Molecules, Compounds, and Nomenclature: Study Notes

Chapter 3: Molecules, Compounds, and Nomenclature

Chapter Objectives

• Understand the difference between ionic and covalent bonds.

• Write and interpret empirical, molecular, and structural formulas.

• Differentiate between atomic/molecular elements and ionic/molecular compounds.

• Apply rules for writing and naming ionic and molecular compounds, acids, and organic compounds.

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

• Determine composition and empirical/molecular formulas from experimental data.

Chemical Bonds

Ionic Bonds

Ionic bonds are formed between metals and non-metals. They involve the transfer of electrons from the metal (which becomes a cation) to the non-metal (which becomes an anion). The resulting electrostatic attraction between oppositely charged ions forms the ionic bond.

• Metals lose electrons to form positively charged cations.

• Non-metals gain electrons to form negatively charged anions.

• The resulting compound is electrically neutral.

Examples:

• $\mathrm{Na}^+ + \mathrm{Cl}^- \rightarrow \mathrm{NaCl}$

• $\mathrm{Mg}^{2+} + 2\mathrm{Cl}^- \rightarrow \mathrm{MgCl}_2$

• $\mathrm{Ca}^{2+} + \mathrm{O}^{2-} \rightarrow \mathrm{CaO}$

The formation of a lattice structure is typical for solid-phase ionic compounds, where ions are arranged in a repeating, three-dimensional pattern.

Covalent Bonds

Covalent bonds are formed between non-metals and involve the sharing of electrons between atoms. The shared electrons interact with the nuclei of both atoms, lowering the potential energy and stabilizing the molecule.

• No transfer of electrons; instead, electrons are shared.

• Results in discrete molecules (not extended lattices).

• Example: In $\mathrm{O}_2$, two oxygen atoms share two pairs of electrons, forming a double bond: $\mathrm{O}=\mathrm{O}$.

Additional info: Covalent bonding is further explored in later chapters, including the concepts of valence electrons and atomic orbitals.

Representing Compounds

Chemical Formulas

• Molecular Formula: Shows the actual number of atoms of each element in a molecule (e.g., $\mathrm{H}_2\mathrm{O}_2$ for hydrogen peroxide).

• Empirical Formula: Shows the simplest whole-number ratio of atoms (e.g., HO for hydrogen peroxide).

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

Categorizing Pure Substances

• Atomic elements: Exist as single atoms (e.g., noble gases like Ne, Ar; metals like Ag, Cu).

• Molecular elements: Exist as molecules of the same element (e.g., $\mathrm{H}_2$, $\mathrm{O}_2$, $\mathrm{N}_2$).

• Molecular compounds: Composed of two or more covalently bonded non-metals (e.g., $\mathrm{H}_2\mathrm{O}$, $\mathrm{CO}_2$).

• Ionic compounds: Composed of cations (usually metals or ammonium) and anions (monoatomic or polyatomic), arranged in a lattice (e.g., $\mathrm{NaCl}$, $\mathrm{MgCl}_2$).

Formulas and Names

Naming Ionic Compounds

• If the metal forms only one possible charge, its name is unchanged (e.g., Potassium chloride: KCl).

• If the metal can form more than one charge, indicate the charge with Roman numerals (e.g., Iron(II) sulfate: FeSO$_4$).

• Non-metal anions are named with the suffix "-ide" (e.g., Oxide, Chloride).

• Polyatomic ions have specific names (e.g., Nitrate: NO$_3^-$, Sulfate: SO$_4^{2-}$).

• Oxyanions with different numbers of oxygens use prefixes and suffixes: "hypo-...ite" (least O), "-ite", "-ate", "per-...ate" (most O).

• Hydrated compounds include the term "hydrate" with a prefix indicating the number of water molecules (e.g., CuSO$_4$·5H$_2$O: Copper(II) sulfate pentahydrate).

Common Polyatomic Ions

Ion	Name
NO$_3^-$	Nitrate
NO$_2^-$	Nitrite
SO$_4^{2-}$	Sulfate
SO$_3^{2-}$	Sulfite
CO$_3^{2-}$	Carbonate
PO$_4^{3-}$	Phosphate
NH$_4^+$	Ammonium
OH$^-$	Hydroxide
ClO$^-$	Hypochlorite
ClO$_2^-$	Chlorite
ClO$_3^-$	Chlorate
ClO$_4^-$	Perchlorate
CH$_3$COO$^-$	Acetate
MnO$_4^-$	Permanganate
HCO$_3^-$	Hydrogen carbonate (bicarbonate)
HSO$_4^-$	Hydrogen sulfate
H$_2$PO$_4^-$	Dihydrogen phosphate

Naming Molecular Compounds

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

• The more metal-like element (lower group/row) is named first.

• "Mono-" is usually omitted for the first element.

• Examples: CO = carbon monoxide, CO$_2$ = carbon dioxide, N$_2$O = dinitrogen monoxide.

Naming Acids

• Binary acids (no oxygen): Use the format "hydro-" + base name + "-ic acid" (e.g., HCl(aq): hydrochloric acid).

• Oxyacids (contain oxygen):

  • If the anion ends in "-ate": acid name ends in "-ic acid" (e.g., H$_2$SO$_4$: sulfuric acid).

  • If the anion ends in "-ite": acid name ends in "-ous acid" (e.g., H$_2$SO$_3$: sulfurous acid).

Organic Compounds

Hydrocarbons

• Alkanes: Only single C–C bonds. General formula: $C_nH_{2n+2}$ (e.g., methane, ethane, propane).

• Alkenes: At least one C=C double bond. General formula: $C_nH_{2n}$ (e.g., ethene, propene).

• Alkynes: At least one C≡C triple bond. General formula: $C_nH_{2n-2}$ (e.g., ethyne, propyne).

• Cyclic hydrocarbons: Ring structures, named with the prefix "cyclo-" (e.g., cyclohexane).

• Aromatic hydrocarbons: Contain delocalized electrons in a ring (e.g., benzene).

Prefixes for chain length: meth- (1), eth- (2), prop- (3), but- (4), pent- (5), hex- (6), hept- (7), oct- (8), non- (9), dec- (10).

Functionalized Hydrocarbons

• Alcohols: –OH group (e.g., ethanol).

• Ethers: –O– linkage (e.g., diethyl ether).

• Aldehydes: –CHO group (e.g., ethanal).

• Ketones: –CO– group (e.g., acetone).

• Carboxylic acids: –COOH group (e.g., acetic acid).

• Esters: –COO– group (e.g., methyl acetate).

Additional info: The position and type of functional group determine the compound's family and its chemical properties.

Formula Mass and the Mole Concept

Formula Mass (Molar Mass)

The formula mass (or molar mass) of a compound is the sum of the atomic masses of all atoms in its formula, expressed in g/mol. It represents the mass of one mole (6.022 × 1023 entities) of the compound.

• Example: For CaCl$_2$, molar mass = atomic mass of Ca + 2 × atomic mass of Cl.

• For Ba$_3$(PO$_4$)$_2$, molar mass = 3 × Ba + 2 × P + 8 × O.

Using Molar Mass

• To convert between mass (g), moles, and number of molecules:

$\text{moles} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

$\text{number of molecules} = \text{moles} \times 6.022 \times 10^{23}$

Composition of Compounds

Mass Percent Composition

The mass percent of an element in a compound is the fraction of the compound's mass contributed by that element, expressed as a percentage.

$\text{Mass percent of element} = \frac{\text{mass of element in 1 mole}}{\text{molar mass of compound}} \times 100\%$

Example: For Cl in CCl$_2$F$_2$:

$\text{Mass percent of Cl} = \frac{2 \times \text{atomic mass of Cl}}{\text{molar mass of CCl}_2\text{F}_2}$

Determining Chemical Formulas

Empirical and Molecular Formulas

1. Calculate the mass percent of each element.

2. Convert mass percent to moles of each element.

3. Divide by the smallest number of moles to get the simplest ratio (empirical formula).

4. If molecular mass is known, divide it by the empirical formula mass to find the multiplier for the molecular formula.

Combustion Analysis

Combustion analysis is used to determine the empirical formula of a compound, especially hydrocarbons and compounds containing C, H, and O.

• Burn a known mass of compound in excess O$_2$.

• Measure the masses of CO$_2$ and H$_2$O produced.

• All C in CO$_2$ and H in H$_2$O come from the original compound.

• If O is present in the compound, use mass conservation to determine its amount.

Additional info: The difference between the mass of the original sample and the sum of C and H masses gives the mass of O in the compound.