 Molten iron is shown being cast as steel. (credit: Clint Budd)
Much of the iron produced is refined and converted into steel. Steel is made from iron by removing impurities and adding substances such as manganese, chromium, nickel, tungsten, molybdenum, and vanadium to produce alloys with properties that make the material suitable for specific uses. Most steels also contain small but definite percentages of carbon ( $0.04 \%-2.5 \%$ ). However, a large part of the carbon contained in iron must be removed in the manufacture of steel; otherwise, the excess carbon would make the iron brittle.


1987. LINK TO LEARNING -
You can watch an animation of steelmaking (http://openstax.org/l/16steelmaking) that walks you through the process.


1988. Isolation of Copper -
The most important ores of copper contain copper sulfides (such as covellite, CuS), although copper oxides (such as tenorite, CuO ) and copper hydroxycarbonates [such as malachite, $\mathrm{Cu}_{2}(\mathrm{OH})_{2} \mathrm{CO}_{3}$ ] are sometimes found. In the production of copper metal, the concentrated sulfide ore is roasted to remove part of the sulfur as sulfur dioxide. The remaining mixture, which consists of $\mathrm{Cu}_{2} \mathrm{~S}, \mathrm{FeS}, \mathrm{FeO}$, and $\mathrm{SiO}_{2}$, is mixed with limestone, which serves as a flux (a material that aids in the removal of impurities), and heated. Molten slag forms as the iron and silica are removed by Lewis acid-base reactions:

$$
\begin{gathered}
\mathrm{CaCO}_{3}(s)+\mathrm{SiO}_{2}(s) \longrightarrow \mathrm{CaSiO}_{3}(l)+\mathrm{CO}_{2}(\mathrm{~g}) \\
\mathrm{FeO}(s)+\mathrm{SiO}_{2}(s) \longrightarrow \mathrm{FeSiO}_{3}(l)
\end{gathered}
$$

In these reactions, the silicon dioxide behaves as a Lewis acid, which accepts a pair of electrons from the Lewis base (the oxide ion).

Reduction of the $\mathrm{Cu}_{2} \mathrm{~S}$ that remains after smelting is accomplished by blowing air through the molten material.

The air converts part of the $\mathrm{Cu}_{2} \mathrm{~S}$ into $\mathrm{Cu}_{2} \mathrm{O}$. As soon as copper(I) oxide is formed, it is reduced by the remaining copper(I) sulfide to metallic copper:


1988. Isolation of Copper -
$$
\begin{gathered}
2 \mathrm{Cu}_{2} \mathrm{~S}(l)+3 \mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{Cu}_{2} \mathrm{O}(l)+2 \mathrm{SO}_{2}(g) \\
2 \mathrm{Cu}_{2} \mathrm{O}(l)+\mathrm{Cu}_{2} \mathrm{~S}(l) \longrightarrow 6 \mathrm{Cu}(l)+\mathrm{SO}_{2}(g)
\end{gathered}
$$

The copper obtained in this way is called blister copper because of its characteristic appearance, which is due to the air blisters it contains (Figure 19.8). This impure copper is cast into large plates, which are used as anodes in the electrolytic refining of the metal (which is described in the chapter on electrochemistry).

FIGURE 19.8 Blister copper is obtained during the conversion of copper-containing ore into pure copper. (credit: "Tortie tude"/Wikimedia Commons)


1989. Isolation of Silver -
Silver sometimes occurs in large nuggets (Figure 19.9) but more frequently in veins and related deposits. At one time, panning was an effective method of isolating both silver and gold nuggets. Due to their low reactivity, these metals, and a few others, occur in deposits as nuggets. The discovery of platinum was due to Spanish explorers in Central America mistaking platinum nuggets for silver. When the metal is not in the form of nuggets, it often useful to employ a process called hydrometallurgy to separate silver from its ores. Hydrology involves the separation of a metal from a mixture by first converting it into soluble ions and then extracting and reducing them to precipitate the pure metal. In the presence of air, alkali metal cyanides readily form the soluble dicyanoargentate(I) ion, $\left[\mathrm{Ag}(\mathrm{CN})_{2}\right]^{-}$, from silver metal or silver-containing compounds such as $\mathrm{Ag}_{2} \mathrm{~S}$ and AgCl . Representative equations are:

$$
\begin{gathered}
4 \mathrm{Ag}(s)+8 \mathrm{CN}^{-}(a q)+\mathrm{O}_{2}(g)+2 \mathrm{H}_{2} \mathrm{O}(l) \longrightarrow 4\left[\mathrm{Ag}(\mathrm{CN})_{2}\right]^{-}(a q)+4 \mathrm{OH}^{-}(a q) \\
2 \mathrm{Ag}_{2} \mathrm{~S}(s)+8 \mathrm{CN}^{-}(a q)+\mathrm{O}_{2}(g)+2 \mathrm{H}_{2} \mathrm{O}(l) \longrightarrow 4\left[\mathrm{Ag}(\mathrm{CN})_{2}\right]^{-}(a q)+2 \mathrm{~S}(s)+4 \mathrm{OH}^{-}(a q) \\
\mathrm{AgCl}(s)+2 \mathrm{CN}^{-}(a q) \longrightarrow\left[\mathrm{Ag}(\mathrm{CN})_{2}\right]^{-}(a q)+\mathrm{Cl}^{-}(a q)
\end{gathered}
$$


1989. Isolation of Silver -
FIGURE 19.9 Naturally occurring free silver may be found as nuggets (a) or in veins (b). (credit a: modification of work by "Teravolt"/Wikimedia Commons; credit b: modification of work by James St. John)

The silver is precipitated from the cyanide solution by the addition of either zinc or iron(II) ions, which serves as the reducing agent:

$$
2\left[\mathrm{Ag}(\mathrm{CN})_{2}\right]^{-}(a q)+\mathrm{Zn}(s) \longrightarrow 2 \mathrm{Ag}(s)+\left[\mathrm{Zn}(\mathrm{CN})_{4}\right]^{2-}(a q)
$$


1991. Refining Redox -
One of the steps for refining silver involves converting silver into dicyanoargenate(I) ions:

$$
4 \mathrm{Ag}(s)+8 \mathrm{CN}^{-}(a q)+\mathrm{O}_{2}(g)+2 \mathrm{H}_{2} \mathrm{O}(l) \longrightarrow 4\left[\mathrm{Ag}(\mathrm{CN})_{2}\right]^{-}(a q)+4 \mathrm{OH}^{-}(a q)
$$

Explain why oxygen must be present to carry out the reaction. Why does the reaction not occur as:

$$
4 \mathrm{Ag}(s)+8 \mathrm{CN}^{-}(a q) \longrightarrow 4\left[\mathrm{Ag}(\mathrm{CN})_{2}\right]^{-}(a q) ?
$$


1992. Solution -
The charges, as well as the atoms, must balance in reactions. The silver atom is being oxidized from the 0 oxidation state to the $1+$ state. Whenever something loses electrons, something must also gain electrons (be reduced) to balance the equation. Oxygen is a good oxidizing agent for these reactions because it can gain electrons to go from the 0 oxidation state to the $2-$ state.


1993. Check Your Learning -
During the refining of iron, carbon must be present in the blast furnace. Why is carbon necessary to convert iron oxide into iron?


1994. Answer: -
The carbon is converted into CO, which is the reducing agent that accepts electrons so that iron(III) can be reduced to iron(0).


1995. Transition Metal Compounds -
The bonding in the simple compounds of the transition elements ranges from ionic to covalent. In their lower oxidation states, the transition elements form ionic compounds; in their higher oxidation states, they form covalent compounds or polyatomic ions. The variation in oxidation states exhibited by the transition elements gives these compounds a metal-based, oxidation-reduction chemistry. The chemistry of several classes of compounds containing elements of the transition series follows.


1996. Halides -
Anhydrous halides of each of the transition elements can be prepared by the direct reaction of the metal with halogens. For example:

$$
2 \mathrm{Fe}(s)+3 \mathrm{Cl}_{2}(g) \longrightarrow 2 \mathrm{FeCl}_{3}(s)
$$

Heating a metal halide with additional metal can be used to form a halide of the metal with a lower oxidation
state:

$$
\mathrm{Fe}(s)+2 \mathrm{FeCl}_{3}(s) \longrightarrow 3 \mathrm{FeCl}_{2}(s)
$$

The stoichiometry of the metal halide that results from the reaction of the metal with a halogen is determined by the relative amounts of metal and halogen and by the strength of the halogen as an oxidizing agent. Generally, fluorine forms fluoride-containing metals in their highest oxidation states. The other halogens may not form analogous compounds.

In general, the preparation of stable water solutions of the halides of the metals of the first transition series is by the addition of a hydrohalic acid to carbonates, hydroxides, oxides, or other compounds that contain basic anions. Sample reactions are:

$$
\begin{gathered}
\mathrm{NiCO}_{3}(s)+2 \mathrm{HF}(a q) \longrightarrow \mathrm{NiF}_{2}(a q)+\mathrm{H}_{2} \mathrm{O}(l)+\mathrm{CO}_{2}(g) \\
\mathrm{Co}(\mathrm{OH})_{2}(s)+2 \mathrm{HBr}(a q) \longrightarrow \mathrm{CoBr}_{2}(a q)+2 \mathrm{H}_{2} \mathrm{O}(l)
\end{gathered}
$$

Most of the first transition series metals also dissolve in acids, forming a solution of the salt and hydrogen gas. For example:

$$
\mathrm{Cr}(s)+2 \mathrm{HCl}(a q) \longrightarrow \mathrm{CrCl}_{2}(a q)+\mathrm{H}_{2}(g)
$$


1996. Halides -
$$
\mathrm{Cr}(s)+2 \mathrm{HCl}(a q) \longrightarrow \mathrm{CrCl}_{2}(a q)+\mathrm{H}_{2}(g)
$$

The polarity of bonds with transition metals varies based not only upon the electronegativities of the atoms involved but also upon the oxidation state of the transition metal. Remember that bond polarity is a continuous spectrum with electrons being shared evenly (covalent bonds) at one extreme and electrons being transferred completely (ionic bonds) at the other. No bond is ever 100\% ionic, and the degree to which the electrons are evenly distributed determines many properties of the compound. Transition metal halides with low oxidation numbers form more ionic bonds. For example, titanium(II) chloride and titanium(III) chloride ( $\mathrm{TiCl}_{2}$ and $\mathrm{TiCl}_{3}$ ) have high melting points that are characteristic of ionic compounds, but titanium(IV) chloride ( $\mathrm{TiCl}_{4}$ ) is a volatile liquid, consistent with having covalent titanium-chlorine bonds. All halides of the heavier $d$-block elements have significant covalent characteristics.

The covalent behavior of the transition metals with higher oxidation states is exemplified by the reaction of the metal tetrahalides with water. Like covalent silicon tetrachloride, both the titanium and vanadium tetrahalides react with water to give solutions containing the corresponding hydrohalic acids and the metal oxides:

$$
\begin{aligned}
& \mathrm{SiCl}_{4}(l)+2 \mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{SiO}_{2}(s)+4 \mathrm{HCl}(a q) \\
& \mathrm{TiCl}_{4}(l)+2 \mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{TiO}_{2}(s)+4 \mathrm{HCl}(a q)
\end{aligned}
$$


1997. Oxides -
As with the halides, the nature of bonding in oxides of the transition elements is determined by the oxidation state of the metal. Oxides with low oxidation states tend to be more ionic, whereas those with higher oxidation states are more covalent. These variations in bonding are because the electronegativities of the elements are not fixed values. The electronegativity of an element increases with increasing oxidation state. Transition metals in low oxidation states have lower electronegativity values than oxygen; therefore, these metal oxides are ionic. Transition metals in very high oxidation states have electronegativity values close to that of oxygen, which leads to these oxides being covalent.

The oxides of the first transition series can be prepared by heating the metals in air. These oxides are $\mathrm{Sc}_{2} \mathrm{O}_{3}$, $\mathrm{TiO}_{2}, \mathrm{~V}_{2} \mathrm{O}_{5}, \mathrm{Cr}_{2} \mathrm{O}_{3}, \mathrm{Mn}_{3} \mathrm{O}_{4}, \mathrm{Fe}_{3} \mathrm{O}_{4}, \mathrm{Co}_{3} \mathrm{O}_{4}$, NiO , and CuO .

Alternatively, these oxides and other oxides (with the metals in different oxidation states) can be produced by heating the corresponding hydroxides, carbonates, or oxalates in an inert atmosphere. Iron(II) oxide can be prepared by heating iron(II) oxalate, and cobalt(II) oxide is produced by heating cobalt(II) hydroxide:

$$
\begin{gathered}
\mathrm{FeC}_{2} \mathrm{O}_{4}(s) \longrightarrow \mathrm{FeO}(s)+\mathrm{CO}(g)+\mathrm{CO}_{2}(g) \\
\mathrm{Co}(\mathrm{OH})_{2}(s) \longrightarrow \mathrm{CoO}(s)+\mathrm{H}_{2} \mathrm{O}(g)
\end{gathered}
$$


1997. Oxides -
With the exception of $\mathrm{CrO}_{3}$ and $\mathrm{Mn}_{2} \mathrm{O}_{7}$, transition metal oxides are not soluble in water. They can react with acids and, in a few cases, with bases. Overall, oxides of transition metals with the lowest oxidation states are basic (and react with acids), the intermediate ones are amphoteric, and the highest oxidation states are
primarily acidic. Basic metal oxides at a low oxidation state react with aqueous acids to form solutions of salts and water. Examples include the reaction of cobalt(II) oxide accepting protons from nitric acid, and scandium(III) oxide accepting protons from hydrochloric acid:

$$
\begin{gathered}
\mathrm{CoO}(s)+2 \mathrm{HNO}_{3}(a q) \longrightarrow \mathrm{Co}\left(\mathrm{NO}_{3}\right)_{2}(a q)+\mathrm{H}_{2} \mathrm{O}(l) \\
\mathrm{Sc}_{2} \mathrm{O}_{3}(s)+6 \mathrm{HCl}(a q) \longrightarrow 2 \mathrm{ScCl}_{3}(a q)+3 \mathrm{H}_{2} \mathrm{O}(l)
\end{gathered}
$$

The oxides of metals with oxidation states of 4+ are amphoteric, and most are not soluble in either acids or bases. Vanadium(V) oxide, chromium(VI) oxide, and manganese(VII) oxide are acidic. They react with solutions of hydroxides to form salts of the oxyanions $\mathrm{VO}_{4}{ }^{3-}, \mathrm{CrO}_{4}{ }^{2-}$, and $\mathrm{MnO}_{4}{ }^{-}$. For example, the complete ionic equation for the reaction of chromium(VI) oxide with a strong base is given by:

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


1997. Oxides -
Chromium(VI) oxide and manganese(VII) oxide react with water to form the acids $\mathrm{H}_{2} \mathrm{CrO}_{4}$ and $\mathrm{HMnO}_{4}$, respectively.


1998. Hydroxides -
When a soluble hydroxide is added to an aqueous solution of a salt of a transition metal of the first transition series, a gelatinous precipitate forms. For example, adding a solution of sodium hydroxide to a solution of cobalt sulfate produces a gelatinous pink or blue precipitate of cobalt(II) hydroxide. The net ionic equation is:

$$
\mathrm{Co}^{2+}(a q)+2 \mathrm{OH}^{-}(a q) \longrightarrow \mathrm{Co}(\mathrm{OH})_{2}(s)
$$

In this and many other cases, these precipitates are hydroxides containing the transition metal ion, hydroxide ions, and water coordinated to the transition metal. In other cases, the precipitates are hydrated oxides composed of the metal ion, oxide ions, and water of hydration:

$$
4 \mathrm{Fe}^{3+}(a q)+6 \mathrm{OH}^{-}(a q)+\mathrm{nH}_{2} \mathrm{O}(l) \longrightarrow 2 \mathrm{Fe}_{2} \mathrm{O}_{3} \cdot(\mathrm{n}+3) \mathrm{H}_{2} \mathrm{O}(s)
$$

These substances do not contain hydroxide ions. However, both the hydroxides and the hydrated oxides react with acids to form salts and water. When precipitating a metal from solution, it is necessary to avoid an excess of hydroxide ion, as this may lead to complex ion formation as discussed later in this chapter. The precipitated metal hydroxides can be separated for further processing or for waste disposal.


1999. Carbonates -
Many of the elements of the first transition series form insoluble carbonates. It is possible to prepare these carbonates by the addition of a soluble carbonate salt to a solution of a transition metal salt. For example, nickel carbonate can be prepared from solutions of nickel nitrate and sodium carbonate according to the following net ionic equation:

$$
\mathrm{Ni}^{2+}(a q)+\mathrm{CO}_{3}{ }^{2-} \longrightarrow \mathrm{NiCO}_{3}(s)
$$

The reactions of the transition metal carbonates are similar to those of the active metal carbonates. They react with acids to form metals salts, carbon dioxide, and water. Upon heating, they decompose, forming the transition metal oxides.


2000. Other Salts -
In many respects, the chemical behavior of the elements of the first transition series is very similar to that of the main group metals. In particular, the same types of reactions that are used to prepare salts of the main group metals can be used to prepare simple ionic salts of these elements.

A variety of salts can be prepared from metals that are more active than hydrogen by reaction with the corresponding acids: Scandium metal reacts with hydrobromic acid to form a solution of scandium bromide:

$$
2 \mathrm{Sc}(s)+6 \mathrm{HBr}(a q) \longrightarrow 2 \mathrm{ScBr}_{3}(a q)+3 \mathrm{H}_{2}(g)
$$

The common compounds that we have just discussed can also be used to prepare salts. The reactions involved include the reactions of oxides, hydroxides, or carbonates with acids. For example:

$$
\mathrm{Ni}(\mathrm{OH})_{2}(s)+2 \mathrm{H}_{3} \mathrm{O}^{+}(a q)+2 \mathrm{ClO}_{4}^{-}(a q) \longrightarrow \mathrm{Ni}^{2+}(a q)+2 \mathrm{ClO}_{4}^{-}(a q)+4 \mathrm{H}_{2} \mathrm{O}(l)
$$

Substitution reactions involving soluble salts may be used to prepare insoluble salts. For example:

$$
\mathrm{Ba}^{2+}(a q)+2 \mathrm{Cl}^{-}(a q)+2 \mathrm{~K}^{+}(a q)+\mathrm{CrO}_{4}{ }^{2-}(a q) \longrightarrow \mathrm{BaCrO}_{4}(s)+2 \mathrm{~K}^{+}(a q)+2 \mathrm{Cl}^{-}(a q)
$$

In our discussion of oxides in this section, we have seen that reactions of the covalent oxides of the transition elements with hydroxides form salts that contain oxyanions of the transition elements.


2002. High Temperature Superconductors -
A superconductor is a substance that conducts electricity with no resistance. This lack of resistance means that there is no energy loss during the transmission of electricity. This would lead to a significant reduction in the cost of electricity.

Most currently used, commercial superconducting materials, such as NbTi and $\mathrm{Nb}_{3} \mathrm{Sn}$, do not become superconducting until they are cooled below $23 \mathrm{~K}\left(-250^{\circ} \mathrm{C}\right)$. This requires the use of liquid helium, which has a boiling temperature of 4 K and is expensive and difficult to handle. The cost of liquid helium has deterred the widespread application of superconductors.

One of the most exciting scientific discoveries of the 1980s was the characterization of compounds that exhibit superconductivity at temperatures above 90 K . (Compared to liquid helium, 90 K is a high temperature.) Typical among the high-temperature superconducting materials are oxides containing yttrium (or one of several rare earth elements), barium, and copper in a 1:2:3 ratio. The formula of the ionic yttrium compound is $\mathrm{YBa}_{2} \mathrm{Cu}_{3} \mathrm{O}_{7}$.

The new materials become superconducting at temperatures close to 90 K (Figure 19.10), temperatures that can be reached by cooling with liquid nitrogen (boiling temperature of 77 K ). Not only are liquid nitrogencooled materials easier to handle, but the cooling costs are also about 1000 times lower than for liquid helium.

Further advances during the same period included materials that became superconducting at even higher temperatures and with a wider array of materials. The DuPont team led by Uma Chowdry and Arthur Sleight identified Bismouth-Strontium-Copper-Oxides that became superconducting at temperatures as high as 110 K and, importantly, did not contain rare earth elements. Advances continued through the subsequent decades until, in 2020, a team led by Ranga Dias at University of Rochester announced the development of a roomtemperature superconductor, opening doors to widespread applications. More research and development is needed to realize the potential of these materials, but the possibilities are very promising.


2002. High Temperature Superconductors -
FIGURE 19.10 The resistance of the high-temperature superconductor $\mathrm{YBa}_{2} \mathrm{Cu}_{3} \mathrm{O}_{7}$ varies with temperature. Note how the resistance falls to zero below 92 K , when the substance becomes superconducting.

Although the brittle, fragile nature of these materials presently hampers their commercial applications, they have tremendous potential that researchers are hard at work improving their processes to help realize.
Superconducting transmission lines would carry current for hundreds of miles with no loss of power due to resistance in the wires. This could allow generating stations to be located in areas remote from population centers and near the natural resources necessary for power production. The first project demonstrating the viability of high-temperature superconductor power transmission was established in New York in 2008.

Researchers are also working on using this technology to develop other applications, such as smaller and more powerful microchips. In addition, high-temperature superconductors can be used to generate magnetic fields for applications such as medical devices, magnetic levitation trains, and containment fields for nuclear fusion reactors (Figure 19.11).

FIGURE 19.11 (a) This magnetic levitation train (or maglev) uses superconductor technology to move along its tracks. (b) A magnet can be levitated using a dish like this as a superconductor. (credit a: modification of work by Alex Needham; credit b: modification of work by Kevin Jarrett)


2003. LINK TO LEARNING -
Watch how a high-temperature superconductor (http://openstax.org/l/16supercond) levitates around a magnetic racetrack in the video.


2004. LEARNING OBJECTIVES -
By the end of this section, you will be able to:

- List the defining traits of coordination compounds
- Describe the structures of complexes containing monodentate and polydentate ligands
- Use standard nomenclature rules to name coordination compounds
- Explain and provide examples of geometric and optical isomerism
- Identify several natural and technological occurrences of coordination compounds

The hemoglobin in your blood, the chlorophyll in green plants, vitamin B-12, and the catalyst used in the manufacture of polyethylene all contain coordination compounds. Ions of the metals, especially the transition metals, are likely to form complexes. Many of these compounds are highly colored (Figure 19.12). In the remainder of this chapter, we will consider the structure and bonding of these remarkable compounds.

FIGURE 19.12 Metal ions that contain partially filled $d$ subshell usually form colored complex ions; ions with empty $d$ subshell $\left(d^{0}\right)$ or with filled $d$ subshells $\left(d^{10}\right)$ usually form colorless complexes. This figure shows, from left to right, solutions containing $\left[M\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]^{\mathrm{n+}}$ ions with $M=\mathrm{Sc}^{3+}\left(d^{0}\right), \mathrm{Cr}^{3+}\left(d^{3}\right), \mathrm{Co}^{2+}\left(d^{7}\right), \mathrm{Ni}^{2+}\left(d^{8}\right), \mathrm{Cu}^{2+}\left(d^{9}\right)$, and $\mathrm{Zn}^{2+}\left(d^{10}\right)$. (credit: Sahar Atwa)