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)