Chapter 4: Molecules and Compounds – Study Notes

Chapter 4: Molecules and Compounds

Overview

This chapter introduces the fundamental concepts of molecules and compounds, focusing on how elements combine to form new substances, the types of chemical bonds, and the representation and naming of compounds. Understanding these concepts is essential for studying the structure, properties, and behavior of matter in chemistry.

Elements to Molecules

Formation of Molecules

• Molecule: A group of two or more atoms bonded together.

• Elements can form molecules by combining with themselves (e.g., O2, H2) or with different elements (e.g., H2O).

• Compound: A substance composed of two or more different elements chemically combined in fixed proportions.

• Compounds can be ionic (formed from metals and nonmetals) or molecular (formed from nonmetals).

• The diversity of substances in nature is due to the ability of elements to form a wide variety of compounds.

4.1 Hydrogen, Oxygen, and Water

Differences Between Elements and Compounds

• When elements combine to form compounds, the resulting substance has entirely new properties.

• Example: Hydrogen and oxygen are both gases with distinct properties, but when combined, they form water, a liquid with very different characteristics.

Substance	Melting Point	State at Room Temp	Reactivity
Hydrogen	-253°C	Gas	Explosive
Oxygen	-183°C	Gas	Necessary for combustion
Water	0°C	Liquid	Used to extinguish flame

4.2 Types of Chemical Bonds

Chemical Bonds and Their Types

• Chemical bond: The force that holds atoms together in a compound, resulting from attractions between charged particles (electrons and protons).

• Three main types of chemical bonds:

  • Ionic bonds: Formed between metals and nonmetals; involve the transfer of electrons from one atom to another, resulting in cations and anions.

  • Covalent bonds: Formed between nonmetals; involve the sharing of electrons between atoms, resulting in molecules.

  • Metallic bonds: Occur between metal atoms; involve a 'sea' of delocalized electrons shared among a lattice of metal cations.

• Example: Table salt (NaCl) is held together by ionic bonds, while water (H2O) is held together by covalent bonds.

4.3 Representing Compounds: Chemical Formulas and Molecular Models

Chemical Formulas

• Chemical formula: Indicates the type and number of each element in a compound.

• Types of chemical formulas:

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

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

  • Structural formula: Shows how atoms are bonded and arranged in a molecule, using lines for bonds.

• Example: For glucose:

  • Empirical formula: CH2O

  • Molecular formula: C6H12O6

  • Structural formula: Shows the arrangement of atoms and bonds (not shown here).

Molecular Models

• Ball-and-stick model: Atoms are represented as balls and bonds as sticks, showing the shape of the molecule.

• Space-filling model: Atoms are shown as spheres that fill the space between them, representing the molecule's approximate size and shape.

4.4 The Lewis Model: Representing Valence Electrons

Lewis Structures

• Valence electrons: Electrons in the outermost shell of an atom, involved in bonding.

• Lewis structure: A diagram showing the arrangement of valence electrons around atoms, using dots for electrons.

• Octet rule: Atoms tend to gain, lose, or share electrons to achieve eight valence electrons (a noble gas configuration).

• Exceptions: Hydrogen (duet, 2 electrons), Boron (often 6 electrons), and elements in period 3 or higher (can have expanded octets).

4.5 Ionic Bonding and the Crystal Lattice

Formation and Properties of Ionic Compounds

• Ionic compounds form a crystal lattice, a repeating three-dimensional arrangement of ions.

• Lattice energy: The energy released when ions come together to form a crystal lattice from the gaseous state.

• Lattice energy depends on the charges of the ions and the distance between them:

$\text{Lattice Energy} \propto \frac{q_1 q_2}{r}$

• Higher charges and smaller ionic radii result in greater lattice energy (stronger ionic bonds).

• Ionic compounds conduct electricity when molten or dissolved in water, as ions are free to move.

4.6 Naming Ionic Compounds

Rules for Naming

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

• Type II: Metal can form more than one type of cation (e.g., Fe2+, Fe3+).

• For binary ionic compounds (two elements):

  • Name the cation (metal) first, then the anion (nonmetal) with the suffix '-ide'.

  • For Type II metals, indicate the charge with Roman numerals in parentheses.

  • Example: FeCl2 is iron(II) chloride; FeCl3 is iron(III) chloride.

• For compounds with polyatomic ions, use the name of the polyatomic ion (e.g., nitrate, sulfate).

• Oxyanions (polyatomic ions with oxygen): '-ate' for more oxygen, '-ite' for less (e.g., sulfate SO42-, sulfite SO32-).

4.7 Covalent Bonding: Bonding and Lone Pair Electrons

Types of Covalent Bonds

• Bonding pairs: Shared electrons between atoms (form bonds).

• Lone pairs: Non-bonding electrons localized on a single atom.

• 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).

4.8 Molecular Compounds: Formulas and Names

Naming Binary Molecular Compounds

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

• The first element keeps its name; the second element ends with '-ide'.

• Prefixes are always used for the second element; 'mono-' is omitted for the first element.

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

4.9 Hydrated Ionic Compounds

Hydrates

• Hydrate: An ionic compound with a specific number of water molecules associated with each formula unit.

• Named using Greek prefixes (mono-, di-, tri-, etc.) to indicate the number of water molecules.

• Example: MgSO4·7H2O is magnesium sulfate heptahydrate.

4.10 Composition of Compounds

Percent Composition

• The chemical formula and molar masses allow calculation of the percent by mass of each element in a compound.

• Percent composition:

$\text{Percent by mass of element} = \frac{\text{mass of element in 1 mol compound}}{\text{molar mass of compound}} \times 100\%$

• Percentages may not total exactly 100% due to rounding.

4.11 Determining a Chemical Formula from Experimental Data

Empirical and Molecular Formulas

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

• Molecular formula: Actual number of atoms of each element in a molecule; a whole-number multiple of the empirical formula.

• To determine the empirical formula from percent composition:

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

  2. Convert grams to moles for each element.

  3. Divide by the smallest number of moles to get the simplest ratio.

  4. Multiply to obtain whole numbers if necessary.

• To find the molecular formula, use the molar mass:

$\text{Molecular formula} = (\text{Empirical formula})_n$

$n = \frac{\text{Molar mass}}{\text{Empirical formula mass}}$

• Example: If the empirical formula is CH2O and the molar mass is 180 g/mol, then $n = 6$ and the molecular formula is C6H12O6.