AP Chem: Unit 8
Bonding: General Concepts
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Types of Chemical Bonds

Electronegativity
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Bond Polarity and Dipole

Ions: Electron Configurations and Sizes

Energy Effects in Binary Ionic Compounds

Partial Ionic Character of Covalent Bonds

The Covalent Chemical Bond

Covalent Bond Energies and Chemical Reactions

The Localized Electron Bonding Model

Lewis Structures

Exceptions to the Octet Rule

Resonance

Molecular Structure: The VSEPR Model
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Types of Chemical Bonds
Ionic Bonds

Ionic Bonds are formed when an atom that loses electrons relatively easily reacts with
an atom that has a high attraction for electrons.
o Ionic Compounds results when a metal bonds with a nonmetal.
Bond Energy


Bond energy is the energy required to break a bond.
The energy of interaction between a pair of ions can be calculated using Coulomb’s law:
o r = the distance between the ions in nm.
o Q1 and Q2 are the numerical ion charges.
o E is in joules
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When the calculated energy between ions is negative, that indicates an attractive force.
A positive energy is a repulsive energy.
The distance where the energy is minimal is called the bond length.
Covalent Bonds

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Covalent bonds form between molecules in which electrons are shared by nuclei.
The bonding electrons are typically positioned between the two positively charged
nuclei.
Polar Covalent Bonds
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
Polar covalent bonds are an intermediate case in which the electrons are not
completely transferred but form unequal sharing.
A δ- or δ+ is used to show a fractional or partial charge on a molecule with unequal
sharing. This is called a dipole.
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Electronegativity
Electronegativity is the ability of an atom in a molecule to attract shared electrons to itself.
(electron love)
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Relative electronegativities are determined by comparing the measured bond energy
with the “expected” bond energy.
Measured in Paulings. After Linus Pauling the American scientist who won the Nobel
Prizes for both chemistry and peace.
Expected H-X bond energy=
Electronegativity values generally increase going left to right across the periodic table
and decrease going top to bottom.
Electronegativity and Bond type
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Bond Polarity and Dipole
Dipoles and Dipole Moments
A molecule that has a center of positive charge and a center of negative charge is said to be
dipolar or to have a dipole moment.
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An arrow is used to show this dipole moment by pointing to the negative charge and the
tail at the positive charge.
Electrostatic potential diagram shows variation in charge. Red is the most electron rich
region and blue is the most electron poor region.
Dipole moments are when opposing bond polarities don’t cancel out.
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Example Problems: For each of the following molecules, show the direction of the bond
polarities and indicate which ones have a dipole moment: HCl, Cl2, SO3, CH4, H2S
o HCl
o Cl2
o SO3
o CH4
o H2S
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Ions: Electron Configurations and Sizes
Electron Configurations of Compounds

When two nonmetals react to form a covalent bond, they share electrons in a way that
completes the valence electron configurations of both atoms. That is, both nonmetals
attain noble gas electron configurations.

When a nonmetal and a representative-group metal react to form a binary ionic
compounds, the ions form so that the valence electron configuration of the nonmetal
achieves the electron configuration of the next noble gas atom and the valence orbitals
of the metal are emptied. In this way both ions achieve noble gas electron
configurations.
Predicting Ionic Formulas

To predict the formula of the ionic compound, we simply recognize that the chemical
compounds are always electrically neutral. They have the same quantities of positive
and negative charges.
Sizes of Ions
Size of an ion generally follows the same trend as atomic radius. The big exception to this
trend is where the metals become nonmetals and the ions switch charge.
A positive ion is formed by removing one or more electrons from a neutral atom, the
resulting cation is smaller than the neutral atom.
o Less electrons allow for less repulsions and the ion gets smaller.
An addition of electrons to a neutral atom produces an anion that is significantly larger
than the neutral atom.
o An addition of an electron causes additional repulsions around the atom and
therefore its size increases.
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Energy Effects in Binary Ionic Compounds
Lattice Energy
Lattice energy is the change in energy that takes place when separated gaseous ions are
packed together to form an ionic solid.
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
The lattice energy is often defined as the energy released when an ionic solid forms
from its ions.
Lattice energy has a negative sign to show that the energy is released.
Lattice Energy Example
Estimate the enthalpy of lithium fluoride and the changes of energy and lattice energy
during formation:
1.Break down LiF into its standard state elements (use formation reaction):
2.Use sublimation and evaporation reactions to get reactants into gas form (since lattice
energy depends on gaseous state). Find the enthalpies to these reactions:
3.Ionize cation to form ions for bonding. Use Ionization energy for the enthalpy of the
reaction.
4.Dissociate diatomic gas to individual atoms:
5.Electron addition to fluorine is the electron affinity of fluorine:
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6.Formation of solid lithium fluoride from the gaseous ions corresponds to its lattice
energy:
The sum of these five processes yields the overall reaction and the sum of the individual
energy changes gives the overall energy change or lattice energy:
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Lattice energy can be calculated with at form of Coulomb’s law:

Q is the charges on the ions and r is the shortest distance between the centers of the
cations and anions. k is a constant that depends on the structure of the solid and the
electron configurations of the ions.
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Partial Ionic Character of Covalent Bonds
Bond Character
Calculations of ionic character:
Even compounds with the maximum possible electronegativity differences are not
100% ionic in the gas phase. Therefore the operational definition of ionic is any
compound that conducts an electric current when melted will be classified as ionic.
The Covalent Chemical Bond
Chemical Bond Model
A chemical bond can be viewed as forces that cause a group of atoms to behave as a unit.
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Bonds result from the tendency of a system to seek its lowest possible energy.
Individual bonds act relatively independent.
Example:
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It takes 1652 kJ of energy required to break the bonds in 1 mole of methane.
1652 kJ of energy is released when 1 mole of methane is formed from gaseous
atoms.
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Therefore, 1 mole of methane in gas phase has 1652 kJ lower energy than the total
of the individual atoms.
One mole of methane is held together with 1652 kJ of energy.

Each of the four C-H bonds contains 413 kJ of energy.
Example:

CH3Cl contains 1578 kJ of energy:
o 1 mol of C-Cl bonds + 3 mol (C-H bonds)=1578 kJ
o C-Cl bond energy + 3 (413 kJ/mol) = 1578 kJ
o C-Cl bond energy = 339 kJ/mol
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Properties of Models

A model doesn’t equal reality; they are used to explain incomplete understanding of
how nature works.
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Models are often oversimplified and are sometimes wrong.
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Models over time tend to get over complicated due to “repairs”.
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Remember that simple models often require restrictive assumptions. Best way to
use models is to understand their strengths and weaknesses.
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We often learn more when models are incorrect than when they are right.
o Cu and Cr.
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Covalent Bond Energies and Chemical Reactions
Bond Energies
Bond energy averages are used for individual bond dissociation energies to give approximate
energies in a particular bond.

Bond energies vary due to several reasons:
o multiple bonds, 4 C-H bonds in methane
o different elements in the molecule, C-H bond in C2H6 or C-H bond in HCCl3
Bond Energy Example:

CH4(g) CH3(g) + H(g)
435 kJ
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CH3(g) CH2(g) + H(g)
453 kJ

CH2(g) CH(g) + H(g)
425 kJ

CH(g) C(g) + H(g)
339 kJ
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Total 1652 kJ
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Average 413 kJ
Bond Energy Example

HCBr3
380 kJ

HCCl3
380 kJ

HCF3
430 kJ

C2H6
410 kJ
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A relationship also exists between the number of shared electron pairs.
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single bond – 2 electrons
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double bond – 4 electrons
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triple bond – 6 electrons
Bond energy values can be used to calculate approximate energies for reactions.

Energy associated with bond breaking have positive signs
o Endothermic process

Energy associated with forming bonds releases energy and is negative.
o Exothermic process
A relationship exists between the number of shared electron pairs and the bond length.
•As the number of electrons shared goes up the bond length shortens.
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ΔH = sum of the energies required to break old bonds (positive signs) plus the sum of the
energies released in the formation of new bonds (negative signs):
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D represents bond energies per mole and always has positive sign
n is number of moles
Bond Energy Example:
H2(g) + F2(g) è2HF(g)

1 H-H bond, F-F bond and 2 H-F bonds

ΔH = DH-H + DF-F – 2DH-F

ΔH= (1mol x 432 kJ/mol) + (1mol x 154 kJ/mol) – (2mol x 565 kJ/mol)

ΔH = -544 kJ
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The Localized Electron Bonding Model

The localized electron 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.
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Electrons are assumed to be localized on a particular atom individually or in the
space between atoms.
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Pairs of electrons that are localized on an atom are called lone pairs.

Pairs of electrons that are found in the space between the atoms are called bonding
pairs
Three parts of the LE Model:
1. Description of the valence electron arrangement in the molecule using Lewis structures.
2. Prediction of the geometry of the molecule using VSEPR model
3. Description of the type of atomic orbitals used by the atoms to share electrons or hold
lone pairs.
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Lewis Structures
The Lewis structure of a molecule show how the valence electrons are arranged among the
atoms in the molecule.

Named after G. N. Lewis

Rules are based on observations of thousands of molecules.

Most important requirement for the formation of a stable compound is that the
atoms achieve noble gas electron configurations.
Only the valence electrons are included.
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The duet rule: diatomic molecules can find stability in the sharing of two electrons.
The octet rule: since eight electrons are required to fill these orbitals, these
elements typically are surrounded by eight electrons.
Lewis Structure Steps
1.Sum the valence electrons from all the atoms. Total valence electrons.
2.Use a pair of electrons to form a bond between each pair of bound atoms.
3.Arrange the remaining electrons to satisfy the duet rule for hydrogen and the octet
rule for the others.
a) Terminal atoms first.
b) Check for happiness
Examples:
HF
N2
NH3
CH4
CF4
NO+
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Exceptions to the Octet Rule
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
Incomplete: An odd number of electrons are available for bonding. One lone
electron is left unpaired.
Suboctet: Less than 4 pairs of electrons are assigned to the central atom
o Suboctets tend to form coordinate covalentbonds
o BH3 + NH3
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Extended: The central atom has more than 4 pairs of electrons.
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At the third energy level and higher, atoms may have empty d orbitals that can
be used for bonding.
General Rules

The second row elements C, N, O, and F always obey the octet rule

The second row elements B and Be often have fewer than eight electrons around
them in their compounds. They are electron deficient and very reactive.
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The second row elements never exceed the octet rule, since their valence orbitals
can only hold 8.
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Third-row and heavier elements often satisfy the octet rule but can exceed the octet
rule by using their empty valence d orbitals.

When writing the Lewis structure for a molecule, satisfy the octet rule for the atoms
first. If electrons remain after the octet rule has been satisfied, then place them on
the elements having available d orbitals
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Resonance

Resonance is when more than on valid Lewis structure can be written for a
particular molecule. The resulting electron structure of the molecule is given by the
average of these resonance structures.

The concept of resonance is necessary because the localized electron model
postulates that electrons are localized between a given pair of atoms. However,
nature does not really operate this way. Electrons are really delocalized- they move
around the entire molecule. The valence electrons in a resonance structure
distribute themselves equally and produce equal bonds.
Formal Charge
Some molecules or polyatomic ions can have several non-equivalent Lewis structures.
•Example: SO42-
Because of this we assign atomic charges to the molecules in order to find the right
structure.
The formal charge of an atom in a molecule is the difference between the number of valence
electrons on the free atom and the number of valence electrons assigned to the atom in the
molecule

Formal charge = (# of valence electrons on neutral ‘free atom’) – (# of valence
electrons assigned to the atom in the molecule)
Assumptions on electron assignment:

Lone pair electrons belong entirely to the atom in question.

Shared electrons are divided equally between the two sharing atoms.
Formal Charge Example: SO42-: All single bonds
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Formal charge on each O is -1
Formal charge on S is 2
Formal Charge Example: SO42-: two double bonds, two single
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Formal charge on single bonded O is -1
Formal charge on double bonded O is 0
Formal charge on S is 0
Formal Charges
1. Atoms in molecules try to achieve formal charges as close to zero as possible.
2. Any negative formal charges are expected to reside on the most electronegative
atoms.
If nonequivalent Lewis structures exist for a species, those with formal charges closest
to zero and with any negative formal charges on the most electronegative atoms are
considered to best describe the bonding in the molecule or ion.
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Molecular Structure: The VSEPR Model
VSEPR
Valence shell electron repulsion model is useful in predicting the geometries of molecules
formed from nonmetals.

The structure around a given atom is determined principally by minimizing electron
– pair repulsion.

From the Lewis structure, count the electron pairs around the central atom.

Lone pairs require more room than bonding pairs and tend to compress the angles
between the bonding pairs.
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Multiple bonds should be counted as one effective pair.
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With a molecule with resonance, all structures should yield the same shape.
Linear
Trigonal Planer
Tetrahedral Arrangements
Tetrahedral
Trigonal Pyramidal
Bent/V
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Bipyramidal Arrangements
Octahedral Arrangements
Molecules without a central atom

The molecular structure of more complicated atoms can be predicted from the
arrangement of pairs around the center atoms. A combination of shapes will
result that allows for minimum repulsion throughout.
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