Study Guide: Chemical Bonding, Nomenclature, and Chemical Calculations

Chapter Six: Chemical Bonding and Nomenclature

Drawing Lewis Dot Structures and Elements in Periods 4 through VIIIA

• Lewis Dot Structures are diagrams that represent the valence electrons of atoms within a molecule. Each dot represents a valence electron, and the arrangement helps predict bonding patterns and molecular structure.

• Elements in Periods 4 through VIIIA refer to elements in the fourth period and groups 4A to 8A (14-18) of the periodic table, which include many main group elements important in GOB Chemistry.

Octet Rule and Formal Charges

• The Octet Rule states that atoms tend to form compounds in ways that give them eight valence electrons, achieving a noble gas configuration.

• Formal Charge is a bookkeeping tool used to determine the most likely arrangement of atoms in a molecule. It is calculated as: $\text{Formal Charge} = \text{Valence Electrons} - (\text{Lone Pair Electrons} + \frac{1}{2} \times \text{Bonding Electrons})$

• Assigning formal charges helps identify the most stable Lewis structure.

Ionic Compounds and Polyatomic Ions

• Ionic Compounds are formed from the electrostatic attraction between cations (positively charged ions) and anions (negatively charged ions).

• Common cations include: $\text{Na}^+$, $\text{K}^+$, $\text{Ca}^{2+}$, $\text{Fe}^{2+}$, $\text{Fe}^{3+}$, $\text{NH}_4^+$.

• Common anions include: $\text{Cl}^-$, $\text{O}^{2-}$, $\text{SO}_4^{2-}$, $\text{NO}_3^-$, $\text{PO}_4^{3-}$.

• Polyatomic Ions are ions composed of more than one atom, such as $\text{SO}_4^{2-}$ (sulfate) and $\text{NH}_4^+$ (ammonium).

Naming Ionic and Covalent Compounds

• Ionic Compounds: Name the cation first, then the anion. For transition metals, indicate the charge with Roman numerals (e.g., iron(III) chloride for $\text{FeCl}_3$).

• Covalent Compounds: Use prefixes to indicate the number of each type of atom (e.g., carbon dioxide for $\text{CO}_2$).

• Common prefixes: mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-.

Bonding Types and Electronegativity

• Ionic Bonds occur between metals and nonmetals, involving the transfer of electrons.

• Covalent Bonds occur between nonmetals, involving the sharing of electrons.

• Electronegativity is the ability of an atom to attract electrons in a bond. The difference in electronegativity determines bond type:

  • Nonpolar covalent: small or no difference (e.g., $\text{Cl}_2$)

  • Polar covalent: moderate difference (e.g., $\text{H}_2\text{O}$)

  • Ionic: large difference (e.g., $\text{NaCl}$)

Lewis Structures and Molecular Geometry

• Lewis structures help determine the arrangement of atoms and lone pairs in a molecule.

• VSEPR Theory (Valence Shell Electron Pair Repulsion) is used to predict molecular geometry based on electron pair repulsions.

• Common geometries: linear, bent, trigonal planar, tetrahedral, trigonal pyramidal.

Polarity and Intermolecular Forces

• Molecular Polarity depends on both bond polarity and molecular geometry.

• Polar molecules have an uneven distribution of charge (e.g., $\text{H}_2\text{O}$), while nonpolar molecules have an even distribution (e.g., $\text{CO}_2$).

• Intermolecular Forces include hydrogen bonding, dipole-dipole interactions, and London dispersion forces.

Examples of Naming and Writing Formulas

• Write the formula for sodium chloride: $\text{NaCl}$

• Name $\text{NH}_4\text{NO}_3$: ammonium nitrate

• Write the formula for calcium phosphate: $\text{Ca}_3(\text{PO}_4)_2$

Chapter Seven: Chemical Calculations

Determining Moles and Molar Mass

• Mole: The amount of substance containing $6.02 \times 10^{23}$ entities (Avogadro's number).

• Molar Mass: The mass of one mole of a substance, expressed in grams per mole (g/mol).

• To calculate moles: $\text{Moles} = \frac{\text{Mass (g)}}{\text{Molar Mass (g/mol)}}$

Empirical and Molecular Formulas

• Empirical Formula: The simplest whole-number ratio of atoms in a compound.

• Molecular Formula: The actual number of atoms of each element in a molecule.

• To determine the empirical formula, divide the number of moles of each element by the smallest number of moles present.

Balancing Chemical Equations

• Law of Conservation of Mass: Atoms are neither created nor destroyed in a chemical reaction.

• To balance an equation, ensure the same number of each type of atom on both sides.

• Example: $\text{CH}_4 + 2\text{O}_2 \rightarrow \text{CO}_2 + 2\text{H}_2\text{O}$

Stoichiometry

• Stoichiometry involves using balanced chemical equations to calculate the amounts of reactants and products.

• Use mole ratios from the balanced equation to convert between substances.

• Example: If 2 moles of $\text{H}_2$ react with 1 mole of $\text{O}_2$, how many moles of $\text{H}_2\text{O}$ are produced? Answer: 2 moles.

Gas Laws

• Pressure (P): The force exerted by gas particles per unit area. Measured in atmospheres (atm), mmHg, or kPa.

• Boyle's Law: $P_1V_1 = P_2V_2$ (at constant temperature)

• Charles' Law: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$ (at constant pressure)

• Combined Gas Law: $\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$

• Ideal Gas Law: $PV = nRT$ Where $P$ = pressure, $V$ = volume, $n$ = moles, $R$ = gas constant ($0.0821 \frac{L \cdot atm}{mol \cdot K}$), $T$ = temperature in Kelvin.

• Dalton's Law of Partial Pressures: $P_{\text{total}} = P_1 + P_2 + P_3 + \ldots$

• Standard Temperature and Pressure (STP): 1 atm and 0°C (273 K).

Table: Common Polyatomic Ions

Name	Formula	Charge
Ammonium	NH4+	+1
Sulfate	SO42-	-2
Nitrate	NO3-	-1
Phosphate	PO43-	-3
Hydroxide	OH-	-1
Bicarbonate	HCO3-	-1
Acetate	CH3COO-	-1
Carbonate	CO32-	-2
Additional info: Other polyatomic ions may be included as needed for course coverage.

Constants and Reference Values

• Avogadro's Number: $6.02 \times 10^{23}$

• Gas Constant (R): $0.0821 \frac{L \cdot atm}{mol \cdot K}$

• STP: 1 atm and 0°C (273 K)