Chapter 4: Molecules and Compounds – Structured Study Notes

Chapter 4: Molecules and Compounds

Introduction

This chapter explores the formation, classification, and properties of molecules and compounds, focusing on the types of chemical bonds, chemical formulas, and methods for determining chemical composition. Understanding these concepts is fundamental to the study of chemistry, as they explain how elements combine to form new substances with unique properties.

Molecules and Compounds

Formation of Molecules and Compounds

When two or more elements combine, they form a molecule. If the elements are the same, such as O2, the molecule is elemental. If the elements are different, such as H2O, the molecule is a compound. Compounds are groups of molecules with distinct chemical properties.

• Molecule: Two or more atoms bonded together.

• Compound: A substance formed from two or more different elements.

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

Properties of Water, Hydrogen, and Oxygen

Compounds have properties distinct from their constituent elements. For example, water is a liquid at room temperature and extinguishes flames, while hydrogen and oxygen are gases and highly reactive.

Selected Properties	Hydrogen	Oxygen	Water
Boiling Point	-253°C	-183°C	100°C
State at Room Temperature	Gas	Gas	Liquid
Flammability	Explosive	Necessary for combustion	Used to extinguish flame

Types of Chemical Bonds

Ionic Bonds

Ionic bonds occur between metals and nonmetals and involve the transfer of electrons. The metal atom loses electrons to become a cation, while the nonmetal gains electrons to become an anion. The resulting oppositely charged ions attract each other, forming a stable ionic compound.

• Cation: Positively charged ion (metal loses electrons).

• Anion: Negatively charged ion (nonmetal gains electrons).

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

Covalent Bonds

Covalent bonds occur between nonmetals and involve the sharing of electrons. The shared electrons are most stable between the bonding atoms, resulting in the formation of individual molecules rather than a lattice.

• Bonding pair: Shared electrons between atoms.

• Lone pair: Electrons not shared, belonging to a single atom.

• Example: Water (H2O) molecules are held together by covalent bonds.

Representing Compounds: Chemical Formulas and Models

Chemical Formulas

A chemical formula indicates the type and number of each atom in a molecule. Subscripts show the relative number of atoms. The order of elements typically places the more metallic element first.

• Empirical formula: Simplest whole-number ratio of atoms (e.g., CH for benzene).

• Molecular formula: Actual number of atoms (e.g., C6H6 for benzene).

• Structural formula: Diagram showing how atoms are bonded.

Empirical vs. Molecular Formulas

The empirical formula is the simplest ratio, while the molecular formula shows the actual number of atoms. For example, glucose has an empirical formula of CH2O and a molecular formula of C6H12O6.

Name of Compound	Empirical Formula	Molecular Formula	Structural Formula	Ball-and-Stick Model	Space-Filling Model
Benzene	CH	C6H6	Ring structure	Ball-and-stick	Space-filling
Acetylene	CH	C2H2	Linear	Ball-and-stick	Space-filling
Glucose	CH2O	C6H12O6	Ring structure	Ball-and-stick	Space-filling
Ammonia	NH3	NH3	Trigonal pyramidal	Ball-and-stick	Space-filling

Lewis Structures and the Octet Rule

Lewis Dot Model

The Lewis dot model represents valence electrons as dots around the element symbol. Chemical bonding involves the transfer or sharing of these electrons to achieve a stable configuration, usually an octet (eight electrons).

• Octet rule: Atoms tend to gain, lose, or share electrons to achieve eight in their outer shell.

• Duet rule: Hydrogen and helium achieve stability with two electrons.

Bonding and Lone Pairs

In Lewis structures, bonding pairs are shared between atoms, while lone pairs belong to a single atom. The arrangement of these pairs determines the molecule's shape and stability.

Ionic Compounds: Structure and Properties

Crystal Lattice and Lattice Energy

Ionic compounds form a three-dimensional network called a crystal lattice. The stability of this lattice is measured by lattice energy, which is the energy released when ions coalesce to form the solid.

• Lattice energy: Proportional to the charge size and inversely proportional to the distance between ions.

• Properties: Ionic solids are hard, brittle, and do not conduct electricity unless dissolved in water.

Electrical Conductivity

Ionic solids do not conduct electricity in the solid state because ions are locked in place. When dissolved in water, ions are free to move and conduct electricity.

Nomenclature of Ionic Compounds

Types of Ionic Compounds

Ionic compounds are classified based on whether the metal forms only one type of ion or multiple types. This affects the naming convention used.

Naming Ionic Compounds

For metals that form only one type of ion, the compound is named by stating the metal followed by the base name of the nonmetal with the suffix -ide. For metals that form more than one type of ion, the charge is indicated in parentheses using Roman numerals.

Metal	Ion	Name	Older Name*
Chromium	Cr2+, Cr3+	Chromium(II), Chromium(III)	Chromous, Chromic
Iron	Fe2+, Fe3+	Iron(II), Iron(III)	Ferrous, Ferric
Cobalt	Co2+, Co3+	Cobalt(II), Cobalt(III)	Cobaltous, Cobaltic
Copper	Cu+, Cu2+	Copper(I), Copper(II)	Cuprous, Cupric
Tin	Sn2+, Sn4+	Tin(II), Tin(IV)	Stannous, Stannic
Mercury	Hg2+, Hg22+	Mercury(II), Mercury(I)	Mercuric, Mercurous
Lead	Pb2+, Pb4+	Lead(II), Lead(IV)	Plumbous, Plumbic

Monatomic and Polyatomic Anions

Monatomic anions are named by adding -ide to the base name. Polyatomic ions have specific names and formulas.

Nonmetal	Symbol for Ion	Base Name	Anion Name
Fluorine	F-	Fluor	Fluoride
Chlorine	Cl-	Chlor	Chloride
Bromine	Br-	Brom	Bromide
Iodine	I-	lod	Iodide
Oxygen	O2-	Ox	Oxide
Sulfur	S2-	Sulf	Sulfide
Nitrogen	N3-	Nitr	Nitride
Phosphorus	P3-	Phosph	Phosphide

Name	Formula	Name	Formula
Acetate	C2H3O2-	Hypochlorite	ClO-
Carbonate	CO32-	Chlorite	ClO2-
Hydrogen carbonate	HCO3-	Chlorate	ClO3-
Hydroxide	OH-	Perchlorate	ClO4-
Nitrite	NO2-	Permanganate	MnO4-
Nitrate	NO3-	Sulfite	SO32-
Chromate	CrO42-	Sulfate	SO42-
Dichromate	Cr2O72-	Hydrogen sulfate	HSO4-
Phosphate	PO43-	Cyanide	CN-
Hydrogen phosphate	HPO42-	Peroxide	O22-
Dihydrogen phosphate	H2PO4-	Ammonium	NH4+

Covalent Bonding and Molecular Compounds

Single, Double, and Triple Bonds

Covalent bonds can involve the sharing of one, two, or three pairs of electrons, resulting in single, double, or triple bonds, respectively. The number of shared pairs affects the bond strength and length.

• Single bond: One pair of electrons shared.

• Double bond: Two pairs of electrons shared.

• Triple bond: Three pairs of electrons shared.

Naming Molecular Compounds

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

• mono = 1

• di = 2

• tri = 3

• tetra = 4

• penta = 5

• hexa = 6

• hepta = 7

• octa = 8

• nona = 9

• deca = 10

Molar Mass and Chemical Composition

Molar Mass Calculation

The molar mass of a compound is the sum of the atomic masses of all atoms in the formula. It is used to convert between grams and moles.

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

• Example: Molar mass of H2O = 2(1.01) + 16.00 = 18.02 g/mol

Percent Composition

The percent composition of an element in a compound is calculated by dividing the mass of the element by the total mass of the compound and multiplying by 100%.

• Formula: $\text{mass\% X} = \frac{\text{mass of X}}{\text{total mass of compound}} \times 100\%$

Empirical and Molecular Formulas from Experimental Data

Determining Empirical Formula

The empirical formula is determined from percent composition, decomposition data, or combustion analysis. The process involves converting mass to moles, finding the simplest ratio, and adjusting for whole numbers.

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

• Convert grams to moles using atomic masses.

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

• Multiply ratios to obtain whole numbers if necessary.

Combustion Analysis

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

Summary

This chapter provides a comprehensive overview of molecules and compounds, including their formation, types of chemical bonds, methods for representing and naming compounds, and techniques for determining chemical composition. Mastery of these concepts is essential for understanding chemical reactions and the properties of substances.