Chapter 20: Metals and Their Compounds

20. Metals and Their Compounds
The two categories of elements, metals and nonmetals,
are so different that most elements can be easily
distinguished as either a metal or a nonmetal.
Compounds can also be classified, although there are
three categories rather than two: combinations of metal
with metal, metal with nonmetal, and nonmetal with
nonmetal. Each class has a distinctive set of characteristics.
The combination of two metals produces an alloy
that is difficult to distinguish from a pure metal because
its characteristics are also metallic. A metal combined
with a nonmetal yields a salt. The majority of such
combinations are difficult to distinguish visually from
table salt, but some are colored and nearly all are toxic.
The combination of two nonmetals produces a material—usually
a transparent gas—that is difficult to distinguish
from a pure nonmetallic element.
These categories show regularities yet to be
explained: Why are salts so different from pure elements?
Why can most metals be melted together in any
proportion and then be solidified into a homogeneous
alloy such that none of either starting metal is left as
excess? Why does a compound composed of a metal
and a nonmetal or of a nonmetal and another nonmetal
nearly always leave some of one reactant in excess,
unchanged at the end of the reaction? This chapter will
begin to answer these questions.
Pure Metals and Alloys
A metal atom has at least one loosely held electron
in a large outer orbital. The atom also has room for
another electron. These facts account for the properties
of metals in the following way. In a solid metal the
atoms are close enough together that the orbitals of the
loosely held valence electrons try to spread out into the
whole piece of metal and overlap significantly (Fig.
20.1). To some degree the orbitals of the valence electrons
of the large number of atoms in the piece of metal
are 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 and
have identical energies. However, the Exclusion
Principle forbids more than two electrons occupying
identical orbitals at the same time. Hence, the orbitals
find a compromise and the energy levels of the electrons
split into very many levels that are very close together
(Fig. 20.2).
The orbitals of these valence electrons spread significantly
beyond the locale of a particular atom so that
a given electron does not “belong” to a particular atom.
The valence electrons can be thought of as a very
mobile kind of “glue” that exists between the atoms to
hold them together. The valence electrons can easily
Figure 20.1. Lithium metal orbitals that overlap permit electrons to move freely throughout the entire lump of metal.
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Figure 20.2. Lithium orbital energies split and form many energy levels that are very close together as more and more
atoms are added.
Figure 20.3. Malleability of a metal. (a) Forces applied to a group of metal atoms. (b) The solid deforms, but will
spring back to its original shape if the force is removed. (c) The solid is deformed permanently by layers sliding to new
positions.
drift from one orbital to another through the overlapping
regions to any part of the lump of metal.
The drifting electrons give metals their unique
properties, the most common of which are discussed
below.
1. Electrical conductivity. If electrons are
forced into one end of a piece of metal, they
push other mobile electrons out the other end.
This is the mechanism by which wires carry
electricity and is called electrical conductivity.
2. Metallic luster. The closely spaced energy
levels shown in Figure 20.2 permit metals to
absorb virtually every visible wavelength of
light and many invisible wavelengths. A photon
of almost any energy can find an energy
gap that just fits. However, the atoms do not
keep the photons, but immediately radiate
them back (reflection). If all of the visible
light is reflected, the metal appears silver, or if
it is roughened, it appears white. If some of
the violet and blue light is permanently
absorbed, the metal appears gold or copper
colored. 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 simply
bend or flatten, particularly if they are hot. For
example, gold can be pounded thinner than
paper (to about a millionth of a centimeter
thickness). This characteristic response to
force is termed malleability. On the microscopic
level, the metal nuclei must slide past
each other into new positions without ever
substantially weakening the metallic bond.
The sea of mobile electrons flows wherever it
is needed to keep the bonds strong during
bending and flattening operations (Fig. 20.3).
4. Thermal conductivity. If thermal energy is
added to a piece of metal, the kinetic energy of
the electrons is increased. The electrons move
rapidly through the metal distributing this
kinetic energy throughout by collisions. Thus,
the whole piece of metal becomes hot much
sooner than an equivalent piece of nonmetal.
This transport of thermal energy is termed
thermal conductivity.
5. Chemical reactivity. The mobile electrons
are easy prey for electron-hungry atoms such
as oxygen, so most metals react easily to form
oxides or other compounds with nonmetals. In
fact, only a few metals (such as gold, silver,
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and copper) are found uncombined in nature.
The others have reacted with oxygen or other
nonmetals and are found as ores.
6. Alloy formation. In metals, the sea of electrons
accommodates one type of metal atom
about as well as another, particularly if their
sizes and orbital energy levels are similar. This
means that metals near one another (especially
vertically) in the Periodic Table should be able
to substitute for one another in solid metals.
Many pairs of metals permit substitution for
one another in all proportions to form alloys
simply by melting the two together. Some
properties of the resulting alloy, such as density,
are intermediate between those of the two
components. Other properties can turn out
quite different than one might expect from the
original components. Some metals, that are far
apart in the Periodic Table, will permit only
limited substitution.
Science attempts to explain these characteristics
with one simple concept: Mobile electrons are the basis
of a metal’s electrical conductivity, metallic luster, malleability,
thermal conductivity, and chemical reactivity.
Oxidation States
According to Figure 17.7, metals have one or more
electrons that are easily removed because they are near
the top of the energy well. These elements are generally
to the left in the Periodic Table. In contrast, nonmetals
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 nonmetal
ion. In forming the ions, the orbitals of the
valence electrons transfer from the metal atom to center
on the nonmetal atom. These oppositely charged ions
are then electrically attracted and become attached to
one another through what is known as an ionic bond.
The oxidation state of an atom in a compound is a
quantitative indication of the number of electrons that
the atom has lost or gained in forming the ion or chemical
bond. An atom that has lost an electron is said to be
in the +1 oxidation state; an example is Li 1+ . An atom
that has lost two electrons is in the +2 oxidation state;
an example is Be 2+ . Similarly, there can be negative
oxidation 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 chemical
bonds 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 energy
wells. The filled shells of the noble gases (helium,
neon, argon, etc.) are particularly low energy arrangements
of the electrons. These elements do not react
with other elements because the reaction does not provide
a means to an overall lowering of the electron energies.
The effective oxidation state of the noble gases is
identically zero. Many ions achieve this “noble gas
electron configuration,” but of course no change has
been made to the nucleus.
In general, metallic elements near the left of the
Periodic Table assume positive oxidation states in
chemical reactions. Nonmetallic elements near the right
side of the Periodic Table assume negative oxidation
states. Near the middle of the Periodic Table things are
somewhat ambiguous. Will an element like carbon
(which has four valence electrons) lose its four valence
electrons (like a metal), or gain four to fill its shell (like
a nonmetal)? The answer depends on the specific compound
that is being formed. An element in the middle,
such as carbon, will do whichever best lowers its electron
energy. Thus, for elements near the middle of the
chart there is a tendency for atoms to assume different
oxidation states in different chemical reactions, depending
on the particular compound which is being formed.
Although many elements can have more than one
oxidation state (depending on the compound), the primary
(likely) oxidation state is related to the atom’s position
in the Periodic Table. All of the elements in column
1 (labeled IA in the Periodic Table) will have an oxidation
state of +1 in ionic compounds. Elements in IIA
have +2 oxidation states. What would you predict for
column IIIA? Check your answer in Figure 20.4.
Elements in columns VIA and VIIA of the Periodic
Table tend to gain electrons to fill their valence shells
and can be predicted with confidence to have primary
oxidation states of –2 and –1, respectively. Elements in
these columns sometimes have additional possible oxidation
states in compounds. See sulfur and chlorine in
Figure 20.4, for example.
Elements in the B columns of the Periodic Table are
also prone to having more than one possible oxidation
state. However, we can use the additional information
that these elements are all metals (and, therefore, tend to
be electron donors) to predict that at least in some compounds
these elements will have positive oxidation
states that correspond directly to the number of valence
electrons indicated by their column number. For example,
chromium ( 24 Cr) in column VIB will in at least
some cases have an oxidation state of +6. The rule is
even less precise for the elements in the centermost
columns of the Periodic Table (labeled VIII [neither A
nor B]). Here we can only predict that, as metals, the
elements will have positive oxidation states, but we
must rely on experimental analysis of the individual
compounds 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 values
can be traced back to the Wave Model and the
Exclusion Principle. If only one electron in an atom has
a low ionization energy, the oxidation state will be +1. If
two electrons have low ionization energies, the oxidation
state will be +2, and so on. If there are no electrons with
low ionization energies, but there is a vacancy in an
orbital deep in the well, the atom will try to fill the vacancy
with an electron. This results in a –1 oxidation state.
Compounds Between Metals and Nonmetals
Imagine a metallic sodium ion, Na + , and a nonmetallic
chloride ion, Cl – , that could form an ionic bond. But
why should the pairing stop there? Could the Na + attract
another negative ion and the Cl – attract another positive
ion? Yes. Long chains of alternating Na + and Cl – ions
form a sheet like a checkerboard (Fig. 20.5), and finally
the checkerboards stack so that positive ions are always
above negative ions and vice versa. The result is a crystal
of salt based on the electrical attraction of the ions.
Figure 20.5. Drawing of a small part of a single layer
of a NaCl (sodium chloride) crystal.
A salt is much different than a metal. Ionic bonding
accounts for the characteristics of salts discussed
below.
1. Electrically nonconducting. The electrons in
both ions of a salt are all in stable, closed
shells, and are difficult to remove from their
respective ions. Because there are no mobile
electrons, solid salts have no way to conduct
electricity. However, if the salt is melted,
whole ions may move about and conduct electricity,
but they are so big and awkward they
are not nearly as conductive as the electrons in
a metal. Also, if the salt is dissolved in a liquid
like water, the ions separate and are able to
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conduct electricity, but in an inefficient manner.
Materials that conduct electricity when
dissolved in water are called electrolytes.
2. Transparent. The energy levels of most salts
are spaced widely apart, so most salts will
absorbs only photons with high energies. Such
high-energy photons are in the invisible, ultraviolet
part of the spectrum. Visible light is not
absorbed by most salts—it passes through and
the salts appear transparent. However, most of
the salts of metals in columns IIIB through IIB
absorb some visible light weakly; this gives
them pale colors. There are also a few highly
colored salts like potassium dichromate; most
of them are composed of more than two elements
(e.g., K 2 Cr 2 O 7 ). If a transparent salt
crystal is ground to powder, the small crystals
bend a light ray many times before it emerges
from the powder. This means light will no
longer pass straight through as it did in a single
crystal, so light coming in from the side may
bend and come out the front. The net result is
that the powder appears white, even though it
is basically transparent or slightly colored. A
block of salt with many cracks or bubbles in it
will also appear white. Many otherwise transparent
minerals in nature appear opaque and
white for this reason.
3. Brittle. Salt crystals break and shatter when
hit with a hammer—they do not bend or flatten.
Imagine sliding one layer of a NaCl crystal
over the underlying layer (Fig. 20.6). When
one sliding layer moves a short distance, all of
the positive ions will be directly over other
positive ions, and likewise for the negative
ions. The repulsion is more than enough to
force 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 considered
super-cooled liquids, because under
mild stress, only a few atoms at a time move
and no breakage occurs. Over a long time
many atoms will have moved, and the object
will have a different shape. Thick rock layers
bend, fold, and take on some surprisingly contorted
shapes if they are not pushed too rapidly.
If they cannot adapt quickly enough, they
snap and cause earthquakes.
Formulas and Names of Salts
The charge in a crystal must be balanced (same
total positive as negative charge). Even a relatively
Figure 20.6. Force applied to layers of ions in a salt
crystal will slide ions with the same charge over one
another and split the crystal.
small excess charge would cause a spark to jump from
the crystal. This implies that there are nearly the same
number of Na + as Cl – ions in a crystal of table salt. We
can use subscripts to indicate how many ions there are
in each crystal. For example, a small crystal containing
70 of each type of ion would be designated Na 70 Cl 70 .
Because the proportion is always one-to-one, it is customary
to write the simplest formula, NaCl. Other oneto-one
compounds are LiF and CaO. Note that the
metal 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 neutralize
the +2 charge on the magnesium ion. The proportion
is now one magnesium to two chlorides, and the
formula is MgCl 2 . By knowing the oxidation numbers,
we can determine the proper proportions and the proper
subscripts in the formulas of compounds composed of a
metal and a nonmetal. The pivotal concept is the cancellation
of charges.
You should be cautioned that this system works
best for the combination of a metal with a nonmetal. It
predicts just one compound—there may be others. If
the oxidation states are known, the rule works for nonmetals
also. For example, nitrogen in the unusual oxidation
state of +4 can combine with oxygen in its –2
oxidation state. The resulting compound of two nonmetals
has the formula NO 2 . However, it is technically
incorrect to discuss such compounds in terms of oxida-
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tion “states” that imply that one of the atoms has entirely
lost one or more electrons and the other atom has
taken full possession of one or more electrons. Ions will
not be formed completely by elements near one another
in the Periodic Table. The smaller nonmetal atom will
pull harder on the electrons than the larger metal atom
does, but the electrons will not leave entirely.
Nevertheless, the pretense that they leave helps us write
down the correct formula, so we keep the concept but
change the name to oxidation number to remind ourselves
that we may not be dealing with completely
formed ions.
In naming these salts, the metal is named first followed
by the nonmetal with its last syllable changed to
-ide: NaCl is sodium chloride, Li 2 O is lithium oxide,
and Al 2 S 3 is aluminum sulfide.
The elements chosen for the preceding examples
usually can have only one oxidation number. Many
other elements have at least two common oxidation
numbers; the name must indicate which is the oxidation
number of interest. One way to do this is to use Roman
numerals in parentheses to indicate the oxidation state
of the metal. For example, FeO would be named
iron(II) oxide, and Fe 2 O 3 would be named iron(III)
oxide.
When two nonmetals are combined, it is not always
clear from the oxidation numbers how many atoms are
in the molecule, so the number of atoms of each element
in the molecule is mentioned using the following prefixes:
mono- for one, di- for two, tri- for three, and
tetra- for four. (Other prefixes are used for higher numbers.)
Some examples are CO, carbon monoxide; CO 2 ,
carbon dioxide; and N 2 O 3 , dinitrogen trioxide.
Summary
Elements can be divided into two categories: metals
and nonmetals. In turn we can divide the types of
chemical bonds into three categories: metals with metals
(metallic bond), metals with nonmetals (ionic bond),
and nonmetals with nonmetals (covalent bond). A distinctive
class of physical characteristics is associated
with each type of bond.
In a solid metal the atoms are so close together that
the orbits of the valence electrons overlap. Because the
outer electrons are so loosely held, they can easily drift
from one orbital to another through the overlapping
regions to any part of the lump of metal. The freedom
of the valence electrons to drift results in the following
properties for metals bound by the metallic bond: good
electrical conductivity, metallic luster, malleability,
good thermal conductivity, high chemical reactivity,
and readiness to form alloys.
It is energetically favorable in metal-nonmetal
reactions for the metal to give up valence electrons to
the nonmetal, thus forming electrically charged ions.
Electrically charged ions with opposite charges are then
attracted to one another and become attached to one
another in an ionic bond. Atoms that have lost electrons
in ionic bonds are said to be in positive oxidation
states. Atoms that have gained electrons in ionic bonds
are said to be in negative oxidation states. The oxidation
states are further characterized by the number of
electrons gained or lost: –1 (for atoms that have gained
one electron), +2 (for atoms that have lost two electrons),
etc.
Ionic compounds of a metal and a nonmetal are
called salts. Salts are generally crystalline solids, transparent,
brittle, white in color (although occasionally
colored), electrically nonconducting as a solid, but conducting
in water solutions or in melted form. These
characteristics can be easily understood in terms of the
nature of the ionic bond.
Bonding between nonmetals (covalent bond) will
be described in the next chapter.
STUDY GUIDE
Chapter 20: Metals and Their Compounds
A. FUNDAMENTAL PRINCIPLES: No new fundamental
principles.
B MODELS, IDEAS, QUESTIONS, OR APPLICA-
TIONS
1. What useful groups are compounds often divided
into?
2. Under what conditions do atoms form metallic
bonds?
3. What are the properties of compounds held in
metallic bonds and why do these compounds have
such properties?
4. Under what conditions do atoms form ionic bonds?
5. What are the simplest rules needed to determine the
primary oxidation states of different atoms?
6. What are the properties of compounds held in ionic
bonds, and why do these compounds have such
properties?
7. What procedure leads to the correct chemical formula
for reactants formed in reactions involving
compounds held together in ionic bonds?
C. GLOSSARY
1. Alloys: Alloys are mixtures of metals. Metals
readily form alloys within certain limits.
2. Brittleness: A characteristic of ionic substances,
such as salts, that readily shatter when struck a
sharp blow.
3. Electrical Conductivity: A measure of the degree
to which a substance conducts an electrical current.
Metals have a high electrical conductivity.
4. Electrical Nonconductivity: A characterization of
ionic substances, such as salts, that do not readily
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conduct electricity.
5. Ionic Bond: The chemical bond that binds a
metallic ion to a nonmetallic ion by electrical
attraction.
6. Ions: A charged object formed when an atom or
molecule loses or gains electrons.
7. Malleability: The characteristic of substances that
allows them to be worked into desirable shapes or
drawn 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 metal
substances.
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 an
ionic bond. The precise oxidation state is represented
by a – sign followed by the number of electrons
gained.
12. Nonmetal: See Chapter 18.
13. Oxidation Number: A signed number that indicates
how many electrons an atom loses (positive
number) or gains (negative number) by forming a
compound. In ionic bonding the transfer is complete
and distinct ions are formed. In covalent
bonding the transfer is only partial and distinct ions
are not formed.
14. Positive Oxidation State: The state of an atom
which has lost one or more electrons to form an
ionic bond. The precise oxidation state is represented
by a + sign followed by the number of electrons
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 degree
to which a substance conducts heat. Metals have a
high thermal conductivity.
17. Transparency: A characteristic of ionic substances,
such as salts, that readily transmit light.
Opposite to opaqueness.
D. FOCUS QUESTIONS
1. Consider atoms held together in metallic bonds:
a. What happens to the valence electrons when
the 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 bonds
are 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-particle
duality that the Wave Model is based on.
2. Consider the following elements held together in
ionic bonds: (a) sodium, chlorine; (b) magnesium,
chlorine. In each case:
a. Determine the positive or negative oxidation
states of the different atoms.
b. Write the chemical formula for the compound.
c. Write and balance the chemical equation for
the reaction.
d. Sketch an energy well for each kind of atom in
the compound. Draw a circle around each electron
soon to be lost in one atom, and an empty circle at
the location soon to be filled in the other atom.
e. List the main parts of the Wave Model of the
atom.
f. State the fundamental principle of wave-particle
duality that the Wave Model is based on.
3. Sketch a diagram showing a possible arrangement
of the ions in a salt. Explain what would happen to
this arrangement if shear forces were exerted on the
salt. Why is the salt brittle? Also, using your understanding
of ionic bonds, explain why table salt dissolved
in water is an ionic conductor. Name and
state the fundamental principle that explains the
forces that are involved.
E. EXERCISES
20.1. When gold and silver are mixed, the appearance
is still that of pure gold. How could you distinguish
the alloy from pure gold?
20.2 Which of the following pairs of metals is most
likely to form alloys of all compositions?
(a) 56Ba and 31 Ga
(b) 50 Sn and 82 Pb
(c) 3Li and 83 Bi?
20.3. Define the following terms: (a) electrical
conductivity, (b) metallic luster, (c) malleability, (d)
thermal conductivity.
20.4. Which of the following pairs of metals is
most likely to form alloys of all proportions?
(a) 55Cs and 49 In
(b) 78 Pt and 79 Au
(c) 82Pb and 19 K
20.5. In the left half of Figure 20.4 notice the
prominent diagonal lines of oxidation states. How far is
it vertically from one line to the next? How far horizontally?
20.6. How many of the 23 elements listed in the
left half of Figure 20.4 have only a single oxidation
state?
20.7. Write down the primary oxidation state of the
following elements on the basis of their positions in the
Periodic Table:
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(a) 19K
(b) 10 Ne
(c) 29Cu
(d) 35 Br
(e) 15P
20.8. Write down the primary oxidation states of
the following elements and check your answers in
Figure 20.4:
(a) 5B
(b) 20 Ca
(c) 17Cl
(d) 1 H
(e) 18Ar
(f) 28Ni
(g) 24 Cr
20.9. Referring to Figure 18.3, select two atoms
that are likely to have negative oxidation states because
they have very high ionization energies. Avoid the
noble gases 2 He, 10 Ne, and 18 Ar. Now look these elements
up in Figure 17.7 to see if there is a vacancy for
an electron in a low-lying orbital.
20.10. Referring to Figure 18.3, select two atoms
that are likely to have positive oxidation states because
they have low ionization energies. Now look these elements
up in Figure 17.7 to see if the top electrons are
near the top of the well.
20.11. Use Figure 20.4 to write the common oxidation
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 is
likely to lose to become Li 1+ . Do the same for Be going
to Be 2+ (two electrons).
20.13. How many I – ions will be required to neutralize
the charge on an Al 3+ ion?
20.14 How many Al 3+ and O 2– ions must be combined
to achieve charge neutralization? (Hint: It will
require more than one of each type of ion.)
20.15. Write the chemical formulas for the following
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 chemical
formulas for the combinations:
(a) 4Be and 35 Br
(b) 31 Ga and 8 O
20.17. Write the names of the following compounds:
(a) BeO
(b) CaCl 2
(c) LiF
(d) Na 2 S
20.18. Write the names of the following compounds:
(a) CrO
(b) CrO 3
(c) Cr 2 O 3
(d) NO 2
(e) SO 3
20.19. Complete Table 20.1.
20.20. Which of the following is an ionic compound?
(a) O 2
(b) NH 3
(c) CO 2
(d) CH 4
(e) MgF 2
is:
20.21. The correct formula for potassium sulfide
(a) KS
(b) K 2 S
(c) KS 2
(d) K 3 S
(e) KS 3
20.22. Determine the primary oxidation state for
Rb. If two Rb atoms combine with one S atom, what is
the oxidation state of S?
(a) +2
(b) +1
(c) 0
(d) –1
(e) –2
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Table 20.1.
NAME
FORMULA
LUSTER OR
COLOR
FORM
CONDUCTIVITY
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
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