Chapter 9 Ionic and Covalent Bonding

Ionic Bonds

Describing Ionic Bonds

Electronic Configurations of Ions(review)

Ionic Radii(review)

Covalent Bonds

Describing Covalent Bonds

Polar covalent bonds; electronegativity

Writing Lewis Dot Formulas (Lewis Structures)

Delocalized Bonding; Resonance

Exceptions to the Octet Rule

Formal Charge and Lewis Structures

Covalent Bond Length, Order and Strength

Objectives

• detemine the number of valence electrons for any atom and write its Lewis symbol.
• describe the concept of ionic bonding.
• recognize and use the octet rule with valence electrons.
• use the Periodic Table to predict the probable formulas of ionic substances formed between common metals and non-metals.
• describe how the radii of ions relate to those of atoms.
• explain the concept of an isoelectronic series and the changes in ionic radius within such a series.
• describe the concept of covalent bonding.
• predict the ionic character of a bond by using the Periodic Table or electronegativity values.
• write Lewis Dot Structures for molecules and ions.
• explain the concept of Resonance and draw resonance structures.
• write Lewis structures for elements with expanded octets or fewer than an octet.
• predict bond strength based upon bond order and bond length.
• Calculate approximate DH values based upon bond energies.

Why don't the Noble gases react readily? Let's look at the electronic configuration.

Something is particularly stable about having eight electrons in the outer (valence) shell.

He 1s²

Ne 1s²2s²2p⁶

Ar 1s²2s²2p⁶3s²3p⁶

Kr 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p⁶

Xe 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p⁶5s²4d¹⁰5p⁶

Ionic Bonding

Octet Rule-

atoms will gain or lose electrons in a chemical reaction in order to achieve an outer shell electron configuration identical to the Noble gases. (ie eight electrons in the valence shell).

Lewis Symbols for Ionic Compounds

Valence electrons = outer shell electrons

Examples:

magnesium and oxygen react to form magnesium oxide, MgO.

calcium reacts with fluorine to form calcium fluoride, CaF₂.

This works well for explaining bonding between metals and non-metals.

Ionic Bonds

The force of electrostatic attraction holding positive and negative charged ions together is called an ionic bond. Most often recognized as a metal ion (or NH₄⁺) and an nonmetal or other anion.

Note: the farther apart two elements are on the Periodic Table, the more ionic is the bond. This is due to the electronegativity of the elements. Electronegativity is the ability of an element to attract electrons in a compound.

Ionic Compounds and Ionic Radii

Review electonic configurations for ions

Ionic Radii

• When atoms gain e^- to form anions they become larger
• When atoms lose e^- to form cations they become smaller
• Isoelectronic series- same number of electrons in the outer shell.

Examples

Lattice Energy(L.E.)- Omit the Born-Haber Cycle

• Energy necessary to break apart ions in an ionic solid.
• No calculations but be able to recognize size relationships of charge, r, and L.E.

• Which has the greatest lattice energy?

Covalent Bonds

Describing Covalent Bonds

Elements, non-metals, which do not form ions in compounds also want to satisfy the octet rule. Electrons are shared more or less equally between the atoms. Bonds formed by sharing electrons are covalent bonds and such molecules are called covalent molecules (non-metal/non-metal compounds)

examples: H₂

CF₄

NF₃

OF₂

F₂

Lewis Dot Structures

1. Count total number of valence electrons for atoms-if you have an ion, add or subtract the appropriate number of electrons

2. Write a skeletal structure putting the least electronegative element in the center, a bond = 2 e^------- subtract the number of electrons used from the total available.

3. Distribute electrons around the outer atoms so each has 8.

4. Put extra electrons on the central atom in pairs

5. If there are not enough electrons for central atom to have 8, may need to use a pair or two from an outer atom to form a double bond. If necessary, move more than one pair.

Note: Only use double bonds if you have to or where formal charges give a better structure.

Examples:

H₂O

N₂

CH₂O

NH₃

CN^-1

Coordinate covalent bonds are formed when both electrons for the bond come from one atom. SO₂

Delocalized Bonding; Resonance

Resonance Structures

Many times it is possible to draw more than one equivalent Lewis Dot structure. Molecules or ions in which more than one equivalent Lewis structure are possible are said to exhibit resonance.

Example- NO₃^- ; CO₃^2-

Exceptions to the Octet Rule

Odd Electron Molecules

nitrogen dioxide

Less than an Octet

boron triiodide

More than an Octet (expanded octet)

phosphorous pentachloride

XeF₄

Polar Covalent Bonds-

Use the Periodic Table to generalize about electronegativity differences or use calculated values to determine type of bonding: ionic, polar covalent, and covalent.

For each of the following pairs of elements, tell whether or not the bond will be ionic or covalent. If covalent, rank their polarity.

Formal Charge and Lewis Structures

The concept of formal charge can be used to decide the best form between alternative Lewis Structures.

The formal charge on an atom is equal to the original number of valence electrons in the atom, minus the number of electrons assigned to the atom in the Lewis Structure.

a. The structure with lowest formal charge on each atom is usually "best". Electrons shared are counted as 1/2, a lone pair counts 2.

b. If 2 structures have the same magnitude of charge, the more electronegative element

Examples:

1. carbon monoxide

2. sulfate ion

3. CNO^-1 or OCN^-1 or CON^-1 ? which structure is best?

Covalent Bond Length, Order and Strength

Bond length depends upon resonance structures and bond energy.

a. The longer the bond, the weaker the bond.

b. The greater the order, the stronger the total bond.

Write Lewis Structures for CH₂O and CH₃OH. Which has the shorter carbon oxygen bond? Which C, O bond is stronger?

Bond Dissociation Energy (Bond Energy)

The enthalpy change, DH, required to break a bond between two atoms in the gaseous state can be related to the enthalpy change of reaction.

WHY?

DH_rxn = Sum reactant bond energies – Sum product bond energies.

Bond breaking is endothermic, bond forming is exothermic.

Example

Use bond energy table in the text to estimate H for the combustion of methane

CH₄ (g) + 2 O₂ (g) ----> CO₂ (g) + 2 H₂O (g)

4 (C-H) + 2 (O=O) 2 (C=O) + 4 (O-H)

[return to top of page]

Updated Feb 13, 2004. Questions or comments on this Web site should go to Robin Terjeson.