Chapter 4: Compounds, Chemical Formulas, and Composition

Types of Compounds and Chemical Formulas

Chemical Bonds and Compound Types

Chemical bonds form because they lower the potential energy of the charged particles that compose atoms. The two main types of chemical bonds are ionic bonds and covalent bonds.

• Ionic Bond: Forms between a metal and a nonmetal. Metals tend to lose electrons, while nonmetals tend to gain them. The transfer of electrons from metal to nonmetal creates ions that are held together by electrostatic forces.

• Covalent Bond: Forms between two or more nonmetals. Nonmetals have high ionization energies, so electrons are shared rather than transferred. The shared electrons interact with the nuclei of both atoms, lowering their potential energy according to Coulomb's law. Covalently bonded atoms form molecules, and the resulting compounds are called molecular compounds.

Examples:

• Ionic: NaCl (table salt), CaF2 (calcium fluoride)

• Covalent: H2O (water), CO2 (carbon dioxide)

Visual Representation of Compounds

Compounds can be represented in several ways:

• Empirical Formula: The simplest whole-number ratio of elements in a compound.

• Molecular Formula: The actual number of atoms of each element in a molecule.

• Structural Formula: Shows how atoms are bonded together.

• Ball-and-Stick Model: 3D representation showing atoms as balls and bonds as sticks.

• Space-Filling Model: Shows the relative sizes of atoms and their spatial arrangement.

Name	Empirical Formula	Molecular Formula	Structural Formula	Ball-and-Stick Model	Space-Filling Model
Benzene	CH	C6H6	Hexagonal ring	Model	Model
Acetylene	CH	C2H2	H–C≡C–H	Model	Model
Glucose	CH2O	C6H12O6	Ring/chain	Model	Model
Ammonia	NH3	NH3	Trigonal pyramidal	Model	Model

Additional info: Models are visual aids for understanding molecular geometry and bonding.

Empirical and Molecular Formulas

Determining Empirical Formulas

The empirical formula represents the simplest whole-number ratio of elements in a compound. To determine it from a molecular formula, divide the subscripts by their greatest common factor.

• Example: Molecular formula of C6H12O6 (glucose) gives the empirical formula CH2O.

Counting Atoms in Chemical Formulas

To determine the number of each atom in a formula, multiply the subscript of each element by the number of formula units if necessary.

• Example: In Ca(NO3)2, there is 1 Ca, 2 N, and 6 O atoms.

Electron Configuration and Lewis Symbols

Electron Configuration of Main Group Elements

Electron configuration describes the arrangement of electrons in an atom. For example, sodium (Na):

• Na: $1s^2 2s^2 2p^6 3s^1$

• Na+: $1s^2 2s^2 2p^6$ (loses one electron from the 3s orbital)

Lewis Symbols

Lewis symbols represent valence electrons as dots around the element symbol. For ions, brackets and charges are used.

• Example: Cl: Cl with 7 dots; Cl-: [Cl]- with 8 dots

• Mg: Mg with 2 dots; Mg2+: [Mg]2+ with no dots

Ionic Bonding and Lattice Energy

Ionic Bond Lewis Structures

Ionic compounds are represented by showing the transfer of electrons from metal to nonmetal, resulting in oppositely charged ions.

• Example: Na• + •Cl• → Na+ + [Cl]-

Lattice Energy

Lattice energy (U) is the energy released when one mole of an ionic compound forms from its free ions in the gas phase.

• General equation: $M^+(g) + X^-(g) \rightarrow MX(s)$

• Lattice energy formula: $U = k \frac{Q_1 Q_2}{d}$

• Higher charges and smaller ion sizes result in higher (more exothermic) lattice energies.

Compound	Lattice Energy (kJ/mol)
NaF	-915
NaCl	-776
NaBr	-742
NaI	-699
MgCl2	-2489
MgO	-3933

Additional info: Lattice energy increases with higher ionic charges and decreases with larger ionic radii.

Formulas and Names for Ionic Compounds

Naming Ionic Compounds

• Monatomic cations: Name of element (e.g., Na+ = sodium)

• Monatomic anions: Root of element + "-ide" (e.g., Cl- = chloride)

• Transition metals: Indicate charge with Roman numerals (e.g., Fe3+ = iron(III))

• Polyatomic ions: Use common names (e.g., SO42- = sulfate, NH4+ = ammonium)

Simple Lewis Structures, Formulas, and Names for Molecular Compounds

Lewis Structures for Diatomic Molecules

Many elements form diatomic molecules (e.g., H2, O2, N2, F2, Cl2, Br2, I2) to achieve stable electron configurations.

• Example: H2 forms by sharing two electrons between two H atoms.

• O2 forms a double bond; N2 forms a triple bond.

Naming Binary Molecular Compounds

• Use prefixes to indicate the number of each atom (mono-, di-, tri-, tetra-, etc.).

• First element keeps its name; second element ends with "-ide".

• Example: CO2 = carbon dioxide, N2O3 = dinitrogen trioxide

Number of Atoms	Prefix	Example
1	mono-	NO: nitrogen monoxide
2	di-	NO2: nitrogen dioxide
3	tri-	N2O3: dinitrogen trioxide
4	tetra-	N2O4: dinitrogen tetraoxide
5	penta-	P2O5: diphosphorus pentoxide

Formula Mass and the Mole Concept

Formula Mass

The formula mass of a compound is the sum of the atomic masses of all atoms in its formula.

• Example: H2O: (2 × 1.01) + 16.00 = 18.02 g/mol

The Mole Concept

A mole is 6.022 × 1023 particles (Avogadro's number). The molar mass (in grams) of a substance contains one mole of its particles.

• 1 mole H2O = 18.02 g = 6.022 × 1023 molecules

Composition of Compounds

Mass Percentage

The mass percent of an element in a compound is calculated as:

$\text{Mass \% of A} = \frac{\text{mass of A in compound}}{\text{total mass of compound}} \times 100$

• Example: In NaCl, mass % Na = (22.99/58.44) × 100 ≈ 39.34%

Calculating Empirical Formulas from Experimental Data

Steps to Determine Empirical Formula

1. Obtain the mass (or percent composition) of each element in the compound.

2. Convert masses to moles using atomic weights.

3. Divide by the smallest number of moles to get the simplest ratio.

4. If necessary, multiply to get whole numbers.

Example: A sample contains 1.651 g Ag and 0.1224 g O. Convert to moles and find the simplest ratio to determine the empirical formula.

Summary Table: Key Concepts

Concept	Definition/Formula
Ionic Bond	Transfer of electrons from metal to nonmetal
Covalent Bond	Sharing of electrons between nonmetals
Empirical Formula	Simplest whole-number ratio of elements
Molecular Formula	Actual number of atoms in a molecule
Lattice Energy	$U = k \frac{Q_1 Q_2}{d}$
Formula Mass	Sum of atomic masses in a formula
Mole	6.022 × 1023 particles
Mass Percent	$\frac{\text{mass of element}}{\text{total mass}} \times 100$

Additional info: These notes cover the foundational concepts of chemical bonding, formulas, and composition, essential for understanding chemical compounds in General Chemistry.