Types of Chemical Bonds

Bond Energy

• Bond Energy – the energy required to break a bond

• Atoms will bond in order to achieve the lowest energy configuration

Two Types of Bonds

• #1: Ionic Bonds – When an atom with a high electron affinity reacts with an atom that loses an electron easily (metal with nonmetal)

• #2: Covalent Bonds – When atoms share electrons (nonmetal with nonmetal)

Ionic Bond creates Ions

• The energy between a pair of ions is calculated using Coulombs Law

• Where• r is the distance between the ion centers • Q1 and Q2 are the ion charges

Size of the Ions

• Remember Isoelectronic Ions?• Example 8.3: Arrange ions Se2-, Br-, Sr2+, Rb+ in

order of increasing size

Bond Length

• A molecule will position itself so the attractive forces are maximized and repulsive forces are minimized (energy is minimized) – this distance is called the bond length (from center of 2 atom to center of the other)

Covalent Bonding

• Polar Covalent Bonds– Electrons are not shared equally due to

electronegativity – Dipole moment is represented by the arrow

pointing toward the negative side – Types of bonds with no dipole moments• Linear molecules with two identical bonds • Planar molecules with three identical bonds• Tetrahedral molecules with 4 identical bonds

Order the following bonds in terms of bond polarity

• H-H, O-H, Cl-H, S-H, F-H

Percent Ionic Character

• How can we tell the difference between a polar covalent bond and an ionic bond?

–Percent ionic character of a bond = (measured dipole moment of X-Y) x 100%(calculated dipole moment of X+Y-)

Percent Ionic Character

• Figure 8.13 – compounds with ionic character greater than 50% are normally considered to be ionic OR any compound that conducts an electric current when melted

% Ionic Character

Lattice Energy

• Lattice Energy – the change in energy that takes place when gas ions are packed together to form a solid (energy released when an ionic solid forms) – Lattice Energy = k(Q1Q2/r) – Where: »k is a proportionality constant

Lattice Energy

Bond Energies

• Scientists can calculate how much energy is required to break down a molecule

• Depending on the bonds of the molecule, they modeled that each bond has a specific amount of energy – Bond Energies (this is purely a scientific invention)

Type of Bonds

–Single Bond – sharing 1 pair of electrons (2 electrons)–Double Bond – sharing 2 pairs of electrons (4

electrons)–Triple Bond – sharing 3 pairs of electrons (6

electrons)

Bond Energy: table 8.4 pg 351

• Bond energy values can be used to calculate approximate energies for reactions – When bonds are broken – energy is added

(endothermic) – When bonds are formed – energy is released

(exothermic) – ΔH = ΣD(bonds broken) – ΣD(bonds formed) – Where: »Σ is the sum of terms »D is the bond energy per mole of bonds

Bond Energies

Use bond energies to determine the following:

• Example: H2 + F2 2 HF

Use bond energies to determine the following:

• Example: Calculate ΔH of methane with chlorine and fluorine to give Freon-12 (CF2Cl2), hydrofluoric acid, and hydrochloric acid.

Localized Electron Bonding Model – a molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms. • Lone pairs – localized on an atom • Bonding pairs – found between atoms

Application Rules

–Description of the valence electron arrangement in the molecule using Lewis Structures –Prediction of the geometry of the molecule using the

valence shell electron pair repulsion (VSEPR) model –Description of the type of atomic orbitals used by the

atoms to share electrons or hold lone pair (Ch. 9)

Lewis Structure

Steps for writing Lewis structures: –Sum the valence electrons from all atoms – this is the

TOTAL number of electrons present

Lewis Structure

Steps for writing Lewis structures: –Sum the valence electrons from all atoms – this is the

TOTAL number of electrons present –Use a pair of electrons to form a bond between each

pair of bound atoms –Arrange the remaining electrons to satisfy the octet

rule for all elements –Compare the TOTAL number electrons to the number

you drew – they must match – if not, add double bonds!

• CH4

• CF4

• NH3

• BH3 (watch out! This one has an exception!)

Exceptions to the Octet rule

–C, N, O, F – always obey the octet rule! –B and Be often have fewer than 8 ve – they are very

reactive –2nd row elements cannot exceed the octet rule

because their orbitals don’t allow it. –3rd row elements can exceed the octet rule by using

their empty valence d orbitals–Satisfy 2nd row elements first – then any left over

electrons should be added to 3rd row elements that have an available d-orbital.

Exceptions to the Octet rule

–C, N, O, F – always obey the octet rule! –B and Be often have fewer than 8 ve – they are very

reactive –2nd row elements cannot exceed the octet rule

because their orbitals don’t allow it. –3rd row elements can exceed the octet rule by using

their empty valence d orbitals–Satisfy 2nd row elements first – then any left over

electrons should be added to 3rd row elements that have an available d-orbital.

Exceptions to the Octet rule

–C, N, O, F – always obey the octet rule! –B and Be often have fewer than 8 ve – they are very

reactive –2nd row elements cannot exceed the octet rule

because their orbitals don’t allow it. –3rd row elements can exceed the octet rule by using

their empty valence d orbitals–Satisfy 2nd row elements first – then any left over

electrons should be added to 3rd row elements that have an available d-orbital.

Exceptions to the Octet rule

–C, N, O, F – always obey the octet rule! –B and Be often have fewer than 8 ve – they are very

reactive –2nd row elements cannot exceed the octet rule

because their orbitals don’t allow it. –3rd row elements can exceed the octet rule by using

their empty valence d orbitals–Satisfy 2nd row elements first – then any left over

electrons should be added to 3rd row elements that have an available d-orbital.

Exceptions to the Octet rule

–C, N, O, F – always obey the octet rule! –B and Be often have fewer than 8 ve – they are very

reactive –2nd row elements cannot exceed the octet rule

because their orbitals don’t allow it. –3rd row elements can exceed the octet rule by using

their empty valence d orbitals–Satisfy 2nd row elements first – then any left over

electrons should be added to 3rd row elements that have an available d-orbital.

Exceptions to the Octet rule

–C, N, O, F – always obey the octet rule! –B and Be often have fewer than 8 ve – they are very

reactive –2nd row elements cannot exceed the octet rule

because their orbitals don’t allow it. –3rd row elements can exceed the octet rule by using

their empty valence d orbitals–Satisfy 2nd row elements first – then any left over

electrons should be added to 3rd row elements that have an available d-orbital.

Draw Lewis Dot Structure for:

• SF6

• ClF3

• XeO3

• RnCl2

• BeCl2

• ICl4-

What do I do with a charge?

• First of all…what does the charge tell us?• So I just add or subtract from the total number

of electrons!– Example: ICl4

-

Formal Charge con’t

• Atoms with a formal charge will..– Try to achieve a charge close to zero–Any formal charges are expectred to reside

on the most electronegative atoms

Resonance

• Is invoked when more than one valid Lewis structure can be written for a particular molecule.

• The resulting structure is an average of these resonance structures. – Ex. Nitrite Ion, Sulfate Ion

Example 8.10:

• Give Possible Lewis Structures for XeO3, an explosive compound of xenon. Which Lewis structure or structures are most appropriate according to the formal charges?