Chapter 4: Molecules and Compounds

Chapter 4: Molecules, Compounds, and Chemical Bonding

Learning Outcomes

• Understand the differences between molecular and empirical formulas.

• Write formulas for ionic compounds and balance charges.

• Name ionic and molecular compounds using systematic conventions.

• Calculate formula mass and use the mole concept for compounds.

• Determine chemical formulas from experimental data.

• Describe the composition of compounds using mass percent and conversion factors.

• Classify and name organic compounds, especially hydrocarbons.

4.1 Hydrogen, Oxygen, and Water

Properties and Importance

• Hydrogen (H2): Explosive gas, used as rocket fuel.

• Oxygen (O2): Essential for combustion and respiration.

• Water (H2O): Vital for life, formed by the reaction of hydrogen and oxygen.

States and Properties

• Hydrogen: Boiling point -253°C, Melting point -259°C

• Oxygen: Boiling point -183°C, Melting point -219°C

• Water: Liquid at room temperature, high boiling point

Elements vs. Compounds

• Elements combine in fixed, definite proportions to form compounds.

• Compounds have distinct properties different from their constituent elements.

• Mixtures can have any proportion of elements, unlike compounds.

Example

• A hydrogen-oxygen molecule can have any ratio of hydrogen to oxygen.

• Water always has a fixed ratio of two hydrogens to one oxygen atom.

4.2 Types of Chemical Bonds

Chemical Bonds Overview

• Ionic bonds: Formed between a metal and a nonmetal by electron transfer.

• Covalent bonds: Formed between nonmetals by electron sharing.

• Metallic bonds: Formed between metals, involving a 'sea' of delocalized electrons.

Bonding and Properties

• Ionic compounds: Crystalline solids, high melting points, conduct electricity when molten.

• Covalent compounds: Can be gases, liquids, or solids; lower melting points; do not conduct electricity.

• Metallic compounds: Malleable, ductile, conduct electricity.

Table: Types of Chemical Bonds

Bond Type	Constituent Elements	Key Properties
Ionic	Metal + Nonmetal	High melting point, crystalline, conducts electricity when molten
Covalent	Nonmetal + Nonmetal	Low melting point, can be gas/liquid/solid, does not conduct electricity
Metallic	Metal + Metal	Malleable, ductile, conducts electricity

4.3 Representing Compounds: Chemical Formulas and Molecular Models

Chemical Formulas

• Empirical formula: Shows the simplest whole-number ratio of atoms.

• Molecular formula: Shows the actual number of atoms of each element.

• Structural formula: Shows how atoms are connected.

Examples

• Empirical formula of hydrogen peroxide: HO

• Molecular formula of hydrogen peroxide: H2O2

• Structural formula: H-O-O-H

Molecular Models

• Ball-and-stick model: Spheres represent atoms, sticks represent bonds.

• Space-filling model: Spheres scaled to atomic radii, showing molecule's size and shape.

4.4 The Lewis Model: Representing Valence Electrons with Dots

Lewis Symbols and Structures

• Valence electrons are represented as dots around the element's symbol.

• Lewis structures show how atoms bond and share electrons.

• Octet rule: Atoms tend to form bonds to achieve eight valence electrons.

Example

• Lewis symbol for silicon: Si with four dots.

• Lewis structure for water: H:O:H (with two lone pairs on O).

4.5 Ionic Bonding: The Lewis Model and Lattice Energies

Ionic Bonding and Electron Transfer

• Electrons are transferred from metals to nonmetals, forming cations and anions.

• Resulting electrostatic attraction forms an ionic bond.

Lattice Energy

• Energy released when ions come together to form a crystalline lattice.

• Higher charges and smaller ions result in higher lattice energies.

Equation

• $\text{Lattice Energy} \propto \frac{Q_1 Q_2}{r}$

Writing and Naming Ionic Compounds

Formulas for Ionic Compounds

• Balance total positive and negative charges.

• Use subscripts to indicate the number of each ion.

Naming Ionic Compounds

• Binary ionic compounds: Name cation first, then anion (ending in -ide).

• Metals with variable charge: Use Roman numerals (e.g., FeCl2 is iron(II) chloride).

• Polyatomic ions: Use standard names (e.g., NaNO3 is sodium nitrate).

• Hydrated ionic compounds: Indicate number of water molecules (e.g., CuSO4·5H2O is copper(II) sulfate pentahydrate).

Table: Common Polyatomic Ions

Ion	Formula	Charge
Nitrate	NO3-	-1
Sulfate	SO42-	-2
Phosphate	PO43-	-3
Ammonium	NH4+	+1

4.7 Covalent Bonding: Simple Lewis Structures

Single, Double, and Triple Bonds

• Single bond: One shared electron pair (e.g., H-H).

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

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

Octet Rule and Bonding

• Atoms share electrons to achieve a stable octet configuration.

• Bond strength increases from single to triple bonds.

4.8 Molecular Compounds: Formulas and Names

Naming Molecular Compounds

• Use prefixes to indicate the number of each atom (mono-, di-, tri-, etc.).

• First element is named fully; second element ends in -ide.

• Example: CO2 is carbon dioxide; N2O is dinitrogen monoxide.

Table: Prefixes for Number of Atoms

Number	Prefix
1	mono-
2	di-
3	tri-
4	tetra-
5	penta-
6	hexa-
7	hepta-
8	octa-

4.9 Formula Mass and the Mole Concept for Compounds

Formula Mass

• The sum of atomic masses of all atoms in a chemical formula.

• Units: grams per mole (g/mol).

Equation

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

Using Molar Mass to Count Molecules by Weighing

• Convert mass to moles using molar mass.

• Convert moles to molecules using Avogadro's number ($6.022 \times 10^{23}$ molecules/mol).

Example Calculation

• 10 g CO2 × (1 mol/44 g) × ($6.022 \times 10^{23}$ molecules/1 mol) = $1.37 \times 10^{23}$ molecules

4.10 Composition of Compounds

Mass Percent Composition

• Indicates the percentage by mass of each element in a compound.

Equation

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

Conversion Factors from Chemical Formulas

• Use mass percent and chemical formulas as conversion factors in stoichiometric calculations.

4.11 Determining a Chemical Formula from Experimental Data

Empirical Formula Determination

• Obtain masses of each element in a sample.

• Convert masses to moles using atomic masses.

• Divide by smallest number of moles to get whole-number ratios.

Example

• Given: 24.9 g N, 70.0 g O. Moles N = 1.78, Moles O = 4.38. Ratio: 1:2.5. Empirical formula: N2O5

Molecular Formula Calculation

• Molecular formula = (empirical formula)n, where n = (molar mass)/(empirical formula mass)

Combustion Analysis

Steps

• Burn sample, measure masses of CO2 and H2O produced.

• Convert masses to moles of C and H.

• Determine empirical formula from mole ratios.

4.12 Organic Compounds

Classification and Examples

• Organic compounds: Contain carbon, often bonded to hydrogen, oxygen, nitrogen.

• Hydrocarbons: Simplest organic compounds, contain only carbon and hydrogen.

Examples

• Propane (C3H8): Chain structure

• Isobutane (C4H10): Branched structure

• Cyclohexane (C6H12): Ring structure

• Ethene (C2H4): Double bond

• Ethyne (C2H2): Triple bond

Key Concepts and Definitions

• Chemical bond: Force holding atoms together in a compound.

• Molecular compound: Atoms share electrons via covalent bonds.

• Empirical formula: Simplest whole-number ratio of atoms.

• Molecular formula: Actual number of atoms in a molecule.

• Formula mass: Sum of atomic masses in a chemical formula.

• Mass percent composition: Percent by mass of each element in a compound.

• Organic compound: Contains carbon, often with hydrogen and other elements.

• Hydrocarbon: Organic compound of only carbon and hydrogen.

Equations and Relationships

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

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

• $\text{Molecular formula} = (\text{empirical formula})^n$

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

Summary Table: Naming Ionic and Molecular Compounds

Compound Type	Naming Rule	Example
Binary Ionic (fixed charge)	Cation + Anion (-ide)	NaCl: sodium chloride
Binary Ionic (variable charge)	Cation (Roman numeral) + Anion (-ide)	FeCl2: iron(II) chloride
Ionic with Polyatomic Ion	Cation + Polyatomic Ion	NaNO3: sodium nitrate
Molecular Compound	Prefix + Element + Prefix + Element (-ide)	CO2: carbon dioxide

Review Questions

• Define and distinguish between empirical and molecular formulas.

• Explain the difference between ionic and covalent bonding.

• Describe how to name ionic compounds with fixed and variable charges.

• Calculate the formula mass and mass percent composition of a compound.

• Determine the empirical formula from experimental data.

• Classify and name simple organic compounds.

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