Chapter 9
Ionic and Covalent Bonding
the chemical bond:
◆ the force that holds atoms or ions together as an aggregate unit
bond energy (or bond dissociation enthalpy, ∆HBDE):
◆ energy required to break a chemical bond Cl–Cl (g) ! 2 Cl (g); ∆H = bond energy
◆ bond energy is always endothermic
generally 3 types of bonds - ionic, covalent, metallic
Lewis Symbols
developed by G.N. Lewis to represent an element and its number of valence electrons
◆ each side of element’s symbol may have 0, 1, or 2 dots
◆ each dot represents a valence electron
◆ for main group elements: # valence e–’s = group #
atoms tend to lose, gain, or share e-’s in such a way that they attain a noble gas configuration (ns2 np6)
◆ ionic compounds: typically metals with low ionization energy lose e-‘s
nonmetals with favorable electron affinity gain e-’s
The Octet Rule
◆ molecular compounds:typically nonmetals will share e-’s to form covalent bonds
The Octet Rule Energy Considerations of Ionic Compounds
Lattice Energy, U: the energy required to separate an ionic solid into its gas phase ions
MX (s) ! M+ (g) + X– (g); endothermic
Lattice Energy and Coulomb’s Law
Coulomb’s Law:◆ describes the energy of electrostatic
interaction between 2 charged particles separated by some distance
Ecoul ∝ ––––––
Q1 and Q2 = charges on ionsr = distance between ions
◆ this relationship is also true for lattice energy
U ∝ charges on ions
U ∝ 1/distance separating ions
Q1●Q2
r
Electron Configurations of Ions
cations
◆ formed by the loss of electrons
◆ e–’s are lost from the highest energy, populated atomic orbital
anions
◆ formed by the gain of electrons
◆ e–’s are added to the lowest energy available atomic orbital
How does the e– configuration of an atom change as an ion is formed?
Electron Configurations of Ions
some examples of cations that do not have a noble gas configuration:
◆ heavy main group metals: ex. Sn2+
◆ transition metals:ex. Fe3+
Ionic Radii
◆ size of cation or anion relative to the parent atom
cations are smaller than their parent, neutral atom
rNa+ < rNa
anions are larger than their parent, neutral atom
rCl– > rCl
Periodic Trend in Ionic Radii
◆ Similar to atomic radii, ionic radii increase from the top to the bottom of the periodic table.
◆ To understand the trend moving left to right across the periodic table, consider isoelectric species.
isoelectric – same number of electrons
O2– F– Na+ Mg2+ Al3+
# e–’s 10 10 10 10 10
# p’s 8 9 11 12 13
ionic radius
140 pm 136 pm 95 pm 65 pm 50 pm
In molecular compounds, atoms share electrons to achieve a noble gas configuration
◆ covalent bonds result from shared pairs of electrons between atoms
◆ electrons are shared in the region between atoms where atomic orbitals overlap
◆ orbital overlap results in a concentration of electron density between 2 nuclei
Molecular Compounds: Covalent Bonds
Nonpolar vs Polar Covalent Bonds:
How is the electron density distributed between atoms in a bond?
Electronegativity, !: the ability of an atom in a polyatomic species to attract electrons to itself
◆ The greater the difference in electronegativity (∆ !) between 2 atoms in a bond, the more polar the bond.
Some Introductory Thoughts About Structure and Bonding in Molecular Compounds
◆ How many covalent bonds is a central atom in a molecule likely to form?
consider the element’s group # and # valence e–’s
use the octet rule as a guideline
◆ bonding vs. nonbonding pairs of electrons
◆ single vs. multiple bonds
1 pair of e–’s shared between atoms (i.e. 2 shared e–’s) ! single bond
2 pairs of e–’s shared between atoms (i.e. 4 shared e–’s) ! double bond
3 pairs of e–’s shared between atoms (i.e. 6 shared e–’s) ! triple bond
Some Introductory Thoughts About Structure and Bonding in Molecular Compounds Drawing Lewis Structures
a 1st step to understanding:
◆ molecular structure◆ atom connectivity◆ arrangement & distribution of valence e–’s◆ numbers and types of bonds◆ numbers of nonbonding pairs of e–’s◆ 3-D shape◆ bond angles◆ molecular polarity
Drawing Lewis Structures
start by thinking about structures of molecules using the localized electron model:
◆ localized electron model assumes molecules are collections of atoms bonded together by covalent bonds
◆ pairs of electrons are either localized on atoms (lone or nonbonding e– pairs), or localized in the space between 2 atoms (bonding e– pairs)
Drawing Lewis Structures
We will use the localized electron model to:
◆ describe the atom arrangement and distribution of valence e–’s in a molecule
Lewis Dot Structures (now)
◆ predict molecular geometry, bond angles, and polarity
VSEPR Theory (Ch. 10)
◆ describe the types of atomic orbitals used by atoms in bonding or to house lone pairs
Valence Bond Theory (Ch. 10)
Drawing Lewis Structures
1. Determine the total number of valence electrons in the molecule:
total # valence e–’s = ∑valence e–’s of atoms in molecule
◆ main group elements: # valence e–’s = group #◆ add 1e– for each unit negative charge on anion◆ subtract 1e– for each positive charge on cation
! this is the total number of electrons you will need to have in your final structure
2. Write symbols for atoms in order of connectivity; connect appropriate atoms with single bonds
Drawing Lewis Structures
3. Complete the octets of atoms bonded to the central atom(s).
H, He, Li, and Be will only have a duplet of e–’s
4. Place any left over electrons on the central atom, even if it results in greater than an octet.
5. If there are not enough electrons to give the central atom a full octet, try multiple bonds.
examples:PCl3 CH2Cl2
HCN ClO2–
NH4+ N2H4
COBr2
Drawing Lewis Structures What if you can draw more than one Lewis Structure that obeys the octet rule?
Which one is the “right” one?
2 concepts to help interpret Lewis structures:
◆ Formal Chargesame atom arrangement; different e– arrangement
◆ Resonance Structuressame atom arrangement; same net e– arrangement; same formal charge distribution
◆ Considering formal charges
to determine formal charges of elements in structure:
1. all unshared e–’s (lone pairs) are assigned to the atom on which they are found
2. bonding e–’s are assumed to be shared evenly between the atoms participating in the bond;
◆ homolytic clevage of the bond
◆ ! of the bonding e–’s are assigned to each atom in the bond
3. formal charge = # valence e–’s of isolated atom – # e–‘s assigned by Lewis structure
In general, the more stable Lewis structure is considered to be the one in which:
◆ atoms bear formal charges closest to 0
◆ any negative formal charge resides on more electronegative element
ex: Consider 2 possible Lewis structures of CO2.
ex: Consider 3 possible Lewis structure of NCS–.
◆ Considering resonance structures
sometimes one Lewis structure does not adequately describe e– arrangement
supported by experimental evidence
example: Consider 2 possible Lewis structures for ozone, O3:
experimental data: ◆ O3 is a bent molecule◆ both O–O bond lengths equivalent
example: carbonate ion, CO32–
resonance hybrid structure
The 2 resonance structures A and B are equivalent contributors to the overall resonance hybrid structure.
Exceptions to the Octet Rule
1. odd number of electrons
ex. NO
2. central atom has less than an octet
ex. BF3
Exceptions to the Octet Rule
3. central atom has more than an octet of e–’s
◆ expanded valence
◆ possible for larger central atoms in the 3rd period and below
ex. PCl5
XeF4
A Way to Think About Expanded Valence
for PCl5:◆ central atom: P◆ valence e– configuration: 3s2 3p3
◆ empty 3d orbitals sit at slightly higher energy
↑ ↑ ↑
↑ ↑ ↑ ↑ ↑
◆ a 3s e– absorbs E and is promoted to a higher E, empty 3d orbital
3s 3p 3d
3s 3p 3d
↑↓
A Way to Think About Expanded Valence
for XeF4:◆ central atom: Xe◆ valence e– configuration: 5s2 5p6
◆ empty 5d orbitals sit at slightly higher energy
↑↓ ↑↓ ↑↓ ↑↓
↑↓ ↑↓ ↑ ↑ ↑ ↑
◆ 2 5p e–’s absorb E and are promoted to higher E, empty 5d orbitals
5s
5s
5p
5p
5d
5d
Using Covalent Radii to Approximate Bond Length
A A
A–A bond length = rA + rA
A B
A–B bond length = rA + rB
◆ knowing periodic trend in atomic radii can help you make predictions about relative lengths of bonds
example: Would you predict that a N–Cl or a P–Br bond would be longer?
Bond Multiplicity (or Bond Order) and Relationship to Bond Length and Bond Energy
bond type
bond order
# e– pairs shared
# e–’s shared
single 1 1 2
double 2 2 4
triple 3 3 6
bond length decreases
bond energy increases
A Word About Tabulated Bond Energy Data
recall: ◆ bond energy (or bond enthalpy) is the energy
required to break a bond
◆ endothermic
◆ units kJ/mol
example:
What is the C–H bond energy?
if: CH4 (g) ! C (g) + 4 H (g); ∆H = 1660 kJ
then: we can approximate the average C–H bond energy as 1660 ÷ 4 = 415 kJ
Using Bond Energies to Approximate ∆Hrxn:
∆Hrxn = ∑ E of bonds broken " ∑ E of bonds formed
Using Bond Energies to Approximate ∆Hrxn:
∆Hrxn = ∑ E of bonds broken " ∑ E of bonds formed
example:
Calculate ∆H for the following reaction using bond energies:
C2H4 (g) + H2O (l) ! C2H5OH (l)
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