Metals and Their Compounds

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The two categories of elements, metals and non-metals, are so different that most elements can be easi-ly distinguished as either a metal or a nonmetal.Compounds can also be classified, although there arethree categories rather than two: combinations of metalwith metal, metal with nonmetal, and nonmetal withnonmetal. Each class has a distinctive set of character-istics. The combination of two metals produces an alloythat is difficult to distinguish from a pure metal becauseits characteristics are also metallic. A metal combined

with a nonmetal yields a salt. The majority of suchcombinations are difficult to distinguish visually fromtable salt, but some are colored and nearly all are toxic.The combination of two nonmetals produces a materi-alusually a transparent gasthat is difficult to distin-guish from a pure nonmetallic element.

These categories show regularities yet to beexplained: Why are salts so different from pure ele-ments? Why can most metals be melted together in anyproportion and then be solidified into a homogeneousalloy such that none of either starting metal is left asexcess? Why does a compound composed of a metaland a nonmetal or of a nonmetal and another nonmetalnearly always leave some of one reactant in excess,unchanged at the end of the reaction? This chapter willbegin to answer these questions.

Pure Metals and Alloys

A metal atom has at least one loosely held electronin a large outer orbital. The atom also has room foranother electron. These facts account for the propertiesof metals in the following way. In a solid metal theatoms are close enough together that the orbitals of theloosely held valence electrons try to spread out into thewhole piece of metal and overlap significantly (Fig.20.1). To some degree the orbitals of the valence elec-

trons of the large number of atoms in the piece of metalare trying to occupy the same space inside the piece of metal at the same time. If they could do so completely,the valence electrons would be in identical orbitals andhave identical energies. However, the ExclusionPrinciple forbids more than two electrons occupyingidentical orbitals at the same time. Hence, the orbitalsfind a compromise and the energy levels of the electronssplit into very many levels that are very close together(Fig. 20.2).

The orbitals of these valence electrons spread sig-nificantly beyond the locale of a particular atom so thata given electron does not belong to a particular atom.The valence electrons can be thought of as a verymobile kind of glue that exists between the atoms tohold them together. The valence electrons can easily

20. Metals and Their Compounds

Figure 20.1. Lithium metal orbitals that overlap permit electrons to move freely throughout the entire lump of metal.


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drift from one orbital to another through the overlap-ping regions to any part of the lump of metal.

The drifting electrons give metals their uniqueproperties, the most common of which are discussedbelow.

1. Electrical conductivity . If electrons areforced into one end of a piece of metal, theypush other mobile electrons out the other end.This is the mechanism by which wires carryelectricity and is called electrical conductivity.

2. Metallic luster . The closely spaced energylevels shown in Figure 20.2 permit metals toabsorb virtually every visible wavelength of light and many invisible wavelengths. A pho-ton of almost any energy can find an energygap that just fits. However, the atoms do not

keep the photons, but immediately radiatethem back (reflection). If all of the visiblelight is reflected, the metal appears silver, or if it is roughened, it appears white. If some of the violet and blue light is permanentlyabsorbed, the metal appears gold or coppercolored. High reflectivity in at least a part of the visible spectrum confers metallic luster.

3. Malleability . Things like glass shatter when

hit with a hammer, but most metals simplybend or flatten, particularly if they are hot. Forexample, gold can be pounded thinner thanpaper (to about a millionth of a centimeterthickness). This characteristic response toforce is termed malleability. On the micro-scopic level, the metal nuclei must slide pasteach other into new positions without eversubstantially weakening the metallic bond.The sea of mobile electrons flows wherever itis needed to keep the bonds strong duringbending and flattening operations (Fig. 20.3).

4. Thermal conductivity . If thermal energy isadded to a piece of metal, the kinetic energy of the electrons is increased. The electrons moverapidly through the metal distributing thiskinetic energy throughout by collisions. Thus,

the whole piece of metal becomes hot muchsooner than an equivalent piece of nonmetal.This transport of thermal energy is termedthermal conductivity.

5. Chemical reactivity . The mobile electronsare easy prey for electron-hungry atoms suchas oxygen, so most metals react easily to formoxides or other compounds with nonmetals. Infact, only a few metals (such as gold, silver,

Figure 20.2. Lithium orbital energies split and form many energy levels that are very close together as more and moreatoms are added.

Figure 20.3. Malleability of a metal. (a) Forces applied to a group of metal atoms. (b) The solid deforms, but willspring back to its original shape if the force is removed. (c) The solid is deformed permanently by layers sliding to newpositions.


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and copper) are found uncombined in nature.The others have reacted with oxygen or othernonmetals and are found as ores.

6. Alloy formation . In metals, the sea of elec-trons accommodates one type of metal atomabout as well as another, particularly if theirsizes and orbital energy levels are similar. Thismeans that metals near one another (especiallyvertically) in the Periodic Table should be ableto substitute for one another in solid metals.Many pairs of metals permit substitution forone another in all proportions to form alloyssimply by melting the two together. Someproperties of the resulting alloy, such as densi-ty, are intermediate between those of the twocomponents. Other properties can turn outquite different than one might expect from theoriginal components. Some metals, that are farapart in the Periodic Table, will permit onlylimited substitution.

Science attempts to explain these characteristicswith one simple concept: Mobile electrons are the basisof a metals electrical conductivity, metallic luster, mal-leability, thermal conductivity, and chemical reactivity.

Oxidation States

According to Figure 17.7, metals have one or moreelectrons that are easily removed because they are nearthe top of the energy well. These elements are general-ly to the left in the Periodic Table. In contrast, non-metals have vacancies for electrons deep in the well.Metals often give one or more electrons to a nonmetal,thus forming a positive metal ion and a negative non-metal ion. In forming the ions, the orbitals of thevalence electrons transfer from the metal atom to centeron the nonmetal atom. These oppositely charged ionsare then electrically attracted and become attached toone another through what is known as an ionic bond .

The oxidation state of an atom in a compound is aquantitative indication of the number of electrons thatthe atom has lost or gained in forming the ion or chem-ical bond. An atom that has lost an electron is said to bein the +1 oxidation state; an example is Li 1+. An atomthat has lost two electrons is in the +2 oxidation state;

an example is Be2+

. Similarly, there can be negativeoxidation states for atoms that acquire extra electrons;examples are 1 for Cl 1 and 2 for O 2.

The important thing to note is that stable chemicalbonds will form if doing so lowers the overall energies

Figure 20.4. Common oxidation states of the elements.


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of the valence electrons by putting them lower in ener-gy wells. The filled shells of the noble gases (helium,neon, argon, etc.) are particularly low energy arrange-ments of the electrons. These elements do not reactwith other elements because the reaction does not pro-vide a means to an overall lowering of the electron ener-gies. The effective oxidation state of the noble gases isidentically zero. Many ions achieve this noble gaselectron configuration, but of course no change hasbeen made to the nucleus.

In general, metallic elements near the left of thePeriodic Table assume positive oxidation states inchemical reactions. Nonmetallic elements near the rightside of the Periodic Table assume negative oxidationstates. Near the middle of the Periodic Table things aresomewhat ambiguous. Will an element like carbon(which has four valence electrons) lose its four valenceelectrons (like a metal), or gain four to fill its shell (likea nonmetal)? The answer depends on the specific com-pound that is being formed. An element in the middle,such as carbon, will do whichever best lowers its elec-tron energy. Thus, for elements near the middle of thechart there is a tendency for atoms to assume differentoxidation states in different chemical reactions, depend-ing on the particular compound which is being formed.

Although many elements can have more than oneoxidation state (depending on the compound), the prim-ary (likely) oxidation state is related to the atoms posi-tion in the Periodic Table. All of the elements in column1 (labeled IA in the Periodic Table) will have an oxida-tion state of +1 in ionic compounds. Elements in IIAhave +2 oxidation states. What would you predict forcolumn IIIA? Check your answer in Figure 20.4.

Elements in columns VIA and VIIA of the PeriodicTable tend to gain electrons to fill their valence shellsand can be predicted with confidence to have primaryoxidation states of 2 and 1, respectively. Elements inthese columns sometimes have additional possible oxi-dation states in compounds. See sulfur and chlorine inFigure 20.4, for example.

Elements in the B columns of the Periodic Table arealso prone to having more than one possible oxidationstate. However, we can use the additional informationthat these elements are all metals (and, therefore, tend tobe electron donors) to predict that at least in some com-pounds these elements will have positive oxidation

states that correspond directly to the number of valenceelectrons indicated by their column number. For exam-ple, chromium ( 24Cr) in column VIB will in at leastsome cases have an oxidation state of +6. The rule iseven less precise for the elements in the centermostcolumns of the Periodic Table (labeled VIII [neither Anor B]). Here we can only predict that, as metals, theelements will have positive oxidation states, but wemust rely on experimental analysis of the individualcompounds to determine the precise oxidation states of

the constituent elements.Why do the regularities exist among the elements?

They are based on ionization energies whose precise val-ues can be traced back to the Wave Model and theExclusion Principle. If only one electron in an atom hasa low ionization energy, the oxidation state will be +1. If two electrons have low ionization energies, the oxidationstate will be +2, and so on. If there are no electrons withlow ionization energies, but there is a vacancy in anorbital deep in the well, the atom will try to fill the vacan-cy with an electron. This results in a 1 oxidation state.

Compounds Between Metals and Nonmetals

Imagine a metallic sodium ion, Na +, and a nonmetal-lic chloride ion, Cl , that could form an ionic bond. Butwhy should the pairing stop there? Could the Na + attractanother negative ion and the Cl attract another positiveion? Yes. Long chains of alternating Na + and Cl ionsform a sheet like a checkerboard (Fig. 20.5), and finallythe checkerboards stack so that positive ions are alwaysabove negative ions and vice versa. The result is a crys-tal of salt based on the electrical attraction of the ions.

Figure 20.5. Drawing of a small part of a single layerof a NaCl (sodium chloride) crystal.

A salt is much different than a metal. Ionic bond-ing accounts for the characteristics of salts discussedbelow.

1. Electrically nonconducting . The electrons in

both ions of a salt are all in stable, closedshells, and are difficult to remove from theirrespective ions. Because there are no mobileelectrons, solid salts have no way to conductelectricity. However, if the salt is melted,whole ions may move about and conduct elec-tricity, but they are so big and awkward theyare not nearly as conductive as the electrons ina metal. Also, if the salt is dissolved in a liq-uid like water, the ions separate and are able to


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conduct electricity, but in an inefficient man-ner. Materials that conduct electricity whendissolved in water are called electrolytes .

2. Transparent . The energy levels of most saltsare spaced widely apart, so most salts willabsorbs only photons with high energies. Suchhigh-energy photons are in the invisible, ultra-violet part of the spectrum. Visible light is notabsorbed by most saltsit passes through andthe salts appear transparent. However, most of the salts of metals in columns IIIB through IIBabsorb some visible light weakly; this givesthem pale colors. There are also a few highlycolored salts like potassium dichromate; mostof them are composed of more than two ele-ments (e.g., K 2Cr 2O7). If a transparent saltcrystal is ground to powder, the small crystalsbend a light ray many times before it emergesfrom the powder. This means light will nolonger pass straight through as it did in a singlecrystal, so light coming in from the side maybend and come out the front. The net result isthat the powder appears white, even though itis basically transparent or slightly colored. Ablock of salt with many cracks or bubbles in itwill also appear white. Many otherwise trans-parent minerals in nature appear opaque andwhite for this reason.

3. Brittle . Salt crystals break and shatter whenhit with a hammerthey do not bend or flat-ten. Imagine sliding one layer of a NaCl crys-

tal over the underlying layer (Fig. 20.6). Whenone sliding layer moves a short distance, all of the positive ions will be directly over otherpositive ions, and likewise for the negativeions. The repulsion is more than enough toforce the two layers apart, which breaks(cleaves) the crystal. Brittleness, however,does not always dominate. Glass, most rocks,and objects made of fired clay can be con-sidered super-cooled liquids, because undermild stress, only a few atoms at a time moveand no breakage occurs. Over a long timemany atoms will have moved, and the object

will have a different shape. Thick rock layersbend, fold, and take on some surprisingly con-torted shapes if they are not pushed too rapid-ly. If they cannot adapt quickly enough, theysnap and cause earthquakes.

Formulas and Names of Salts

The charge in a crystal must be balanced (sametotal positive as negative charge). Even a relatively

small excess charge would cause a spark to jump fromthe crystal. This implies that there are nearly the samenumber of Na + as Cl ions in a crystal of table salt. Wecan use subscripts to indicate how many ions there arein each crystal. For example, a small crystal containing

70 of each type of ion would be designated Na 70Cl70.Because the proportion is always one-to-one, it is cus-tomary to write the simplest formula, NaCl. Other one-to-one compounds are LiF and CaO. Note that themetal is written first, the nonmetal last.

Now consider Mg 2+ ions coming together with Cl

ions to form a crystal. Two Cl ions are required to neu-tralize the +2 charge on the magnesium ion. The pro-portion is now one magnesium to two chlorides, and theformula is MgCl 2. By knowing the oxidation numbers,we can determine the proper proportions and the propersubscripts in the formulas of compounds composed of ametal and a nonmetal. The pivotal concept is the can-

cellation of charges.You should be cautioned that this system works

best for the combination of a metal with a nonmetal. Itpredicts just one compoundthere may be others. If the oxidation states are known, the rule works for non-metals also. For example, nitrogen in the unusual oxi-dation state of +4 can combine with oxygen in its 2oxidation state. The resulting compound of two non-metals has the formula NO 2. However, it is technicallyincorrect to discuss such compounds in terms of oxida-

Figure 20.6. Force applied to layers of ions in a saltcrystal will slide ions with the same charge over oneanother and split the crystal.


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conduct electricity.5. Ionic Bond: The chemical bond that binds a

metallic ion to a nonmetallic ion by electricalattraction.

6. Ions: A charged object formed when an atom ormolecule loses or gains electrons.

7. Malleability: The characteristic of substances thatallows them to be worked into desirable shapes ordrawn out into wires. Metals are malleable.

8. Metal: See Chapter 18.9. Metallic Bond: The chemical bond that binds

metal atoms to other metal atoms in forming metalsubstances.

10. Metallic Luster: The shiny appearance of metals.11. Negative Oxidation State: The state of an atom

which has gained one or more electrons to form anionic bond. The precise oxidation state is repre-sented by a sign followed by the number of elec-trons gained.

12. Nonmetal: See Chapter 18.13. Oxidation Number: A signed number that indi-

cates how many electrons an atom loses (positivenumber) or gains (negative number) by forming acompound. In ionic bonding the transfer is com-plete and distinct ions are formed. In covalentbonding the transfer is only partial and distinct ionsare not formed.

14. Positive Oxidation State: The state of an atomwhich has lost one or more electrons to form anionic bond. The precise oxidation state is repre-sented by a + sign followed by the number of elec-trons lost.

15. Salt: An substance formed from the ionic bond of

a metal with a nonmetal. NaCl is a salt.16. Thermal Conductivity: A measure of the degreeto which a substance conducts heat. Metals have ahigh thermal conductivity.

17. Transparency: A characteristic of ionic sub-stances, such as salts, that readily transmit light.Opposite to opaqueness.

D. FOCUS QUESTIONS1. Consider atoms held together in metallic bonds:

a. What happens to the valence electrons whenthe bonds are formed?b. What happens to the energy of the system

when the bonds are formed?c. What happens to the orbitals when the bondsare formed?d. How does this account for the luster of metals?e. What is the Wave Model of the atom?f. State the fundamental principle of wave-parti-cle duality that the Wave Model is based on.

2. Consider the following elements held together inionic bonds: (a) sodium, chlorine; (b) magnesium,chlorine. In each case:

a. Determine the positive or negative oxidationstates of the different atoms.b. Write the chemical formula for the compound.c. Write and balance the chemical equation forthe reaction.d. Sketch an energy well for each kind of atom inthe compound. Draw a circle around each electronsoon to be lost in one atom, and an empty circle atthe location soon to be filled in the other atom.e. List the main parts of the Wave Model of theatom.f. State the fundamental principle of wave-parti-cle duality that the Wave Model is based on.

3. Sketch a diagram showing a possible arrangementof the ions in a salt. Explain what would happen tothis arrangement if shear forces were exerted on thesalt. Why is the salt brittle? Also, using your under-standing of ionic bonds, explain why table salt dis-solved in water is an ionic conductor. Name andstate the fundamental principle that explains the forces that are involved.

E. EXERCISES20.1. When gold and silver are mixed, the appear-

ance is still that of pure gold. How could you distin-guish the alloy from pure gold?

20.2 Which of the following pairs of metals is mostlikely to form alloys of all compositions?

(a) 56Ba and 31Ga(b) 50Sn and 82Pb(c) 3Li and 83Bi?

20.3. Define the following terms: (a) electricalconductivity, (b) metallic luster, (c) malleability, (d)thermal conductivity.

20.4. Which of the following pairs of metals ismost likely to form alloys of all proportions?

(a) 55Cs and 49In(b) 78Pt and 79Au(c) 82Pb and 19K

20.5. In the left half of Figure 20.4 notice theprominent diagonal lines of oxidation states. How far isit vertically from one line to the next? How far hori-zontally?

20.6. How many of the 23 elements listed in theleft half of Figure 20.4 have only a single oxidationstate?

20.7. Write down the primary oxidation state of thefollowing elements on the basis of their positions in thePeriodic Table:


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(a) 19K(b) 10Ne(c) 29Cu(d) 35Br(e) 15P

20.8. Write down the primary oxidation states of

the following elements and check your answers inFigure 20.4:(a) 5B(b) 20Ca(c) 17Cl(d) 1H(e) 18Ar(f) 28Ni(g) 24Cr

20.9. Referring to Figure 18.3, select two atomsthat are likely to have negative oxidation states becausethey have very high ionization energies. Avoid thenoble gases 2He, 10Ne, and 18Ar. Now look these ele-ments up in Figure 17.7 to see if there is a vacancy foran electron in a low-lying orbital.

20.10. Referring to Figure 18.3, select two atomsthat are likely to have positive oxidation states becausethey have low ionization energies. Now look these ele-ments up in Figure 17.7 to see if the top electrons arenear the top of the well.

20.11. Use Figure 20.4 to write the common oxi-dation states of the following elements (e.g., He: 0; and

C: 2,0,2, and 4):(a) N(b) P(c) O(d) S(e) Ne(f) Ar

20.12. From Figure 17.6, state which electron Li islikely to lose to become Li 1+. Do the same for Be goingto Be 2+ (two electrons).

20.13. How many I

ions will be required to neu-tralize the charge on an Al 3+ ion?

20.14 How many Al 3+ and O 2 ions must be com-bined to achieve charge neutralization? (Hint: It willrequire more than one of each type of ion.)

20.15. Write the chemical formulas for the follow-ing combinations:

(a) K + and F

(b) Na + and S 2

(c) Mg 2+ and N 3

20.16. Determine the principal oxidation states of the following elements and then write the correct chem-ical formulas for the combinations:

(a) 4Be and 35Br(b) 31Ga and 8O

20.17. Write the names of the following com-pounds:

(a) BeO(b) CaCl 2(c) LiF(d) Na 2S

20.18. Write the names of the following com-pounds:

(a) CrO(b) CrO 3(c) Cr 2O3(d) NO 2(e) SO 3

20.19. Complete Table 20.1.

20.20. Which of the following is an ionic com-pound?

(a) O 2(b) NH 3(c) CO 2(d) CH 4(e) MgF 2

20.21. The correct formula for potassium sulfideis:

(a) KS(b) K 2S(c) KS 2(d) K 3S(e) KS 3

20.22. Determine the primary oxidation state forRb. If two Rb atoms combine with one S atom, what isthe oxidation state of S?

(a) +2(b) +1(c) 0(d) 1(e) 2


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Table 20.1.

Sodium fluoride solid salt

Iron-cobalt alloy metallic luster

CuCl 2

Calcium oxide

Calcium chloride CaCl 2

Silver-gold alloy indefinite

Chromium (III) oxide

FeO

Magnesium bromide

Na

CONDUCTIVITYFORMULALUSTER OR

COLOR FORMNAME


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Metals and Their Compounds

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