Bonding
Ms. Pollock
Chem 1
2014 - 2015
Ions
• Elements in same group – same
properties
• Dictated by configuration of
valence electrons
• Gain or loss of electron =
creation of ion
Octet Rule
• Noble gases unreactive due to
electron configuration – full
outer shell
• Elements form compounds in
ways that give them eight
valence electrons – like noble
gases.
• Exception first row – only two
electrons possible
Electron Dot Diagrams
• Common way to keep track of
valence electrons
• Chemical symbol with one dot
for each valence electron
• Dots match up to group
numbers on periodic table
Cations and Anions
• Metals tend to lose electrons.
Nonmetals tend to gain
electrons.
• Isoelectronic – same numbers of
electrons
Cations
• Positively charged ion
• More than one species possible
with transition metals
• Examples Na+, Ca2+, Al3+, Ag+,
Au3+
Anions
• Negatively charged ion
• Cl -, O2-, N3-
Ionic Bonds and Ionic Compounds
• Gain or loss of electrons result of
interaction between two
different chemical species
• Ionic bond – attraction between
positively charged cations and
negatively charged anions
• Occur in valence (outer) shell
Crystal Lattices
• Ionic compounds composed of
large numbers of anions and
cations
• Cations attracted to anions and
repelled by other cations
• Crystal lattice – 3D structure
formed between interactions of
cations and anions
• Regular, repeating pattern
Lattice Energy
• Energy needed to completely
break apart ionic substance into
isolated ions
• Not directly measurable
• Relative strengths of bonds
Properties of Ionic Compounds
• Form crystals
• High mp, high bp
• Hard and brittle
• Conduct electricity when
dissolved in water
Metals and Metallic Bonds
• Metals about 25% of earth’s
crust
• Typically not found alone in
nature
• Generally cation portion of
compound
Properties of Metals
• Good conductors of electricity
and heat
• Malleable
• Ductile
• Lustrous
• Resistant to corrosion
• High melting point, high boiling
point
The “Sea of Electrons”
• Behavior of metals explained by
bonding of metal atoms to make
solid
• Crystalline solids, but every atom
in lattice identical
• Electrons in outer energy level
mobile – drift from one atom to
another – delocalized electrons
• Explains properties of metals
Types of Metals
• Precious
• Very resistant to corrosion
• Scarce
• High IE, high EN
• Rare Earth
• Difficult extraction
• Used in many industrial processes
• Alloys
• Mixtures of metals
• Properties often different from base
elements
• Mercury-based alloys called
amalgams
Lewis Electron Dot Structures
• Covalent bonds formed when
two or more elements share
electrons
• Electrons not fully possessed by
single atom – shared between
both
• First proposed in 1916 by G.N.
Lewis, expanded by Linus
Pauling
Single Covalent Bonds
• Two atoms sharing two valence
electrons
• Simplified representation in Lewis
structure
• Each dot one valence electron
• Dots between atoms shared in
covalent bond; also represented by
single straight line
• Number of covalent bonds formed
correlates to number of valence
electrons needed to make octet
Single Covalent Bonds
• Pairs of electrons not involved in
covalent bonding lone pairs
• Contribute to overall shape of
molecule
Double and Triple Bonds
• Double bonds – sharing of two
pairs of electrons
• Triple bonds – sharing of three
pairs of electrons
• Represented with two or three
lines between bonded atoms
Steps for Drawing Lewis Structures
• 1. Identify atoms participating in covalent bond.
• 2. Draw element symbols and represent valence electrons with dots.
• 3. Use the octet rule to predict the number of covalent bonds that
will form.
• 4. Draw bonding atoms next to each other, indicating shared pairs
with lines.
Molecular Geometry
• Three-dimensional arrangement
of atoms in a molecule
• Affects physical and chemical
properties
• Predictable based on Lewis
electron-dot structure
VSEPR Theory
• Valence shell electron-pair
repulsion theory
• Molecule will arrange self so that
outer shell electron pairs stay as far
apart as possible
• Opposites repel
• Classification based on number of
bonding and nonbonding pairs
around central atom
• Single/double/triple bonds same in
terms of repulsion
Central Atom with No Lone Pairs
• A = central atom
• B = atoms surrounding central
atom
• Subscripts after B to denote
number of B atoms bonded to
central atom
AB2
• Beryllium hydride, carbon
dioxide
• Central atom with two single
bonds
• Bonded atoms directly opposite
on both sides of central atom
• Linear molecule
• 180 bond angle
AB3
• Central atom with three single
bonds
• Trigonal planar
• Bonded atoms positioned at
vertices of equilateral triangle
• 120 bond angles
• All atoms in same plane
AB4
• Central atom with four single
bonds
• Bonds unable to lie in same
plane to maximize distance from
each other
• Tetrahedral shape
• Face equilateral triangle
• Bond angles 109.5
AB5
• Central atom with five single
bonds
• Violates octet rule – ten
electrons on central atom;
access to d orbitals
• Bond angles not equivalent –
some 120, some 90
• Trigonal bipyramidal geometry
AB6
• Central atom with six single
bonds
• Exceeds octet rule
• All bonded atoms equivalent
• All bond angles 90
• Octahedral geometry
Central Atom with One or More Lone Pairs
• Lone pairs change molecular
geometry
• Electron domain geometry –
number of bonds on central
atom + number of lone pairs on
central atom
• Based on molecular geometry,
just modified for lone pairs
Ammonia
• Three single bonds, one lone
pair
• Domain pair geometry
tetrahedral (four electron pairs)
• Molecular geometry trigonal
pyramidal (one unbonded pair)
• Bond angles 107 - lone pairs
slightly more repulsive than
bonded pairs
Water
• Two bonding pairs, two lone
pairs
• Tetrahedral domain geometry
• Molecular geometry bent
• Bonding angle 104.5
Polarity in Chemical Bonds
• Result of greater attraction for
electrons by one atom in a bond
than by another atom
• Electrons unequally shared
Electronegativity
• Ability to attract shared
electrons
• Higher value = greater pull for
electrons
• Trend: increase left to right
(greater number of protons),
decrease top to bottom
(electron shielding)
• Most electronegative element
fluorine
Polar Bonds
• Difference in electronegativity =
polar bond
• More electronegative element
stronger attraction for electrons
• Each atom has apparent partial
negative charge
Classifying Chemical Bonds
• 1. Pure nonpolar covalent bonds only between two identical atoms
(no difference in electronegativity).
• 2. Electronegativity difference 0.4 or less polarity minimal
• 3. Electronegativity difference 0.5 – 1.6 electrons shared but
significantly polarized
• 4. Ionic bonds electronegativity 2.0 and above, generally one metal
and one nonmetal
Intermolecular Forces
• Collection of gas molecules
simple interaction of matter
• Individual molecular far apart –
only fleeting interaction
• Liquid and solid particles
clumped, so constantly
interacting
Ion-Ion Interactions
• Most significant in solid state
• Formation of lattices
• Can also involve polyatomic ions
Dipole-Dipole Interactions
• Dipole – two opposite charges
separated by some distance
(polar bond example)
• Molecular dipole – geometric
sum of all individual bond
dipoles in a molecule
• Must have at least one polar
bond
• Net zero molecular dipole in
symmetrical molecules
Dipole-Dipole Force
• Two polar molecules interact
• Partial positive region of one
molecule attracted to partial
negative region of adjacent
molecule
• Both charges partial, so not as
strong as ion-ion interaction
Ion-Dipole Interactions
• Interaction between fully
charged entity and polar
molecule
• Involves both cations and anions
• Cation attracted to partial
negative
• Anion attracted to partial
positive
Dispersion Forces
• Attractive forces that arise as a
result of temporary dipoles
induced in atoms or molecules
• London dispersion forces
• Relatively weak, no requirement of
permanent polarity
• Caused by local and temporary
environmental changes
• Necessary for nonpolar substances
to form solids or liquids
The Hydrogen Bond
• Only when hydrogen bonded to
fluorine, oxygen, or nitrogen (all
highly electronegative)
• Strong attraction of hydrogen
electrons by bonded atom –
hydrogen nucleus exposed,
strong attraction to nearby lone
pairs
• Very important to water
• Not always drawn, but always
present
The Hydrogen Bond
Valence Bond Theory
• Covalent bond = overlap of
electron clouds in two atoms
• Result of combination of atomic
orbitals
• Valence bond theory: quantum
mechanical theory; covalent
bonds formed by overlap of
partially filled atomic orbitals
Hybrid Orbitals
• Bonding scheme related to
molecular geometries predicted
by VSEPR theory
• Need hybrid orbitals
• Overlap of existing orbitals not
sufficient to explain some
bonding and molecular
geometries
3
sp
Hybridization
• Hybridization – mixing of atomic
orbitals in an atom to produce
set of hybrid orbitals
• Must involve nonequivalent
orbitals
• tetrahedral
2
sp
Hybridization
• Three occupied orbitals in
valence shell
• Paired s electron promoted to p
orbital
• Trigonal planar
sp Hybridization
• Linear, bent
• S electron promoted to p orbital
Hybridization in Molecules with Double or
Triple Bonds
• Hybridization important to double
or triple bond geometry
• Must designate between types of
covalent bonds
• Sigma – overlap of orbitals end-toend
• Pi – overlap of orbitals side-to-side
• Single bond = sigma
• Double bond = 1 sigma, 1 pi
• Triple bond = 1 sigma, 2 pi
Molecular Orbitals
• Atomic orbitals combined to
make molecular orbitals
• Important to understanding of
molecular properties