CHEMICAL BONDING &
MOLECULAR STRUCTURE
WHAT IS A CHEMICAL BOND?
 Bonds are attractive forces that hold groups of atoms together and make them function as a unit.
 Bonding relates to physical properties such as melting point, hardness and electrical and thermal
conductivity as well as solubility characteristics.
 The system is achieving the lowest possible energy state by bonding.
 Most of the chemical substances you can name or identify are NOT elements. They are
compounds. That means being bound requires less energy than existing in the elemental form.
 It also means that energy was released from the system — HUGE misconception — it takes
energy to break a bond, not make a bond! Energy is RELEASED when a bond is formed,
therefore, it REQUIRES energy to break a bond.
TYPES OF CHEMICAL BONDS
 IONIC BONDS — electrons are transferred; metals react with nonmetals; ions paired have
lower energy (greater stability than separated ions).
 COULOMB'S LAW — used to calculate the Energy of an ionic bond; the energy interaction
between a pair of ions.
- J is the energy in Joules
- Q1 and Q2 are the numerical
ion charges; don’t neglect their
signs
- r is the distance between the ion
centers in nanometers
There is a negative sign on the Energy once calculated—it indicates an attractive force so that the
ion pair has lower energy than the separated ions.

BOND LENGTH — the distance between the 2 nuclei where the system energy is at a
minimum between the two nuclei; energy is given off when two atoms achieve greater stability
together than apart.
 Attractive forces — proton - electron
 Repulsive forces — electron - electron
 Small energy decrease - van der Waals IMFs
 Large energy decrease - chemical bonds

COVALENT BONDS — electrons are shared by nuclei
 Pure covalent (non-polar covalent) - Electrons are shared evenly
 Polar covalent bonds - Electrons are shared unequally; Atoms end up with fractional
charges; δ+ or δ-.



ELECTRONEGATIVITY (Eneg)
The ability of an atom in a molecule (BONDED) to attract shared electrons to
itself. Think tug of war.
TREND: Increases across a period and decreases down a group: Fluorine is the
most Eneg and Francium is the least Eneg.
Characterizing Bonds - Use the difference in Eneg to determine the type of bond
formed.
 ionic — Eneg difference > 1.7; When the electrons are shared unequally to
a greater extent (c).
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

polar covalent — Eneg difference ≥ 0.4 and ≤ 1.7; When the two nuclei are different the
electrons are not shared equally, setting up slight +/− poles (b).
nonpolar covalent — Eneg difference < 0.4 - When two nuclei are the same, the sharing
is equal (a).
PRACTICE ONE Relative Bond Polarities
Order the following bonds according to polarity: H—H, O—H, Cl—H, S—H, and F—H.
BOND POLARITY AND DIPOLE MOMENTS


DIPOLAR MOLECULES - A molecule with a somewhat positive end and a somewhat negative
end; a dipole moment; also molecules with preferential orientation in an electric field; all diatomic
molecules with a polar covalent bond are dipolar.
REPRESENTATION
 by a arrow pointing to the negative charge center with the tail of the arrow indicating the
positive center of change.
 by an electrostatic potential diagram.
The diagram shows the charge distribution in the water molecule, the water molecule in an electric field, and the
electrostatic potential diagram of a water molecule.

MOLECULES WITH POLAR BONDS BUT NO DIPOLE MOMENT - Linear, radial or
tetrahedral symmetry of charge distribution. Examples: CO2 - linear and CCl4 - tetrahedral.
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IONS: ELECTRON CONFIGURATIONS AND SIZES





Goal is for atoms in a stable compound to achieve a noble gas configuration.
COVALENT: two nonmetals share the valence electrons so each has a noble gas configuration.
IONIC: Metal and representative group metal: form a binary ionic compound where electrons
are transferred so the nonmetal has the valence configuration of the noble gas on its period and
the metal has the noble gas configuration from the period above.
Placement of elements on the periodic table suggests how many electrons are lost or gained to
achieve a noble-gas configuration
a. Group I loses one electron, Group II loses two, Group VI gains two, Group VII gains
one….
Formulas for compounds are balanced so that the total positive ionic charge is equal to the total
negative ionic charge
Al+32 O-2 3
Total positive = +6
Total negative = -6
IONIC COMPOUNDS
 The final result of ionic bonding is a solid, regular array of cations and anions called a crystal
lattice; This configuration limits negative ion - negative ion and positive ion - positive ion
interactions and maximizes positive ion - negative ion interactions.
 Ion size plays a role in determining the structure and stability of ionic solids and the properties of
ions in aqueous solutions.
PRACTICE TWO Relative ion size
Arrange the ions Se2-, Br-, Rb+, and Sr+2 in order of decreasing size.
PRACTICE THREE Relative ion size again
Choose the largest ion in each of the following groups.
a. Li+, Na+, K+, Rb+, Cs+
b. Ba2+, Cs+, I-, Te2-
ENERGY EFFECTS IN BINARY IONIC COMPOUNDS



Remember that ionic solids form because the resulting conglomeration of ions has a lower energy
than the original elements.
LATTICE ENERGY is the change in energy that takes place when separated gaseous ions are
packed together to form an ionic solid. It is the energy released when an ionic solid forms exothermic; it will have a (-) sign.
Example: Formation of lithium fluoride
Process
Li(s) → Li(g)
Li(g) → Li+1(g) + e-
Description
Sublimation energy
Ionization energy
Energy Change (kJ)
161
520
3
½F2(g) → F(g)
Bond energy (1/2 mole)
77
F(g) + e- → F-1(g)
Electron affinity
-328
Li+1(g) + F-1(g) → LiF(s)
Lattice energy
-1047
Li(s) + ½F2(g) → LiF(s)
ΔH
-617
- The formation of ionic compounds is endothermic until the formation of the lattice
- The lattice formed by alkali metals and halogens (1:1 ratio) is cubic except for cesium salt.

Lattice energy calculations
k = a proportionality constant dependent on the solid structure and the electron configuration;
Q1 and Q2 are the ion charges; r = shortest distance between centers of the cations and the anions.

Lattice energy increases as the ionic charge increases and the distance between anions and
cations decreases.
PARTIAL IONIC CHARACTER OF COVALENT BONDS

Calculating % ionic character

Ionic vs. Covalent
 Ionic compounds generally have greater than 50% ionic character
 Ionic compounds generally have electronegativity differences greater than 1.7
 Percent ionic character is difficult to calculate for compounds containing polyatomic ions.
COVALENT BONDING
Most compounds are covalently bonded, especially carbon compounds.
 Strengths of the Bond Model - associates quantities of energy with the formation of bonds
between elements; allows the drawing of structures showing the spatial relationship between
atoms in a molecule; provides a visual tool to understanding chemical structure
 Weaknesses of the Bond Model - bonds are not actual physical structures; bonds cannot
adequately explain some phenomena like resonance.
COVALENT BOND ENERGIES AND CHEMICAL REACTIONS

Average bond energies
Process
Energy Required (kJ/mol)
CH4(g) → CH3(g) + H(g)
CH3(g) → CH2(g) + H(g)
CH2(g) → CH(g) + H(g)
CH(g) → C(g) + H(g)
435
453
425
339
Total 1652
Average 413
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
Single and multiple bonds
 single bond - one pair of electrons shared → sigma (σ) bond
 Multiple bonds are most often formed by C,N,O,P and S atoms — C-NOPS
 double bond - two pairs of electrons shared → one σ bond and one π bond
 triple bond - three pairs of electrons shared → one σ bond and two π bonds
Multiple Bonds, Average Energy (kJ/mole)
C=C
C≡C
O=O
C=O
C≡O


614
839
495
745
1072
N=O
N=N
N≡N
C≡N
C=N
607
418
941
891
615
as the number of shared electrons increases, the bond length shortens.
Bond energy and enthalpy (using bond energy to calculate approximate energies for rxns)
 ΔH = sum of the energies required to break old bonds(endothermic) + sum of the energies
released in forming new bonds (exothermic).
 ΔH =∑D(Bonds broken) -∑D(Bonds formed) (D always has a positive sign)
PRACTICE FOUR ΔH from Bond Energies
Using the bond energies in in the table in your text, calculate Δ H for the reaction of methane with
chlorine and fluorine to give Freon, CF2Cl2.
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THE LOCALIZED ELECTRON (LE) BONDING MODEL
Assumes that a molecule is composed of atoms that are bound together by sharing pairs of electrons
using the atomic orbitals of the bound atoms. Electron pairs are assumed to be localized on a particular
atom -lone pairs - or in the space between two atoms - bonding pairs.
 Lone electron pairs - electrons localized on an atom (unshared)
 Bonding electron pairs - electrons found in the space between atoms (shared pairs)
 Derivations of the Localized Model
1. Lewis Structures describe the valence electron arrangement
2. Geometry of the molecule is predicted with VSEPR
3. Description of the type of atomic orbitals “blended” by the atoms to share electrons or hold
lone pairs (hybrids—next chapter).
LEWIS STRUCTURES


Electrons and Stability
 "the most important requirement for the formation of a stable compound is that the atoms
achieve noble gas configurations
 Duet rule - hydrogen, lithium, beryllium, and boron form stable molecules when they share two
electrons (helium configuration)
 Octet Rule - elements carbon and beyond form stable molecules when they are surrounded by
eight electrons
Writing Lewis Structures - rules
1. H is always a terminal atom and connected to only one
other atom.
2. Lowest electronegativity elements is central atom in
molecule.
3. Add the TOTAL number of valence electrons from all
atoms.
4. Place one pair of electrons, a σ bond, between each pair of bonded atoms.
5. Arrange the remaining atoms to satisfy the duet rule for hydrogen and the octet rule for the
second row elements.
PRACTICE FIVE Writing Lewis Structures
Give the Lewis structure for each of the following:
a. HF
d. CH4
b. N2
e. CF4
c. NH3
f. NO+
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EXCEPTIONS TO THE OCTET RULE
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Fewer than eight — H at most only two electrons (one bond); BeH2, only 4 valence electrons
around Be (only two bonds); Boron compounds, only 6 valence electrons (three bonds);
ammonia + boron trifluoride is a classic Lewis A/B reaction.
Example: Boron Trifluoride
1. Note that boron only has six electrons around it
2. BF3 is electron deficient and acts as a Lewis acid
(electron pair acceptor)
3. Boron often forms molecules that obey the octet rule
Expanded Valence — can only happen if the central element has d-orbitals which means it is
from the 3rd period or greater and can thus be surrounded by more than four valence pairs in
certain compounds. The number of bonds depends on the balance between the ability of the
nucleus to attract electrons and the repulsion between the pairs.
Example: Sulfur Hexafluoride
1. Note that sulfur has 12 electrons around it, exceeding
the octet rule
2. Sulfur hexafluoride is very stable
3. SF6 fills the 3s and 3p orbitals with 8 of the valence
electrons, and places the other 4 in the higher energy 3d
orbital
Odd-electron compounds — A few stable compounds contain an odd number of valence
electrons and thus cannot obey the octet rule. NO, NO2 , and ClO2.
More About the Octet Rule
1. Second row elements C, N, O and F should always obey the octet rule
2. B and Be (second row) often have fewer than eight electrons around them, and form
electron deficient, highly reactive molecules
3. Second row elements never exceed the octet rule
4. Third row and heavier elements often satisfy (or exceed) the octet rule
5. Satisfy the octet rule first. If extra electrons remain, place them on elements having
available d orbitals. When necessary to exceed the octet rule for one of several third
row elements, assume that the extra electrons be placed on the central atom.
Coordinate covalent bonds
 Some atoms, such as N and P, tend to share a lone pair with another atom that is short of
electrons, leading to the formation of a coordinate covalent bond. These bonds are in all
coordination compounds and Lewis Acids/Bases.
 N is sharing the lone PAIR of electrons by drawing an arrow from it to the H+, remember
H+ has NO electrons to contribute to the bond. Note that all four bonds are actually
identical in length and strength.
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PRACTICE SIX Lewis Structures for Molecules That Violate the Octet Rule I
Write the Lewis structure for PCl5.
PRACTICE SEVEN Lewis Structures for Molecules That Violate the Octet Rule II
Write the Lewis structure for each molecule or ion.
a. ClF3
d. BeCl2
b. XeO3
e. ICl4
c. RnCl2
RESONANCE STRUCTURES
Experiments show that ozone, O3 has equal bond lengths and equal bond strengths, implying that there is
an equal number of bond pairs on each side of the central O atom. When you draw the Lewis structure,
that situation is NOT what you draw. So, we draw a resonance structure:
BUT since there are equal bond lengths and strengths, they are not as pictured above. We use the double
edged arrows to indicate resonance.
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The central atom is surrounded by three sites of electron density: a
lone pair, a single σ bond and a double bond (consisting of one σ
bond and one π bond). Three electron-dense sites surrounding a
central point in a 3-dimensional space will orient themselves at
120° so that electron /electron repulsions are minimized.
Carbonate ion:
KEY POINTS
1. Resonance structures differ only in the assignment of electron pair positions, NEVER atom
positions.
2. Resonance structures differ in the number of bond pairs between a given pair of atoms.
3. The actual structure is an average of the depicted resonance structures.
PRACTICE EIGHT Resonance Structures
Describe the electron arrangement in the nitrite anion, NO2-, using the localized electron model.
FORMAL CHARGE
Use formal charge to determine the most favored structure.
 Formal Charge - The number of valence electrons on the free element minus the number of
electrons assigned to the atom in the molecule.
- lone pair (unshared electrons belong completely to the atom in question.
- shared electrons are divided equally between the sharing atoms
 Atom’s formal charge = group number − [# of lone electrons − 2 (# of bonding electrons)]
 The sum of the formal charges must equal an ion’s charge.
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Use formal charges along with the following to determine resonance structure
 Atoms in molecules (or ions) should have formal charges as small as possible—as close to zero
as possible - called principle of electroneutrality.
 A molecule (or ion) is most stable when any negative formal charge resides on the most
electronegative atom.
PRACTICE NINE Determining formal charges
Draw all possible structures for the sulfate ion. Decide which is the most plausible using formal
charges.
REMINDER: Although formal charges are considered closer to the atomic charges than the oxidation
states, they are still only estimates and should not be taken as the actual atomic charges. Second, using
formal charges can often lead to erroneous structures, so tests based on experiments must be used to
make the final decisions on the correct description of bonding.
PRACTICE TEN Formal Charges II
Give possible Lewis structures for XeO3, an explosive compound of xenon. Which Lewis structure or
structures are most appropriate according to the formal charges?
MOLECULAR STRUCTURE - VSEPR
VSEPR: valence shell electron pair repulsion theory
 the structure around a given atom is determined principally by minimizing electron-pair
repulsions.
 Non-bonding and bonding electron pairs will be as far apart as possible.
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# of
electron
pairs
2
Electronic Geometry
(NOT molecular
geometry!)
linear
3
Trigonal planar
4
Tetrahedral (or
pyramidal)
5
Trigonal bipyramidal
6
Octahedral





Arrangement of electron
pairs
Arrangement of electron pairs
Electronic geometry – all the electron pairs surrounding a central atom are considered.
Molecular geometry - the arrangement in space of the atoms bonded to a central atom and
nonbonding electrons become “invisible”. This is not necessarily the same as the structural pair
geometry.
ANY time lone pairs are present, the electronic and molecular geometries will be different.
Each lone pair or bond pair repels all other lone pairs and bond pairs - they try to avoid each
other making as wide an angle as possible.
Lone pairs have a different repulsion since they are experiencing an attraction or “pull” from
only one nucleus as opposed to two nuclei. They also take up more space around an atom as you
can see on the left.
- works well for elements of the s and p-blocks
- VSEPR does not apply to transition element compounds (exceptions).
- They do not cause distortion when bond angles are 120° or greater
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Effect of Unshared Electron Pairs
TO DETERMINE THE GEOMETRY:
1. Sketch the Lewis dot structure – do not skip this step!
2. Determine the number of electron pair groups surrounding the central atom(s). Remember that
double and triple bonds are treated as a single group.
3. Determine the geometric shape that maximizes the distance between the electron groups. This is the
geometry of the electron groups.
4. Mentally allow nonbonding pairs to become invisible. Determine the molecular geometry by looking
at the remaining arrangement of atoms (as determined by the bonding electron groups) around the
central atom.
Molecular shapes for central atoms with normal valence
 This will be no more than four structural pairs if the atom obeys the octet rule.
 Since no lone pairs are present, the molecular and structural pair [or electronic] geometry is the
same. 109.5° bond angle.
 Ignore lone pairs AFTER you’ve determined the angles and only the relative positions of the
atoms are important in molecular geometry
 Examples are ammonia (107.5° bond angle) and water (104.5° bond angle).
Molecular shapes for central atoms with expanded valence
 only elements with a principal energy level of 3 or higher can expand their valence and violate
the octet
 rule on the high side. Why?
 d orbitals are needed for the expansion to a 5th or 6th bonding location—the combination of 1 s
and 3 p’s provides the four bonding sites that make up the octet rule.
 seems to be a limit of 3 lone pairs about the central atom
 Examples are XeF4 (equatorial and axial)
PRACTICE ELEVEN Prediction of Molecular Structure I
Describe the molecular structure of the water molecule.
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PRACTICE TWELVE Prediction of Molecular Structure II
When phosphorus reacts with excess chlorine gas, the compound phosphorus pentachloride (PCl5) is
formed. In the gaseous and liquid states, this substance consists of PCl5 molecules, but in the solid state
it consists of a 1:1 mixture of PCl4+ and PCl6- ions. Predict the geometric structures of PCl5, PCl4+, and
PCl6-.
Recap including bond angles for all
• structural pairs - σ bond (π bond pairs occupy the same space) pairs about an atom
A. 2 at 180° → linear (and planar)
B. 3 at 120° → trigonal planar
C. 4 at 109.5° → tetrahedral
D. 3 at 120° & 2 at 90° with each other through central atom → trigonal bipyramidal
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E. 6 at 90° → octahedral
*The presence of lone pairs alters the six basic MOLECULAR geometries, but the electronic or
structural pair geometry remains one of these six basic types.
MOLECULAR POLARITY


polar--bonds can be polar while the entire molecule isn’t and vice versa.
dipole moment--separation of the charge in a molecule; product of the size of the charge and the
distance of separation – learned this earlier.
• align themselves with an electric field.
• align with each other as well in the absence of an electric
field
• water — two lone pairs establish a strong negative pole
• ammonia — one lone pair which establishes a negative
pole
• note that the direction of the arrow indicating the dipole
moment always points to the negative pole with the cross
hatch on the arrow at the positive pole.
Nonpolar molecule - octet rule is obeyed AND all the surrounding bonds are the same; the molecule is
nonpolar since all the dipole moments cancel each other out. Example: CO2
Polar Molecule - octet rule is obeyed and all the surrounding bonds are NOT the same; the molecule is
polar since all the dipole moments do not cancel each other out. Example: NH3. Methane is a great
example. Replace one H with a halogen and it becomes polar. Replace all and it iss nonpolar again!
PRACTICE THIRTEEN Drawing molecules
Draw CH4 , CH3Cl, CH2Cl2, CHCl3, CCl4. Indicate dipole moment(s) where necessary.
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PRACTICE FOURTEEN Prediction of Molecular Structure III
Because the noble gases have filled s and p valence orbitals, they were not expected to be chemically
reactive. In fact, for many years these elements were called inert gases because of this supposed inability
to form any compounds. However, in the early 1960s several compounds of krypton, xenon, and radon
were synthesized. For example, a team at the Argonne National Laboratory produced the stable colorless
compound xenon tetrafluoride (XeF4). Predict its structure and whether it has a dipole moment.
PRACTICE FIFTEEN Structures of Molecules with Multiple Bonds
Predict the molecular structure of the sulfur dioxide molecule. Is this molecule expected to have a dipole
moment?
SUMMARY
CHEMICAL BONDS—Forces of attraction that hold groups of atoms together within a molecule or
crystal lattice and make them function as a unit.
IONIC
Characteristics of ionic substances usually include:
 electrons that are transferred between atoms having high
differences in electronegativity (greater than 1.7)
 compounds containing a metal and a nonmetal (Remember that
metals are located on the left side of the periodic table and
nonmetals are to the right of the “stairs”.)
 strong electrostatic attractions between positive and negative ions
 formulas given in the simplest ratio of elements (empirical
formula; NaCl, MgCl2)
 crystalline structures that are solids at room temperature
 ions that form a crystal lattice structure as pictured in Figure 1
 compounds that melt at high temperatures
 Substances that are good conductors of electricity in the molten
or dissolved state
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COVALENT
Characteristics of covalent substances usually include:
 the sharing of electrons between atoms having small differences in electronegativities (less than
1.7)
 nonmetals attracted to other nonmetals
 formulas that are given in the true ratios of atoms (molecular formulas; C6H12O6)
 substances that may exist in any state of matter at room temperature (solid, liquid, or gas)
 compounds that melt at low temperatures
 substances that are nonconductors of electricity
METALLIC
Characteristics of metallic substances usually include:
 substances that are metals
 a “sea” of mobile or delocalized electrons surrounding a positively charged metal center
 an attraction between metal ions and surrounding electrons
 formulas written as a neutral atom (Mg, Pb)
 solids with a crystalline structure at room temperature
 a range of melting points—usually depending on the number of valence electrons
 substances that are excellent conductors of electricity since the electrons in the “sea” are free to
move.
 Most chemical bonds are in fact somewhere between purely ionic and purely covalent.
 When answering bonding multiple choice or free-response questions, draw the Lewis structure.
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