Chapter 9: Chemical Bonding

Chemical Bonds

Valence electrons are the outermost electrons (all electrons in the outermost quantum shell) – these are the electrons involved in chemical bonding.

And recall:

Ionic compounds (metals with nonmetals) are held together by the electrostatic attraction of opposite charges - this is also known as coulombic attraction. This attraction is an ionic bond.

When nonmetals bond with other nonmetals, they do not bond as ions. Instead, they remain uncharged and share electrons to form covalent bonds.

Lewis Structures

Lewis structures are a symbol notation used to describe bonded/unbonded atoms:

The Octet Rule

Atoms are especially stable when all of the valence orbitals are either completely filled or completely empty (the "noble gas" configuration).

This has been adapted to the octet rule: (most) atoms are stable when there are 8 electrons in their outermost (valence) shell.

Note: For this course: 1st & 2nd row atoms can never have more than 8 valence electrons and/or 4 valence orbitals!

Electronegativity

Ionic bonds: one atom transfers electrons to another.

Bonds with strong ionic character are formed between atoms of very different electonegativities (>2.0 electronegativity difference)

Ionic Compounds

The Ionic-Covalent Continuum

Lattice Energy

Lattice Energy is the energy required to completely separate a mole of a solid ionic compound into individual ions in the gas phase.

Covalent bonds: Two atoms share electrons.

Both atoms can count the shared electrons toward their octet. This type of bond is formed between two atoms of similar electonegativities (<2.0 electronegativity difference)

Covalent Compounds

Polar Covalent Compounds

Some atoms share electrons, but not 100% equally.

This creates polarization of the bond.

Dipole Moment

Dipole moment (μ) is a measure of charge separation. The units of dipole moment are debye (D).

μ= Qr

Where Q is the (absolute value) of the charge on each particle and r is the distance of separation

Dipole moment is often identified using a vector arrow.

Writing Lewis Structures

•All valence electrons must be used (as bonds or lone pairs)

•All electrons are paired (except for odd-electron species)

•All atoms get a full octet (except hydrogen/helium w/ 2 e-, group 13)

•Hydrogen atoms will only ever have 1 bond

•In general, more electronegative atoms (and hydrogen atoms) go on the periphery, and less electronegative atoms are usually more central to the structure.

•The Need-Have-Shared-Nonbonded counting algorithm can be helpful for compounds that obey the octet rule.

Formal Charge

Formal charge is a measure of how electron rich/poor a bonded atom is if all bonds were assumed to be 100% covalent.

As a general rule: formal charges should be minimized when writing Lewis structures (without violating other rules).

Formal charge = (# valence e-) – [nonbonded e- + ½ bonded e-]

Octet Exceptions

•Compounds with an odd number of electrons

•Group 13 compounds

•“Expanded” Octets

Resonance Structures: Multi-center Bonding

In many compounds, single Lewis structures are inadequate to describe the electron distribution and bonding. We use resonance structures to better describe these systems.

Resonance structures are individual Lewis structures that must be “combined” to provide a full picture of the compound.

Some important “rules”:

Resonance structures do not change what atoms are bonded to each other (and never more than 8 electrons!).

Resonance results from the reorganization of electrons in multiple bonds & nonbonded electrons.

An Introduction to Arrow Pushing

Chemists use curved arrows to show electron movement. For our purposes, this is particularly helpful for resonance structures:

Bond Enthalpies (Bond Dissociation Energies)

ΔHrxn = Σ(BDEbonds broken) – Σ(BDEbonds

formed)

A bond enthalpy (also called bond dissociation energy) is the energy required to break a bond (1 mole, in the gas phase) and give half the electrons in the bond to each of the atoms (homolytic cleavage)

Bond enthalpies can be used to determine heat of reaction (without resorting to heats of formation):