Chapter 3: Molecules and Compounds – Study Notes

Chapter 3: Molecules and Compounds

Rocket Fuel and Chemical Reactions

This section introduces the concept of chemical reactions using rocket fuel as an example, highlighting the transformation of reactants into products and the properties of the substances involved.

• Example Reaction: Combustion of methane with oxygen produces carbon dioxide and water.

• Equation: $\mathrm{CH_4 + 2\ O_2 \rightarrow CO_2 + 2\ H_2O}$

• Properties Table:

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

Mixtures vs. Compounds:

• Mixture: Hydrogen and oxygen gases can be mixed in any ratio.

• Compound: Water molecules have a fixed ratio of hydrogen to oxygen (2:1).

Chemical Bonds

Types of Chemical Bonds

Chemical bonds are the forces that hold atoms together in compounds. There are two main types:

• Ionic Bond: Attraction between a cation (positive ion) and an anion (negative ion).

• Covalent Bond: Sharing of a pair of electrons, with one electron contributed by each atom.

• Bonding Continuum: Ionic and covalent bonds represent extremes; many bonds have characteristics in between.

Ionic Bonding in Solids

In solid ionic compounds, ions of opposite charge are arranged in a three-dimensional lattice structure, maximizing attractive forces and minimizing repulsion.

• Example: Sodium chloride (NaCl) forms a crystalline lattice of Na+ and Cl- ions.

Covalent Bonds Share Electrons

Covalent bonds form when atoms share electrons to achieve a stable electron configuration, often resulting in the lowest potential energy for the system.

• Bonding Pair: Each atom contributes one electron to the shared pair.

• Stability: Covalent bonds are energetically favorable due to electron density between nuclei.

Molecular Formulae

Types of Chemical Formulas

Chemical formulas provide information about the composition and structure of compounds.

• Empirical Formula: Simplest whole-number ratio of atoms in a compound (e.g., HO for hydrogen peroxide).

• Molecular Formula: Actual number of atoms of each element in a molecule (e.g., H2O2 for hydrogen peroxide).

• Structural Formula: Shows how atoms are bonded together, including connectivity and sometimes geometry.

Information Hierarchy: Empirical < Molecular < Structural

Examples of Formulas

• Hexene: C6H12 (empirical: CH2)

• Pentene: C5H10 (empirical: CH2)

• Butene: C4H8 (empirical: CH2)

• Methane: CH4 (empirical and molecular formula are the same)

• Structural Formula Example: D-Glucose shows connectivity and stereochemistry using wedge bonds.

Number of Possible Compounds: The more atoms in a molecular formula, the greater the number of possible isomers and compounds (e.g., C6H12O6 has over 1,500 possible compounds).

Molecular Models of Structure

Different models are used to visualize molecules for various purposes, such as drug discovery or teaching.

• Molecular Formula: Shows composition only.

• Structural Formula: Shows connectivity.

• Ball-and-Stick Model: Shows 3D arrangement and bonds.

• Space-Filling Model: Shows realistic atom sizes and spatial relationships.

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 structure	Ball-and-stick	Space-filling
Glucose	C6H12O6	C6H12O6	Ring structure	Ball-and-stick	Space-filling
Ammonia	NH3	NH3	Trigonal pyramidal	Ball-and-stick	Space-filling

Compounds

Molecular Compounds

Molecular compounds consist of two or more nonmetals bonded covalently. Their basic units are molecules.

• Examples: Water (H2O), dry ice (CO2), propane (C3H8).

Ionic Compounds

Ionic compounds are formed from cations (usually metals) and anions (usually nonmetals) held together by ionic bonds. The basic unit is the formula unit.

• Example: Table salt (NaCl) consists of Na+ and Cl- ions in a 1:1 ratio.

Naming Ionic Compounds

Cations and Anions

Cations are positively charged ions, typically formed by metals. Anions are negatively charged ions, typically formed by nonmetals.

• General Rules:

  • Metals form cations.

  • Halides, oxygen, and sulfur form anions.

  • Polyatomic ions contain more than one atom.

  • Oxyanions contain one or more oxygen atoms.

Rules for Writing Formulas

• Cation is listed before anion.

• Polyatomic ions are written as a unit; if more than one is present, use parentheses and a subscript.

• Recognizing ionic compounds: Contains a metal from Group I or II, or a polyatomic ion.

• Examples: NaOH (sodium hydroxide), Ba(OH)2 (barium hydroxide).

Visualizing Ionic Compounds

• Formation involves electron transfer: Na atom loses an electron to become Na+; Cl atom gains an electron to become Cl-.

Naming Cations

• Monatomic cations (one atom) use the element name: e.g., lithium ion (Li+), calcium ion (Ca2+).

• Cations with multiple charges use Roman numerals: e.g., iron(II) (Fe2+), iron(III) (Fe3+).

• Alternative naming uses -ous and -ic: ferrous (Fe2+), ferric (Fe3+).

• Polyatomic cations of nonmetals end in -ium: ammonium (NH4+), hydronium (H3O+).

Common Cations Table

Charge	Formula	Name
1+	H+	hydrogen ion
1+	Na+	sodium ion
2+	Mg2+	magnesium ion
2+	Ca2+	calcium ion
3+	Al3+	aluminum ion
1+	NH4+	ammonium ion
1+	H3O+	hydronium ion
2+	Fe2+	iron(II) ion (ferrous)
3+	Fe3+	iron(III) ion (ferric)

Naming Anions

• Monatomic anions end in -ide: chloride (Cl-), sulfide (S2-), oxide (O2-).

• Some polyatomic anions also end in -ide: hydroxide (OH-), peroxide (O22-).

• Oxyanions (contain oxygen) end in -ate or -ite: nitrate (NO3-), nitrite (NO2-), sulfate (SO42-), sulfite (SO32-).

Common Anions Table

Charge	Formula	Name
1-	F-	fluoride ion
1-	Cl-	chloride ion
1-	Br-	bromide ion
2-	O2-	oxide ion
2-	S2-	sulfide ion
1-	OH-	hydroxide ion
1-	NO3-	nitrate ion
2-	CO32-	carbonate ion
2-	SO42-	sulfate ion
3-	PO43-	phosphate ion

Charges on Elemental Ions

The charge of ions formed by elements can be predicted from their position in the periodic table. Group I metals form 1+ ions, Group II metals form 2+ ions, and nonmetals typically form negative ions.

Naming Acids

Rules for Naming Acids

Acids are named based on the anion they contain:

• If the anion ends in -ide, the acid is named as hydro-___-ic acid (e.g., chloride → hydrochloric acid, HCl).

• If the anion ends in -ate, the acid is named as ___-ic acid (e.g., chlorate → chloric acid, HClO3).

• If the anion ends in -ite, the acid is named as ___-ous acid (e.g., chlorite → chlorous acid, HClO2).

Naming Molecular Compounds

Rules for Naming Binary Molecular Compounds

Molecular compounds (binary compounds) are formed between two nonmetals. Their names reflect the ratio and identity of the elements.

• The element farther left in the periodic table is named first.

• If both elements are in the same group, the lower element is named first.

• The second element is named with an -ide ending.

• Prefixes indicate the number of atoms of each element (never use 'mono' for the first element).

Prefixes Table

Prefix	Meaning
Mono-	1
Di-	2
Tri-	3
Tetra-	4
Penta-	5
Hexa-	6
Hepta-	7
Octa-	8
Nona-	9
Deca-	10

Note: If the prefix ends in 'a' or 'o' and the element name begins with a vowel, the 'a' or 'o' is dropped.

Inorganic Naming Flowchart

A flowchart can help determine the correct naming convention for inorganic compounds, distinguishing between ionic, molecular, and acid nomenclature.

Calculating Empirical Formulae

Steps to Find an Empirical Formula

The empirical formula represents the simplest whole-number ratio of atoms in a compound. To determine it experimentally:

• Measure the mass of each element in the compound.

• Express each mass as a percentage of the total mass.

• Assume a 100 g sample for ease of calculation.

• Calculate the number of moles of each element.

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

• Multiply to obtain whole numbers if necessary.

Calculating Formula Mass

Formula Mass Calculation

The formula mass (or molecular mass/weight) is the sum of the atomic masses of all atoms in a molecule or formula unit.

• Example: Mass of 1 molecule of H2O:

$\text{Formula mass of H}_2\text{O} = (2 \times 1.01\ \text{amu}) + (1 \times 16.00\ \text{amu}) = 18.02\ \text{amu}$

Calculating Molar Mass

Molar Mass Calculation

The molar mass is the mass in grams of one mole of a substance, numerically equal to the formula mass in amu but expressed in g/mol.

• Example: 1 mole of H2O contains 2 moles of H and 1 mole of O.

$\text{Molar mass of H}_2\text{O} = 2 \times 1.01\ \text{g/mol} + 1 \times 16.00\ \text{g/mol} = 18.02\ \text{g/mol}$

Molar Mass and Counting Molecules

Molar mass and Avogadro's number ($6.022 \times 10^{23}$) are used to convert between mass, moles, and number of molecules.

• Convert mass to moles using molar mass.

• Convert moles to number of molecules using Avogadro's number.

$\text{Number of molecules} = \frac{\text{grams of A}}{\text{molar mass of A}} \times 6.022 \times 10^{23}$

Additional info: These notes cover all major aspects of Chapter 3: Molecules and Compounds, including chemical bonding, formula types, nomenclature, and calculations relevant to general chemistry.