2014-­‐05-­‐08 Periodic Trends
™  There are various trends on the periodic table that need
to be understood to explain chemical bonding.
™  These include:
™ 
™ 
™ 
™ 
™ 
Homework:
™  Do:
™  Periodic Trends Handout
Elements of his theory:
™  Valence electrons play a fundamental role in chemical
bonding.
™  Sometimes bonding involves the transfer of one or more
electrons from one atom to another. This leads to the ion
formation and IONIC BONDS.
™  Sometimes bonding involves sharing electrons between
atoms, this leads to COVALENT BONDS.
Atomic/Ionic Radius
Ionization Energy
Electronegativity
Electron Affinity
Effective Nuclear Charge
Lewis Theory
Now, where
did I leave my
keys…
™  From 1916-1919, Gilbert N.
Lewis made several
important proposals on
bonding which lead to the
development of Lewis
Bonding Theory.
™  LEWIS SYMBOLS:
™  A common chemical symbol surrounded by up to 8
dots.
™  The symbol represents the nucleus and the electrons of
the filled inner shell orbitals.
™  The dots represent the valence electrons.
™  For Example:
™  Electrons are transferred or shared such that each atom
gains a more stable electron configuration.
™  Usually this is that of a noble gas (having 8 outer shell
electrons).
™  This arrangement is called an OCTET.
1 2014-­‐05-­‐08 Hold the phone…
Types of Bonding
™  Ionic (metal/non-metal)
This only works well for the representative
elements.
Transition metals, actinides and lanthanides
have incompletely filled inner shells - we can't
write simple Lewis structures for them.
Ionic Bonds
™  Ionic bonds are forces that hold ionic compounds
together.
™  Forming the ionic bond:
™  Step 1:
™  A cation forms by the LOSS of 1 or more e-.
™  Representative elements become cations, isoelectronic
with the nearest Noble gas.
™  For others (transition metals), not necessarily.
™  Covalent (non-metal/non-metal)
™  Intermolecular Bonding (between molecules):
™  Hydrogen Bonding
™  London Dispersion Forces
™  Step 2:
™  A anion forms by GAINING sufficient e- to become
isoelectronic with the nearest Noble gas.
™  The Lewis symbol will show an OCTET (8) of electrons
for the anion.
™  Step 3:
™  The oppositely charged ions come together to form an
ionic compound.
™  The electrostatic attraction between the ions forms the
bond.
™  In the solid state, each anion surrounds itself with
cations, and each cation with anions, forming an ionic
crystal.
™  Step 4:
™  A formula unit of an ionic compound is the smallest
collection of ions that would be electrically NEUTRAL.
™  The formula unit is automatically obtained when the
Lewis structure of the compound is written.
™  The ionic crystal then consists of each constituent ion
bound together in the crystal, not of individual
molecules.
™  ATOMS → COMPOUNDS, properties change…
™  Their reactivity decreases considerably.
™  They become neutral overall.
™  There are no unique molecules in many ionic
solids.
™  Ionic compounds become electrically
conductive when melted or dissolved.
2 2014-­‐05-­‐08 Covalent Bonds
IONIC BONDS are STRONG so that
compounds held together by ionic bonds
have HIGH MELTING
TEMPERATURES.
™  The valence electrons involved in the bond are called
the BONDING ELECTRONS or the BOND PAIR.
™  Those not involved in the bond are called the
NONBONDING ELECTRONS or the LONE
PAIRS.
Bond
Pair
Lone
Pair
™  Covalent bonds arise from the sharing of electrons
between atoms (generally of groups IVA, VA, VIA, and
VIIA).
™  Each electron in a shared pair is attracted to both
nuclei involved in the bond.
COVALENT BONDS are VERY STRONG.
™  OCTET RULE :
™  An atom other than hydrogen tends to form bonds until
it is surrounded by eight valence e-.
™  The pairs repel each other and thus tend to stay as far
away as possible.
Multiple Bonds
™  Covalent compounds can form multiple bonds.
™  Depending upon how many pairs of electrons are
shared, different types of bonds are made.
™  Single Bond: Share ONE pair of electrons.
™  Double Bond: Share TWO pairs of electrons.
™  Triple Bond: Share THREE pairs of electrons.
Polar Covalent Bonds
™  Imagine we have one atom with a somewhat higher
electronegativity than the other in a covalent bond:
™  This will cause the electrons to be shared unevenly, such
that the shared electrons will spend more time (on
average) closer to the atom that has the higher
ELECTRONEGATIVITY.
™  This is done to fill their valence shell (Octet Rule).
3 2014-­‐05-­‐08 ™  The greater the difference in electronegativity in the
bonding atoms, the greater the polarity of the bond.
™  Atoms with widely different electronegativity values
(ΔE ≥ 2.0) tend to form IONIC BONDS.
™  True NON-POLAR COVALENT BONDS form only
when diatomic molecules are formed with two identical
atoms (ΔE ≤ 0.4)
™  Everything else will form a POLAR COVALENT
BOND (ΔE 0.5 – 1.9) .
Homework:
™  Do:
™  W.S. 7-1
Hydrogen Bonding
™  The force of attraction, shown here as a dotted line, is
called a HYDROGEN BOND:
™  Polar molecules, such as water molecules, have a weak,
partial negative charge at one region of the molecule
(the oxygen atom in water) and a partial positive
charge elsewhere (the hydrogen atoms in water).
™  When water molecules are close together, their positive
and negative regions are attracted to the oppositelycharged regions of nearby molecules.
4 2014-­‐05-­‐08 ™  Each water molecule is hydrogen bonded to four
others:
London Dispersion Forces
™  The London dispersion force is the weakest
intermolecular force.
™  It is a temporary attractive force that results when the
electrons in two adjacent atoms occupy positions that
make the atoms form temporary dipoles.
HYDROGEN BONDS are WEAK.
™  Because of the constant motion of the electrons, an
atom or molecule can develop a temporary
(instantaneous) dipole when its electrons are
distributed unsymmetrically about the nucleus.
™  Dispersion forces are present between any two
molecules (even polar molecules) when they are almost
touching.
LONDON DISPERSION FORCES are the
WEAKEST.
Writing Lewis Structures
Simple Ionic Compounds
™  Easy to do…
™  For Example:
™  KBr
™  Li2S
™  The overall charge on the compound must equal zero,
that is, the number of electrons lost by one atom MUST
EQUAL the number of electrons gained by the other
atom.
™  The Lewis Structure of each ion is used to construct
the Lewis Structure for the ionic compound.
™  K3P
Ionic compounds end up having
ZERO valence electrons.
5 2014-­‐05-­‐08 THE LEWIS STRUCTURES OF COVALENT
COMPOUNDS THAT OBEY THE OCTET RULE
Rules of the Road…!
1.  Find the total number of valence e-.
™  Lewis structures show how the VALENCE electrons are
distributed in a molecule.
™  Covalent compounds share electrons to fill their
valence shells.
™  Therefore, the Lewis structures for these compounds
are drawn a little differently.
™  Go by the column or group that it’s in.
2.  After connecting your central atom to the terminal atoms with single
bonds, begin adding the remaining valence e- as lone pairs.
™  First around the terminal atoms.
™  Then around the central atom (if you have any left over).
3.  Satisfy the octet rule for your central atom by either:
™  Replacing a lone pair on your terminal atoms with a bond (to
make a double bond).
OR:
™  By replacing two lone pairs with two bonds (to make a triple
bond).
Remember:
™  You can’t just add double bonds without
first removing a lone pair.
™  Not only are you adding more electrons
than you started with, but you’re
probably breaking the octet rule for the
terminal atoms.
™  NH4+
™  O2
™  NO2™  C2H4
™  CO
THE LEWIS STRUCTURES OF COVALENT
COMPOUNDS THAT VIOLATE THE OCTET RULE
Incomplete Octets:
™  In addition to H, Be, B, and Al are exceptions to the
octet rule.
Hypervalent Compounds:
™  Elements in the third and fourth periods MAY attain
more than an octet of valence e- when they form
covalent compounds (the electrons are placed in low lying
d-orbitals).
™  Since they have very low electronegativities, they can
only accept one electron for every one they donate.
™  Other than the fact that the central atom will end up
with more than eight valence electrons, the same rules
apply.
™  For Example:
™  Examples include Xe, P, S, and I.
BF3
6 2014-­‐05-­‐08 Formal Charge
™  For Example:
™  When we draw Lewis structures, we may end up with
several possible structures for the molecule.
™  PCl5
™  SF4
™  The question then becomes…
™  XeF4
™  Calculating the FORMAL CHARGE allows us to figure
this out:
⎡1
⎤
F.C. = (Valence e − ) - ⎢ (number of e − in covalent bonds) + (number of e − in lone pairs)⎥
⎣2
⎦
Note: The sum of the formal charges equals the overall
charge on the ion or molecule.
™  The “guidelines”:
™  We generally choose the Lewis structure in which the
atoms bear formal charges closest to ZERO.
™  We generally choose the Lewis structure in which any
negative formal charge resides on the more electronegative
atom.
Resonance
™  When we draw Lewis structures in which we must make a
choice as to what gets a double bond, the structure is
actually a blend of two or three structures.
For Example:
™  Draw the three possible Lewis structures for the
thiocyanate ion, NCS-.
™  Determine the formal charges of the atoms in each
structure.
™  Which Lewis structure is the preferred one?
For Example:
™  Draw three resonance structures for the polyatomic ion
CO3 2-.
™  We “say” that the structure RESONATES (contains
contributions from each of the resonance structures).
™  Resonance occurs simply because the electron-dot model is
too limited to show how electrons are being shared between
the atoms.
7 2014-­‐05-­‐08 Homework:
™  Do:
™  Lewis Structures W.S.
™  Study for your quiz!
™  Lewis Diagrams
VSEPR
™  The Valence Shell Electron Pair Repulsion model is not so much
a model of chemical bonding, as a scheme for explaining the
shapes of molecules.
™  It is based on the quantum mechanical view that bonds represent
electron clouds that repel each other and thus try to stay as far
apart as possible.
™  This minimizes the energy of repulsion, represents the lowest
energy configuration of the molecule, and gives it a distinctive
shape.
Electron Group Geometries
™  An ELECTRON GROUP is any collection of valence
electrons, localized in a region around a central atom, that
repels other groups of valence electrons.
™  Can include:
™  A single unpaired electron.
™  A lone pair of electrons.
™  One bonding pair of electrons in a single covalent bond.
™  The overall shape of the molecule is determined by its
BOND ANGLES (the angles made by the lines joining
the nuclei of the atoms in the molecule).
™  The central idea is that bonding and non-bonding pairs
around a given atom will be positioned as far apart as
possible.
™  Mutual repulsions among electron groups lead to their
electron group geometry.
It’s all about
repulsion,
baby!
™  Two bonding pairs of electrons in a double covalent bond.
™  Three bonding pairs of electrons in a triple covalent bond.
™  Three e- groups:
Let’s reason our way through
a few…
™  Four e- groups:
CO2
™  The electron group geometries are:
™  Two e- groups:
™  Five
e-
BCl3
CH4
groups:
™  Six e- groups:
™  Seven e- groups:
PCl5
IF7
SF6
8 2014-­‐05-­‐08 It does get a
little more
complicated…
How will we do this?
Step 1:
Draw a Lewis structure of the molecule.
Step 2:
Determine the number of electron groups
around the central atom, and identify each as
either a bonding group or a lone pair.
Step 3:
Identify the electron group geometry.
Step 4:
Identify the molecular geometry.
Ok… here goes nothing!!
™  BeCl2
Homework:
™  Do:
™  VSEPR W.S.
™  BF3
™  SO2
™  NH3
™  XeF2
Covalent Bonding and Orbital Overlap
™  VSEPR provides a simplistic model for predicting
molecular shape, but does not explain why bonds exist.
™  Lewis Theory + Quantum Mechanics =
™  Atoms sharing electrons concentrates electron density
between the two nuclei.
™  These orbitals are then said to share a region of space, or
to OVERLAP.
™  Take H2 for example:
Electron overlap creates
the covalent bond.
™  Another involves HCl:
™  Chlorine has an electron configuration of [Ne]3s23p5.
™  All its orbitals are full except for
one 3p orbital with a single electron.
™  This electron pairs up with the single
electron of H (1s1) to form the
covalent bond.
™  Yet another… Cl2.
9 2014-­‐05-­‐08 ™  There is always an optimum distance between the two
bonded nuclei in any covalent bond.
™  Overlapping of orbitals helps us to understand the
formation of covalent bonds, but its not easy to extend
these ideas to polyatomic molecules.
™  In these cases, we must explain both the formation of
electron-pair bonds and the observed geometries of the
molecules.
™  To explain geometries, we assume that the atomic orbitals
on an atom mix to form new orbitals called HYBRID
ORBITALS.
™  During the process of HYBRIDIZATION, the total
number of atomic orbitals on an atom remains constant,
so, the number of hybrid orbitals on an atom equals the
number of atomic orbitals mixed.
The Hybrid Orbital Model
™  Consider BeF2:
™  Developed by Linus Pauling in 1931.
™  VSEPR predicts that the molecule is linear with two
identical Be-F bonds.
™  Each F has an electron configuration of 1s22s22p5.
™  There are several types of hybridization to
know:
,
,
,
and
™  Each one of these types of hybridization
are connected to the five basic electron
geometries:
sp Hybrid Orbitals
™  This 2p e- can be paired with an unpaired electronUnpaired e in
from the Be atom to form a polar covalent bond. a 2p orbital.
-
Linus Pauling
(1901 - 1994)
™  Linear, trigonal planar, tetrahedral,
trigonal bipyramidal, and octahedral.
™  The orbital diagram for a ground state Be atom is:
™  Notice that there are no unpaired electrons available,
therefore, it is incapable of forming bonds with Fluorine
atoms.
™  With this e- configuration, the Be atom now has two
unpaired e- it can form two polar covalent bonds with the
F atoms.
™  But… these concepts still do not explain the structure of
BeF2 since:
™  BUT… it does, by promoting one of the 2s electrons to a
2p orbital:
10 2014-­‐05-­‐08 ™  This is accomplished by “mixing” the 2s and 2p orbitals to
generate two new, equal orbitals as follows:
™  The e- in the sp hybrid orbital can form two e- bonds with
the two Fluorine atoms.
™  These two orbitals are identical in shape, but point in
opposite directions.
™  Since these orbitals are equivalent, but in opposite
directions, BeF2 has two identical bonds and a linear
geometry.
™  This is what we call a sp HYBRID ORBITAL.
™  A linear arrangement of electron domains implies sp
hybridization.
sp2 and sp3 Hybrid Orbitals
™  This results in three equivalent sp2 hybrid orbitals:
™  Mixing a certain number of atomic orbitals always results in
the same number of hybrid orbitals.
™  For example:
™  Mixing one 2s and one 2p orbital results in two equivalent sp
hybrid orbitals pointing in opposite directions.
™  There are other possible combinations.
™  Using BF3 as an example, a 2s e- on the B atom can be
promoted to a vacant 2p orbital.
™  These three sp2 hybrid orbitals lie in the same plane, 120o apart
from one another.
™  This leads to the TRIGONAL PLANAR geometry of BF3.
™  Note that an unfilled 2p orbital remains unhybridized... more
to come later (hint: Carbon)!
™  In sp3 hybridization, such as in the C in CH4, an s orbital
can mix with all three p orbitals in the same subshell.
™  The result is four equivalent sp3 hybrid orbitals.
™  Each one of these hybrid orbitals has a large lobe that
points towards a vertex of a tetrahedron.
™  These hybrid orbitals can form two electron bonds by
overlapping with the atomic orbitals of H to form four
equivalent sp3 hybrid orbitals on the C.
11 2014-­‐05-­‐08 Hybridization Involving d Orbitals
™  Atoms in the third period and beyond can also use d
orbitals to form hybridized orbitals.
™  For example:
™  Mixing one s orbital, three p orbitals, and one d orbital
leads to FIVE sp3d hybrid orbitals directed towards the
vertices of a trigonal bipyramid (eg: PF5).
™  Mixing one s orbital, three p orbitals, and two d orbitals
leads to SIX sp3d2 hybrid orbitals directed towards the
vertices of an octahedron (eg: SF6).
To Summarize:
™  Hybridization offers a convenient model for using valence
bond theory to describe covalent bonds whose geometries
are predicted by the VSEPR.
™  To predict hybrid orbitals used by an atom in bonding:
1.  Draw the Lewis structure for the molecule or ion.
2.  Determine the electron-domain geometry using the VSEPR
model.
3.  Specify the hybrid orbitals needed to accommodate the electron
pairs based on their geometric arrangement.
An example:
™  Indicate the hybridization of orbitals employed by the
central atom in:
™  NH2™  SF4
Multiple Bonds
™  In all the hybridization examples so far, e- density is
concentrated symmetrically about the line connecting the
nuclei (the intermolecular axis).
™  The line joining the two nuclei passes through the middle of
the overlap region.
™  These bonds are called SIGMA (σ) BONDS.
12 2014-­‐05-­‐08 ™  Multiple bonds involve a second type of bond resulting
from perpendicular overlap between two p orbitals.
™  This produces a PI (π)BOND.
™  In pi bonds, the overlap regions are above and below the
intermolecular axis.
× 
The sideways orientation of
p orbitals in a π bond makes
for weaker overlap.
× 
As a result, π bonds are
generally weaker than σ
bonds.
™  A bond angle of 120o suggests that each C atom uses sp2
hybrid orbitals to form σ bonds with the other Carbon
and two Hydrogens.
e-,
™  Since Carbon has four valence after
hybridization,
one e- in each of the Carbon atoms remains in the
unhybridized 2p orbital:
sp2
I’m so
lonely…
™  Single, double, and triple bonds are all different:
™  In almost all cases, single bonds are σ bonds.
™  A double bond consists of one σ bond and one π bond.
™  A triple bond consists of one σ bond and two π bonds.
™  Consider Ethylene (C2H4) which possesses a C=C double
bond:
™  This unhybridized 2p
orbital is directed
perpendicular to the
plane that contains the
three sp2 hybrid orbitals.
™  The same principle can
be used to explain triple
bonds.
™  Consider Acetylene (C2H2):
™  Acetylene is also linear and uses sp hybrid orbitals to form
σ bonds with the other Carbon and one Hydrogen.
The two remaining
unhybridized 2p orbitals
overlap to form a pair of
π bonds.
13 2014-­‐05-­‐08 What is on the exam?
Homework:
™  Periodic Trends
™  Do:
™  Bonding W.S.
™  Study for your quiz!
™  VSEPR and Hybridization
™ 
™ 
™ 
Atomic and Ionic Radii
Ionization Energy
Electron Affinity and
Electronegativity
™  Lewis Theory
™ 
™ 
™ 
Elements of the Theory
Ionic Bonds
Covalent Bonds
™ 
Multiple Bonds
™ 
Polarity
™  Drawing Lewis structures
™ 
™ 
™ 
™ 
™ 
Simple Ionic Compounds
Structures that Obey the Octet
Rule
Structures that Violate the
Octet Rule
Formal Charge
Resonance
™  VSEPR
™  Hybridization
™  Intermolecular Bonding
™ 
™ 
Hydrogen Bonding
London Dispersion Forces
14