What is a Bond?
• A force that holds atoms together
How do you determine the strength of the bond?
• Bond Energy- the energy required to break a bond
Why are compounds formed?
• Because it gives the system the lowest potential energy
3 Types of Bonds
1. Metallic bonding
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Found between metal cations
Attraction between the nuclei of the cations with a sea of valence electrons
The “sea of electrons model explains the malleability and ductility of metals as well as
their great ability to conduct electricity
2. Ionic Bonding
• An atom with a low ionization energy reacts with an atom with high electron affinity,
also known as a metal reacting with a nonmetal. The electron is transferred from the
metal to the nonmetal opposite charges hold the atoms together (electrostatic
attraction)
Ions
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Atoms tend to react to form noble gas configurations
Metals lose electrons to form cations (+) charged particles; charge is the group # next to
letter A
Nonmetals gain electrons to form anions (-) charged particles; charge is 8- group # next
to letter A
Examples
Predict the ion generally formed by Sr, S, and Al
Predicting Formulas for Ionic Compounds
• Referred to as in solid state
• Ions align themselves to maximize attractions between opposite charges and minimize
repulsion between like ions
• React to achieve noble gas configuration
• Chemical formula is always neutral and is a formula unit (the simplest way in which ions
react together
Examples: calcium and oxygen
Sizes of Ions
magnesium and nitrogen
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Ion sizes increase down a group
Cations are smaller than the atoms they came from, anions are larger
Ion sizes decrease across a period
Size of isoelectronic ions
• Iso- means the same
• Isoelectronic atoms have the same number of electrons
Al+3 Mg+2 Na+1 Ne F-1 O-2 N-3 - all have 10 electrons ( 1s2 2s2 2p6 )
• The ion that has the highest proton number is the smallest in size
Examples
Arrange these ions in order of decreasing size Se-2, Br-1, Rb+1, Sr+2
Arrange these atoms and ions in order of decreasing size: Mg+2, Ca+2, Ca
Choose the largest ion in each of the following categories
a. Li+1 , Na+1 , K+1 , Rb+1 , Cs+1
b.
Ba+2 , Cs+1 , I-1 , Te-2
Formation of binary Ionic Compounds
• Lattice energy- the energy associated with making a solid ionic compound from its
gaseous ions M+ (g) + X- (g) --> MX(s)
• This is the energy that “pays” for making ionic compounds.
• (-) means exothermic- releases energy (+) means endothermic- absorbs energy
• Energy is a state function so we can get from reactants to products in a round about way.
Example:
Na(s) + ½ F2(g) --> NaF(s)
First sublime Na:
Na(s) --> Na(g)
Enthalpy of sublimation
Ionize Na(g)
Na(g) --> Na+(g) + e- Ionization energy
Δ H = 495 kJ/mol
Break F-F bond
½ F2(g)
Δ H = 154/2
=77 kJ/mol
Add electron to F
F(g) + e- --> F- (g)
Electron affinity
Δ H = -328 kJ/mol
Formation of solid
sodium fluoride
Na+(g) + F- (g) --> NaF(s) Lattice Energy
Δ H = -928 kJ/mol
--> F(g)
Energy Change is -575 kJ/mol
Lattice Energy Calculations
Δ H = 109 kJ/mol
To measure the energy of interaction between ions, we use Coulomb’s Law
Eion pair = 2.31 x 10-19 J · nm (Q1Q2)/r
• Q is the charge number
• R is the distance between the centers in (nm)
• K is a constant that depends on the structure of the crystal
Example: NaCl
Lattice E = 2.31 x 10-19 J · nm (+1)(-1)/.276nm
= -8.37 x 10-19 J
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If the charges are opposite, E is negative and thus considered exothermic- they release E
and thus have a lower energy than the separated ions
Same charge, positive E, endothermic energy is required to bring them together
Lattice energy is greater with more highly charged ions and as the distances between the
ions decrease
The bigger the lattice energy “ pays” for the extra ionization energy
Also “pays” for unfavorable electron affinity
Example
Arrange the following ionic compounds in order of increasing lattice energy
NaF, CsI , CaO
Which one has the greatest lattice energy?
AgCl, CuO, CrN
Characteristics of Ionic Compounds
• An ionic compound is defined as a substance that contains ionic bonds due to a transfer
of electrons; strong forces of attraction
• Found in solid state
• conducts electricity when melted or dissolved in water
• also known as a salt
• have rigid, crystalline structures that can fracture easily
• high melting points
2. Covalent Bonding
• The electrons in each atom are attracted to the nucleus of the other. The electrons repel
each other. The nuclei repel each other
• They reach a distance with the lowest possible energy- this distance is called the bond
length .
See figure 8.1 b in book
• Electrons are shared by atoms
• Nonpolar covalent bonds- two identical atoms share electrons equally
• Polar covalent bonds- electrons are not shared equally; one end is slightly positive and
the other is slightly negative, indicated using a small delta (δ)
δ+ δ-
H–F
In an electric field, HF orients itself accordingly, implying it has partial charge distribution.
Electronegativity
• The ability of an atom to attract shared electrons to itself
• Trend tends to increase from left to right across a period and decreases down a group
• Noble gases are not discussed
• Difference in electronegativity between atoms tells us type of bond(ionic v covalent)
(polar covalent v nonpolar covalent)
Electronegativity difference
Zero
Bond type
0- .3
Nonpolar Covalent
Intermediate .3-1.67
Polar covalent
Large
Ionic
> 1.67
Example
Rank the bonds according to polarity : H-H, O-H, Cl-H, S-H, F-H
Dipole Moments
• A molecule with a center of negative charge and a center of positive charge is dipolar
(two poles) Or
has a dipole moment
• Center of charge does not have to be on an atom
• How its drawn: the arrow head is towards the negative side
δ+

δ-
H–F
Example
Which bond is more polar (a) B-Cl or C-Cl (b)
Example
Which molecules have dipole moments? HCl- linear,
CH4-tetrahedral and
H2S-bent
Characteristics of Covalent Compounds
P-F or
P-Cl
Cl2- linear,
SO3-trigonal planar,
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A covalent or molecular compound is defined as a substance that contains covalnet bonds
due to a tsharing of electrons; weaker forces of attraction in general when compared to
ionic bonds
Found mostly in liquid or gas state; some soft solids
Does not conduct electricity when melted or dissolved in water
Low melting points/boiling points
Drawing Lewis Structures
Lewis Structure: shows how the valence electrons are arranged, one dot for each valence
electron
• A stable compound has all of its atoms with a noble gas configuration
• Hydrogen follows the duet rule, all other nonmetals follow the octet rule.
• A bonding pair of electrons are between the atoms
• A lone pair are not between atoms, they are also called unshared pair
Rules to drawing a structure
1. Add up all the valence electrons for the atoms within the chemical formula
2. Arrange the atoms in a skeleton structure- Carbon is always central, H is never
central, & the least electronegative is central if carbon is not present- When
looking at a chemical formula, the way it’s written usually is the order it is drawn
or if there is a central atom, it is usually written first.
3. Connect the atoms with a bonding pair of electrons with a dash.
4. Complete the octets of the atoms bonded to the central atom by distributing pairs
of electrons to external atoms( remember hydrogen can only have 2 electrons
(duet rule)
5. Place any leftover electrons on the central atom even if doing so results in more
than an octet
6. If there are not enough electrons to give the central atom and octet, try multiple
bonds ( double or triple bond)
Examples:
PCl3
H2 O
HCN
NO+1
C2H4
CH2Cl2
Covalent bonds can be:
• Single bonds-one pair of electrons is shared
• Double and triple bonds- two and three pair of electrons are shared
• The more bonds, the shorter the average bond length
• The more bonds between atoms, the greater the bond energy
Bond type
Average Bond Energy
Bond length
C-C
C=C
347 kJ/mol
614 kJ/mol
154 pm
134 pm
Calculating ΔH for a reaction using bond energies (review)
• It takes energy to break bonds and end up with atoms (+)
• We get energy when we use atoms to form bonds ( -)
• ΔH = bonds broken – Bonds Formed
= reactant bonds – product bonds
Example Find the energy for this:
2 CH2 =CHCH3 + 2 NH3 + O2  2 CH2 = CHC=N + 6 H2O
C-H
C=C
N-H
C-C
413 kJ/mol
614 kJ/mol
391 kJ/mol
347 kJ/mol
O-H 467 kJ/mol
O=O 495 kJ/mol
C=N 891 kJ/mol
Localized Electron Model
• Simple model stating that a molecule is composed of atoms that are bound together by
sharing pairs of electrons using the atomic orbitals of the bound atoms
• 3 parts:
1) Using valence electrons to create Lewis Structures
2) Prediction of geometry using VSEPR theory
3) Description of the types of orbitals (Chapter 9)
Exceptions to the Octet Rule:
• Be and B often do not achieve an octet; beryllium is stable with 2 electrons and boron
often has 6 valence electrons and is considered stable; it acts electron deficient and form
coordinate covalent bonds
BH3
BeCl2
How BH3 reacts with NH3 ammonia or other compounds with a free lone pair
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3rd row and higher elements can satisfy or exceed the octet-use empty3d orbitals to hold
electrons; place electrons around atoms to acquire 8 and then place extra electrons on the
central atom which should be a 3rd period element
SF6
I3-1
ClF3
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ICl4-1
Free radicals: atoms and molecules that are considered stable with 7 valence electrons;
very reactive because of 1 unpaired electron
NO
NO2
Resonance
• sometimes there is more than one valid structure for a molecule or ion
• In order to draw resonance, the atom arrangement does not change, just the location of
the electrons is different.
• Use double arrows to indicate it is the “average” of the structures-IT DOES NOT
SWITCH BETWEEN THEM!!
• NO2-1
NO3-1
Formal Charges
• For molecules and polyatomic ions that exceed the octet, there are several different
structures so we evaluate which structure if several can be drawn is the most accurate
• How to calculate:
FC = valence number of electrons – (total # of lone electrons around the atom + # of bonds connected to that atom
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Most valid structure will have
1. formal charges as close to zero for all atoms
2. negative formal charges will be placed on the most electronegative atom
To check your formal charge calculations: the sum of the formal charges must equal the
overall charge of molecule
Examples
SO4-2
HCN
ClO3-1
VSEPR theory -Valence Shell Electron Pair Repulsion
• The shape of a molecule can greatly affect its properties
• Use VSEPR to predict shape: molecules take a shape that puts electron pairs as far away
from each other as possible
• Have to draw the Lewis structure to determine the electron pairs
• Lone pair takes more space
• Multiple bonds count as one lone pair
• The number of electron pairs determines: bond angles and underlying shape= Electron
pair geometry
• The number & position of atoms determines the Molecular Geometry
Electron Pairs
2
3
4
5
6
Bond Angles
180º
120º
109.5º
90º & 120º
90º
Actual Shape
Electron pair/Molecular Geometry
linear
trigonal planar
tetrahedral
trigonal bipyramidal
octogonal
Electron Pairs
2
3
3
4
4
4
5
5
5
5
6
6
6
6
6
Bonding Pairs
2
3
2
4
3
2
5
4
3
2
6
5
4
3
2
Lone pairs off central
0
0
1
0
1
2
0
1
2
3
0
1
2
3
1
Molecular Geometry
linear
trigonal planar
bent(120º)
tetrahedral
trigonal pyramidal (107º)
bent(104.5º)
trigonal bipyramidal
See-saw (90º & 120º)
T-shaped(90º)
linear(180º & 90º)
octohedral
square pyramidal (90º)
square planar(90º)
T-shaped(90º)
linear(180º & 90º)
No lone pair off the central:
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BeCl2
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AB or AB2 both the electron pair geometry and the molecular geometry are linear
HCl
AB3 both the electron pair geometry and the molecular geometry are trigonal planar
BF3
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CH4
AB4 both the electron pair geometry and the molecular geometry are tetrahedral
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AB5 both the electron pair geometry and the molecular geometry are trigonal
bipyramidal
PCl5
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AB6 both the electron pair geometry and the molecular geometry are octohedral
PCl6-1
Lone pair off the central:
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AB3E the electron pair geometry is tetrahedral but the molecular geometry is trigonal
pyramidal
NH3
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H2 O
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SeF4
AB2E or AB2E2 the electron pair geometry is tetrahedral but the molecular geometry is
bent or v-shaped
SO2
AB4E the electron pair geometry is trigonal bipyramidal but the molecular geometry is
see saw
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AB3E2 the electron pair geometry is trigonal bipyramidal but the molecular geometry is
T- shaped
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AB2E3
linear
the electron pair geometry is trigonal bipyramidal but the molecular geometry is
AB4E2
planar
the electron pair geometry is octohedral but the molecular geometry is square
IF3
ClF2-1
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KrF4
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AB5E the electron pair geometry is octohedral but the molecular geometry is square
pyramidal
Predicting Polarity of Molecule
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Any molecule that retains a dipole moment is considered to be a polar molecule.
Any two atom molecule with a polar bond has a dipole moment and is considered a polar
molecule
With 3 or more atoms there are 2 considerations to be considered a polar molecule
- there must be a polar bond
- geometry cannot cancel it out
Draw dipoles of the polar bonds into the Lewis structure that has been drawn in the
correct shape
 If the dipoles move in equal but opposite directions, the dipoles cancel out, the
molecule is NONPOLAR
 If the dipoles move in the same direction, the dipoles remain, the molecule is
POLAR
 Hint: Think of an electric field and see how molecules orient themselves- if they
can’t orient, the dipoles cancel out due to geometry and the molecule is
NONPOLAR, if they do orient themselves, the molecule is POLAR
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When in doubt, memorize!
 Linear, planar and tetrahedral, trigonal bypyramidal and octohedral shapes with
identical bonds will cancel out the dipoles: NONPOLAR
 Square planar and triatomic linears are usually NONPOALR
 Usually a lone pair off the central atom makes a molecule POLAR:
Trigonal Pyramidal, Bent, T-Shaped, Square pyramidal, See saw
To Calculate Bond Order
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Bond order is the number of pairs of electrons in a bond
H2
bond order = 1
O2
bond order = 2
N2
bond order = 3
If a resonance structure: # electron pairs/ attachments
CO3-2
bond order of 1.3
Remember
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not all bonds are the same length. That is why we take an average bond length.
The more tight the atoms are, the stronger the bond, the shorter the bond and the larger
the bond dissociation energy
Sigma(σ) and Pi (Π )Bonds
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A sigma bond = a single bond; orbitals directly overlap each other
A pi bond is found in a double or triple bond; p orbitals on neighboring atoms align in a
parallel fashion
 A double bond would consist of a sigma and a pi bond
 A triple bond would consist of a sigma and 2 pi bonds
Sigma bonds are weaker than pi bonds
Hybridization
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Individual, isolated atoms have their own atomic orbitals. (S, P, D, F)
When atoms bond, electrons change their arrangement due to the influence of other atoms
in “bonding” orbitals by a process called hybridization
Orbitals mix together to rearrange to form new HYBRIDIZED orbitals
Linear= sp2
Trigonal planar= sp3
Tetrahedral = sp4
Trigonal Bipyramidal = sp3d
Octohedral = sp3d2
*********Watch podcast 9.1 attached to website regarding hybridization and fill in missing
pieces to the notes pages that follow the podcast************