Chapter 4: Forces Between Particles – Ionic and Covalent Bonding, Lewis Structures, and Molecular Shapes

Forces Between Particles

Introduction

This chapter explores the fundamental forces that hold atoms and molecules together, focusing on ionic and covalent bonding, the use of Lewis structures, the octet rule, and the prediction of molecular shapes using VSEPR theory. Understanding these concepts is essential for explaining the properties and behaviors of chemical substances.

Noble Gas Configurations

Definition and Importance

• Noble gas configuration: The stable electron arrangement found in noble gases, characterized by a full valence shell (2 electrons for helium, 8 for others).

• Other elements tend to gain, lose, or share electrons to achieve this stable configuration.

Element Symbol	Conventional Form	Abbreviated Form
Li	1s22s1	[He]2s1
Na	1s22s22p63s1	[Ne]3s1
K	1s22s22p63s23p64s1	[Ar]4s1
Rb	1s22s22p63s23p64s23d104p65s1	[Kr]5s1

Lewis Structures

Representation of Atoms and Ions

• Lewis structure: A diagram showing the element symbol surrounded by dots representing valence electrons.

• Valence electrons are the outermost electrons and are key to chemical bonding.

• Each side of the symbol can hold up to two dots (electrons).

Determining Valence-Shell Electrons

• Write the electronic configuration and identify electrons with the highest principal quantum number (n).

• For representative elements, the number of valence electrons equals the group number (Roman numeral) in the periodic table.

Example: Drawing Lewis Structures

• Beryllium (Be, group IIA/2): Two valence electrons → Lewis structure: Be with two dots.

• Cesium (Cs, group IA/1): One valence electron → Lewis structure: Cs with one dot.

The Octet Rule

Definition and Application

• Atoms tend to gain, lose, or share electrons to achieve eight electrons in their valence shell (octet), similar to noble gases.

• Simple ion: An atom that has gained or lost electrons to acquire a net charge.

Example equations:

• $\mathrm{Na} \rightarrow \mathrm{Na}^+ + e^-$

• $\mathrm{Cl} + e^- \rightarrow \mathrm{Cl}^-$

Determining Ionic Charges for Representative Elements

• Metals (left side of periodic table) form positive ions (cations) with a charge equal to their group number.

• Nonmetals (right side) form negative ions (anions) with a charge equal to 8 minus their group number.

• Example: Sr (group IIA/2) forms $\mathrm{Sr}^{2+}$; P (group VA/5) forms $\mathrm{P}^{3-}$.

Ionic Bond Formation

Process and Result

• Ions with opposite charges attract, forming an ionic bond.

• Metals lose electrons (become cations), nonmetals gain electrons (become anions).

• Both achieve noble gas configurations and are held together by electrostatic attraction.

Isoelectronic Species

Definition

• Isoelectronic: Atoms or ions with the same number of electrons and identical electronic configurations.

Particle	Symbol	Protons	Electrons	Net Charge
Magnesium atom	Mg	12	12	0
Magnesium ion	Mg2+	12	10	+2
Neon atom	Ne	10	10	0

Ionic Compounds

Definition and Types

• Ionic compounds: Substances formed from the electrostatic attraction between cations and anions.

• Binary ionic compounds: Ionic compounds formed from only two elements (a metal and a nonmetal).

Formulas for Ionic Compounds

• Formulas represent the simplest ratio of ions that results in electrical neutrality.

• Example: Na+ and F- combine to form NaF.

• Example: Na+ and S2- combine to form Na2S.

Naming Binary Ionic Compounds

• Name = Metal name + nonmetal stem + -ide

• For metals with multiple possible charges (transition metals), indicate the charge with a Roman numeral in parentheses.

Element	Stem	Formula of Ion
Chlorine	chlor-	Cl-
Oxygen	ox-	O2-
Sulfur	sulf-	S2-
Nitrogen	nitr-	N3-

Examples

• K2O: Potassium oxide

• Mg3N2: Magnesium nitride

• BeS: Beryllium sulfide

• AlBr3: Aluminum bromide

Ionic Compound Structure

Crystal Lattice

• Ionic compounds form a rigid, three-dimensional crystal lattice rather than discrete molecules.

• Lattice site: The position occupied by an ion in the crystal lattice.

Ionic Compound Formulas and Weights

• Formula weight: The sum of the atomic weights of the ions in the empirical formula of an ionic compound.

• 1 mole of an ionic compound contains Avogadro's number ($6.022 \times 10^{23}$) of formula units.

Example calculation:

• NaCl: $\text{Formula weight} = 22.99\,u + 35.45\,u = 58.44\,u$

Covalent Bonds

Definition and Formation

• Covalent bond: A bond formed when two atoms share one or more pairs of valence electrons to satisfy the octet rule.

• Shared electrons are counted in the octet of both atoms involved.

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

Drawing Lewis Structures for Molecules

1. Determine the number of atoms of each type in the molecule.

2. Arrange atoms and connect them with single bonds.

3. Count total valence electrons.

4. Place electron pairs between bonded atoms, then distribute remaining electrons to complete octets.

5. If octets are incomplete, form double or triple bonds as needed.

Example: SO3 Lewis Structure

• Sulfur (S) and three oxygens (O) are arranged with S in the center.

• Total valence electrons: 24 (6 from S, 6 from each O).

• Single bonds are formed first, then remaining electrons complete octets.

• Double bonds are used if necessary to satisfy the octet rule for the central atom.

Shapes of Molecules: VSEPR Theory

Valence-Shell Electron-Pair Repulsion (VSEPR) Theory

• Predicts the three-dimensional shapes of molecules based on the repulsion between electron pairs around a central atom.

• Bonding and nonbonding pairs are counted equally; double/triple bonds count as one pair.

Common Electron Pair Arrangements

• 2 pairs: linear

• 3 pairs: trigonal planar

• 4 pairs: tetrahedral

Examples

• CO2: Linear (O=C=O)

• NH3: Tetrahedral electron pair geometry, but molecular shape is trigonal pyramidal due to one lone pair on N.

Polarity of Molecules and Electronegativity

Bond Polarity

• Polar covalent bond: Electrons are shared unequally due to differences in electronegativity.

• Nonpolar covalent bond: Electrons are shared equally.

• Electronegativity: The ability of an atom to attract shared electrons in a bond.

Molecular Polarity

• Polar molecule: Contains polar bonds arranged asymmetrically, resulting in a net dipole moment.

• Nonpolar molecule: Either contains only nonpolar bonds or polar bonds arranged symmetrically.

Polyatomic Ions

Definition and Examples

• Polyatomic ion: A covalently bonded group of atoms with a net charge.

• Common examples: $\mathrm{NO}_3^-$ (nitrate), $\mathrm{SO}_4^{2-}$ (sulfate), $\mathrm{NH}_4^+$ (ammonium).

Naming Compounds with Polyatomic Ions

• Rules are similar to those for binary ionic compounds.

• Examples: Na3PO4 (sodium phosphate), CaCO3 (calcium carbonate).

Interparticle Forces

Types of Forces

• Ionic and covalent bonds: Hold atoms together within compounds.

• Intermolecular forces: Hold molecules together in the solid and liquid states.

• Hydrogen bonding: Strong dipole-dipole attraction between molecules with H bonded to N, O, or F.

• Dispersion forces: Weak attractions present in all molecules, especially significant in nonpolar substances.

• Network solids: Solids where atoms are covalently bonded in a continuous network (e.g., diamond).

• Metallic bonding: Attraction between metal cations and delocalized electrons.

Effect on Properties

• The strength and type of interparticle forces determine melting and boiling points, solubility, and other physical properties.