Chemical Formulas, Compounds, and Nomenclature: Study Notes for General Chemistry I

Chemical Formulas and Molecular Models

Introduction to Chemical Formulas

Chemical formulas are symbolic representations that convey the types and relative numbers of atoms in a molecule or compound. They are fundamental tools in chemistry for describing substances and predicting their behavior.

• Molecular Formula: Shows the exact number of each type of atom in a molecule (e.g., H2O, CO2, C6H12O6).

• Empirical Formula: Shows the simplest whole-number ratio of atoms in a compound (e.g., CH2O for glucose).

• What molecular formulas tell you: The relative numbers and types of atoms in a molecule.

• What molecular formulas do NOT tell you: The arrangement of atoms (structure) or how atoms are bonded.

Example: The molecular formula for glucose is C6H12O6, while its empirical formula is CH2O.

Ionic Compounds

Definition and Formation

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

• Example: NaCl is formed from Na+ (cation) and Cl- (anion).

Formation of Ions:

• Atoms that lose electrons form cations.

• Atoms that gain electrons form anions.

Comparison: Ionic vs. Molecular Compounds

• Molecular Compounds: Composed of nonmetals, held together by covalent bonds, represented by molecular formulas.

• Ionic Compounds: Composed of ions, held together by electrostatic forces, represented by empirical formulas (simplest ratio).

Example Table:

Property	Molecular Compounds	Ionic Compounds
Constituents	Nonmetals	Metals & Nonmetals
Bonding	Covalent	Ionic (electrostatic)
Formula Type	Molecular formula	Empirical formula
Example	H2O	NaCl

Elements vs. Compounds

Definitions and Differences

An element is a pure substance consisting of only one type of atom, while a compound is a substance formed from two or more different elements chemically bonded together. Compounds have properties distinct from their constituent elements.

• Element: Cannot be broken down into simpler substances by chemical means.

• Compound: Can be broken down into elements by chemical reactions.

• Molecule: The smallest unit of a compound that retains its chemical properties.

Example: Water (H2O) is a compound formed from hydrogen and oxygen elements.

Classification of Pure Substances

• Elements: Can exist as individual atoms (e.g., Ne) or as molecules (e.g., O2).

• Compounds: Can be molecular (e.g., H2O) or ionic (e.g., NaCl).

Example Table:

Type	Example	Representation
Atomic Element	Ne	Ne
Molecular Element	O2	O2
Molecular Compound	H2O	H2O
Ionic Compound	NaCl	NaCl (formula unit)

Formula Mass and the Mole Concept for Compounds

Calculating Formula Mass and Moles

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: Calculated by adding the atomic masses of each atom in the formula.

• Molar Mass: The mass of one mole of a substance, numerically equal to the formula mass in grams per mole (g/mol).

Example Calculation for Glucose (C6H12O6):

	As a single molecule	As 1 mole of molecules
C atoms	6	6 × Avogadro's number
H atoms	12	12 × Avogadro's number
O atoms	6	6 × Avogadro's number
Formula mass	180.16 amu	180.16 g

Key Equations:

• Number of moles: $n = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

• Number of particles: $\text{Number of particles} = n \times N_A$ where $N_A$ is Avogadro's number ($6.022 \times 10^{23}$ mol-1).

Example: How many atoms of carbon are in 0.40 mole of procaine (C13H20N2O2)?

• Number of C atoms = $0.40 \text{ mol} \times 6.022 \times 10^{23} \text{ molecules/mol} \times 13 \text{ C atoms/molecule}$

Composition of Compounds

Percent Composition

The percent composition of a compound is the percentage by mass of each element in the compound. It is calculated using the formula mass and the atomic masses of the constituent elements.

• Percent by mass of element X: $\%X = \frac{\text{mass of X in 1 mol of compound}}{\text{molar mass of compound}} \times 100\%$

Example: Calculate the percent by mass of each element in lead(II) chromate (PbCrO4).

• Find the molar mass of PbCrO4.

• Calculate the mass of each element in one mole.

• Apply the percent composition formula for each element.

Application: Determining how many grams of oxygen are in 50.0 g of CO2 involves using the percent composition of oxygen in CO2 and multiplying by the sample mass.

Summary Table: Steps for Percent Composition

Step	Description
1	Calculate molar mass of the compound
2	Find total mass of each element in one mole
3	Divide mass of each element by molar mass, multiply by 100%

Additional info:

• Section 3.4 (Formulas and Names, including polyatomic anions) is important for nomenclature but not covered in detail here.

• Section 3.5 (Organic Compounds) is not required for this section.