Chemical Bonds
• Forces that hold atoms together
• Ionic bonds: the forces of attraction
between ions; involves a transfer of
electrons.
• Covalent bonds: the forces of attraction
between two atoms that are sharing
electrons
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Ionic Bonds
• Result from reaction between metal and
nonmetal to form a cation and an anion
• An ionic bond is the attraction between a
positive ion and negative ion.
• The ions are arranged in a pattern called a
crystal lattice.
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Crystal Structures
sodium chloride
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Covalent Bonds
• Sharing pairs of electrons
• Molecules attracted to each other weakly
• Often found between nonmetal atoms
N2
Cl2
HF
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Bond Polarity
• Covalent bonding between unlike atoms
results in unequal sharing of the electrons.
• The result is bond polarity; HF
 H •• F 
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Electron Sharing in HF
If electrons were equally shared
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Actual molecule
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Electronegativity
• Measure of the inherent ability of an atom to attract shared
electrons
• A larger electronegativity means that an atom attracts shared
electrons more strongly; fluorine in HF
• The difference in electronegativity between two atom in a bond is
a measure of bond polarity. A larger difference in electronegativity means a more polar bond.
If the difference is ≥2, the bond is ionic
If the difference is >0 and <2, the bond is polar covalent
If the difference is 0, the bond is covalent
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Electronegativity (cont.)
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Choose the bond in each pair that will be more polar.
• H-P or H-C:
• O-F or O-I:
• N-O or S-O:
• N-H or Si-H:
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Classify the following bonds as ionic,
covalent, or polar covalent
•
•
•
•
•
HCl
NaF
Cl2
Cl-F
NH3
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Dipole Moment
• Any molecule that has a center of positive
charge and a center of negative charge at
different points in space is said to have a
dipole moment.
H-F
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Dipole Moment (cont.)
• In polyatomic molecules the dipole
moment depends upon the atoms involved
and the three-dimensional structure.

H
H
O  

H

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O


H
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Dipole Moment (cont.)
The dipole moment affects the attractive forces between
molecules, and therefore the physical properties of the
substance. For example, water readily dissolves NaCl.
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Electron Arrangements
and Ionic Bonding
• Metals lose their valence electrons to form cations.
• Nonmetals gain electrons to form anions.
• Both try to acquire the electron configuration of a
noble gas; for example, NaCl
Na: [Ne]3s1
→
[Ne] + 1eCl: [Ne]3s23p5 + 1e- → [Ar]
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What are the electron configurations of the
ions in the following compounds?
• Al2S3
• MgO
• SrF2
• LiCl
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Electron Arrangements
and Ionic Bonding (cont.)
• Formulas are predicted by achieving electrical
neutrality using the component cations and anions.
What is the formula of the ionic compound formed from
Al and S?
• In polyatomic ions, the atoms in the ion are
connected with covalent bonds. The ions are
attracted to oppositely charged ions to form an
ionic compound.
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Properties of Ionic Compounds
• All are solids at room temperature.
– Melting points are greater than 300°C.
• Many are soluble in water. When dissolved
the solution becomes an electrical conductor.
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Bonding & Structure
of Ionic Compounds
• Crystal lattice: geometric pattern determined by the
size and charge of the ions.
• Anions are almost always larger than cations.
• Anions are generally considered “hard” spheres
packed as efficiently as possible, with the cations
occupying the “holes” in the packing.
• Compounds with polyatomic anions contain covalent
bonds within the anion structure, which are ionically
bonded to the cation; CuSO4
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Copper(II) sulfate
Cu2+
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SO4212 | 19
Lithium fluoride
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Lewis Structures
• Use the symbol of the element to represent the
nucleus and inner electrons.
• Use dots around the symbol to represent valence
electrons.
• Elements try to achieve a noble gas configuration.
sodium chloride
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Lewis Structures (molecules with covalent bonds)
•
Hydrogen shares two electrons (duet rule).
• Helium already has two, so it does not form bonds.
• The second row nonmetals require eight electrons to
fill the 2s and 2p orbitals (octet rule).
example: Cl2
•
Neon has eight electrons, so it does not form bonds.
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Writing Lewis Structures of Molecules
• Count the total number of valence electrons from all
the atoms.
• Attach the atoms together with one pair of electrons.
• Arrange the remaining electrons in pairs so that all
hydrogen atoms have two electrons (one bond) and
other atoms have eight electrons (combination of
bonding and nonbonding). Nonbonding pairs of
electrons are also know as lone pairs.
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Draw Lewis structures for the following molecules:
• NH3
• CCl4
• LiBr
• CH3OH
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Multiple Covalent Bonds
• Single covalent bond: atoms share two
electrons (one pair); HBr
• Double covalent bond: atoms share four
electrons (two pairs); C2H4
• Triple covalent bond: atoms share six
electrons (three pairs); HCN
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Lewis Structures of Molecules with
Multiple Bonds
• Determine the total number of valence
electrons.
• Form a single bond around each atom.
• Distribute the remaining electrons.
• Make double or triple bonds as needed until
each atom (except H) acquires an octet of
electrons.
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Draw Lewis structures for the following molecules:
• CO2
• NO3-1
• HCO3-1
• SO3
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Resonance
• When there are multiple Lewis structures for a
molecule that differ only in the positions of the
electrons, they are called resonance structures.
(Lone pairs and multiple bonds are in different
positions.) The real structure is a hybrid.
NO3-1
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Draw resonance structures for HCO3-1
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Exceptions to the Octet Rule
Some atoms may violate the octet rule.
beryllium: forms two bonds; BeCl2
boron and aluminum: normally form three bonds; BH3, AlCl3
phosphorus: can form up to five bonds; PCl5
sulfur: can form up to six bonds; SF6
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Molecular Structure
• Refers to the three-dimensional geometry of a molecule.
• Can specify a bond angle.
120°
• We will consider five different structures:
- linear
- trigonal planar
- tetrahedral
- trigonal pyramid
- bent
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Molecular Structure (cont.)
180°
• Linear: BeCl2
• Trigonal planar: BH3
120°
109.5°
• Tetrahedral: CCl4
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• Trigonal pyramid: NH3
• Bent: H2O
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Predicting Molecular Structure
• VSEPR Theory
– Valence Shell Electron Pair Repulsion
• The shape around the central atom can be predicted
by assuming that the electrons surrounding the
central atom will position themselves as far apart as
possible.
BeCl2
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Steps for Predicting Molecular Structure
• Draw the correct Lewis structure.
• Count the electron pairs and place them as far
apart as possible.
• Determine the positions of the atoms.
• Determine the name of the molecular structure
based on the positions of the attached atoms.
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Determine the molecular structures of the
following molecules:
• NH3
• CCl4
• AlCl3
• BeF2
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• Some molecules contain a combination of
more than one shape.
CH3OH
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Molecules with Multiple Bonds
• When using the VSEPR model to predict the
molecular geometry, multiple bonds are
counted the same as a single electron pair.
CO2
HCN
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