CHEM 1120 Exam III Study Guide: Chemical Bonding, Intermolecular Forces, Organic Chemistry, Solutions, and Acids & Bases

Chemical Bonding

Octet Rule and Its Exceptions

The octet rule states that atoms tend to gain, lose, or share electrons to achieve a stable configuration of eight valence electrons, similar to the noble gases.

• Definition: Atoms are most stable when surrounded by eight electrons in their valence shell.

• Exceptions: Hydrogen (2 electrons), Helium (2 electrons), Beryllium (4 electrons), Boron (6 electrons), and elements in period 3 and beyond can have expanded octets.

• Example: BF3 (boron trifluoride) has only 6 electrons around boron.

Lewis Diagrams for Atoms and Molecules

• Lewis diagrams (or Lewis structures) represent valence electrons as dots around atomic symbols.

• First and second row atoms: Place dots to represent valence electrons, pairing them as needed.

• For molecules, connect atoms with lines (bonds) and show lone pairs as dots.

• Example: The Lewis structure for water (H2O) shows two bonding pairs and two lone pairs on oxygen.

Bonds Sought by Second Period Nonmetals

• Second period nonmetals "seek" enough bonds to complete their octet:

  • Carbon (C): 4 bonds

  • Nitrogen (N): 3 bonds

  • Oxygen (O): 2 bonds

  • Fluorine (F): 1 bond

Lone and Bonding Electron Pairs

• In Lewis structures, bonding pairs are shared between atoms, while lone pairs are non-bonding electrons localized on a single atom.

• Count the number of each to predict molecular shape and reactivity.

• Example: Ammonia (NH3) has three bonding pairs and one lone pair on nitrogen.

Steric Number and Molecular Geometry

• The steric number is the sum of bonded atoms and lone pairs on the central atom.

• Use the steric number to determine electron geometry (e.g., linear, trigonal planar, tetrahedral).

• Molecular geometry depends on the arrangement of atoms (not lone pairs).

• Example: Methane (CH4) has a steric number of 4 (tetrahedral geometry).

Electronegativity and Periodic Trends

• Electronegativity is the ability of an atom to attract shared electrons in a bond.

• Trends: Increases across a period (left to right), decreases down a group (top to bottom).

• Example: Fluorine is the most electronegative element.

Molecular Polarity

• A molecule is polar if it has an uneven distribution of electron density, resulting in a dipole moment.

• Polarity depends on both bond polarity (difference in electronegativity) and molecular geometry.

• Example: Carbon dioxide (CO2) has polar bonds but is nonpolar overall due to its linear shape.

Intermolecular Forces

Types of Intermolecular Forces

Intermolecular forces are attractions between molecules that influence physical properties like boiling and melting points.

• Dispersion Forces (London Forces): Weakest; present in all molecules due to temporary dipoles.

• Dipole-Dipole Forces: Occur between polar molecules with permanent dipoles.

• Hydrogen Bonds: Strongest; occur when H is bonded to N, O, or F and interacts with lone pairs on these atoms in other molecules.

Identifying Intermolecular Forces

• Check for H bonded to N, O, or F for hydrogen bonding.

• Polar molecules exhibit dipole-dipole forces.

• All molecules have dispersion forces.

• Example: Water exhibits all three types; methane only has dispersion forces.

Intermolecular Forces Between Molecules

• Identical molecules: Analyze their structure to determine possible forces.

• Different molecules: Consider both structures to identify possible interactions.

• Example: Ethanol and water can form hydrogen bonds with each other.

Organic Chemistry

Hydrocarbons: Alkanes, Alkenes, Alkynes

• Alkanes: Saturated hydrocarbons with only single bonds (general formula: CnH2n+2).

• Alkenes: Unsaturated hydrocarbons with at least one double bond (CnH2n).

• Alkynes: Unsaturated hydrocarbons with at least one triple bond (CnH2n-2).

• Naming: Use prefixes (meth-, eth-, prop-, etc.) and suffixes (-ane, -ene, -yne) to indicate structure.

• Example: Ethene (C2H4) is an alkene; ethyne (C2H2) is an alkyne.

Functional Groups

Functional groups are specific groups of atoms within molecules that determine chemical reactivity.

Functional Group	Structure	Example
Alcohol	-OH	Ethanol
Ether	-O-	Diethyl ether
Amine	-NH2	Methylamine
Ketone	R-CO-R'	Acetone
Carboxylic Acid	-COOH	Acetic acid
Ester	-COOR	Ethyl acetate
Aldehyde	R-CHO	Formaldehyde

Solutions

Mass and Molar Quantity Conversions

• Use molar mass (g/mol) to convert between mass and moles:

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

• Example: 18 g of H2O is 1 mole (molar mass = 18 g/mol).

Solute vs. Solvent

• Solute: Substance dissolved in a solution (present in lesser amount).

• Solvent: Substance that dissolves the solute (present in greater amount).

• Example: In saltwater, NaCl is the solute, water is the solvent.

Mass Percent Calculations

• Mass percent of solute or solvent:

$\text{Mass percent} = \frac{\text{mass of component}}{\text{total mass of solution}} \times 100\%$

• Example: 10 g NaCl in 90 g water: mass percent NaCl = $\frac{10}{100} \times 100\% = 10\%$

Molarity Calculations

• Molarity (M): Moles of solute per liter of solution.

$M = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Example: 0.5 moles NaCl in 1 L solution = 0.5 M

Moles of Solute from Volume and Concentration

• Calculate moles using:

$\text{moles} = M \times \text{volume (L)}$

• Example: 2.0 L of 0.1 M HCl contains 0.2 moles HCl.

Dilution Calculations

• Use the dilution equation:

$M_1 V_1 = M_2 V_2$

• Where $M_1$ and $V_1$ are the initial concentration and volume, $M_2$ and $V_2$ are the final concentration and volume.

• Example: To make 100 mL of 0.5 M solution from 1.0 M stock: $V_1 = \frac{M_2 V_2}{M_1} = \frac{0.5 \times 0.100}{1.0} = 0.050$ L = 50 mL.

Acids and Bases

Definitions and Behavior in Aqueous Solution

• Arrhenius Acid: Produces H+ ions in water.

• Arrhenius Base: Produces OH- ions in water.

• Bronsted-Lowry Acid: Proton (H+) donor.

• Bronsted-Lowry Base: Proton (H+) acceptor.

• Example: HCl is an acid (donates H+), NH3 is a base (accepts H+).

pH Calculations

• pH is a measure of hydrogen ion concentration:

$\text{pH} = -\log_{10}[H^+]$

• For hydroxide ion concentration:

$\text{pOH} = -\log_{10}[OH^-]$

$\text{pH} + \text{pOH} = 14$

• Example: [H+] = $1 \times 10^{-3}$ M, pH = 3.

Titration Calculations

• Used to determine the concentration of an acid or base by reacting it with a solution of known concentration.

• At the equivalence point, moles of acid = moles of base.

$M_a V_a = M_b V_b$

• Where $M_a$ and $V_a$ are the molarity and volume of acid, $M_b$ and $V_b$ are those of the base.

• Example: 25.0 mL of 0.10 M NaOH neutralizes 50.0 mL of acid. Acid concentration = $\frac{0.10 \times 25.0}{50.0} = 0.050$ M.