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Bonding: General Concepts
Chemical Bonds
Electronegativity, Polarity
Ionic Bonds
Covalent Bonds: Lewis Structures, VSEPR
CHEMICAL BONDS
• Forces that hold groups of atoms together to form
molecules.
• The driving force is the lowering of energy due to
electrostatic attractions between the positive nuclei
and the negative electrons exceeding repulsions
between nuclei and between electrons..
• Separated atoms have zero energy and chemically
bonded atoms have negative (lower) energy. (Fig
8.1).
• The minimum energy or well corresponds to the
bond length
Figure 8.1 a & b (a) The Interaction of Two
Hydrogen Atoms (b) Energy Profile as a
Function of the Distance Between the
Nuclei of the Hydrogen Atoms
CHEMICAL BONDS (2)
• This lowering of energy is achieved when
atoms achieve a noble gas electron
configuration or an octet.
• We will see that bonds form in order that
each participating atom achieves an octet.
• We will also see that there are exceptions.
CHEMICAL BONDS (3)
• Form between atoms resulting in molecules
(covalent bonds, sharing of electrons).
• Form between ions resulting in ionic cmps
(ionic bonds, electron transfer).
• Chemical bonding model assumes molecule
consists of individual chemical bonds.
• Bond strength varies and is measured by
bond energy (kJ/mol) = energy required to
break a mole of bonds.
ELECTRONEGATIVITY
• Defined as the ability of an atom to attract
shared electrons in a covalent bond to itself.
• EN > 0; Fig 8.3
• EN largest in upper right hand corner of PT.
• This unequally sharing leads to unequal
charges on the atoms.
• Use δ+ and δ- to indicate partial charges on
the atoms.
Figure 8.3 The Pauling
Electronegativity Vaules
BOND POLARITY
• Polar covalent bond forms when electron
pair is not shared equally due to bonded
atoms having different EN values.
• ΔEN = difference in EN
– ~ 0, nonpolar covalent bond. E.g. H2, O2
– < 2, polar covalent bond; e-pair is held more
closely by atom with greater EN
– > 2, bond is ionic and electron is transferred to
form anion and cation (vs Sec 8.6)
Figure 8.12
a-c The
Three
Possible
Types of
Bonds
DIPOLE MOMENT
• When there is a separation of electron
charge leading to polar bonds, the molecule
may have a dipole moment.
– All diatomics with polar bonds have a dipole
moment. (HCl, NO, CO)
– Polyatomics with polar bonds MAY have a
dipole moment. (Fig 8.2). H2O, NH3, SO2)
Table 8.2 Types of Molecules with
Polar Bonds but No Resulting Dipole
Moment
Figure 8.6 a-c The Structure and
Charge Distribution of the
Ammonia Molecule
IONIC BONDS (8.4)
• (Metal) Cation + (Nonmetal) Anion 
Ionic Solid held together with ionic bonds.
• This solid has a continuous network of
cations surrounded by anions and anions
surrounded by cations.
• The formation of ionic bonds is driven by
favorable energy considerations: this is
illustrated by the Born-Haber cycle.
ATOMIC ION SIZE
• Cations shrink and anions expand as
electrons are removed or added to the
neutral atom.
• In an isoelectronic series, the number of
electrons stays the same, but Z is constant.
– As Z increases, the ion size decreases.
– Fig 8.8
• Note that
Figure 8.8 Sizes of Ions Related to
Positions of the Elements on the
Periodic Table
Born-Haber Cycle (Fig 8.9, 8.11)
•
•
•
•
•
•
•
Li(s)  Li(g)
Sublimation energy > 0
Li(g)  Li+(g) + eIE, T7.6
½ F2(g)  F(g) Dissociation energy > 0
F(g) + e-  F-(g) EA, T7.7
Li+(g) + F-(g)  LiF(s) Lattice energy
Sum all of these rxns to get energy for
Li(s) + ½ F2(g)  LiF(s) ΔHfo = -617
kJ/mol
Figure 8.9 The Energy Changes
Involved in the Formation of Lithium
Fluoride from Its Elements
Lattice Energy, U
• KF(s)  K+(g) + F-(g)
U>0
• Electrostatic attraction between Cation and
Anion.
• As charge increases, U increases.
COVALENT BONDS (8.7)
• Most common type of chemical bond.
• Involve electrons shared by two nuclei.
• The covalent bond model assumes that a
molecule is an arrangement of individual
bonds that form between 2 atoms because
the molecule is energetically favored (i.e.
energy is at a minimum) compared to the
separated atoms.
DISSOCIATION BOND
ENERGY
• Chemical bonds can be assigned average (±10%)
dissociation bond energies (T8.4) and bond
lengths (T8.5)
• D > 0 kJ/mol; measure of bond strength.
• AB(g)  A(g) + B(g)
• Note single vs double vs triple bonds D values.
• ΔHrxn ≈ Σ D(bonds in R) – ΣD (bonds in P)
because bond breaking is endothermic and bond
formation is exothermic.
Table 8.4 Average Bond Energies
(kj/mol)
Table 8.5 Bond Lengths for
Selected Bonds
COVALENT BONDS (2)
• Determine physical and chemical properties
of cmps.
• Determine the likelihood and products of
chemical reactions.
• Determine molecular shape (Sec 8.13).
LOCALIZED ELECTRON (LE)
BONDING MODEL
• Valence electrons participate in the
formation of chemical bonds.
• Electron pairs are localized between (shared
or bonding pair) or on (lone pair) atoms
such that each atom has an octet or duet of
electrons.
• VSEPR model predicts molecular geometry
based on LE bonding model.
LEWIS SYMBOLS and
STRUCTURES
• Lewis symbol: picture of molecule showing
arrangement of its valence electrons around
atoms.
• Lewis structure: picture of molecule
showing bonding electrons as lines and
nonbonding electrons as dots or lines.
• Especially used for main group elements (p
357)
COVALENT BONDS (3)
• Form when electron pairs are shared so that
each atom achieves an octet (duet).
• Coordinate covalent bond forms when one
atom provides both bonding electrons.
• Multiple covalent bond forms when more
than one electron pair is shared between two
atoms (double bond, bond order 2 [CO2]
and triple bond, bond order 3 [N2]).
WRITING LEWIS
STRUCTURES
• Determine total # of valence electrons.
• Write skeletal structure with central atom
[lowest EN]; terminal atoms [H, higher EN]
• Use electron pairs to form bonds.
• Achieve octet rule for terminal atoms
• Add the remaining to the central atom.
• Form multiple bonds if needed.
WRITING LEWIS
STRUCTURES (2)
• Exceptions to octet rule [odd # of valence
electrons (NO), free radicals, incomplete
octets (B), more than 8 electrons (expanded
valence shell SF6)].
• Resonance structures showing different but
equivalent distributions of electrons; note
delocalization (vs localization) of electrons.
• Be guided by experimental observations.
FORMAL CHARGE (FC)
• FC = [VE in free atom] - [VE asigned in
molecule]
• FC is a hypothetical charge for electron loss
(+) or gain (-) due to bond formation.
• [VE]free = # valence e’s for Group A atoms
• [VE] assigned = all lone pair electrons on atom
+ 1/2 shared electrons
FORMAL CHARGE (2)
• Best Lewis structure has minimum FC
(zero).
• Formal Charge method is not perfect and
can lead to incorrect “best” Lewis
structures.
• The best Lewis structure is consistent with
exptal evidence (bond lengths, EN data, etc)
VSEPR MODEL
• VALENCE-SHELL ELECTRON-PAIR
REPULSION (VSEPR) Method helps us
determine molecular geometry.
• Molecular geometry: 3-D shape of the
molecule.
• This method assumes that the final positions
of nuclei are the ones that minimizes
electron repulsions because this is the one
associated with the lowest energy.
VSEPR METHOD (2)
• Determine Lewis structure of molecule.
• Count electron “pairs” around the central
atom where a “pair” may be a single e, lone
pair, single bond, double bond, triple bond.
• Determine geometry of electron pairs.
• Determine molecular group geometry with
A = central atom; X = terminal atom; E =
lone pair of electrons. T8.6, 8.7, 8.8
Table 8.6
Arrangements
of Electron
Pairs Around
an Atom
Yielding
Minimum
Repulsion
MOLECULAR GEOMETRY
#e
pairs
2
e pair geometry
molecular geometry
Linear
Linear
3
Trigonal planar
Trigonal planar, bent
4
Tetrahedral
Tetrah, trig pyram, bent
5
Trig bipyramidal Trig bipyra, seesaw, Tshaped, linear
Octahedral
Octah, sq pyrami, sq
planar
6
MOLECULAR GEOMETRY (2)
• Electron pair geometry differs from
molecular geometry when there are lone
electron pairs (E).
• Electron-electron repulsions decrease as
E-A-E> E-A-X> X-A-X; X = bonded pair
• Resonance structures
• Note bond angles