CHM103_Chapter_4.2

Chapter 4: Forces Between Particles

Overview

• This chapter discusses the interaction forces between particles, focusing on ionic and covalent bonding.

Ionic Compound Structure

• Stable ionic compounds form crystals, not molecules.

• Crystals consist of many oppositely charged ions arranged in a rigid three-dimensional structure called a crystal lattice.

Ionic Compound Formulas and Weights

• Formulas for ionic compounds represent the simplest ratio of ions, not specific atom counts found in a molecule.

• Formula weight is the total sum of the atomic weights of the elements in the formula, analogous to molecular weight.

• One mole of an ionic compound contains Avogadro’s number (6.022 x 10²³) of the simplest combining ratio.

Example Calculation of Formula Weight

For gem emeralds (Al₂Be₃Si₆O₁₈)

• Aluminum (Al): 2 * (26.98) = 53.96 g/mole

• Beryllium (Be): 3 * (9.012) = 27.036 g/mole

• Silicon (Si): 6 * (28.09) = 168.54 g/mole

• Oxygen (O): 18 * (16.00) = 288.00 g/mole

• Total: 537.54 g/mole

Writing Ionic Compound Formulas

Example: Tin(IV) Oxide and Lead(II) Sulfide

• Correct formulas: SnO₂ (tin(IV) oxide) and PbS (lead(II) sulfide).

Covalent Bonding

• Covalent bonds are formed when atoms share valence electrons, satisfying the octet rule.

• The shared electrons count toward the octets of the involved atoms.

• Represented by shared pairs or a line between bonded atoms.

Electron Sharing in Covalent Bonds

Types

• Sharing can occur between identical atoms (e.g., Cl₂, O₂, N₂) or different atoms (e.g., H₂O, CH₄).

Examples of Covalent Bonding

Molecular Examples

• Nitrogen gas (N₂): N≡N (triple bond)

• Carbon dioxide (CO₂): O=C=O (double bonds)

• Water (H₂O): H-O-H (single bonds)

• Methane (CH₄): H-C-H (four single bonds with hydrogen)

Drawing Lewis Structures for Covalent Molecules

Steps

1. Use the molecular formula to determine the number of each type of atom.

2. Draw an initial molecular structure based on the atom connections.

3. Determine the total valence-shell electrons in the molecule.

4. Place pairs of electrons between each bonded atom and complete the octets accordingly.

5. Move lone pairs to form double or triple bonds if octets cannot be satisfied.

Example: Lewis Structure for SO₃

1. Identify: One S and three O atoms

2. Draw: O-S-O-O

3. Total valence electrons: S (6) + 3O (18) = 24 total

4. Create bonds: S-O bonds take 6 electrons, octets filled for O's

5. Move lone pair from an O to S to form a double bond

Key Notes on Lewis Structures

1. Hydrogen forms only 1 bond for 2 electrons.

2. Boron and aluminum form only 3 single bonds with no lone pairs.

3. Carbon can form up to 4 single bonds with no lone pairs.

4. Nitrogen can form up to 3 single bonds with 1 lone pair.

5. Oxygen can form up to 2 single bonds with 2 lone pairs.

6. Fluorine forms 1 single bond with 3 lone pairs.

7. Noble gases typically form no bonds as they have complete valence shells.

8. Elements in the 3rd period or below can exceed octets through super-octet structures.

9. Resonance structures can occur with excess lone pairs.

Practice Exercises for Lewis Dot Structures

• NH₃

• CH₄

• H₂S

• HCN

• C₂H₄

• SiF₄

• SO₄²-

• NO₃-

• PCl₃

Conclusion

• Understanding ionic and covalent bonding principles is essential for studying molecular interactions and structures.