CH3-1: Molecules, Compounds, and Nomenclature: Chemical Formulas, Ionic Compounds, and the Mole Concept

Chapter 3: Molecules, Compounds, and Nomenclature

Chemical Formulas and Molecular Models

Chemical formulas and molecular models are essential tools in chemistry for representing the composition and structure of molecules. They provide information about the types and numbers of atoms present in a compound.

• Molecule: A definite group of atoms that are chemically bonded together.

• Molecular formula: An abbreviated description of the composition of a molecule, indicating the actual numbers and types of atoms present. Examples: $H_2$, $H_2O$, $CO_2$

• Empirical formula: Gives the relative numbers and types of atoms in a molecule, representing the simplest whole-number ratio. Examples: $H_2O$ (molecular and empirical), $C_2H_4$ (molecular), $CH_2$ (empirical for $C_2H_4$)

• Structural formula: Shows how atoms are bonded together in a molecule.

• Ball-and-stick and space-filling models: Three-dimensional representations of molecules, illustrating the spatial arrangement of atoms.

Key Questions:

• Where do we encounter empirical formulas?

• Which is better, empirical or molecular formulas?

Ionic Compounds

Ionic compounds are formed from the electrostatic attraction between cations (positively charged ions) and anions (negatively charged ions), resulting in an overall neutral species. Typically, these are combinations of metals and nonmetals.

• Cation: A positively charged ion formed when a neutral atom loses one or more electrons.

• Anion: A negatively charged ion formed when a neutral atom gains one or more electrons.

• Example: $NH_4^+$ (ammonium cation), $Cl^-$ (chloride anion), forming $NH_4Cl$.

• Difference from molecular compounds: Ionic compounds are composed of ions and held together by electrostatic forces, while molecular compounds consist of neutral molecules held together by covalent bonds.

Elements vs. Compounds

Understanding the distinction between elements and compounds is fundamental in chemistry.

• Element: A substance consisting of only one kind of atom, which cannot be decomposed into simpler substances by ordinary chemical means.

• Compound: A substance consisting of two or more kinds of atoms chemically combined in fixed proportions, forming a molecule or molecular entity with properties different from its components.

• Most chemicals exist as compounds rather than as pure elements.

Classification of Elements

• Atomic elements: Exist as single atoms (e.g., Fe, Na).

• Molecular elements: Exist as molecules composed of two or more atoms of the same element (e.g., $O_2$, $N_2$, $Cl_2$, $S_8$, $P_4$).

Molecular vs. Ionic Compounds

	Molecular	Ionic
Constituents	Elements, not ions	Elements present as ions
Bonding	Held together by covalent bonds	Held by electrostatic forces
Formula type	Molecular formula	Empirical formula (formula unit)

Summary of Pure Substances and Mixtures

Pure substances can be elements or compounds, while mixtures contain two or more substances physically combined.

• Atoms: e.g., He

• Molecules: e.g., $O_2$

• Compounds: e.g., $H_2O$ (molecular), $NaCl$ (ionic)

• Mixtures: Physical combinations of elements and/or compounds

Formula Mass and the Mole Concept for Compounds

The formula mass (or molecular mass) is the sum of the atomic masses of all atoms in a chemical formula. The mole concept allows chemists to relate the mass of a substance to the number of particles (atoms, molecules, or formula units) it contains.

• Formula mass: The sum of the atomic masses of all atoms in a formula unit (for ionic compounds) or molecule (for molecular compounds).

• Mole: The amount of substance containing $6.022 \times 10^{23}$ entities (Avogadro's number).

• Molar mass: The mass of one mole of a substance, expressed in grams per mole ($g \cdot mol^{-1}$).

Example (Glucose, $C_6H_{12}O_6$):

As a single molecule	As 1 mole of glucose molecules
6 carbon atoms	6 moles of C atoms
12 H atoms	12 moles of H atoms
6 O atoms	6 moles of O atoms
Formula mass = 180.16 amu	Molar mass = 180.16 g/mol

Calculating Formula Mass

• Sum the atomic masses of all atoms in the formula.

• Example (Calcium hydroxide, $Ca(OH)_2$):

$1 \times Ca = 40.078\ g\ mol^{-1}$ $2 \times O = 2 \times 15.999\ g\ mol^{-1}$ $2 \times H = 2 \times 1.008\ g\ mol^{-1}$ Total = $74.093\ g\ mol^{-1}$

Using the Mole Concept

• To find the number of molecules: $\text{Number of molecules} = \text{moles} \times 6.022 \times 10^{23}$

• To find the number of atoms of a specific element: $\text{Number of atoms} = \text{number of molecules} \times \text{number of that atom per molecule}$

• Example: How many carbon atoms are in 0.400 mole of procaine ($C_{13}H_{20}N_2O_2$)? $0.400\ \text{mol} \times 6.022 \times 10^{23}\ \text{molecules/mol} = 2.409 \times 10^{23}\ \text{molecules}$ $2.409 \times 10^{23}\ \text{molecules} \times 13 = 3.13 \times 10^{24}\ \text{C atoms}$

Percent Composition by Mass

The percent composition of a compound expresses the mass percentage of each element in the compound.

$\text{mass \% of element X} = \frac{\text{mass of element X in 1 mol of compound}}{\text{mass of 1 mol of compound}} \times 100$

• Example (Lead(II) chromate, $PbCrO_4$): Molar mass = $323.20\ g\ mol^{-1}$

• % Cr = $\frac{52.00\ g}{323.20\ g} \times 100 = 16.09\%$

• % Pb = $\frac{207.20\ g}{323.20\ g} \times 100 = 64.11\%$

• % O = $\frac{4 \times 16.00\ g}{323.20\ g} \times 100 = 19.80\%$

Approaches to Determining Composition

1. Using percent composition

2. Using mole relationships

Additional info: These notes provide foundational concepts for understanding chemical formulas, the classification of substances, and quantitative relationships in chemistry, which are essential for further study in General Chemistry.