Chapter 9
Chemical Bonding I:
Lewis Theory
Bonding Theories
• explain how and why atoms attach together
• explain why some combinations of atoms are
stable and others are not
9 why is water H2O, not HO or H3O
• One of the simplest bonding theories was
•
•
developed by G.N. Lewis (1875–1946) and is
called Lewis theory.
Lewis theory emphasizes valence electrons to
explain bonding.
Using Lewis theory, we can draw models—called
Lewis structures—that allow us to predict many
properties of molecules.
9 aka electron dot structures
9 such as molecular stability, shape, size, polarity
Tro, Principles of Chemistry: A Molecular Approach
Tro, Principles of Chemistry: A Molecular Approach
Bonding
Types of Bonds
• A chemical bond forms when the potential
energy of the bonded atoms is less than the
potential energy of the separate atoms.
• To calculate, you need to consider the
following interactions:
9nucleus-to-nucleus repulsions
9electron-to-electron repulsions
9nucleus-to-electron attractions
Tro, Principles of Chemistry: A Molecular Approach
3
2
Tro, Principles of Chemistry: A Molecular Approach
Types of Bonding
Ionic Bonds
• When a metal atom loses electrons, it
becomes a cation.
9metals have low ionization energy; are relatively
easy to remove an electron from
• When a nonmetal atom gains electrons, it
becomes an anion.
9nonmetals have high electron affinities; they
release energy when they gain electrons
• The oppositely charged ions are then
attracted to each other, resulting in an ionic
bond.
Tro, Principles of Chemistry: A Molecular Approach
5
Covalent Bonds
Tro, Principles of Chemistry: A Molecular Approach
6
Metallic Bonds
• Nonmetals have relatively high ionization energies,
• The low ionization energy of metals allows
•
• The simplest theory of metallic bonding
so it is difficult to remove electrons from them.
When nonmetals bond together, it is better in terms
of potential energy for the atoms to share valence
electrons.
9 Potential energy is lowest when the electrons are
between the nuclei.
• Shared electrons hold the atoms together by
attracting nuclei of both atoms.
them to lose electrons easily.
involves the metals atoms releasing their
valence electrons to be shared by all
atoms/ions in the metal.
9an organization of metal cation islands in a sea
of electrons
9electrons delocalized throughout the metal
structure
• Bonding results from the attraction of metal
cations for the delocalized electrons.
Tro, Principles of Chemistry: A Molecular Approach
7
Tro, Principles of Chemistry: A Molecular Approach
8
Metallic Bonding
Determining the Number of
Valence Electrons in an Atom
• The column number on the periodic table will tell you
how many valence electrons a main group atom
has.
Tro, Principles of Chemistry: A Molecular Approach
9
Lewis Symbols of Atoms
• aka electron dot symbols
• Use symbol of the element to represent the
•
Tro, Principles of Chemistry: A Molecular Approach
Practice—Write the Lewis symbol
for 1) arsenic
2) barium
nucleus and inner electrons.
Use dots around the symbol to represent
valence electrons.
9 1) Pair first two electrons for the s orbital.
9 2) Put one electron on each open side for p electrons.
9 3) Then pair the rest of the p electrons.
Tro, Principles of Chemistry: A Molecular Approach
Tro, Principles of Chemistry: A Molecular Approach
12
Lewis Symbols of Ions
• (Group A) Cations have Lewis symbols
without valence electrons.
9electrons lost in the cation formation
• Anions have Lewis symbols with eight
valence electrons.
9electrons gained in the anion formation
Tro, Principles of Chemistry: A Molecular Approach
Stable Electron Arrangements
and Ion Charge
• Metals form cations by
losing enough electrons to
get the same electron
configuration as the
previous noble gas.
• Nonmetals form anions by
gaining enough electrons
to get the same electron
configuration as the next
noble gas.
• The noble gas electron
configuration must be very
stable.
Tro, Principles of Chemistry: A Molecular Approach
Octet Rule
When atoms bond, they tend to gain, lose, or share electrons
to result in eight valence electrons.
• ns2np6 (noble gas configuration)
• exceptions
9 H, Li, Be, B attain an electron configuration like He
¾ H shares or gains one electron.
– though it commonly loses its one electron to become H+
¾ Be loses two electrons to become Be2+.
– though it commonly shares its two electrons in covalent bonds,
resulting in four valence electrons
¾ B loses three electrons to become B3+.
– though it commonly shares its three electrons in covalent bonds,
resulting in six valence electrons
9 expanded octets for elements in period 3 or below (later in chapter)
¾ using empty valence d orbitals
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15
Lewis Theory and Ionic Bonding
• Lewis symbols can be used to represent the
transfer of electrons from metal atom to
nonmetal atom, resulting in ions that are
attracted to each other and therefore bond.
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16
Energetics of Ionic Bond Formation
Predicting Ionic Formulas
Using Lewis Symbols
• The ionization energy of the metal is endothermic.
• Electrons are transferred until the metal loses all
•
9 Na(s) → Na+(g) + 1 e ─
ΔH° = +603 kJ/mol
• The electron affinity of the nonmetal is exothermic.
its valence electrons and the nonmetal has an
octet.
Numbers of atoms are adjusted so the electron
transfer comes out even.
9 ½Cl2(g) + 1 e ─ → Cl─(g)
ΔH° = −227 kJ/mol
• Generally, the ionization energy of the metal is
•
Li2O
larger than the electron affinity of the nonmetal;
therefore, the formation of the ionic compound
should be endothermic.
But the heat of formation of most ionic compounds
is exothermic and generally large. Why?
9 Na(s) + ½Cl2(g) → NaCl(s) ΔH°f = −410 kJ/mol
Tro, Principles of Chemistry: A Molecular Approach
17
WHY?
• The extra energy that is released comes
from the formation of a structure in which
every cation is surrounded by anions, and
vice versa.
• Electrostatic attraction is nondirectional!
Tro, Principles of Chemistry: A Molecular Approach
18
Crystal Lattice
• The ions are arranged in a pattern called a crystal
•
lattice.
The crystal lattice maximizes the attractions
between cations and anions, leading to the most
stable arrangement.
9no direct anion–cation pair
• Therefore, there is no ionic molecule.
9The chemical formula is an empirical formula,
simply giving the ratio of ions based on charge
balance.
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Tro, Principles of Chemistry: A Molecular Approach
20
Trends in Lattice Energy—
1) Ion Size
Lattice Energy
• The extra stability that accompanies the
formation of the crystal lattice is measured
as the lattice energy.
• The lattice energy is the energy released
when the solid crystal forms from separate
ions in the gas state,
9always exothermic
9hard to measure directly, but can be
calculated from knowledge of other processes
• Lattice energy depends directly on size of
• The force of attraction between charged
particles is inversely proportional to the
distance between them.
• Larger ions mean that the center of
positive charge (nucleus of the cation) is
farther away from negative charge
(electrons of the anion).
9larger ion = weaker attraction
9weaker attraction = smaller lattice energy
charges and inversely on distance
between ions.
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21
Lattice Energy vs.
Ion Size
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22
Trends in Lattice Energy—
2) Ion Charge
• The force of attraction between
•
oppositely charged particles is
directly proportional to the product
of the charges.
Larger charge means the ions are
more strongly attracted.
9 larger charge = stronger attraction
9 stronger attraction = larger lattice
energy
• Of the two factors, ion charge is
generally more important.
Tro, Principles of Chemistry: A Molecular Approach
23
Tro, Principles of Chemistry: A Molecular Approach
Lattice Energy =
−910 kJ/mol
Lattice Energy =
−3414 kJ/mol
HUGE!
24
Covalent Bonds
Covalent Bonds
• When two nonmetals bond, they often share
electrons since they have similar attractions
for them. This sharing of valence electrons
is called the covalent bond.
9These atoms will share sufficient numbers of
electrons in order to achieve a noble gas
electron configuration (that is, eight valence
electrons).
Tro, Principles
of Chemistry:
A Molecular
Approach
Copyright
© Houghton
Mifflin Company.
All rights reserved.
Presentation of Lecture Outlines, 9–25
• The tendency of atoms in a molecule to
have eight electrons in their outer shell (two
for hydrogen) is called the octet rule.
9Next slide illustrates how two electrons can be
shared by bonded hydrogen atoms.
Tro, Principles
of Chemistry:
A Molecular
Approach
Copyright
© Houghton
Mifflin Company.
All rights reserved.
Presentation of Lecture Outlines, 9–26
Lewis Structures
Lewis Structures
For binary molecular compounds
For binary molecular compounds
• You can represent the formation of the
covalent bond in H2 as follows:
H
. + .H
time in the region around each atom.
:
H H
– This uses the Lewis dot symbols for the
hydrogen atom and represents the
covalent bond by a pair of dots.
Tro, Principles
of Chemistry:
A Molecular
Approach
Copyright
© Houghton
Mifflin Company.
All rights reserved.
• The shared electrons in H2 spend part of the
Presentation of Lecture Outlines, 9–27
:
H H
– In this sense, each atom in H2 has a
helium configuration.
Tro, Principles
of Chemistry:
A Molecular
Approach
Copyright
© Houghton
Mifflin Company.
All rights reserved.
Presentation of Lecture Outlines, 9–28
• Electrons that are shared by atoms are
• When two atoms share one pair of electrons it is
called a single covalent bond.
9 two electrons
• One atom may use more than one single bond
to fulfill its octet.
9 to different atoms
9 H-only duet
9aka nonbonding pairs
••
F •
• F
••
•••
Bonding Pairs
••
••
Lone Pairs
••
F
Tro, Principles of Chemistry: A Molecular Approach
29
F
A single line can be used to show a bond
Tro, Principles of Chemistry: A Molecular Approach
Double Covalent Bond
• When two atoms share two pairs of
electrons, the result is called a double
covalent bond.
Triple Covalent Bond
electrons, the result is called a triple
covalent bond.
9six electrons
••
•N
•
••
•N
•
•
•
•
•
N ••
•• N
••
••
••
•O
•O
••
••
• •• ••
•O
•• •• O••
30
• When two atoms share three pairs of
••
9four electrons
F
•
••
F
••
••
Using lines as a pair
of electrons
Tro, Principles of Chemistry: A Molecular Approach
31
••
• O •H
••
••
H O H
••
••
••
H•
••
••
••
•
called bonding pairs.
• Electrons that are not shared by atoms but
belong to a particular atom are called lone
pairs.
Single Covalent Bonds
••
“Bonding” and “Lone Pair”
Electrons
Tro, Principles of Chemistry: A Molecular Approach
32
Predictions of Molecular
Formulas by Lewis Theory
Predictions of Molecular
Formulas by Lewis Theory
Hydrogen is more stable when it is singly bonded to another atom.
Oxygen is more stable when it is singly bonded to two other atoms
+
H2
+
+
or doubly bonded to one other atom.
H2O
+
HCl
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33
Covalent Bonding—
Model vs. Reality
O2
+
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34
Intermolecular Attractions
vs. Bonding
• Lewis theory predicts the melting and boiling points
of molecular compounds should be relatively low.
9 involve breaking the attractions between the molecules,
but not the bonds between the atoms
9 The covalent bonds are strong, but the attractions
between the molecules are generally weak.
• Molecular compounds have low melting points and
boiling points.
9 MP generally <300 °C
9 Molecular compounds are found in all three states at
room temperature.
Tro, Principles of Chemistry: A Molecular Approach
35
Tro, Principles of Chemistry: A Molecular Approach
36
Covalent Bonding—
Model vs. Reality
• Lewis theory predicts that the more electrons two
•
•
atoms share, the stronger the bond should be.
Bond strength is measured by how much energy
must be added into the bond to break it in half.
In general, triple bonds are stronger than double
bonds, and double bonds are stronger than single
bonds.
9 Lewis theory would predict that double bonds are twice as
strong as single bonds; however, double bonds are
generally less than twice as strong.
Tro, Principles of Chemistry: A Molecular Approach
37
Covalent Bonding—
Model vs. Reality
• Lewis theory predicts that the more electrons two
atoms share, the shorter the bond.
9 when comparing bonds to like atoms
• Bond length is determined by measuring the
•
distance between the nuclei of bonded atoms.
In general, triple bonds are shorter than double
bonds, and double bonds are shorter than single
bonds.
Tro, Principles of Chemistry: A Molecular Approach
Bond Polarity
HF
• Covalent bonding between unlike atoms
results in unequal sharing of the electrons.
9One atom pulls the electrons in the bond closer
to its side.
9One end of the bond has a larger electron
density than the other.
38
EN 2.1
EN 4.0
• The result is a polar covalent bond.
9bond polarity
9The end with the larger electron density gets a
partial negative charge.
9The end that is electron deficient gets a partial
positive charge.
Tro, Principles of Chemistry: A Molecular Approach
39
δ+ H •• F δ−
EN = electronegativity
Tro, Principles of Chemistry: A Molecular Approach
40
Ionic Character
• Most bonds have some degree of sharing
and some degree of ion formation to them.
• Bonds are classified as covalent if the
amount of electron transfer is insufficient for
the material to display the classic properties
of ionic compounds.
• If the sharing is unequal enough to produce
a dipole in the bond, the bond is classified
as polar covalent.
Tro, Principles of Chemistry: A Molecular Approach
41
Electronegativity Scale
Electronegativity
• measure of the pull an atom has on bonding
electrons
• Increases across period (left to right).
• Decreases down group (top to bottom).
9Fluorine is the most electronegative element.
9Francium is the least electronegative element.
9opposite of atomic size trend
• The larger the difference in electronegativity,
the more polar the bond.
9negative end toward more electronegative atom
Tro, Principles of Chemistry: A Molecular Approach
42
Electronegativity and Bond Polarity
• If the difference in electronegativity between bonded
atoms is 0, the bond is pure covalent.
9 equal sharing
• If the difference in electronegativity between bonded
atoms is 0.1 to 0.4, the bond is nonpolar covalent.
• If the difference in electronegativity between bonded
atoms is 0.5 to 1.9, the bond is polar covalent.
• If the difference in electronegativity between bonded
atoms is ≥2.0, the bond is ionic.
0
Tro, Principles of Chemistry: A Molecular Approach
43
4%
Percent ionic character
51%
“100%”
0.4
2.0
Electronegativity difference
4.0
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44
Bond Polarity
ENCl = 3.0
3.0 − 3.0 = 0
Pure covalent
ENCl = 3.0
ENH = 2.1
3.0 – 2.1 = 0.9
Polar covalent
Tro, Principles of Chemistry: A Molecular Approach
ENCl = 3.0
ENNa = 0.9
3.0 – 0.9 = 2.1
Ionic
45
Example 9.3(c)
Determine whether an N–O bond is ionic,
covalent, or polar covalent.
Tro, Principles of Chemistry: A Molecular Approach
46
Lewis Structures
of Molecules
• Lewis theory allows us to predict the
distribution of valence electrons in a
molecule.
• useful for understanding the bonding in
many compounds
• allows us to predict shapes of molecules
• allows us to predict properties of molecules
and how they will interact together
Tro, Principles of Chemistry: A Molecular Approach
47
Tro, Principles of Chemistry: A Molecular Approach
48
Lewis Structures
Beware!
• Lewis theory predicts that atoms will be most stable when
they have their octet of valence electrons.
• It does not require that atoms have the same number of
lone pair electrons they had before bonding.
9 first use the octet rule
• Some atoms commonly violate the octet rule.
9 Be generally has two bonds and zero lone pairs in its compounds.
9 B generally has three bonds and zero lone pairs in its compounds.
9 Many elements may end up with more than eight valence
electrons in their structure if they can use their empty d orbitals for
bonding.
• generally try to follow the common bonding patterns
9 C = 4 bonds & 0 lone pairs, N = 3 bonds & 1 lone pair,
O = 2 bonds & 2 lone pairs, H and halogens = 1 bond,
Be = 2 bonds & 0 lone pairs, B = 3 bonds & 0 lone pairs
9 Often, Lewis structures with line bonds do not include the
lone pairs.
¾ Their presence is assumed from common bonding patterns.
• Structures that result in bonding patterns different
from the common patterns may have formal charges.
¾ Called the expanded octet
Tro, Principles of Chemistry: A Molecular Approach
49
Tro, Principles of Chemistry: A Molecular Approach
Example—Writing Lewis structures
of molecules: HNO3
1) Write the skeletal structure.
in oxyacid, H outside attached to O’s
atoms with pairs of electrons and subtract
from the total.
9 Make the least electronegative
atom central.
¾
9 Don’t forget: A line represents two electrons.
N is central
2) Count valence electrons.
9 Sum the valence electrons for
each atom.
9 Add 1 electron for each − charge.
9 Subtract 1 electron for each +
charge.
Tro, Principles of Chemistry: A Molecular Approach
51
Example—Writing Lewis structures
of molecules: HNO3
3) Attach the central atom to the surrounding
9 H always terminal
¾
50
N=5
H=1
O3 = 3x6 = 18
Total = 24 e−
Start
Used
Left
Tro, Principles of Chemistry: A Molecular Approach
52
Electrons
24
8
16
Example—Writing Lewis structures
of molecules: HNO3
4) Complete octets, outside-in
9 H is already complete with 2
¾
1 bond
¾
Bonding electrons only
N=5
H=1
O3 = 3·6 = 18
Total = 24 e−
5) If all octets are complete,
give extra electrons to central
atom.
9 Elements with d orbitals can
have more than 8 electrons.
and recount electrons
tally
Lewis structures of molecules:
HNO3
Electrons
Start
24
Used
8
Left
16
Tro, Principles of Chemistry: A Molecular Approach
plus non-bonding
Start
Used
Left
Electrons
16
16
0
53
period 3 and below
6) If central atom does not have
octet, bring in electrons from
outside atoms to share.
9 Follow common bonding
patterns if possible.
Correct structure
Tro, Principles of Chemistry: A Molecular Approach
Practice—Lewis structures
54
Formal Charge
During bonding, atoms may end with more or less
electrons than the valence electrons they brought
in order to fulfill octets.
This results in atoms having a formal charge (FC).
CO2
H3PO4
SeOF2
SO32−
FC = valence e− − nonbonding e− − ½ bonding e−
left O …FC = 6 − 4 − ½ (4) = 0
for S … FC = 6 − 2 − ½ (6) = +1
right O….FC = 6 − 6 − ½ (2) = −1
NO2−
P2H4
• sum of all the formal charges in a molecule = 0
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55
9 in an ion, the sum equals the ion charge
Tro, Principles of Chemistry: A Molecular Approach
56
Continued- Lewis Structures
Practice—Assign formal charges.
7) Assign formal charges to the atoms.
CO2
a) FC = valence e− − lone pair e− − ½ bonding e−
b) Follow the common bonding patterns. (See slide 50)
all 0
SeOF2
−1
0
+1
−1
H3PO4
P = +1
rest 0
−1
−1
SO32−
S = +1
Se = +1
−1
−1
−1
0
Tro, Principles of Chemistry: A Molecular Approach
0
+1
57
Lewis structures- Resonance
• When there is more than one Lewis structure
for a molecule that differ only in the position
of the electrons, they are called resonance
structures.
• The actual molecule is a combination of the
resonance forms—a resonance hybrid.
9It does not resonate between the two forms,
though we often draw it that way.
• Look for multiple bonds or lone pairs.
Tro, Principles of Chemistry: A Molecular Approach
59
NO2−
0
−1
Tro, Principles of Chemistry: A Molecular Approach
•
•
•
•
•
•
•
P2H4
all 0
58
Rules of Resonance Structures
Resonance structures must have the same
connectivity.
9 Only electron positions can change.
Resonance structures must have the same number of
electrons.
Second row elements have a maximum of eight
electrons.
9 bonding and nonbonding
9 third row elements or below can have an expanded octet
Resonance structures must have the same total of
the formal charges.
Better structures have fewer atoms with formal
charges.
Better structures have smaller formal charges.
Better structures have the negative formal charge on
the more electronegative atom.
Tro, Principles of Chemistry: A Molecular Approach
60
Drawing Resonance Structures
1. First draw Lewis structure that
maximizes octets.
2. Assign formal charges.
3. Move electron pairs from
atoms with − formal charge
toward atoms with + formal
charge.
4. if + FC atom 2nd row, only
move in electrons if you can
move out electron pairs from
a multiple bond
5. if + atom 3rd row or below,
keep bringing in electron pairs
to reduce the formal charge,
even if get expanded octet.
Tro, Principles of Chemistry: A Molecular Approach
61
−1
−1
+1
−1
−1
+1
Drawing Resonance Structures
−1
1. Draw first Lewis structure that
maximizes octets.
2. Assign formal charges.
3. Move electron pairs from
atoms with − formal charge
toward atoms with + formal
charge.
4. if + FC atom 2nd row, only
move in electrons if you can
move out electron pairs from
multiple bond
5. if + FC atom 3rd row or below,
keep bringing in electron pairs
to reduce the formal charge,
even if get expanded octet
Tro, Principles of Chemistry: A Molecular Approach
+2
−1
62
Practice—Identify structures with better or
equal resonance forms and draw them.
all 0
SeOF2
P = +1
rest 0
−1
Tro, Principles of Chemistry: A Molecular Approach
−1
SO32−
S = +1
Se = +1
NO2−
−1
H3PO4
CO2
−1
Tro, Principles of Chemistry: A Molecular Approach
P2H4
all 0
64
−1
−1
Practice—Identify structures with better or
equal resonance forms and draw them.
CO2
−1
none
H3PO4
all 0
−1
2−
SO
−1 3
−1
−1
−1
all 0
−1
−1
−1
−1
Tro, Principles of Chemistry: A Molecular Approach
−1
+1
+1
−1
• expanded octets
9Elements with empty d orbitals can have more
than eight electrons.
+1
SeOF2
NO2−
Exceptions to the Octet Rule
−1
S=0
in all new
resonance forms
• odd-numbered electron species, e.g., NO
9will have one unpaired electron
9free radical
9very reactive
• incomplete octets
9B, Al
65
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66
Expanded octet around the central atom
B has more than 6 electrons
Tro, Principles of Chemistry: A Molecular Approach
Tro, Principles of Chemistry: A Molecular Approach
Metallic Bonding
Metallic Bonds
• Low ionization energy of metals allows
them to lose electrons easily.
• The simplest theory of metallic bonding
involves the metals atoms releasing their
valence electrons to be shared by all the
core atoms in the structure.
9an organization of metal cation islands in a
sea of electrons
9electrons delocalized throughout the metal
structure
• Bonding results from attraction of the atom
core for all the delocalized electrons.
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69
Properties of Metals
• generally high melting points, and all have very
high boiling points
9 All are solids at room temperature except Hg.
9 Melting points of metals generally increase left to right
across a period.
9 Melting points of metals generally decrease down a
column.
Tro, Principles of Chemistry: A Molecular Approach
Metallic Bonding—Model vs. Reality
• This model says that the attractions of the
•
• Solid and liquid states conduct heat and electricity.
9 The ability of metals to conduct electricity decreases
with increasing temperature.
• Solid is malleable and ductile.
• Solid surface reflects light.
Tro, Principles of Chemistry: A Molecular Approach
71
70
•
•
atom cores for the delocalized electrons is
strong because it involves full charges.
In order to melt, some of the attractions
holding the metallic crystal together must be
broken. In order to boil, all the attractions must
be broken.
This model predicts that metals should have
high melting points and boiling points.
Metals generally have high melting points and
boiling points.
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72
Metallic Bonding—Model vs. Reality
• Because the electrons are delocalized,
they should be able to move through the
metallic crystal.
• Since electrical conductivity takes place
when charged particles (like electrons) are
able to move through the structure, this
model predicts metallic solids should
conduct electricity well.
• Metallic solids do conduct electricity
well.
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73
Metallic Bonding—Model vs. Reality
• Metal ions get larger as you traverse down a
column.
• This model says that the attractions of the atom
•
•
•
•
cores for the delocalized electrons will be stronger
when the ions are smaller.
This model predicts that metals with smaller atom
cores should have higher melting points.
This model predicts that the melting points of
metals should decrease down a column.
Melting points of metals generally decrease down
a column.
Li (180.54ºC) > Na (97.72ºC) > K (63.38ºC)
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75
Metallic Bonding—Model vs. Reality
• Atoms emit light when electrons jump from higher
energy levels to lower energy levels.
• This model says that the delocalized electrons will
•
•
•
share a set of orbitals that belong to the entire
metallic crystal.
This model says that the delocalized electrons on
the surface can absorb the outside light and then
emit it at the same frequency.
This model predicts that the surface of a metallic
solid should reflect light.
Metallic solids do reflect light.
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74