Chapter 03. Molecules and Compounds

Chapter 03. Molecules and Compounds

Introduction to Molecules and Compounds

The diversity of substances in nature arises from the ability of elements to combine and form compounds. While there are about 91 naturally occurring elements, their combinations result in an almost limitless number of compounds, each with unique properties. Understanding molecules and compounds is essential for comprehending the chemical basis of life and matter.

Hydrogen, Oxygen, and Water

Properties and Formation of Water

Hydrogen (H2) is an explosive gas, and oxygen (O2) is essential for combustion. Both are gases at room temperature with low boiling points. When they react, they form water (H2O), a liquid at room temperature with a much higher boiling point and very different properties from its constituent elements. This transformation illustrates how compounds have properties distinct from the elements that compose them.

• Hydrogen: Explosive, gas at room temperature, boiling point -253°C.

• Oxygen: Supports combustion, gas at room temperature, boiling point -183°C.

• Water: Smothers flames, liquid at room temperature, boiling point 100°C.

In compounds, elements combine in fixed, definite proportions, unlike mixtures where proportions can vary.

Chemical Bonds

Ionic and Covalent Bonds

Chemical bonds are the forces that hold atoms together in compounds, arising from electrostatic attractions between charged particles (electrons and protons). There are two main types of chemical bonds:

• Ionic Bonds: Formed by the transfer of electrons from a metal to a nonmetal, resulting in the formation of cations (positively charged ions) and anions (negatively charged ions). These ions are held together in a lattice by electrostatic forces.

• Covalent Bonds: Formed by the sharing of electrons between two or more nonmetals, resulting in the formation of molecules. The shared electrons lower the potential energy and hold the atoms together.

Example: Sodium (Na, a metal) reacts with chlorine (Cl, a nonmetal) to form sodium chloride (NaCl), an ionic compound.

Representing Compounds: Chemical Formulas and Molecular Models

Types of Chemical Formulas

• Empirical Formula: Shows the simplest whole-number ratio of atoms in a compound (e.g., CH2 for C4H8).

• Molecular Formula: Shows the actual number of atoms of each element in a molecule (e.g., H2O2).

• Structural Formula: Shows how atoms are connected or bonded (e.g., H–O–O–H for hydrogen peroxide).

Molecular Models

• Ball-and-Stick Model: Atoms are balls, bonds are sticks; shows geometry.

• Space-Filling Model: Atoms fill the space between each other, representing the molecule's actual size and shape.

An Atomic-Level View of Elements and Compounds

Classification of Substances

• Atomic Elements: Exist as single atoms (e.g., Ne, He).

• Molecular Elements: Exist as molecules of two or more atoms of the same element (e.g., O2, N2, Cl2).

• Molecular Compounds: Composed of molecules formed by covalently bonded nonmetals (e.g., H2O, CO2).

• Ionic Compounds: Composed of a lattice of cations and anions (e.g., NaCl).

Ionic Compounds: Formulas and Names

Writing Formulas for Ionic Compounds

Ionic compounds are charge-neutral and contain positive (cation) and negative (anion) ions. The formula reflects the smallest whole-number ratio of ions that results in charge neutrality.

• Example: Na+ and Cl- combine to form NaCl.

• Example: Ca2+ and Cl- combine to form CaCl2.

Naming Ionic Compounds

• Type 1: Metal forms only one type of cation (e.g., Na+, Ca2+). Name: [metal] [base name of nonmetal]-ide (e.g., sodium chloride).

• Type 2: Metal forms more than one type of cation (often transition metals). Name: [metal] (Roman numeral for charge) [base name of nonmetal]-ide (e.g., iron(III) chloride).

• Polyatomic Ions: Use the name of the polyatomic ion (e.g., sodium nitrate, NaNO3).

Molecular Compounds: Formulas and Names

Naming Molecular Compounds

• Composed of two or more nonmetals.

• Prefixes indicate the number of each atom: mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca-.

• Example: CO2 is carbon dioxide; N2O5 is dinitrogen pentoxide.

Naming Acids

Binary Acids

• Composed of hydrogen and a nonmetal.

• Name: hydro-[base name of nonmetal]-ic acid (e.g., HCl(aq) is hydrochloric acid).

Oxyacids

• Contain hydrogen, oxygen, and another element (usually a nonmetal).

• If the oxyanion ends in -ate: [base name of oxyanion]-ic acid (e.g., HNO3 is nitric acid).

• If the oxyanion ends in -ite: [base name of oxyanion]-ous acid (e.g., HNO2 is nitrous acid).

Formula Mass and the Mole Concept for Compounds

Formula Mass

The formula mass of a compound is the sum of the atomic masses of all atoms in its chemical formula. The molar mass (g/mol) is numerically equal to the formula mass (amu).

• Example: Formula mass of CO2 = 12.01 + 2(16.00) = 44.01 amu.

Using Molar Mass

• Converts between mass (g), amount (mol), and number of molecules (using Avogadro's number, $6.022 \times 10^{23}$).

Composition of Compounds

Mass Percent Composition

The mass percent composition of an element in a compound is the percentage of the compound's total mass contributed by that element.

Formula:

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

• Example: In CCl2F2, calculate the mass percent of Cl.

Determining a Chemical Formula from Experimental Data

Empirical and Molecular Formulas

• Empirical formula: Simplest whole-number ratio of atoms in a compound, determined from mass or percent composition data.

• Molecular formula: Actual number of atoms, found by multiplying the empirical formula by a whole number $n$ such that $\text{molar mass} = n \times \text{empirical formula mass}$.

Steps to determine empirical formula:

1. Convert mass (or percent) of each element to moles.

2. Write a pseudoformula using mole values as subscripts.

3. Divide all subscripts by the smallest number of moles.

4. If necessary, multiply to get whole numbers.

Organic Compounds

Introduction to Organic Chemistry

Organic compounds are primarily composed of carbon and hydrogen, often with oxygen, nitrogen, or sulfur. The simplest organic compounds are hydrocarbons, which contain only carbon and hydrogen. Hydrocarbons are classified as:

• Alkanes: Only single bonds (suffix -ane).

• Alkenes: At least one double bond (suffix -ene).

• Alkynes: At least one triple bond (suffix -yne).

Functionalized hydrocarbons contain additional characteristic groups (functional groups) such as alcohols (-OH), ethers (R-O-R'), aldehydes (-CHO), ketones (C=O), carboxylic acids (-COOH), esters (-COOR), and amines (-NH2).

Summary Table: Types of Substances

Type	Basic Unit	Example
Atomic Element	Single atom	Ne, He
Molecular Element	Molecule (same atom)	O2, N2
Molecular Compound	Molecule (different atoms)	H2O, CO2
Ionic Compound	Formula unit (lattice)	NaCl, CaCl2

Key Equations and Relationships

• Formula Mass: $\text{Formula mass} = \sum (\text{number of atoms} \times \text{atomic mass})$

• Mass Percent Composition: $\text{Mass percent of X} = \frac{\text{mass of X in 1 mol compound}}{\text{mass of 1 mol compound}} \times 100\%$

• Molecular Formula: $\text{Molecular formula} = n \times (\text{empirical formula})$ where $n = \frac{\text{molar mass}}{\text{empirical formula mass}}$

Practice and Application

• Classify substances as atomic element, molecular element, molecular compound, or ionic compound.

• Write and name formulas for ionic and molecular compounds, including those with polyatomic ions and hydrates.

• Calculate formula mass, molar mass, and mass percent composition.

• Determine empirical and molecular formulas from experimental data.