Download Assignment 20 ELECTRONIC STRUCTURE OF ATOMS AND IONS I

Assignment 20
ELECTRONIC STRUCTURE OF ATOMS AND IONS
I. BACKGROUND
The modern description of an atom is that of a positively charged nucleus, which contains nearly all
the mass of the atom, surrounded by negatively charged electrons , which distribute themselves in an
electron cloud that accounts for most of the volume of the atom. The nucleus is composed of positively
charged particles called protons and neutrally charged particles called neutrons. Though the proton mass
and neutron mass are not exactly the same, each has a mass three orders of magnitude heavier than that of
an electron. The number of protons in the nucleus is called the atomic number and is equal to the nuclear
charge. The atomic number determines an element’s chemical identity. The volume of the electron cloud is
five orders of magnitude greater than that of the nucleus. Since a typical atom is 10-8 cm in diameter, this
volume relationship implies that the average nuclear diameter is 10-13 cm. To put this in perspective, if the
space occupied by an atom’s electrons was the size of a professional baseball stadium, then the nucleus
would be about the size of a fly at the center of the field. In a neutral atom, the nuclear charge is balanced
by the total electrical charge of the electron cloud.
II. AN ACCURATE DESCRIPTION OF ATOMS—QUANTUM NUMBERS
The theory of quantum mechanics will be discussed in class and covered in your textbook. The
vocabulary of quantum mechanics, which is described below, is unusual and probably unfamiliar to you but
is needed to completely describe the atom and its electrons. According to quantum mechanics, the spatial
distribution of the electron cloud and the internal energy of an atom are determined by the electron
occupancy of various atomic orbitals. Quantum mechanics tells us how to identify the electrons and their
atomic energy using a set of labels called quantum numbers. Quantum mechanics also dictates the rules
that one must follow when determining the occupancy the available orbitals.
The energy associated with any single electron is most strongly dependent on the principal
quantum number, n, where n is an integer greater than zero. All electrons whose description contains the
same n-quantum number are referred to as being in the same shell. Both the electron energy and the
volume which that electron is most likely to occupy increases with increasing value of n.
The energy associated with an electron is dependent to a lesser extent on the angular momentum
quantum number, l, whose values range from 0 to n-1 in integral steps, e.g., for n = 3; l = 0, 1, or 2.
Again, for a given value of n, the energy associated with an electron increases with increasing l. Electrons
whose description contains the same l quantum number belong to the same sub-shell, e.g., s- (l=0), p(l=1), d- (l=2), and f- (l=3) sub-shells. Each of these sub-shells has associated with it 2l+1 atomic orbitals
that are degenerate, i.e., at the same energy. Thus an s- sub-shell is 1-fold degenerate (i.e., has one
atomic orbital associated with it) since l=0 and 2(0) + 1 = 1, whereas a p- sub-shell is 3-fold degenerate
(i.e., has three atomic orbitals associated with it) since l=1 and 2(1) + 1 = 3.
The atomic orbitals in a given sub-shell are identified by the magnetic quantum number, ml,
whose values range from +l to – l in increments of 1, e.g., for a p- sub-shell, ml may have values of –1, 0,
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or 1. Each of these values corresponds to a different atomic orbital that electrons may occupy. For this
class we do not treat these values as containing any information about the energy of the electron associated
with this orbital. (The influence of the ml quantum number on the energy structure of an atom is covered in
Chemistry 262). For this class, information about the relative electron energies is expressed in the
combination of n and l quantum numbers associated with the electron.
In terms of occupancy of the atomic-orbitals, once again quantum mechanics guides us. The Pauli
Principle dictates that only two electrons can exist in a given atomic orbital, and that the two electrons must
have opposite spin quantum numbers, ms, where ms = +1/2 or –1/2. Thus, the capacity for electrons is 2
in the s- sub-shell (1 atomic orbital that can hold two electrons), 6 in the p- sub-shell (3 atomic orbitals that
can hold two electrons each, 3x2 = 6), 10 in the d- sub-shell, and 14 in the f- sub-shell. The actual number
of electrons in any particular kind of sub-shell is shown as a superscript after the letter designating the subshell; e.g., s1, s2, p3, d6, f 5, etc.
In summary, each electron in an atom can be identified by its own, unique set of quantum numbers,
n, l, ml, and ms. This set of quantum numbers contains information about the relative energy associated with
an atom’s electron and restricts how the electrons may occupy the atomic orbitals. Much of the information
covered to this point is summarized in the following table, where the shells and their available electron
capacities are listed, as are the sub-shells into which these electrons may be placed. What is not contained
in this table is a relative ordering of the energy of the sub-shells. This will be discussed below.
n (Shell)
1
2
3
4
5
6
7
l (Sub-shells)
Electron Capacity of the Shell
s
2
s p
8 (=2+6)
s p d
18 ( = 2 + 6 + 10 )
s p d f
32 ( = 2 + 6 + 10 + 14 )
s p d f g
more than 32
s p d f g h
"
s p d f g h i
"
?The italicized sub-shells are not needed for normal (ground state) atoms.]
Almost everything in chemistry is explained in terms of the behavior of electrons on atoms, thus
making it very important for us to know about the electrons. The Periodic Table is a convenient summary
of the electron properties in an atom. In this Assignment we will learn the shorthand conventions which
chemists use when discussing electrons in atoms, while reflecting upon the structure of the Periodic Table.
Orbitals and the Periodic Table's Shape
The Periodic Table pre-dates the quantum mechanics revolution in chemistry and was first devised
to show inherent relationships among the properties of the chemical elements. However, it turns out that
each position of an element on the Periodic Table also can be matched to a specific quantum mechanical
electron configuration in a systematic way that we will explore.
Each horizontal row in the Periodic Table is called a period. The first period contains only two
elements (H and He), but the second period contains eight elements, the fourth period expands to include
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eighteen elements, and the sixth period expands even more to include thirty-two elements. A vertical column
on the Periodic Table is called a column or group. The Periodic Table contains a total of 18 different
columns or groups, within which there are recurrences of chemical characteristics.
In each period the first two columns correspond to atoms whose highest energy electrons reside in
the s- sub shell, i.e., they are s-type electrons. Thus the highest energy electrons in H and He in the first
period are in the 1s atomic orbital (note that s- type sub-shell can also be referred to as an atomic orbital
because there is only a single atomic orbital associated with the s- sub-shell), while the highest energy
electrons in Li and Be in the second period are in the 2s atomic orbital, and so on down the rows. The n
value in “ns” matches the period number.
The elements in the last six columns have their highest energy electrons in a p- sub-shell, i.e., they
are p-type electrons. Thus, B through Ne (in the second period) correspond to 2p sub-shell, with the n
value in “np” matching the period number. Note that there is no 1p location on the Periodic Table. This is
because l = n-1 and for n = 1, l can only be equal to zero, i.e., only an s- sub-shell is available.
Transition elements—located in vertical columns 3 through 12—add electrons to the d- subshell, i.e., they are d-type electrons. Here the n value in “nd” is one unit less than that of s- and p- in the
same period. Thus, in the fourth period, the positions from left to right correspond to electrons being placed
in the 4s- sub-shell for the first two positions (K, Ca), the 3d- sub-shell for the next ten positions (Sc
through Zn), and the 4p- sub-shell for the last six positions (Ga through Kr).
In effect, the SHAPE of the Periodic Table reflects all that has been stated and thus provides a
guide to how electrons occupy atomic energy levels, i.e., the Periodic Table reveals the electron
configuration. For example, the shape of the Periodic Table allows 1s locations but no 1p. Similarly, there
are 2s and 2p but not 2d. The shape also sets a maximum number of electrons of each type. For
example, in the Periodic Table, only two positions are 3s; six positions are 3p; ten positions are 3d (these in
row 4). This idea of associating each location in the Periodic Table with a specific electron configuration
undergirds the Aufbau Principle, which is the topic that follows.
III. AUFBAU PRINCIPLE
Electron Configurations of Neutral Atoms
An atom's electron configuration is a summary of the electron shell and sub-shell population
within that neutral atom. The major emphasis of this Assignment is to recognize the relationship of the
electron distributions to the Periodic Table.
The elements in the Periodic Table are arranged according to increasing atomic number. Thus each
new element in the sequence has a neutral atom containing not only one more proton but also one more
electron than the element immediately preceding it. (Unfortunately, there is no similar convenient information
regarding the change in the number of neutrons in the nucleus as we move through the periodic table. This
information must be gleaned from the atomic mass.) The Aufbau ("build-up") Principle assumes that each
new element contains the same configuration as the previous element, except that one additional electron
must now be placed into the next available sub-shell. The type of sub-shell for each "new electron" is
summarized diagrammatically below.
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1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
s
s
s
d
p
Transition Metals
f
Lanthanides and Actinides
Recall that each new period in the Periodic Table begins a new principal quantum number, n, for the
valence electrons and that each period ends with a noble gas element. In the first period, n = 1; only the ssub-shell is available to the electrons in the ground state species. The two elements and their configurations
are H(1s) and the noble gas He(1s2). In the second period, n = 2 for the valence electrons and the
electrons are added to the s- sub-shell (for Li and Be) and then to the p- sub-shell (for B through noble gas
Ne).
Rather than listing all the electrons of all of an atom's shells, chemists find it convenient to write an
abbreviated version of the electron configuration. Starting with the noble gas found at the end of the
preceding period (its symbol being enclosed in parentheses), the abbreviated version lists, in filling order, all
the electrons added thereafter, from the beginning of the period containing the element. Thus:
AUFBAU
Li = (He) 2s1
C = (He) 2s22p2
Ne = (He) 2s22p6
(one 2s electron in the period beyond noble gas He)
(four electrons in the period beyond noble gas He)
(completion of second period with 8 valence electrons)
In the third period, n = 3 for the valence electrons and both s-type and p- type sub-shells are available to
the electrons in the ground state, the lowest energy state:
AUFBAU
Mg = (Ne) 3s2
P = (Ne) 3s23p3
Ar = (Ne) 3s23p6
(two 3s electrons in the period beyond noble gas Ne)
(five electrons in the period beyond noble gas Ne)
(completion of third period with 8 valence electrons)
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In the fourth period, n = 4 for the valence electrons. The next lowest energy sub-shells into which
electrons are placed in this period are s-type (for K and Ca), d-type (for Sc through Zn), and p-type (for
Ga through Kr), respectively. The electrons occupying the s- and p- sub-shells are in the valence shell of
that period, n = 4. The electrons occupying the d-sub-shell are in the n- 1 = 3 shell, although still in period
4. Note that d-type electrons are not considered valence electrons. Some examples showing this filling
order are:
AUFBAU
Ca = (Ar) 4s2
Zn = (Ar) 4s23d10
(only the two 4s electrons are in the valence shell)
2
10
3
As = (Ar) 4s 3d 4p (the valence shell now has five electrons)
Kr = (Ar) 4s23d104p6 (completion of the fourth period)
In the fifth period, n = 5 for the valence electrons. The next lowest energy sub-shells into which electrons
are placed in this period are the 5s- (for Rb and Sr), the 4d- (for Y through Cd), and the 5p- (for In
through Xe) sub-shells, respectively.
AUFBAU
Tc = (Kr) 5s24d5
In = (Kr) 5s24d105p1
Xe = (Kr) 5s24d105p6
In the sixth period, n = 6 for the valence electrons and the next lowest energy sub-shells into which
electrons are placed in this period are the 6s- (for Cs and Ba), the 5d- (for La, and Hf through Hg), the 4f(for the rare earth elements following La, that is, Ce through Lu), and the 6p- (Tl through Rn) sub-shells,
respectively. The electrons occupying the s- and p- sub-shells are in the valence shell of that period, n = 6
in this case. The electrons occupying the d- sub-shell are in the n- 1 = 5 shell, although still in period 6.
The electrons occupying the f- sub-shell are in the n- 2 = 4 shell, although still in period 6. Note that f-type
electrons are not considered to be valence electrons. The f-type electrons are first encountered in this sixth
period of the Periodic Table. The 4f- sub-shell begins to fill at element Ce after the 5d- sub-shell has
received its first electron at element La. The 4f- sub-shell then becomes filled before any additional
electrons are added to the 5d- sub-shell. (In the writing of aufbau configurations, all the d-type electrons
should be listed together. The order of listing the (n- 1)d- and the (n- 2)f- sub-shells in the configuration is
arbitrary.)
AUFBAU
2
56Ba = (Xe) 6s
(two 6s electrons in the valence shell)
La = (Xe) 6s25d1
(the 1 of 5d1 should be written)
57
Ce = (Xe) 6s24f15d1
or
(Xe) 6s25d14f1
(4f and 5d arbitrarily reversed—see above)
Lu = (Xe) 6s24f145d1
or
(Xe) 6s25d14f14
(the f sub-shell is now full)
Hg = (Xe) 6s24f145d10
or
(Xe) 6s25d104f14 (from Hf to Hg, 5d is being filled)
58
71
80
Tl = (Xe) 6s24f145d106p1
(now valence shell contains three electrons)
Rn = (Xe) 6s24f145d106p6
(completion of sixth period with 8 valence
81
86
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electrons)
In the seventh period, n = 7 for the valence electrons and the next lowest energy sub-shells into which
electrons are placed are 7s-, 6d-, 5f-, and 7p- sub-shells, respectively. For example,
AUFBAU
Ac = (Rn) 7s26d1
(two 7s electrons in the valence shell)
89
Lr = (Rn) 7s25f146d1
or
103
(Rn) 7s26d15f14
(the f- sub-shell is full)
Discrepancies from Aufbau
The aufbau method is a quick way to predict electron configurations of ground state, isolated neutral
atoms. For most chemical purposes these predictions are adequate, even where the predictions do not
exactly match the actual distributions found by lab experiment; but some elements do indeed show unusual
electron distributions. Aufbau predictions are not particularly useful for anticipating which will be unusual
cases. It is fortunate for chemists that the aufbau predictions work so well.
Unusual electron order occurs with certain transition metals (involving rearrangements between the
(n)s- and (n- 1)d- sub-shells) and many rare earth elements (involving rearrangements between (n- 1)dand (n- 2)f- sub-shells). The rare earth atom peculiarities are beyond the scope of this course, and the
unusual cases among the transition metals mostly need not concern us either. But there are four cases
among the transition elements that you must memorize:
Cr (in column 6 of the Periodic Table) and
Cu, Ag, and Au (in column 11).
These neutral atoms actually have only one s electron in the valence shell. Note the configuration of the
silver atom, Ag,
AUFBAU PREDICTION
2
9
Ag = (Kr) 5s 4d
ACTUAL
Ag = (Kr) 5s14d10.
From its position in the Periodic Table, Ag is predicted to have an electron configuration with two valence
electrons and a nearly full d- sub-shell. However, the actual distribution for Ag shows an electron repositioned, completing the d- sub-shell. The Ag atom actually contains only one electron in the valence
shell. Copper and gold (which lie in the same vertical column of the Periodic Table as Ag does) show
similarity to silver, by completing the inner d- sub-shell, also;
AUFBAU PREDICTION
2
9
Cu = (Ar) 4s 3d
Au = (Xe) 6s24f145d9
ACTUAL
Cu = (Ar) 4s13d10
Au = (Xe) 6s14f145d10.
Chromium atom, by re-positioning one electron, half-fills the d- sub-shell.
AUFBAU PREDICTION
ACTUAL
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Cr = (Ar) 4s23d4
Cr = (Ar) 4s13d5
These unusual configurations have important chemical ramifications for these elements.
Aufbau and Electron Energy
Note that the aufbau configuration does not necessarily tell the relative energies of an
atom’s electrons . The aufbau electron configuration tells only what electrons are present on a particular
atom and where. The entire cluster of orbitals mentioned in an aufbau listing identifies the kinds of electrons
most energetic in the atom, but the order of the listing gives no indication of the relative energies of those
listed electrons. To briefly review the comments about energy:
•
•
In any aufbau electron configuration, the electron with the highest n value has the highest energy.
For orbitals with the same n value, the energy increases with increasing orbital complexity, s
electrons being lower energy than p electrons, which are lower than d electrons, which are lower
than f electrons.
IV. ELECTRON CONFIGURATIONS OF SIMPLE IONS
Simple ions differ from neutral atoms by containing more or fewer electrons than the corresponding
neutral atoms. Since an electron is negatively charged, the loss of electron(s) from any neutral atom results
in a positive ion, whereas the gain of electron(s) results in a negative ion. Rigorously speaking, simple ions
only exist under special gas phase conditions. However, the term “simple ion” is often a reasonable
approximation and will be used in this course. For atoms in periods 1-3 and for atoms in the first two
columns, the ion's existence as a simple ion came from either (1) removal of all electrons from the atom's
valence shell, or (2) adding of electrons to complete a shell of 8 valence electrons (2 for hydrogen). It is
rare that a single atom can gain or lose more than two electrons to form an ion in this manner.
To predict (or deduce) the aufbau configuration of simple ions, use this mental construct:
?? Always state the neutral atom by aufbau first—this configuration establishes the orbitals of
concern.
?? Next, devise the ion by altering the neutral atom, so as to account for loss or gain of electrons.
?? When a neutral atom loses electrons in becoming a cation (a positively charged ion), the most
energetic electrons (of highest n value) are lost first.
?? When a neutral atom gains electrons in becoming an anion (a negatively charged ion), the extra
electrons are added where a sub-shell is only partly-filled or to the next highest sub-shell.
For many common ions the resulting electron distribution matches that of a noble gas. Some examples of
predicting configurations of simple ions (where the ion configuration corresponds to a noble gas):
Mg (Ne)3s2
can become
Mg2+ (Ne)3s0
(by loss of 2e- )
comparable to a Ne configuration
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Cl (Ne)3s23p5
can become
Cl– (Ne)3s23p6
(by gain of 1e- )
comparable to an Ar configuration
Br (Ar)4s23d104p5
can become
Br– (Ar)4s23d104p6
(by gain of 1e- )
comparable to a Kr configuration
Ce (Xe)6s25d14f1
can become
Ce4+ (Xe)6s05d04f0
(by loss of 4e- )
comparable to a Xe configuration
When ions are formed by transition metals and post-transition metals, the first electrons to be lost are
the electrons of the valence shell (largest n value). If additional electrons are lost, they come from the subshell of next highest energy—usually , these are d electrons in the n-1 shell. Some examples of predicting
configurations of transition metal ions are:
Co (Ar)4s23d7
can become
Co2+ (Ar)4s03d7
(by loss of 2 e- )
(Note that electrons are first removed from the 4s- sub-shell rather than the 3d- sub-shell because
the 4s electrons are higher energy. Of the nine electrons that a cobalt atom has beyond the noble
gas configuration, the Co2+ ion retains seven, and these are d-type electrons.)
Cr (Ar)4s23d4 (aufbau structure) can become
Cr3+ (Ar)4s03d3
(by loss of 3e- )
OR (accounting for the actual structure of the neutral Cr atom)
Cr (Ar)4s13d5 (actual structure)
can become
Cr3+ (Ar)4s03d3
(by loss of 3e- )
(Note that either neutral configuration yields the same conclusion for the Cr3+ ion. Of the six
electrons a chromium atom has beyond the noble gas configuration, the Cr3+ ion retains three.
Besides loss of 4s-type electron(s), 3d-type electron(s) are also lost to give the final +3 ionic
charge.)
Pr (Xe)6s25d14f2
can become
Pr3+ (Xe)6s05d04f2
(by loss of 3e- ; two 4f e- remain)
Two important unusual cases should be mentioned involving Periodic Table column 11. Silver, when present
as an ion, almost always occurs as a Ag+ ion. In this unusual case, one must begin with the actual
configuration for Ago in order to get the correct ion configuration.
Ag (Kr)5s14d10
(one valence e- )
can become
Ag+ (Kr)5s04d10
(no valence e- )
(by loss of le- )
Copper is usually encountered as a Cu2+ ion, but sometimes Cu+ ion occurs. The Cu+ configuration must be
derived like Ag+ ion:
Cu (Ar)4s3d10
can become
Cu+ (Ar)4s03d10
(by loss of 1e-)
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V. ELECTRON DOT STRUCTURES
The electron dot structure of an atom is given by the symbol of the atom surrounded by a number of
dots equal to the number of electrons in the valence electron shell. By convention, ONLY valence electrons
appear as dots. For atoms of the first eight-membered period, these dot structures are:
Li
Be
B
C
N
O
F
Ne
As a "mental construct" (merely for convenience in devising consistent and informative dot structures for
individual, isolated atoms), imagine these valence electrons at the sides of an imaginary square around the
symbol.
s
p ? p
p
By means of this simplistic mental construct, the build-up and pairing of electrons in an isolated atom's
valence shell is readily represented as in the dot pictures of the second period elements depicted above.
Notice that the first two electrons in a valence shell will occupy the lower energy s-type orbital. The p- subshell consists of three p-type, degenerate atomic orbitals. Consistent with Hund’s rule, any electrons added
to the p- sub-shell will fill the atomic orbitals evenly and with the same electron spin orientation (i.e., ms
value) while there is only one electron in an atomic orbital, i.e., the electrons are unpaired. The number of
valence electrons, and thus the number of dots, is suggested by the Roman numeral at the top of the
Periodic Table.
Dot Pictures for Simple Ions
Dot pictures for simple ions are designed by depicting the neutral atom first, and then altering the dot
picture appropriately to account for the ionic charge. An ion's dot picture MUST have the ion's electrical
charge stated explicitly as a necessary aspect of the dot structure. Examples include the following:
Problem:
deduce electron dot structures for these ions (from neutral atoms).
Method:
It is convenient—as a visual representation of the reasoning process—to express each
answer by depicting the dot structure of the neutral atom first and then of the requested
ion.
a) for sodium ion?
Na
becomes
Na+
Al
becomes
Al
Cl
becomes
( by loss of 1 electron)
b) for aluminum ion?
+3
( by loss of 3 electrons)
c) for chloride ion?
Cl
( by gain of 1 electron)
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d) for oxide ion?
2O
becomes
O
( by gain of 2electrons)
e) for V3+ ion?
V
becomes
V
3+
f) for Nd3+ ion?
Nd
becomes
Nd 3+
( by loss of 3 electrons, including one of the d electrons,
which was not represented as a dot on the V atom at left)
( by loss of 3 electrons, including one of the d electrons,
which was not represented as a dot on the V atom at left)
As seen in examples (e) and (f) above, electron dot pictures have some limitations as representations for
transition metals and their ions, because, by convention, (n- 1)d electrons are not depicted as dots. By
aufbau, each neutral transition metal atom contains only two valence electrons, ns2, therefore those dot
pictures are written with only 2 dots (there are exceptions such as Ag and Cu). Their ions will have lost
these s-type electrons, thus there typically are no dots on transition metal ions. However, (n- 1)d electrons,
not shown as dots, may have been lost also, as is demonstrated in examples (e) and (f). Only the notation of
the charge (on the "bare" dot picture) provides the clue that the ion is missing a third electron (one of its
(n- 1)d electrons).
In the case of V3+, a neutral vanadium atom contains 5 electrons beyond that of the noble gas argon—
the configuration is (Ar)4s23d3. It loses the two valence 4s electrons and one of the 3d electrons; the V3+
ion still retains two 3d electrons beyond the noble gas structure. None of the 3d electrons appear as dots in
the dot pictures.
Dot Pictures for Ionic Compounds
A dot picture for an ionic compound should represent each ion as a separate, charged entity since the
ions in an ionic compound do not share any electrons. Instead of joining ions together, use a comma or
parentheses to emphasize the separateness of the ions in the compound.
Problem:
Deduce electron dot structures of these ionic compounds.
Method:
As a means of showing your reasoning visually, depict dot pictures of the neutral atoms
first. Then, deduce and depict the structures for the ionic compound—show each ion
individually with its electrical charge explicitly stated as a fundamental part of its dot picture.
a) Dot structure for NaCl (which is an ionic compound, Na+Cl- ) ?
Na
and
Cl
will occur as
Na , Cl
-
This construction is a summary of two changes: the removal of an electron from Na and the addition of
an electron to Cl.
Na
becomes
Na
by loss of 1 electron and
Cl
becomes Cl
by gain of 1 electron
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This ionic compound ultimately becomes electrostatically neutral by having equal numbers of
(independent) positive Na+ ions and negative Cl- ions as neighbors within the crystal.
b) Dot structure for ZnF 2 (which is an ionic compound, Zn2+(F- )2 ) ?
2
Zn
and
2 F
will occur as
Zn , 2 F
Zn has a complete 3d10 sub-shell, but only the valence 4s electrons are involved in the electron dot
diagrams. The zinc atom loses both of its two valence electrons when it becomes an ion; only one
electron is gained by each fluorine atom in becoming a fluoride ion. In order for the overall compound
to become electrically balanced, two fluoride ions are needed per one zinc ion, but the ions continue to
be independent species.
c) Dot structure of Na2S (which is an ionic compound, (Na+)2S2- ) ?
2 Na
and
S
will occur as
2 Na ,
S
2
The use of parentheses around dot structures of ions is optional, but convenient.
d) Dot structure of CaO (which is an ionic compound, Ca2+O2- ) ?
Ca
and
O
will occur as
2
Ca ,
O
2
Remember you must denote explicitly the separateness of the ions comprising ionic compounds.
***
***
***
***
***
NOTE: A word of caution is relevant, as you construct answers to problems in this Assignment and on
quizzes and exams. Whenever you write electron dot pictures:
?? Draw the dots darkly and precisely—with careful attention to the number and placement of the
dots—revealing whether they are paired or single (unpaired).
?? Always provide exact labeling of the charge of any ion or ionic component.
?? In ionic compounds, the ions MUST NOT be linked by the electrons (there is NO SHARING
OF AN ELECTRON PAIR between the ions)—each ion exists as a separate entity.
Emphasize each ion's individuality, through good symbolic use of commas as separators and/or
parentheses as enclosures.
Casualness in the drawing of the dots or in the writing of the charges will result in serious mistakes in these
problem sets, quizzes and exams.
?
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