Chapter 3: Molecules and Compounds – General Chemistry Study Notes

Chapter 3: Molecules and Compounds

Rocket Fuel: Chemical Reactions and Properties

This section introduces the concept of chemical reactions using rocket fuel as an example, focusing on the reaction between methane and oxygen to produce carbon dioxide and water. It also compares the physical and chemical properties of hydrogen, oxygen, and water.

• Chemical Reaction Example: The combustion of methane ($CH_4$) with oxygen ($O_2$) produces carbon dioxide ($CO_2$) and water ($H_2O$): $CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O$

• Properties Comparison:

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

• Mixtures vs. Compounds: A mixture of hydrogen and oxygen can have any ratio of the two gases, while water (a compound) always has a fixed ratio of two hydrogens to one oxygen.

Chemical Bonds

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

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

Atoms form covalent bonds by sharing electrons, creating a stable electron pair between nuclei. This results in a concentration of electron density between the atoms, lowering potential energy and stabilizing the molecule.

• Example: The $H_2$ 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, molecular, and structural formulae.

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

• Molecular Formula: Actual number of atoms of each element in a molecule (e.g., for hydrogen peroxide: $H_2O_2$).

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

Examples of Formulae

• Hexene: $C_6H_{12}$ (empirical formula: CH$_2$)

• Pentene: $C_5H_{10}$ (empirical formula: CH$_2$)

• Butene: $C_4H_8$ (empirical formula: CH$_2$)

• Methane: $CH_4$ (empirical and molecular formula are the same)

• Structural Formula Example: D-Glucose, showing wedge bonds to indicate 3D structure.

The number of possible compounds increases with the number of atoms in the molecular formula (e.g., $C_6H_{12}O_6$ 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 understanding molecular geometry.

• Molecular Formula: Shows the number and type 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.

Name of Compound	Empirical Formula	Molecular Formula	Structural Formula	Ball-and-Stick Model	Space-Filling Model
Benzene	CH	C$_6$H$_6$	Ring structure	Ball-and-stick diagram	Space-filling diagram
Acetylene	CH	C$_2$H$_2$	Linear structure	Ball-and-stick diagram	Space-filling diagram
Glucose	CH$_2$O	C$_6$H$_{12}$O$_6$	Ring structure	Ball-and-stick diagram	Space-filling diagram
Ammonia	NH$_3$	NH$_3$	Trigonal pyramidal	Ball-and-stick diagram	Space-filling diagram

Compounds

Molecular Compounds

Molecular compounds are typically composed of two or more covalently bonded nonmetals. The basic unit is the molecule, which consists of the constituent atoms bonded together.

• Examples: Water ($H_2O$), dry ice ($CO_2$), propane ($C_3H_8$).

Ionic Compounds

Ionic compounds are composed of cations (usually metals) and anions (usually nonmetals) held together by ionic bonds. The basic unit is the formula unit, the smallest electrically neutral group of ions.

• 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: Use Roman numerals to indicate charge (e.g., Iron(II): $Fe^{2+}$, Iron(III): $Fe^{3+}$).

• -ous and -ic: -ous for lower charge, -ic for higher charge (e.g., Ferrous ion: $Fe^{2+}$, Ferric ion: $Fe^{3+}$).

• Nonmetal Cations: End in -ium (e.g., Ammonium: $NH_4^+$, Hydronium: $H_3O^+$).

• Anions: Negative ions. Monatomic anions end in -ide (e.g., Chloride: $Cl^-$, Sulfide: $S^{2-}$, Oxide: $O^{2-}$).

• Polyatomic Anions: With oxygen, end in -ate or -ite (e.g., Nitrate: $NO_3^-$, Nitrite: $NO_2^-$, Sulfate: $SO_4^{2-}$, Sulfite: $SO_3^{2-}$).

• Formula Writing Rules: Cation is listed first, polyatomic ions are written as units, and parentheses are used for multiples (e.g., $Ba(OH)_2$).

• Recognizing Ionic Compounds: Contains a metal from Group I or II, or a polyatomic ion.

Naming Acids

Acids are named based on the anion they contain:

• -ide anion: Add 'hydro-' prefix and '-ic acid' suffix (e.g., Chloride $Cl^-$ → Hydrochloric acid $HCl$).

• -ate anion: Add '-ic acid' suffix (e.g., Chlorate $ClO_3^-$ → Chloric acid $HClO_3$).

• -ite anion: Add '-ous acid' suffix (e.g., Chlorite $ClO_2^-$ → Chlorous acid $HClO_2$).

Naming Molecular (Binary) Compounds

Binary molecular compounds are formed between two nonmetals. The naming system uses prefixes to indicate the number of atoms of each element.

• 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 Adjustment: If a prefix ends in 'a' or 'o' and the element name begins with a vowel, drop the 'a' or 'o'.

Calculating Empirical Formulae

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

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

2. Express as a percentage of total mass.

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

4. Calculate moles of each element.

5. Divide 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/weight) is the sum of the atomic masses of all atoms in a molecule or formula unit.

• Example: For water ($H_2O$): $\text{Formula mass} = (2 \times 1.01\,\text{amu}) + (1 \times 16.00\,\text{amu}) = 18.02\,\text{amu}$

Calculating Molar Mass

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

• Example: For water ($H_2O$): $\text{Molar mass} = 2 \times 1.01\,\text{g/mol} + 16.00\,\text{g/mol} = 18.02\,\text{g/mol}$

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.

• 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 $CO_2$: $\text{Moles of } CO_2 = \frac{9\,\text{g}}{44.01\,\text{g/mol}}$ $\text{Number of molecules} = \text{moles} \times 6.022 \times 10^{23}$

Additional info: Some context and examples have been expanded for clarity and completeness, including stepwise procedures and formula calculations.