Chapter 6
The Chemical Bond
Some questions
• Why do noble gases rarely bond to other
elements?
• How does this relate to why the atoms of
other elements do form bonds?
• Why do certain elements combine to form
ionic and others molecular compounds?
• Why is the formula of water H2O and not
H3O or HO2
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Bond formation and noble
gases
• Noble gases rarely form compounds
• They have filled s and p outer subshells
• This is a total of eight electrons, referred to
as an octet
• Eight electrons in the outer s and p orbitals
is a particularly stable configuration
• The energy required to remove an electron
from these full subshells is particularly
high
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Lewis dot symbols for elements
• Since only the
valence electrons are
involved in bonding
we can concentrate
on those
• Lewis dot symbols
are used to represent
the valence electrons
of an atom
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How atoms achieve an octet
• Metals can lose one to three electrons to
form a cation with the electron
configuration of the previous noble gas
• Nonmetals can gain one to three
electrons to form an anion with the
electron configuration of the next noble
gas
• Atoms can also share electrons
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Formation of ions
• Metals can lose electrons to form ions
– Na ([Ne]2s1) → Na+ ([Ne]) + e– If a metal loses all of its outer electrons, it
acquires the octet of the previous noble gas
• Nonmetals can gain electrons to form ions
– Cl ([Ne]2s22p5 + e- → Cl ([Ne]2s22p6
• Lewis dot structures of the atoms can be
very helpful here
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Forming ionic compounds
•
•
•
•
Reaction of Na with Cl
Na donates an electron to Cl
Na+ has the previous noble gas structure (Ne)
Cl- has the next noble gas structure (Ar)
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Binary ionic compounds
• In NaCl, each Na+ is surrounded by six Cl-,
and each Cl- is surrounded by six Na+
• Ionic lattice is a three-dimensional array of
ions
• These electrostatic attractions are called ionic
bonds
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But… remember!
• Atoms can also share electrons…
-
when this happens, the electrons form
covalent bonds…
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Chemical bonds
(let’s review)
• A chemical bond is the force that holds
two or more atoms together
• Chemical bonds involve the electrons
• A bond results if a more stable electron
configuration results
• The valence electrons are the electrons in
the outer s and p subshells
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The Covalent Bond
• Covalent bonds result from electron sharing between two
atoms
• We use Lewis dot structures to show the order and
arrangement of the atoms in a molecule and all of the
valence electrons
H
.
+H
.
→ H:H
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The Octet Rule
• The s block and p block elements (often termed
the representative elements) will form bonds
such that there are eight electrons surrounding
each atom (the octet rule)
• Obtaining this configuration is the driving force
for bond formation for many compounds formed
by the representative elements
• The exceptions are H, Li and Be, which tend
to follow a “duet” rule (filling the ns subshell)
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Types of covalent bonds
• Two nonmetals can share one, two or
three electron pairs
• The bonds resulting from this sharing are
referred to as single, double or triple
bonds respectively
• Multiple bonds are frequently observed in
compounds of 2nd period elements
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Writing Lewis dot structures
• The octet rule and Lewis dot structures
allow us to justify the formulas that we
know
• We can also predict the formulas of new
compounds
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Lewis dot structures
1. Count the number of valence electrons for the
atoms in the molecule, include charge.
2. Place the most electropositive atom in the
center (the inner atom). Draw simple diagram;
single atom in middle, one bond to each outer
atom.
3. Subtract two electrons from the total number of
valence electrons for each bond.
4. Satisfy octet rule for all outer atoms.
5. Subtract these atoms from total.
6. Place remaining pairs on center atom.
7. Check octet rule for center atom.
8. CELEBRATE!
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Lewis Dot Diagrams
• Examples: CF4; NH3; NH4+; SCl2; PF3
• Other examples: NO2-;
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Resonance structures
• In compounds with multiple bonds,
sometimes you can draw structures which
vary only by placement of the double
bonds
• The structures are called resonance
structures, and are an approximation of
the true structure of the molecule
• Actually, the molecule is a superposition of
all of the resonance structures
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Nitrate ion
• Nitrate has three resonance structures
• Each is identical except for the placement of the
double bond and associated lone pairs
• Experimentally, all N-O bonds are identical
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More examples Lewis dot diagrams
• NO2
-- A bit more complicated:
C2H6; (hydrocarbons in general =C3H8);
H2O2; H2CO, BF3 (no multiple bonds for
halogens!)
Start to anticipate base/acid chemistry:
NH3 , OH-, etc…
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Electronegativity
• the ability of an atom to attract electrons in a
bond to itself
• Differences in electronegativities of atoms that
are bonded together results in a partial transfer
of electron charge to the more electronegative
atom.
• The bond is therefore a polar covalent bond
• The polar bond has a negative end and a
positive end (a so-called dipole; which we
indicate with a δ with the appropriate sign
added)
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Nonpolar, polar and ionic bonds
http://web.mit.edu/3.091/www/pt/pert8.html
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Polarity of bonds
• Bonds that involve atoms of differing
electronegativities have a concentration of
negative charge at the more
electronegative atom, and a deficiency of
charge at the less electronegative atom
• This unequal distribution of negative
charge creates a dipole, where one end of
the bond is slightly negative and the other
is slightly positive
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Geometry of Simple Molecules
• Electron pairs will repel each other, and
will govern the structure of the molecule,
all other things being equal
• Electron pairs will arrange themselves to
be as far apart as possible
• Note that molecular geometry is described
by the bonded atoms and does not include
the lone pairs
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Electron pair repulsion
• Degree of repulsion depends on the electron
pair types; in order of decreasing repulsion
– lone pair-lone pair
– lone pair-bonding pair
– bonding pair-bonding pair
• We will also treat all of the electrons that bond
together two atoms as one electron group
regardless of whether the bond is single, double
or triple
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Parent structures
• We will consider three parent structures to
begin
• The central or inner atom is designated A,
outer atoms are designated X, and lone
pairs are designated E
• The parent structures are based on the
number of electron pairs that surround the
central atom
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Parent structures
Parent structure Name
Bond angles
AX4
tetrahedron
109.5°
AX3
trigonal planar
120°
AX2
linear
180°
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Parent structure derivatives
• Each parent structure can give rise to a
family of derivatives, simply by replacing
bonding pairs with lone pairs
• For AX4, there are two derivatives
– AX3E - the trigonal pyramid (NH3)
– AX2E2 - bent (H2O)
• Replacement of the bonding pairs with the
lone pair(s) compresses the bond angles
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Tetrahedral - AX4
• Each molecule has four electron pairs around
them
• Replacement of a bonding pair with a lone pair
yields the AX3E (NH3), AX2E2 (H2O)
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Trigonal - AX3
• The examples shown here have three electron pairs
around the central atom
• Note that when one of the bonding pairs is replaced by a
lone pair, the bond angle is smaller
• Structure is designated by AX2E, where E is a lone pair
“electron deficient”
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Linear examples - AX2
• The bonding pairs in the following molecules
arrange themselves to be as far apart as
possible
• These examples have two electron pairs
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Expanded Valence Shells
Valence Shell Electron Pair Repulsion (VSEPR)
• 3rd Row and beyond,
valence shell MAY have
more than octet… up to
12 electrons, 6 electron
pairs.
• Use VSEPR chart from
website:
http://www.sci.uidaho.edu
/chem101/Lecture
supplements/VSEPR1.P
DF
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Examples
• SCl4; KrCl2 ; BrF4
• IF5; SF6; XeBr4
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Polarity of bonds can lead to polar
molecules
33
Polarity of molecules
• If the forces are equal but opposite, the polarity
of the bonds cancel
• Such molecules have polar bonds, but are
themselves nonpolar (e. g. CO2)
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Nonpolar molecules
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Polar molecules
• Polar molecules result from an
arrangement of polar bonds such that the
entire molecule has a dipole
• The polar bonds can be arranged to
cancel the polarity (CO2)
• The best way to predict this in complex
cases is vector algebra, but it is important
to learn to recognize this in clear cut
structures
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Polar molecules
• If the two bonds are not equally polar, then a net
dipole exists (a)
• If the two bonds are equally or not equally polar,
yet are not opposite, a net molecular dipole
exists
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Linear example
• The Be-Cl and Be-H are both polar, but not to
the same extent
• Even though BeClH is linear, the polarity of the
bonds are not equal so the molecule is polar
http://web.mit.edu/3.091/www/pt/pert8.html
38
Noncanceling bonds
• Water is bent, so the dipolar O-H bonds cannot
cancel each other
Same for SO2 or
any bent
molecule
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That’s all folks!
On to Chapter 10
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