Chapter 3: Molecules and Compounds – General Chemistry Study Notes

Chapter 3: Molecules and Compounds

Rocket Fuel and Chemical Reactions

This section introduces the concept of chemical reactions using the example of rocket fuel combustion. It highlights the transformation of reactants into products and compares the properties of substances involved.

• Chemical Reaction Example: Combustion of methane (CH4) with oxygen (O2) produces carbon dioxide (CO2) and water (H2O).

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

• Properties Comparison:

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: Mixtures (e.g., hydrogen and oxygen gases) can have any ratio of components, while compounds (e.g., water) have a fixed ratio of atoms (2 H : 1 O).

Chemical Bonds

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

• Ionic Bond: Attraction between a cation (positively charged ion) and an anion (negatively charged ion).

• Covalent Bond: Sharing of a pair of electrons between atoms, 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 ionic solids, 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 and Electron Sharing

Covalent bonds form when atoms share electrons, creating a stable electron pair between nuclei. This results in a concentration of electron density between the atoms.

• Example: The H2 molecule forms when two hydrogen atoms share electrons.

Molecular Formulae

Molecular formulae provide information about the composition and structure of compounds. There are three main types:

• 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 the arrangement of atoms and bonds (e.g., H–O–O–H for hydrogen peroxide).

Examples of Formula Types

• Hexene: Empirical formula CH2, molecular formula C6H12, structural formula CH3CH2CH2CH=CHCH3

• Pentene: Empirical formula CH2, molecular formula C5H10, structural formula CH3CH2CH=CHCH3

• Butene: Empirical formula CH2, molecular formula C4H8, structural formula CH3CH=CHCH3

• Methane: Empirical and molecular formula are both CH4.

• Structural Formula with Wedge Bonds: D-Glucose is often depicted with wedge and dash bonds to show 3D orientation.

Additional info: The number of possible compounds increases with the number of atoms in the molecular formula (e.g., C6H12O6 has over 1,500 isomers).

Molecular Models of Structure

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

• Molecular Formula: Shows the types and numbers of atoms.

• Structural Formula: Shows connectivity of atoms.

• Ball-and-Stick Model: Represents atoms as balls and bonds as sticks, useful for visualizing geometry.

• Space-Filling Model: Shows the relative sizes of atoms and how they fill space.

Comparison Table: Different Views for Different Needs

Name of Compound	Empirical Formula	Molecular Formula	Structural Formula	Ball-and-Stick Model	Space-Filling Model
Benzene	CH	C6H6	Ring structure	Ball-and-stick image	Space-filling image
Acetylene	CH	C2H2	Linear structure	Ball-and-stick image	Space-filling image
Glucose	CH2O	C6H12O6	Ring structure	Ball-and-stick image	Space-filling image
Ammonia	NH3	NH3	Trigonal pyramidal	Ball-and-stick image	Space-filling image

Additional info: Ball-and-stick and space-filling models are especially useful for visualizing molecular geometry and atomic sizes.

Compounds

Compounds are substances formed from two or more elements chemically bonded in fixed ratios. They can be classified as molecular or ionic compounds.

• Molecular Compounds: Composed of covalently bonded nonmetals. The basic unit is the molecule (e.g., H2O, CO2, C3H8).

• Ionic Compounds: Composed of cations (usually metals) and anions (usually nonmetals) held together by ionic bonds. The basic unit is the formula unit (e.g., NaCl).

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

Naming Ionic Compounds

Naming conventions for ionic compounds depend on the types of ions present.

• Cations: Positive ions, usually metals. Named after the element (e.g., Lithium ion, Li+).

• Multiple Charges: Some metals form cations with different charges; use Roman numerals (e.g., Iron(II), Fe2+; Iron(III), Fe3+).

• -ous and -ic: Alternative naming for cations with multiple charges (-ous for lower, -ic for higher charge).

• Nonmetal Cations: End in -ium (e.g., ammonium NH4+, hydronium H3O+).

• Anions: Negative ions. Monatomic anions end in -ide (e.g., chloride Cl-, sulfide S2-).

• Polyatomic Anions: With oxygen, end in -ate or -ite (e.g., nitrate NO3-, nitrite NO2-).

Common Anions Table

Ion	Name
Cl-	Chloride ion
O2-	Oxide ion
NO3-	Nitrate ion
SO42-	Sulfate ion
PO43-	Phosphate ion
OH-	Hydroxide ion

Additional info: Polyatomic ions are ions composed of more than one atom (e.g., NH4+, SO42-).

Naming Acids

Acids are named based on the anion they contain:

• -ide anion: Add 'hydro-' prefix and '-ic acid' suffix (e.g., chloride → hydrochloric acid, HCl).

• -ate anion: Add '-ic acid' suffix (e.g., chlorate → chloric acid, HClO3).

• -ite anion: Add '-ous acid' suffix (e.g., chlorite → chlorous acid, HClO2).

Naming Molecular (Binary) Compounds

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

• Order: Element farther left in the periodic table is named first; if in the same group, the lower element is named first.

• Second Element: Name ends in '-ide'.

• Prefixes: Indicate the number of atoms (mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca-). 'Mono-' is never used for the first element.

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

Additional info: When a prefix ends in 'a' or 'o' and the element name begins with a vowel, the final vowel of the prefix is dropped.

Calculating Empirical Formulae

The empirical formula represents the simplest whole-number ratio of atoms in a compound. It can be determined from experimental data.

1. Measure the mass of each element in the compound.

2. Express as a percentage of the total mass.

3. Assume a 100 g sample for ease of calculation.

4. Calculate the number of moles of each element.

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

6. Multiply to obtain whole numbers if necessary.

Calculating Formula Mass

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

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

• Example: For H2O: $\text{Formula mass} = (2 \times 1.01\ \text{amu}) + (1 \times 16.00\ \text{amu}) = 18.02\ \text{amu}$

Calculating Molar Mass

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

• Formula: $\text{Molar mass} = \sum \text{(number of moles of each element)} \times \text{molar mass of element}$

• Example: For H2O: $\text{Molar mass} = 2 \times 1.01\ \text{g/mol} + 1 \times 16.00\ \text{g/mol} = 18.02\ \text{g/mol}$

Molar Mass and Counting Molecules

Molar mass and Avogadro's number are used to relate mass, moles, and number of molecules.

• Avogadro's Number: $6.022 \times 10^{23}$ molecules/mol

• Conversion Steps:

  1. Use molar mass to convert grams to moles.

  2. Use Avogadro's number to convert moles to number of molecules.

• Example: To find molecules in 9 g of CO2:

$\text{Moles of CO}_2 = \frac{9\ \text{g}}{44.01\ \text{g/mol}}$ $\text{Number of molecules} = \text{moles} \times 6.022 \times 10^{23}$

Additional info: These conversions are fundamental for quantitative chemical analysis and stoichiometry.