Chapter 4: Molecules and Compounds – Structured Study Notes

Chapter 4: Molecules and Compounds

Introduction

This chapter explores the nature of molecules and compounds, focusing on how elements combine to form new substances, the types of chemical bonds, and the representation and calculation of chemical formulas. Understanding these concepts is fundamental to the study of chemistry, as it provides the basis for predicting compound properties and behaviors.

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 properties from their constituent elements.

• 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 Compared to Its Elements

Water is a compound with properties distinct from its constituent elements, hydrogen and oxygen.

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 an ionic bond.

• 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 between atoms. The shared electrons are most stable between the bonding atoms, resulting in the formation of individual molecules.

• Bonding pair: Shared electrons between atoms.

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

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

Representing Compounds: Chemical Formulas and Models

Chemical Formulas

A chemical formula indicates the type and number of each atom present in a molecule. Subscripts show the relative number of each atom, with the order of elements typically starting with the most metallic.

• Empirical formula: Simplest whole-number ratio of atoms (e.g., CH2O for glucose).

• Molecular formula: Actual number of atoms in a molecule (e.g., C6H12O6 for glucose).

• Structural formula: Diagram showing how atoms are bonded.

Empirical vs. Molecular Formulas

The empirical formula gives the simplest ratio of atoms, while the molecular formula gives the actual number of atoms. For example:

• Molecular formula: C4H8 → Empirical formula: CH2

• Molecular formula: B2H6 → Empirical formula: BH3

• Molecular formula: CCl4 → Empirical formula: CCl4

Molecular Models

Different models are used to represent molecules:

• Ball-and-stick model: Atoms as balls, bonds as sticks, showing shape.

• Space-filling model: Atoms fill the space between one another.

The Lewis Model and Electron Dot Structures

Lewis Dot Model

The Lewis electron-dot structure focuses on valence electrons, which are involved in chemical bonding. The octet rule states that atoms tend to gain, lose, or share electrons to achieve a noble gas configuration (eight electrons in the outer shell).

• Duet rule: Applies to hydrogen and helium (two electrons).

• Octet rule: Applies to most other elements (eight electrons).

Exceptions to the Octet Rule

Some elements (Z = 1–5) do not follow the octet rule strictly. For example, hydrogen and helium follow the duet rule, while boron and beryllium may have fewer than eight electrons.

Ionic Compounds

Structure and Properties

Ionic compounds are composed of cations and anions bound by ionic bonds. The sum of the charges must equal zero, and the formula represents the smallest whole-number ratio of ions (formula unit).

• Example: NaBr (Na+ and Br-), CaHPO4 (Ca2+ and HPO42-).

• Crystal lattice: 3D network of ions held by electrostatic attraction.

Lattice Energy

Lattice energy is the extra stability gained when ions form a crystal lattice. It is proportional to the charge size and inversely proportional to the distance between ions.

• Energy is released when the crystal forms.

• Ionic solids are hard and brittle due to strong electrostatic forces.

Electrical Conductivity

Ionic solids do not conduct electricity in the solid state because ions are locked in position. However, when dissolved in water or melted, ions can move and conduct electricity.

Naming Ionic Compounds

Rules for Ionic Compound Formulas

• Contain both positive and negative ions.

• Sum of positive charges equals sum of negative charges.

• Formula shows smallest whole-number ratio of ions.

Types of Metals in Ionic Compounds

Metals may form only one type of ion or more than one type. This affects naming conventions.

Naming Ionic Compounds

• If the metal forms only one type of ion: name of cation (metal) + base name of anion (nonmetal) + -ide

• If the metal forms more than one type of ion: name of cation (metal) + charge in Roman numerals + base name of anion (nonmetal) + -ide

Common Monoatomic and Polyatomic Ions

Monoatomic ions are single atoms with a charge, while polyatomic ions are groups of atoms with a charge.

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+

Hydrates

Hydrates are ionic compounds containing water molecules within their crystal structure. Prefixes indicate the number of water molecules.

• hemi = ½, mono = 1, di = 2, tri = 3, tetra = 4, penta = 5, hexa = 6, hepta = 7, octa = 8

• Example: CuSO4·5H2O is copper sulfate pentahydrate.

Covalent Bonding and Molecular Compounds

Bonding Pairs and Lone Pairs

Electrons shared between atoms are called bonding pairs, while those not shared are lone pairs. Covalent bonds can be single, double, or triple, depending on the number of shared electron pairs.

Diatomic Elements and Multiple Bonds

Some elements exist naturally as diatomic molecules (e.g., H2, O2, N2). Double and triple bonds involve sharing two or three pairs of electrons, respectively.

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

• Example: N2O4 is dinitrogen tetroxide.

Molar Mass and Composition Calculations

Molar Mass

The molar mass of a compound is the mass of one mole of its molecules, calculated by summing the atomic masses of all atoms in the formula.

• 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 as:

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

• Example: Calculate percent calcium by mass in CaCO3.

Empirical and Molecular Formulas from Experimental Data

Determining Empirical Formula

The empirical formula is the simplest whole-number ratio of atoms in a compound. It can be determined from percent composition, decomposition data, or combustion analysis.

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

2. Convert grams to moles using atomic masses.

3. Divide all moles by the smallest number to get ratios.

4. Multiply ratios to obtain whole numbers (see table below).

Combustion Analysis

Combustion analysis is used to determine the empirical formula by burning a known mass of compound and measuring the products (CO2 and H2O).

From Empirical to Molecular Formula

The molecular formula is a multiple of the empirical formula:

• $\text{Molecular formula} = (\text{Empirical formula}) \times n$

• $n = \frac{\text{molar mass}}{\text{empirical formula molar mass}}$

Summary

This chapter covers the essential concepts of molecules and compounds, including types of chemical bonds, formula representation, naming conventions, and calculation of molar mass and percent composition. Mastery of these topics is crucial for understanding chemical reactions and properties of substances.