Molecules, Compounds, and Nomenclature: Chemical Formulas, Models, and Composition

Molecules, Compounds, and Nomenclature

Introduction

This section introduces the fundamental concepts of molecules, compounds, and their nomenclature, which are essential for understanding chemical formulas, molecular models, and the composition of substances in general chemistry.

Chemical Formulas and Molecular Models

Types of Chemical Formulas

• Molecular Formula: Shows the exact number and type of atoms in a molecule. For example, the molecular formula for hydrogen peroxide is H2O2.

• Empirical Formula: Indicates the simplest whole-number ratio of atoms in a compound. For example, the empirical formula for H2O2 is HO.

Key Point: Molecular formulas provide the actual number of atoms, while empirical formulas show only the simplest ratio.

• What do molecular formulas tell you? The exact number and type of atoms in a molecule.

• What do they NOT tell you? The arrangement of atoms (structure) or the simplest ratio (unless the molecule is already in its simplest form).

Example: For glucose (C6H12O6), the empirical formula is CH2O.

Ionic Compounds

Definition and Formation

• Ionic Compound: A chemical compound composed of cations (positively charged ions) and anions (negatively charged ions) that combine to form a neutral substance.

• Typically formed from the reaction of metals (which lose electrons to become cations) and nonmetals (which gain electrons to become anions).

• Cation: An atom or molecule that has lost one or more electrons, gaining a positive charge.

• Anion: An atom or molecule that has gained one or more electrons, gaining a negative charge.

Key Point: Ionic compounds are held together by electrostatic forces between oppositely charged ions.

Example: Sodium chloride (NaCl) is formed from Na+ and Cl- ions.

Comparison: Ionic vs. Molecular Compounds

• Molecular Compounds: Consist of nonmetals bonded together by covalent bonds (sharing electrons).

• Ionic Compounds: Consist of metals and nonmetals bonded by the transfer of electrons (ionic bonds).

Elements vs. Compounds

Definitions

• Element: A pure substance consisting of only one type of atom. Examples include O2 (oxygen), Ne (neon).

• Compound: A substance formed when two or more different elements are chemically bonded together, having properties different from its constituent elements. Examples include H2O (water), NaCl (sodium chloride).

Key Point: Most chemicals exist as compounds, not as pure elements.

Classification of Pure Substances

Type	Example
Element (Atomic)	Ne (neon)
Element (Molecular)	O2 (oxygen)
Compound (Molecular)	H2O (water)
Compound (Ionic)	NaCl (sodium chloride)

Formula Mass and the Mole Concept for Compounds

Definitions and Calculations

• Formula Mass: The sum of the atomic masses of all atoms in a chemical formula (for both molecular and ionic compounds).

• Mole: A unit representing 6.022 × 1023 entities (Avogadro's number).

• Molar Mass: The mass of one mole of a substance, usually expressed in grams per mole (g/mol).

Key Point: Formula mass and molecular weight are often used interchangeably.

Ways to Look at Glucose (C6H12O6)

As a Single Molecule	As 1 Mole of Glucose Molecules
Contains 6 C, 12 H, 6 O atoms	Contains 6 × Avogadro's number of C atoms, 12 × Avogadro's number of H atoms, 6 × Avogadro's number of O atoms
Molecular mass = sum of atomic masses	Molar mass = molecular mass in grams

Example Calculation:

• Calculate the formula mass and molar mass of calcium hydroxide, Ca(OH)2:

$\text{Formula mass} = 40.08\ (Ca) + 2 \times [16.00\ (O) + 1.01\ (H)] = 74.10\ \text{g/mol}$

• How many atoms of carbon are there in 0.40 mole of procaine, C13H20N2O2?

$0.40\ \text{mol} \times 6.022 \times 10^{23}\ \text{molecules/mol} \times 13\ \text{C atoms/molecule} = 3.13 \times 10^{24}\ \text{C atoms}$

Composition of Compounds

Percent Composition

• Percent Composition: The percentage by mass of each element in a compound.

• Calculated using the formula:

$\%\ \text{Element} = \frac{\text{Mass of element in 1 mole of compound}}{\text{Molar mass of compound}} \times 100\%$

• Example: Lead(II) chromate (PbCrO4) is used in paint pigment. To find the percent by mass of each element, sum the atomic masses and divide each element's contribution by the total molar mass.

• Example: How many grams of oxygen are in 50.0 g of carbon dioxide (CO2)?

$\text{Molar mass of CO}_2 = 44.01\ \text{g/mol}$ $\text{Mass of O in 1 mol CO}_2 = 2 \times 16.00 = 32.00\ \text{g}$ $\%\ \text{O} = \frac{32.00}{44.01} \times 100\% = 72.7\%$ $\text{Mass of O in 50.0 g CO}_2 = 50.0\ \text{g} \times 0.727 = 36.4\ \text{g}$

Summary of Key Steps

• Calculate the formula mass of a compound.

• Determine the mass percentage of each element.

• Apply mole relationships to convert between mass, moles, and number of atoms.

Additional info: The notes reference sections not covered (e.g., organic compounds and polyatomic ions), but the core content here is foundational for understanding chemical formulas, types of compounds, and composition calculations in general chemistry.