Earlier we examined the individual properties of atoms and ions. But, in nature, most
substances occur with these atoms or ions bonded together.
When you examine these
atoms/ions you can see how they influence that types of chemical bonds that are formed, and
the type of bonding ultimately is responsible for the behavior of the compound.
The term chemical bond describes a strong link between the atoms in a molecule or crystal, as
opposed to weaker forces which attract one molecule to another. The energies required to break
a chemical bond lie between 100 and 1000 kJ per mole of bond formed. This is MUCH greater
than the energy required to overcome attractions between molecules (which we will talk about
in CH132!). In nearly all stable substances, each atom is chemically bonded to at least one other
atom – only the inert gases exist normally as isolated atoms.
Why should atoms bond together? Bonding lowers the potential energy between positive and
negatively charged particles – whether the particles are actually ions which retain charges or are
composed of atomic nuclei and the negative electrons spinning around that nucleus. The
electron configurations that we spent time writing helped us to determine properties of the
atoms (will they gain or lose electrons?, what ions will they form?, what species will bond
together?, etc . . .), while the type and strength of the bond will help us to determine the
properties of a substance.
There are three principle types of chemical bonds that you should be aware of: ionic, covalent,
and metallic.
Ionic bonds: an electrostatic attraction between oppositely charged ions. Ionic bonds are
formed when one atom transfers its electrons to another atom. We examined the formation of
ionic compounds being formed from a metal and a non-metal. We also examined how to name
these species. The ions generally arrange themselves into nice ionic structures. For example,
NaCl consists of alternating Na+1 ions with Cl-1 ions. Each ion is surrounded by, and attracted
to, neighbors of opposite charge. There are no molecules in an ionic solid! It is composed of
IONS!! The chemical formula represents the cation-anion ratio formed in the ionic solid.
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Covalent bonding: Electron transfer (gain/loss) is unlikely between two identical atoms (e.g.
H2, O2, N2, F2, Cl2, Br2, I2, S8), or between atoms with similar electron-attracting tendencies. Such
atoms are held together by covalent bonds or metallic bonds. Bonding between non-metals
consists of two electrons shared between two atoms. Using the Wave Theory, the covalent bond
involves an overlap of the electron clouds from each atom. The electrons are concentrated in the
region between the two atoms (termed localized). In covalent bonding, the two electrons shared
by the atoms are attracted to the nucleus of both atoms. Neither atom completely loses or gains
electrons as in ionic bonding. We have previously discussed how to name covalent compounds
with regard to two non-metals bonded together.
There are two types of covalent bonding:
1. Non-polar bonding with an equal sharing of electrons. (DIATOMICS)
NON-POLAR BONDING results when two identical non-metals equally share
electrons between them. One well known exception to the identical atom rule is the
combination of carbon and hydrogen in all organic compounds.
Octet Rule: Elemental atoms generally lose, gain, or share electrons with other atoms
in order to achieve the same electron structure as the nearest rare gas with eight
electrons in the outer level.
The proper application of the Octet Rule provides valuable assistance in predicting and
explaining various aspects of chemical formulas
2. Polar bonding with an unequal sharing of electrons. The number of shared electrons depends
on the number of electrons needed to complete the octet. (DIFFERENT ATOMS)
POLAR BONDING is the unequal sharing of electrons between two different non metal
atoms. A proper understanding of polar bonding is gained by viewing the types of
bonding on a continuum. Ionic bonding is on one extreme with a complete transfer of
electrons forming charged ions. Non-polar covalent bonding with equal sharing of
electrons is at the other extreme. Somewhere in the middle but favoring the covalent side
is polar bonding with unequal sharing of electrons with partial but incomplete transfer of
electrons. A polar bond results when different atoms share electrons. One atom will
attract the bonding electrons more strongly than the other atom and will acquire more
than a half share of these electrons. This leaves the other atom with less than a half share
and makes the electron distribution unsymmetrical. On a time-average basis the electrons
spending more time with one atom cause it to have a partial negative charge. The other
atom deficient in electrons acquires a partial positive charge.
The simplest type of covalent bond is one in which a pair of electrons is shared between two
atomic nuclei. More complex covalent bonds consist of two or even three pairs of electrons.
Formation of a covalent bond results in a shift of the electron density. The shared electrons will
tend to spend more of their time between the two nuclei involved in sharing the electrons. The
negative electron density (electron cloud) is held between the two positive nuclei and results in
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an electrostatic attraction that holds the molecule together and gives the covalent bond its
strength.
Polar Covalent Bonding: unequal sharing of electrons
Again, the electrons move around the nuclei with the electrons spending the majority of the
time near the more electronegative element. This generates a partial negative charge near the
more electronegative element and a partial positive charge near the less electronegative
element.
Non-Polar Covalent Bonding: equal sharing of electrons
Metallic bonding: atoms in a metal are held together by metallic bonding in which the valence
electrons are not confined to individual atoms or to individual bonds but are allowed to flow
freely through the metal (termed delocalized). Metallic bonding is sometimes described as a
lattice of positive ions bathed in a sea of mobile electrons; these mobile electrons provide metals
with their conductivity and other distinctive properties.
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The strength of the bonding depends on the number of electrons released by each atom to the
electron sea, and the charge density of the cations (i.e. the more valence electrons, the stronger
the bonding is likely to be).
A method used to indicate the location of the valence electrons for an atom. Remember that
atoms are using only their valence electrons to participate in the formation of a chemical bond.
First you must determine the number of valence electrons a species has:
1. For main group elements (nsxnpy) it is the number of electrons in the HIGHEST principle
quantum level. It also happens to be the group number! Group 1A has 1 valence
electron (ns1) Group 2A has 2 valence electrons (ns2), Group 3A has 3 valence electrons
(ns2np1), Group 4A has 4 valence electrons (ns2np2), Group 5A has 5 valence electrons
(ns2np3), Group 6A has 6 valence electrons (ns2np4), Group 7A has 7 valence electrons
(ns2np5), Group 8A has 8 valence electrons (ns2np6). PLEASE remember that even
though you might sandwich in the d block electrons, main group elements ONLY have
nsxnpy valence electrons!!
Calculate the number of valence electrons for Se using its electron configuration:
Calculate the number of valence electrons for I using its electron configuration:
Calculate the number of valence electrons for Sn using the electron configuration:
2. For transition group elements (nsx(n-1)dy) it is the number of electrons in the HIGHEST
principle quantum level and sometimes the n-1 d electrons as well. It is impossible to
predict the number of valence electrons from looking a the periodic table, either you
must memorize the ion types that are formed, or sometimes it can be determined by
examining the electron configurations too. It is a little more ambiguous. Generally
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speaking, you will be given charges to memorize and thus can determine which electrons
from the e- configuration are involved as the valence electrons.
Knowing that Mn can form +7, +6, +4, +3, and +2 ions indicate the valence electrons for Mn
The Lewis symbol for an atom consists of the atomic symbol surrounded by a number of dots
equal to the number of valence electrons:
Rules for writing Lewis dot structures for main group elements:
1.) Write the elemental symbol for the element. This represents the element and the inner or
core electrons
2.) Notice the number of valence electrons that the element has (indicated by the group
number)
3.) Place the valence electrons as dots on the outside of the element. There will be a
maximum of 4 locations for the electron. Again, electrons want to be single until it
becomes necessary to pair them. It does NOT matter which electrons you pair up first.
Notice Nitrogen, you could have paired up the electrons on the left instead of the
electrons at the top, or the electrons on the right, or the bottom instead. N just has one
pair of electrons and the others are singles.
4.) The number of dots provides information about the element: for metals it is the number
of electrons that the atom will lose to become positive and to achieve a noble gas like
electron configuration. It is also the number of electrons that will be lost when forming
an ionic bond (thus the charge that the metal will adopt in an ionic bond). For a nonmetal, the number of unpaired dots are the number of electrons that will become paired
when that atom either gains electrons or forms a covalent bond. Thus, the number of
unpaired dots indicates the charge of the anion, or the number of covalent bonds that the
atom will form.
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From the table above, we see that C will form four bonds, Phosphorus will form 3 bonds, Cl will
form 1 bond, and Ar will form 0 bonds.
Draw Lewis dot diagrams for Se, Br, K, and Ca
The central idea behind the ionic bond is the gain/loss of electrons; that a metal will lose its
electrons and transfer them to a non-metal which will gain those electrons. For nearly every
main-group element that forms an ion, the ionic electron configuration has a filled outer level: 8
electrons (or 2 electrons . . .), the same number as the nearest noble gas. This idea, of achieving a
stable ionic state was used to develop the octet rule: when atoms bond, they lose or gain
electrons to attain a filled outer shell of eight electrons (or two . . .).
Electrons lost from a metal and then gained by a non-metal can be depicted using orbital box
diagrams: The electrons lost by the metal will be added to the empty or half-empty boxes of the
non-metal
Draw the orbital box diagram for Ca: Draw the orbital box diagram for O. Show the movement
of electrons from Ca as it loses its valence electrons to O as it gains those electrons and indicate
where in the orbitals the electrons “go”:
Ca: 1s22s22p63s2: valence electrons are the 3s electrons
O: 1s22s22p4: valence electrons are 2s22p4
O becomes:
as it takes the electrons from the 3s shell of Ca.
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The 3s subshell of calcium will now be empty, but the electron configuration for the Ca+2 ion is
isoelectronic with Argon, the nearest noble gas.
In all cases for ionic bonding, the number of electrons lost by the metal must indicate the
number of electrons gained by the non-metal.
Ionic compounds tend to be hard, rigid, and brittle due to the powerful attractive forces that
hold the ions in specific positions through the crystal (remember the ions arrange themselves
positive to negative in a nice crystal array – it is highly organized!!). Moving the ions out of
position requires a lot of energy, a lot of force to overcome the attractions. Thus, ionic
compounds tend not to dent or bend. Instead, they break or shear. Most ionic compounds do
not conduct electricity in the solid state but DO conduct electricity when molten or in the liquid
state (dissolved in water). The ions become free to move about and thus are able to carry an
electric current). It takes extremely high temperature to melt or boil ionic compounds because
that would require freeing the ions from their set and happy  positions inside the crystal
lattice. The attractions are very strong, giving rise to boiling points in the 1000 oC!
The number of known covalent compounds greatly outnumbers the number of known ionic
compounds. Simple molecules such as H2, N2,O2, H2O, NH3 are covalent, and so are more
complex substances like sugar, rubber, Teflon, silk, silicates, which make up most of the rocks in
the Earth’s crust contain covalent networks of oxygen and silicon atoms. Living matter consists
of even more intricate covalently bonded molecules made up of carbon, hydrogen, oxygen,
sulfur and phosphorus (DNA, RNA, proteins, amino acids . . .) Metals, especially the transition
metals also form covalent bonds. There are only 22 nonmetals on the periodic table yet over
90% of the known compounds consist solely of these elements!! Sharing electrons is the
principle way that atoms chemically interact with one another.
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Forming a covalent bond
Why do atoms exist bonded together? Imagine two isolated H atoms, some distance apart from
one another (1). At this distance neither atom is really aware of the other . . .As the distance
between the atoms decreases, they move closer together (2), eventually, they get close enough
that the nucleus starts to attract the electrons of the other atom (like gravitational pull). The
attraction of the nucleus to the electrons pulls the atoms closer and closer together, which
lowers the potential energy of the system (3) (remember nature wants to be at the LOWEST
energy state possible because it’s LAZY!!). Eventually, the atoms get so close that the electron
clouds start to overlap and the electrons repulse one another, also, the nucleus of each atom
starts to get too close together (4) and the atoms repel one another back to the ideal distance and
lower energy state seen at position 3. At this position, the optimal distance between the nuclei
is achieved and is known as the bond length. It is also the lowest energy state for the two
atoms. When the bond is formed, energy is released (-432 kJ). When the bond breaks, energy
must be used to overcome the nucleus-electron attractions and it takes +432 kJ of energy to
break the H-H bond. This mutual attraction of the nucleus with the electrons from another
atom constitutes the covalent bond and holds the molecule together.
In covalent bonding, unlike ionic bonding, the electrons are shared between the two atoms.
Again, each atom wants to achieve its full outer level of electrons (the octet). In order to achieve
this, each atom will count the electrons surrounding it, as if they only belonged to themselves.
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Thus, the electrons simultaneously participate in outer levels of both atoms in the bond. This is
known as the shared pair, or shared pairs, or bonding electrons. When two electrons are shared
between atoms, this is known as a bond, and is drawn with a single
between atoms. The
electrons that are present as part of an atom’s valence shell but do NOT participate in bonding
are known as the lone pairs or unshared pairs.
Example: Show the shared electrons of Cl2 indicating the electrons that are involved in the bond
and the electrons that are lone pairs.
Step 1: Write the element with its valence electrons around it
Step 2: Place the atoms next to one another, so that each fulfills the octet
Step 3: Account for the rest of the electrons
Above is a covalent molecule represented by a Lewis structure diagram. There are several types
of bonds to be aware of which will be used when drawing Lewis structures:
1.) Single bonds: one pair of electrons shared between two atoms
2.) Double bonds: two pairs of electrons shared between two atoms
3.) Triple bonds: three pairs of electrons shared between two atoms
The covalent bond that is formed results from the nucleus attracting the electrons balanced out
with e-/e- repulsions and nucleus/nucleus repulsions. Each atom has its own size, which was
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discussed previously. The strength of the bond depends on the magnitude of the attraction
between the nuclei bonded to the shared electrons. Bond Energy (BE) (bond strength) is the
energy required to overcome this attraction, or the energy required to break that bond. It takes
energy to break the bond (endothermic), thus Bond Energy is defined as being positive. This
bond energy is the difference in energy between the separated atoms and the bonded atoms (the
potential energy difference between points 1 and 3 in the diagram above). The same amount of
energy that is required to break the bond is given off when then bond is formed (exothermic).
Since energy is released when the bond is formed we say that Bond Formation Energy is
negative.
A
A
B
+
A
B
A
+
H:
B
B
BE  0
H: BE  0
Bond energies depend on the atoms that are involved in the bonding, specifically think about
each atom’s electron configuration which is directly related to the size of each atomic radius for
each atom. Thus, each type of bond has its own bond length and its own bond energy. Even
worse . . . the energy of a given bond varies slightly from molecule to molecule, even within the
same molecule, so the values for bond energy given are average values. Bonds between smaller
atoms tend to be stronger (therefore harder to break) than bonds between larger atoms.
Comparing Cl-Cl to I-I we know that from our atomic size trend, atoms at the top of a group are
smaller than atoms at the bottom of the group. This means that Cl is smaller than I. Thus, two
Cl atoms can get closer together than two I atoms. This makes the bond length shorter and
shorter bonds are stronger bonds, therefore harder to break. This is a general rule, however, as
in all aspects of chemistry, there are exceptions (see tables below).
A bond also has a length associated with it. The bond length is the distance between the nuclei
of two bonded atoms at a minimum energy point. Remember, nature is lazy, so the bond length
will be a balance between attraction of the nucleus with the electrons and repulsion of the
electrons and repulsion of the two nuclei. More simply put, it is the equilibrium distance
between the centers of two bonded atoms. It is the distance in a molecule where the net
attractive and repulsive forces are at zero.
Again, the bond lengths given are average values for a given bond in different substance. Bond
length is related to the sum of the radii of the bonded atoms. Most atomic radii are then
calculated from measured bond lengths.
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There is a strong relationship between bond length and bond strength. The shorter the bond,
the stronger the bond. We previously discussed multiple shared pairs of electrons (we called
them double and triple bonds). Examining the tables below, compare the bond energies with
to the bond lengths.
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12
Covalent bonding proposes that the electrons are shared between the atoms which results in a
strong, localized bond holding the atoms together in a molecule. But if they are so strong, then
how do we explain the properties that covalent compounds have? Many gases are covalent
molecules (methane gas), as are liquids (water, organics, such as octanol, hexanes, etc . . .) or
even low melting point solids such as paraffin wax. Covalent bonds are strong – it takes
anywhere between 200-500 kJ/mole energy to break them, but why do they melt and boil at
such low temperatures? The reason is that covalent molecules have both strong covalent bonds,
but also intermolecular forces that attract the molecules together. These intermolecular forces
are weaker than the covalent bonds themselves and thus, are easily overcome to turn liquids
into gases or solids into liquids.
There are covalent substances that do not rely on intermolecular forces to hold them together.
These are known as network covalent solids. These compounds are held together by covalent
bonds which go throughout the sample, in three dimensions. Thus, their properties will reflect
the strength of only the covalent bonds that hold the atoms together. Two examples of
extremely strong covalent bonds are diamonds and quartz.
Covalent compounds are poor electrical conductors, even when melted or when dissolved in
water since electrical current is carried by either mobile electrons or mobile ions, and covalent
compounds do not have free electrons. Their electrons are localized in bonds or are localized as
lone pairs on the atoms. Also, no ions are present which would allow for mobility of the
electrons.
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We have seen example after example of “exceptions” to the rule. Now we get to see more! We
have discussion ionic bonding as being completely 100% gain/loss of electrons. We have
discussed non-polar covalent binding as being completely 100% equal sharing of the electrons
in a bond. However, nature mostly exists on the continuum between those two extremes, where
the bond is neither 100% ionic nor 100% pure covalent. We termed these molecules polar
covalent molecules. They exist in-between ionic and pure covalent bonds.
Often abbreviated  and defined as the ability of a bonded atom to attract the shared electrons.
It is NOT the same as electron affinity, which is defined as the addition of 1 mole of electrons to
1 mole of gaseous state atoms or ions. With electron affinity we are talking about the GAIN of
an electron (or two or three). With electronegativity we are talking about still sharing the
electrons, but that one of the atoms attracts the electrons more than the other atom. Although
they are different, many elements that have a high EA value (realllllly want the electron) also
have a high  value. Linus Pauling came up with his electronegativity scale, with the MOST
electronegative atom being F and given a value of 4.0. It was noticed that many bonds took
more energy to break them. This is due to this extra charge contribution that results when two
atoms of different electronegativities are bonded together. The attraction of the partial charges
that were previously mentioned when we discussed polar bonds, means that there is an “extra”
attraction that must also be overcome in order to break the bond. A certain amount of energy
must go into disrupting the charge attraction, the more energy must be used to break the
covalent bond. Pauling discovered that it took almost 270 kJ/mole extra energy to break the
H-F bond than expected or calculated. The electronegativity values are relative values and
based off Pauling’s highest assignment which went to F.
Bond polarity results from a difference in electronegativity between the bonded atoms. The
more electronegative atom draws the electron density towards itself, away from the less
electronegative atom.
With some exceptions (of course!!) electronegativities increase from the left to right on the
periodic table. Thus, non-metals are more electronegative than metals. Highly electronegative
atoms tend to want to hold onto the electrons that you do have and thus have high ionization
energies (they do NOT want to form positive ions!!). They also tend to acquired electrons very
easily and thus have very negative electron affinities (they DO want to form negative ions!!)
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Electronegativity and Oxidation Number
Electronegativity can be used to determine an atom’s oxidation number, or the charge that the
atom will adopt when it becomes an ion. For the metals, we determined the oxidation numbers
simply by looking on the periodic table, Group 1A always has a charge of +1, Group 2A has a
charge of +2, etc. . . But what about covalent compounds? Determining oxidation numbers
becomes a bit more difficult.
1.) The more electronegative atom involved in the bond is assigned all the shared electrons,
while the less electronegative atom is assigned none.
2.) Each atom involved in the bond is assigned all of its lone pair electrons
3.) The oxidation number is given as the # of valence electrons – (# shared + # lone pairs)
We will do more with this after we learn about drawing Lewis structures for compounds, just
keep these rules in mind for a bit longer . . .
Electronegativity and Bond Type
The difference in electronegativity between two atoms tells us something about the type of bond
that will be formed. Atoms with identical electronegativities were called pure covalent
molecules – because they share the electrons equally. Examining their electronegativities
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explains why – neither atom in the bond wants the electrons any more than the other atom –
their electronegativity values are identical! If their electronegativity values are different, then
the sharing will be unequal, one atom in the bond will selfishly hog the electron and the bond is
said to be polar. In general, the greater the electronegativity difference between the two atoms,
the more polar the bond.
When the electronegativity difference is zero, as in H2, O2 and N2, the electrons are
perfectly shared, equally between each atom and the bond is non-polar.
When the electronegativity difference is between 0 and 1.7, the bond is polar, but
predominantly polar covalent.
When the electronegativity difference is greater than 1.7 the bond is ionic
Unfortunately, again, these rules are not infallible. For example, the electronegativity difference
between Li-H is less than 1.7 but the bond happens to be ionic. Lithium hydride is a crystalline
compound that melts at 688oC and conduct electricity (as ionic compounds do and remember
covalent compounds do not – because there are no ions!!). The electronegativity difference
between H-F is 1.8 but the bond is polar covalent. Hydrogen Fluoride is a liquid at room
temperature and a poor conductor of electricity – indication that no ions are present in solution.
When atoms with different electronegativity values form a bond, the electrons are not shared
equally between them. This results in a partially positive region and a partially negative region
on the molecule.
-
+
This type of bond is called a polar covalent bond and is depicted by an arrow pointing in the
direction of the most electronegative atom or by + - indicating the partial charges that result
when electrons are shared unequally. Another term to describe a molecule whose ends carry
equal but opposite charges is a dipole. By knowing the electronegativity values OR by knowing
the general trend of electronegativity for atoms on the periodic table one can determine the
direction of the bond polarity.
The existence of
partially ionic.
electronegativity
between the two
bond becomes.
partial charges means that a polar covalent molecule behaves as if it were
Partial ionic character in a bond is directly related to the difference in
between the two atoms. The greater the difference in electronegativity
atoms the larger the partial charges become and the more ionic in nature the
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Various attempts to classify bond types into ionic, polar covalent, and nonpolar have been made
but they all use arbitrary cutoff values when examining the difference in electronegativity
between the atoms. Another method calculates the percent ionic character of a bond. A value
of 50% ionic character is chosen to divide ionic from covalent substances. How % character is
determined is not really important, but what is, is that in reality, there is no bond that is 100% in
ionic character. Ultimately, sharing occurs to some extent in every bond, even between a metal
and a halogen.
A metal and a non-metal interact and form a bond by electron transfer. A non-metal and a nonmetal interact and form a bond by sharing the electrons between them. Thus there is a gradual
shifting across the periodic table from ionic to polar covalent to non-polar covalent bonding.
NaCl to MgCl2 to AlCl3 to SiCl4 to PCl3 to S2Cl2 and finally to Cl2 shows the transition from ionic
to nonpolar covalent. NaCl is ionic, it is a hard brittle solid, it conducts electricity when
dissolved in water, and has a high melting point. MgCl2 is still considered to be ionic ( =1.8)
but it has a lower melting point and lower conductivity. AlCl3 has a  value of 1.5 which
actually falls in the polar covalent bond. It actually is not a lattice of Al +3 and Cl-1 ions, instead
it consists of linked Al and Cl atoms (but still named like an ionic compound). AlCl 3 has a
lower melting point and does not conduct electricity well – a sign that there is a lack of ions
present. SiCl4 does not conduct electricity and has a very low melting point. It is still a polar
covalent bond. PCl3 is not as polar in nature, does not conduct electricity and has a very low
melting point. S2Cl2 approaches the bottom end of polar covalent bonds ( =0.5). It also does
not conduct electricity. Cl2 is purely nonpolar covalent bonding between identical atoms and
exists as a gas at room temperature. As we go across the period there is a shift from compounds
as ionic solids to molecules what exist as a gas.
Metals are known for transferring their electrons to non-metals and creating ionic bonds. They
can also share their electrons forming a covalent bond. In the electron sea model of metallic
bonding the valence electrons from the metal surround the positive metal nuclei in a delocalized
cloud and can travel throughout the substance. The positive nuclei are submerged in the sea in
an orderly manner.
The metal ions are not held in place as rigidly as in an ionic solid as the electrons are mobile.
The electrons in the delocalized sea are shared amongst all the positive nuclei in the substance.
The metal is held together by the mutual attraction of the metal cations with the mobile
electrons.
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Most metals are solids with moderate to high melting points (solids to liquids) and much higher
boiling points (liquids to gas). Metals typically bend or dent rather than break or shatter (ionic
compounds), many can be flattened into sheets (malleable), can be pulled into thin wires
(ductile), and they conduct electricity and heat better than ionic or covalent compounds in the
solid and liquid states.
Two features of the metallic bonding account for this: the electron sea allows for the movement
and shifting of the nuclei and the sea allowing the solids to bend and dent. The moderately
high melting points result from the attraction from the metal nuclei to the electrons, but shifting
from the solid state to the liquid state does not require the charges to be separated from one
another. The high boiling points are also related to the strength of the bonding or the attraction
of the metal nuclei with the sea of electrons. Boiling requires the positive metal nuclei to
separate from the negative electrons. Gallium is an ideal example of metal behavior. It will
turn from a solid to a liquid slightly above room temperature (29.8oC) but turning from a liquid
to a gas requires a temperature in excess of 2400oC.
Metals can be dent or bent or pulled into wires as the electrons are not locked into place this
allows the metal ions to slide past one another and generate a new lattice position and the
electron sea moves with them. The metal ion cores do not get so close together that they repel
one another and cause the solid to crack (as in ionic solids).
Metals are good conductors of electricity because of the mobile sea of electrons. Electrons can
flow freely from the electrical source through the metal. If the metal contains an irregular
arrangement of ions in the lattice, the conductivity of the metal is reduced. For example, if
another metal is mixed with copper it reduces the conductivity of the metal.
We have examined the individual properties of atoms and ions and substances with these
atoms or ions bonded together. We can now examine the three-dimensional reality of what a
molecule looks like with all the atoms arranged in some specific orientation in space. Whether
we are studying biology or chemistry or physics it is the 3 dimensional organization of atoms in
space that governs chemical reactions, properties of synthetic materials and all of life.
The first step to discovering the 3 dimensional shape of a molecule is to write the Lewis
structure for the entire molecule. A Lewis structure is a two-dimensional structural formula
which shows valence electrons as dots and bonds as
which show each atom and how
those atoms are bonded together. Drawing Lewis structures generally will mean following the
octet rule: placing and arranging the atoms such that each atom is surrounded by 8 electrons.
However, some atoms, like Hydrogen, do not want to be surrounded by 8 electrons. There is
another rule, known as formal charge, which sometimes will be followed instead.
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Using the Octet Rule to Write Lewis Structures:
1.) Find the total number of valence electrons which will be involved in the structure by
adding up the number of valence electrons for each atom in the molecule or ion; adding
electrons for negatively charged species or subtracting electrons for positively charged
species.
For main group elements the number of valence electrons is determined by the
group number.
For transition metals the number of valence electrons is the highest n value and
also sometimes the (n-1)d electrons as well.
2.) Draw a skeleton structure of the molecule by placing the LEAST electronegative atom in
the center and arranging the atoms with single bonds connecting them together.
Hydrogen can only form 1 bond, therefore it will never be the central atom. Oxygen
generally will only have a maximum of 2 bonds – either two single bonds or 1 double
bond. It is ok for oxygen to have just one bond, but it is very rare for oxygen to
participate in a structure and have 3 bonds (either three singles or a double and a single
bond). Carbon will always always always have 4 bonds, NO lone pairs and NO 5
bonded carbon atoms!! NEVER NEVER NEVER!!
3.) Distribute the remaining valence electrons around the outside atoms as lone pairs, first
attempting to satisfy the octet rule for all the atoms. If you have remaining electrons,
then place them around the central atom as lone pairs. Keep in mind you may need to
form double or triple bonds to satisfy the octet rule. The double and triple bonds come
from involving what would be a lone pair into a bond.
Draw Lewis structures for NH3, CH4, CO2, and HCN
Calculate the number of valence electrons for NH3:
N=
H=
total number of valence electrons =
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Calculate the number of valence electrons for CH4:
C=
H=
total number of valence electrons =
Calculate the number of valence electrons for CO2:
C=
O=
total number of valence electrons =
Calculate the number of valence electrons for HCN:
C=
H=
N=
total number of valence electrons =
Keep in mind that there are several structures sometimes that show that you have followed all
the rules for drawing Lewis structures. IF you have obeyed the rules then you have drawn a
valid Lewis structure. Generally speaking, when drawing Lewis structures you will be asked to
do so following the octet rule or formal charge. Sometimes you will violate one rule to satisfy
another. That is ok, as long as you are following the rule that you were asked to follow.
What this means, is that following the octet rule there is actually another Lewis structure that
could have been drawn for CO2:
O C O
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Notice that when you started drawing the Lewis structure for CO2 and HCN that there are not
enough electrons on the central atom such that it becomes very difficult to figure out how to get
that central atom its octet. This is usually a sign that a double or triple bond must be used.
Remember that when asked to obey the octet rule you need to find a way to do so whenever
possible (remember some atoms do not want to – like Hydrogen!!)
Two or more Lewis structures, each with the same arrangement of atoms and the same number
of electrons pairs can often be drawn for a molecule or ion (see CO2 above!). For example, draw
the Lewis structure for NO3-1 (no O-O bonds).
-1
-1
O
O
N
-1
O
O
O
N
O
O
O
N
O
Notice that for ions, the Lewis structure is surrounded with brackets and the charge is indicated
as a superscript. The double headed arrow does not mean that the molecule changes back and
forth between one structure to another, it means that the actual structure of the molecule is
intermediate to those represented by these contributing structures. This is the proper Lewis
structure depiction for ions and should be used.
When drawing resonance structures you will want to pretend as if the atoms are frozen in
space. Notice the atoms and their location remains the same. Instead of moving atoms, we
move the bonding electrons and the lone pairs. If it helps, number the atoms in your
drawing and then shift the bonds around.
Remember back to bond energies: were double bonds stronger or weaker than single bonds?
Were they longer or shorter than single bonds? Double bonds were stronger and they were also
shorter than single bonds. Since double bonded atoms are more tightly bonded than singly
bonded atoms, the above Lewis structures would suggest that one of the bonds in the nitrate ion
is shorter and stronger than the other two bonds. Experimental evidence however, shows that
the nitrate ion is completely symmetrical – meaning all of the bonds are equally strong and each
bond is 124 pm in length. This means that the three bonds share the double bond character
equally. This double bond is an example of a delocalized bond, once in which certain bonding
electrons are spread out over several bonds and help to bond all of them. Delocalization of
bonding electrons can not really be shown in a single Lewis structure so we draw all three.
These structures are called resonance structures and each contribute to the molecules overall
structure known as the resonance hybrid. It is sometimes drawn with a dashed line indicating
the atoms which share these delocalized electrons. Electrons delocalization diffuses the electron
density over a greater area which reduced electron-electron repulsions and thus stabilizes the
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molecule. What is important to remember is that the hybrid is a blend of all resonance
structures that can be drawn and contribute to the overall TRUE structure. Individual
resonance structures do NOT exist. They are drawn for our method of bookkeeping in our
attempt to freeze electrons in space and time. Think of it this way. The overall resonance
hybrid is like a mule. What makes a mule? A horse and a donkey. But a mule is always a mule
100% of the time. It is not a donkey somedays and a horse on others. Neither are resonance
hybrids! They are the mules and the structures that we draw to capture the overall essence of
the molecule/ion are the donkeys and horses.
Resonance structures for atoms ONLY differ in the location of the electrons NOT the atoms – if
you start moving the atoms around this is NOT a resonance structure!!
All resonance structures that represent a given molecule MUST have the same arrangement of
atoms and the same number of electrons (either in bonds or as lone pairs). However, all Lewis
structures that you can draw for a particular molecule that satisfy the requirements do not
necessarily contribute equally to the stable arrangement of atoms. The actual molecule, the
resonance hybrid, is more stable than any of the individual Lewis structures that you could
draw for the molecule. Therefore, the most stable Lewis structure is the one that contributes
the most to the overall structure of the molecule. Less stable structures, even though they might
be valid Lewis structures (meaning you followed the octet rule) might actually not contribute at
all to the resonance hybrid.
One way to determine which Lewis structure contributes more the other overall resonance
hybrid is to determine the formal charge of each atom in the molecule.
Two rules to follow when examining formal charge:
1.) The most stable structure has the lowest formal charge
2.) Structures in which the adjacent atoms have formal charges of the same sign are
especially unstable.
Examine the two Lewis structures that were drawn for CO2.
A
B
O C O
(1)
(2)
O C O
(1)
(2)
When calculating formal charge, you want to pretend as if the atoms are frozen in time again.
Number the atoms that are identical. It is not necessary to number the carbon atoms since there
is only one present, but since there are two oxygens, you want to be clear about which oxygen
atoms you are talking about.
Calculating formal charge: # valence electrons – the number of electrons that the atom “sees”
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An atoms sees ALL of its own lone pairs and ½ of the electrons involved in a bond.
Calculating formal charge for structure A:
O1: has 6 electrons as lone pair electrons and 1 electron from the bond
O1: 6 valence – 7 electrons = -1
C: has 4 electrons from the bonds
C: 4 valence – 4 electrons = 0
O2: has 2 electrons from the lone pair and 3 electrons from the bonds
O2: 6 valence – 5 electrons = +1
When you sum the formal charges for a neutral molecule they better add up to zero!!
-1 + 0 + 1 = 0 !!
Calculating the formal charge for structure B:
O1: has 4 electrons from the lone pairs and 2 electrons from the bond
O1: 6 valence – 6 electrons = 0
C: has 4 electrons from the bonds
C: 4 valence – 4 electrons = 0
O2: has 4 electrons from the lone pairs and 2 electrons from the bond
O2: 6 valence – 6 electrons = 0
Summing the formal charges:
0+ 0 + 0 = 0 !!
There is one more resonance structure that can be drawn for CO2. Can you see it? Draw it and
calculate the formal charge for that structure as well. Notice you will get the same charges as
Structure A (except O1 will have the +1 and O2 will have the -1 charge associated with it).
Given a situation like this where one resonance hybrid/structure has a formal charge of 0 and
the others actually have formal charge values associated to them means that the structures with
formal charges will contribute much less to the overall structure of the molecule. In the case of
CO2 their contributions to the structure is negligible.
Electronegativity and Oxidation Number
Electronegativity can be used to determine an atom’s oxidation number, or the charge that the
atom will adopt when it becomes an ion. For the metals, we determined the oxidation numbers
simply by looking on the periodic table, Group 1A always has a charge of +1, Group 2A has a
charge of +2, etc. . . But what about covalent compounds? Determining oxidation numbers
becomes a bit more difficult.
4.) The more electronegative atom involved in the bond is assigned all the shared electrons,
while the less electronegative atom is assigned none.
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5.) Each atom involved in the bond is assigned all of its lone pair electrons
6.) The oxidation number is given as the # of valence electrons – (# shared + # lone pairs)
Calculate the Oxidation number of each atom in CO2
Oxygen is the more electronegative atom in the molecule. Examining the true Lewis structure
we have :
O C O
(1)
(2)
Oxygen 1 is more electronegative than the central C atom, so we pretend that the oxygen gets its
lone pair electrons AND all the electrons in the bonds between itself and carbon:
O1: 4 lone pair electrons and 4 electrons in the bond = 8 total electrons
The same is true for O2
Oxidation number for O1 (and O2):
O = 6 valence electrons – 8 electrons = -2
For carbon, since it is less electronegative than the atoms it is bonded to, it does not get ANY of
the electrons in its bonds.
Thus the oxidation number for carbon is 4 valence – 0 = +4
Notice each Oxygen is a -2 (thus a total of -4) and the carbon atom is a +4 leading to a neutral
molecule. If you ever end up with extra charge somewhere it is a good sign you have either
calculated something incorrectly or your Lewis structure is not correct!
The octet rule is a very useful guide, and in general, atoms tend to follow it, but as in all aspects
of chemistry, there are exceptions. Sometimes central atoms have fewer than 8 electrons
(electron deficient atoms) and some central atoms can have more than 8 electrons (expanded
octets). We will not deal with radicals, which means a lone electron at this time.
More than 4 Electron Pairs
The atoms in the second row (period) have only 4 orbitals in their valence shells: 1 2s orbital
and 3 2p orbitals (2px, 2py, and 2pz). These atoms only have room for electron pairs to fill these
orbitals and these orbitals only. Thus, they achieve the s2p6 configuration and have a full octet.
Second period atoms B, C, N, O, F, and Ne never exhibit octet expansion. However, once the atom
gets enough electrons to enter into the 3rd level we know from writing electron configurations
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that this is the first indication of the 3d subshell. Even if the atom does not have any electrons in
the 3d subshell, it has the d orbitals if it has the 3s and 3p! They are a package deal, if you get
the 3s and 3p you have the 3d – it might be empty, but it is there. It is because of this 3d
subshell that a particular atom can accept more than 8 electrons – it has a place to put them –
into the 3 d orbitals! Remarkably, some molecules with expanded octets are extremely stable, so
this situation is ok! SF6 is an example of a molecule that exists in a stable form with an
expanded octet. It is through expanded octets that some of the noble gases have been found to
covalently bond with other atoms XePtF6 was synthesized in 1962, and XeF6 was synthesized
shortly thereafter.
Octet expansion generally occurs in a molecule to minimize formal charges. This means that
atoms in the 3rd row will often form double and triple bonds with atoms, expanding their octets
but eliminating formal charges on the atoms. Sulfur and Phosphorus can accommodate up to 12
electrons and Iodine as many as 14.
Fewer than 4 Electron Pairs
Beryllium, Boron, and Aluminum form covalent compounds, however, these atoms cannot
complete their octets without acquiring a negative formal charge:
Example: Draw a Lewis structure for BF3 which satisfies the octet rule. Calculate formal
charge for the atoms:
Example: Draw a Lewis structure for BF3 which satisfies formal charge
The main way that electron deficient molecules interact chemically is to attain an octet by
forming additional bonds in chemical reactions. Usually, when this happens, an entire pair of
electrons must be donated to the atom to form the bond (notice B is using its three available
valence electrons in the bonds to F – additional bonds would have to result from another atoms
donating a pair of electrons – we’ll get to that later!!)
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Now that we can draw Lewis structures, we can go a step farther and determine the shape of
the molecule – generally about the central atom. Sometimes, however, molecules do not have a
defined “central” atom – as we saw in the Lewis structure lab when you were asked to draw
CnHx structures – there were several carbons – and ultimately, you must examine all of them to
determine the shape about each.
There are two different ways to designate the shape of the molecule. The first way is to take
into account the presence of the lone electron pairs in the molecule (known as VSEPR or the
electronic shape) the second is to pretend that any non-bonded electrons do not contribute to
the shape of the molecule (known as the molecular structure).
The two methods are definitely related to one another and you must be able to determine the
shape of a molecule from its Lewis structure. Therefore, you must know the shapes!!
Ultimately it is easiest to memorize or learn the shapes attributed to VSEPR and then deduce the
molecular shape.
The Lewis structure is the flat, 2 dimensional representation of a molecule that we know has a 3
dimensional shape. To construct the shape, we start off using VSEPR – Valence Shell Electron
Pair Repulsion theory. It is exactly what it sounds like. VSEPR takes into account lone pair or
non-bonding electrons and states that they DO contribute to the overall shape of the molecule.
Lone pairs have mass – albeit small, and they take up space. So just because an atom has a lone
pair of electrons does not mean that the shape of the molecule behaves as if they were not there.
A stable structure will result which takes into account the size of the bonded atoms and the lone
pairs of electrons. Ultimately, from the lab, you learned that the angles for the different shaped
molecules stay fairly constant – but each molecule has some wiggle room – and it will lose those
few degrees to maintain the lowest energy distance between the electron pairs and the bonded
atoms. Ultimately, it is all about the lowest energy state and the atoms and lone pairs will
arrange themselves around the central atoms such that they are as faaaar apart as possible
(repulsion!)
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Ultimately, for VSEPR, it does NOT matter if there is a lone pair or an actual atom bonded to the
central atom, the shape will be the same. So to determine the shape of a molecule using VSEPR,
count up the number of bonded atoms and lone pairs ON the central atom (the central atom
does NOT get counted!!!!!!) if the number is 2 then the molecule is linear, if it is 3 then it is
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trigonal planar, if it is 4 it is tetrahedral, 5 it is trigonal bipyramidal, if it is 6 then the molecule is
octahedral. Once again, VSEPR includes both non-bonding electron pairs AND bonded atoms!!!
Molecular shapes are determined off the VSEPR shape, however, they do NOT take into account
the presence of the lone pair of electrons when determining the overall shape of the molecule.
Think of the molecular shape as being a photograph of the molecule. What are the chances that
you would “see” the electrons in the photograph? Not very good. Sure they contribute to the
shape that the atoms adopt around the central atom, but they will appear as invisible object in
the photograph. All we will see in our picture are the actual atoms. Thus, we must classify the
shape as something “new” compared to the VSEPR shapes (which only account for 5 different
shaped molecules!)
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Let’s examine some molecules and determine their VSEPR shape and their molecular shape.
Draw Lewis structures and determine the shapes (VSEPR and molecular) for: BeCl2, SO3, SO2,
CH4, NH2Cl, H2O, PBr5, SF4, ClF3, XeF2, SF6, BrF5, XeF4.
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Practice drawing Lewis structures and determining shape about the central atom for each
structure – both the VSEPR shape and the molecular shape.
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Previously we have examined whether or not a bond was polar simply by looking for a
difference in electronegativity between the two atoms bonded together. But just because a
molecule CONTAINS polar bonds does not mean that the molecule as a whole is polar. Polar
molecules influence the melting and boiling points of the compound, solubility, chemical
reactivity and biological function.
Remember we talked about a polar bond resulting in partial charges. For simple molecules that
only have one bond (e.g. H-F) a difference in electronegativity results in a polar bond AND a
polar molecule. And remember these charges act as forces the must be overcome with extra
energy in order to disrupt the interaction which results in higher melting and boiling points.
Determining molecular polarity can be tricky sometimes. In order for a molecule to be polar
overall there must be a difference in electronegativity values between the atoms in question and
the molecule must by asymmetrical. Generally speaking, for all molecules that do NOT have
lone pairs, determining polarity follows the rules. When there are lone pairs in the molecule, it
actually becomes a bit more difficult and ambiguous to determine (for example, NH 3 is a polar
molecule with a dipole moment of 1.47D while NF3 has a very small dipole moment and is not
nearly as polar at 0.24 D) Thus, we will be determining if a molecular is polar only if it has no
lone pairs.
Compare CCl4 to CHCl3 and determine if the molecules are polar (CCl4 is not – symmetrical
shape, CHCl3 is polar as it is asymmetrical)
Examine BCl3 is the molecule polar? (no, symmetrical shape)
Examine HCN, is the molecule polar? (yes, it is linear but the difference in electronegativity
pulls the electrons towards the N atom)
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