Molecules and Compounds

Elements to Molecules

When two or more elements combine, they form a molecule. Molecules can consist of the same element (e.g., O2) or different elements (e.g., H2O). The diversity of substances in nature arises from the ability of elements to form compounds, which are collections of molecules.

• Molecule: A group of atoms bonded together, representing the smallest fundamental unit of a chemical compound.

• Compound: A substance formed when two or more elements are chemically bonded together.

Hydrogen, Oxygen, and Water: Properties Comparison

The properties of compounds are often dramatically different from the elements that compose them. For example, hydrogen and oxygen are both gases with distinct properties, but when combined to form water, the resulting compound is a liquid with unique characteristics.

• Boiling Point: Hydrogen (-253°C), Oxygen (-183°C), Water (100°C)

• State at Room Temperature: Hydrogen and oxygen are gases; water is a liquid.

• Flammability: Hydrogen is explosive, oxygen is necessary for combustion, and water is used to extinguish flames.

Law of Definite Proportion

The Law of Definite Proportion states that a chemical compound always contains exactly the same proportion of elements by mass. For example, water (H2O) always consists of two hydrogen atoms for every one oxygen atom, regardless of the sample size.

• Mixtures can have variable proportions of components.

• Compounds have fixed, definite proportions of their constituent elements.

Types of Chemical Bonds

Chemical Bonds Overview

Atoms in compounds are held together by chemical bonds, which arise from the attractions between charged particles (electrons and protons). There are three main types of chemical bonds:

• Ionic Bonds: Formed by the transfer of electrons from a metal to a nonmetal.

• Covalent Bonds: Formed by the sharing of electrons between nonmetals.

• Metallic Bonds: Involve the pooling of electrons among a lattice of metal atoms (not covered in detail here).

Ionic Bonds

Ionic bonds occur between metals and nonmetals. Metals lose electrons to become positively charged cations, while nonmetals gain electrons to become negatively charged anions. The resulting electrostatic attraction forms an ionic compound, which typically exists as a crystalline lattice.

• Example: Sodium chloride (NaCl) forms from sodium (Na) and chlorine (Cl) atoms.

Covalent Bonds

Covalent bonds occur between nonmetal atoms and involve the sharing of electrons. The atoms are held together as a discrete molecule, known as a molecular compound.

• Example: Water (H2O) is formed by covalent bonding between hydrogen and oxygen atoms.

Representing Compounds

Chemical Formulas and Molecular Models

Chemical compounds are represented by chemical formulas, which indicate the types and numbers of atoms present. There are several ways to represent compounds:

• Empirical Formula: Simplest whole-number ratio of elements (e.g., CH2).

• Molecular Formula: Actual number of atoms of each element (e.g., C2H4).

• Structural Formula: Shows how atoms are bonded (e.g., H–O–H for water).

• Molecular Models: Ball-and-stick and space-filling models provide 3D representations.

Summary Table: Empirical, Molecular, Structural, and Model Representations

The following table summarizes the different ways to represent compounds such as benzene, acetylene, glucose, and ammonia:

Lewis Structures and the Octet Rule

Lewis Structure Model

The Lewis structure is a way to represent the valence electrons of atoms within a molecule using dots. It is especially useful for illustrating covalent bonding and predicting molecular structure.

• Valence electrons are shown as dots around the chemical symbol.

• Lewis structures help visualize the transfer or sharing of electrons.

The Octet Rule

Atoms tend to gain, lose, or share electrons to achieve a noble gas configuration, typically eight valence electrons (the octet rule). Hydrogen, however, follows the duet rule (two electrons).

• Period 2 nonmetals must obey the octet rule.

• Exceptions: H, Li, Be, and B can have fewer than eight electrons.

Lewis Structures for Ionic Compounds

Lewis symbols can also represent the transfer of electrons in ionic bonding, such as between calcium and chlorine:

Ionic Compounds: Properties and Conductivity

Crystal Lattice and Conductivity

Ionic compounds form a crystal lattice structure, maximizing the attraction between cations and anions. In the solid state, ions are fixed in place and do not conduct electricity. When melted or dissolved in water, ions are free to move, allowing the compound to conduct electricity.

Writing and Naming Ionic Compounds

Ionic compounds are named based on the cation and anion present. The sum of the charges must equal zero, and the formula reflects the smallest whole-number ratio of ions.

• Type I: Metal forms only one type of ion (e.g., Na+, Ca2+).

• Type II: Metal forms more than one type of ion (e.g., Fe2+, Fe3+).

Type II Cations and Polyatomic Ions

Some metals can form cations with different charges. Their names include a Roman numeral to indicate the charge. Polyatomic ions are groups of atoms that carry a charge and act as a single ion in compounds.

Covalent Bonding and Molecular Compounds

Bonding and Lone Pairs

In covalent bonds, electrons are shared between atoms. Shared electrons are called bonding pairs, while unshared electrons are lone pairs.

Single, Double, and Triple Bonds

Atoms can share one, two, or three pairs of electrons, forming single, double, or triple bonds, respectively.

• Single bond: One pair of shared electrons (e.g., H–H).

• Double bond: Two pairs of shared electrons (e.g., O=O).

• Triple bond: Three pairs of shared electrons (e.g., N≡N).

Naming Molecular Compounds

Molecular compounds are named using prefixes to indicate the number of each type of atom. The first element is named first, followed by the second element with the suffix -ide.

• Prefixes: mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca-

• Example: CO2 is carbon dioxide; PCl5 is phosphorus pentachloride.

Formula Mass and Percent Composition

Calculating Formula Mass

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

• Example: For H2O, amu

Using Molar Mass and Avogadro's Number

Molar mass allows conversion between mass and number of molecules using Avogadro's number ( particles/mol).

Percent Composition

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

Empirical and Molecular Formulas

Determining Empirical Formulas

The empirical formula gives the simplest whole-number ratio of atoms in a compound. It can be determined from percent composition or experimental data.

• Convert percentages to grams (assume 100 g sample).

• Convert grams to moles using molar mass.

• Divide by the smallest number of moles to get ratios.

• Multiply to obtain whole numbers if necessary.

From Empirical to Molecular Formula

The molecular formula is a whole-number multiple of the empirical formula. To find it, divide the compound's molar mass by the empirical formula mass:

Molecular formula = (empirical formula)n

Combustion Analysis

Combustion Analysis Technique

Combustion analysis is used to determine the empirical formula of organic compounds. The compound is burned, and the masses of CO2 and H2O produced are measured to calculate the amounts of C, H, and O in the original sample.

Organic Compounds

Definition and Characteristics

Organic compounds are primarily composed of carbon and hydrogen, often with oxygen, nitrogen, phosphorus, sulfur, and other elements. Carbon's ability to form four covalent bonds and create chains and rings is central to organic chemistry.

• Organic compounds are now synthesized in the lab and found throughout nature.

• Carbon forms the backbone of organic molecules.