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

Chemical Bonding

Octet Rule and Its Exceptions

The octet rule states that atoms tend to gain, lose, or share electrons to achieve a full set of eight valence electrons, resembling the electron configuration of noble gases. However, some elements have exceptions to this rule.

• Octet Rule: Atoms (especially main group elements) seek eight electrons in their valence shell.

• Exceptions: Hydrogen (2 electrons), Helium (2 electrons), Boron (often 6 electrons), and elements in period 3 or higher can have expanded octets.

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

Lewis Diagrams for First and Second Row Atoms

Lewis diagrams represent valence electrons as dots around the atomic symbol. For first and second row atoms, these diagrams help visualize bonding and lone pairs.

• Drawing Lewis Diagrams: Place dots for each valence electron around the element symbol.

• Example: Oxygen (O) has 6 valence electrons: two pairs and two single electrons.

Bonds Sought by Second Period Non-Metals

Second period non-metals (C, N, O, F) seek a specific number of bonds to satisfy the octet rule.

• Carbon: 4 bonds

• Nitrogen: 3 bonds

• Oxygen: 2 bonds

• Fluorine: 1 bond

• Example: Water (H2O)—oxygen forms 2 bonds.

Lewis Diagrams for Simple Molecules: Lone and Bonding Pairs

Lewis diagrams show both bonding pairs (shared electrons) and lone pairs (unshared electrons).

• Bonding Pairs: Electrons shared between atoms.

• Lone Pairs: Electrons not involved in bonding.

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

Steric Number, Electron and Molecular Geometries

The steric number is the sum of bonded atoms and lone pairs around a central atom. It determines the molecule's geometry.

• Steric Number: Number of atoms bonded + number of lone pairs.

• Electron Geometry: Arrangement of electron groups (bonding and lone pairs).

• Molecular Geometry: Arrangement of atoms (ignoring lone pairs).

• Example: Methane (CH4)—steric number 4, tetrahedral geometry.

Electronegativity and Periodic Trends

Electronegativity is the ability of an atom to attract electrons in a bond. It increases across a period and decreases down a group.

• Trend: Highest in top right (fluorine), lowest in bottom left (francium).

• Example: Oxygen is more electronegative than carbon.

Polarity of Molecules

A molecule is polar if it has an uneven distribution of charge due to differences in electronegativity and molecular shape.

• Determining Polarity: Draw Lewis diagram, check for symmetry and electronegativity differences.

• Example: Water is polar; carbon dioxide is non-polar.

Intermolecular Forces

Types of Intermolecular Forces

Intermolecular forces are attractions between molecules, affecting physical properties like boiling and melting points.

• Dispersion Forces (London): Present in all molecules, strongest in large, nonpolar molecules.

• Dipole-Dipole Forces: Occur between polar molecules.

• Hydrogen Bonds: Strongest; occur when H is bonded to N, O, or F.

• Example: Water exhibits hydrogen bonding.

Identifying Types of Intermolecular Bonds

• Hydrogen Bond: Requires H bonded to N, O, or F.

• Dipole-Dipole: Requires polar molecules.

• Dispersion: Present in all molecules.

• Example: Methane (CH4)—only dispersion forces.

Intermolecular Bonds Between Molecules

• Identical Molecules: Same type of intermolecular force as the molecule's structure allows.

• Different Molecules: The strongest force present in either molecule will dominate.

• Example: Mixing water and ethanol—hydrogen bonding occurs.

Organic Chemistry

Hydrocarbons: Alkanes, Alkenes, Alkynes

Hydrocarbons are compounds made of carbon and hydrogen. They are classified by the types of bonds between carbon atoms.

• Alkanes: Only single bonds (saturated hydrocarbons).

• Alkenes: At least one double bond.

• Alkynes: At least one triple bond.

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

• Example: Ethene (C2H4) is an alkene.

Functional Groups

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

• Alcohol: -OH group

• Ether: R-O-R'

• Amine: -NH2 group

• Ketone: C=O within carbon chain

• Carboxylic Acid: -COOH group

• Ester: -COOR group

• Aldehyde: C=O at end of carbon chain

• Example: Acetone is a ketone; ethanol is an alcohol.

Solutions

Mass and Molar Quantity Conversion

To convert between mass and moles, use the molar mass of the substance.

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

• Example: 18 g of water ($\text{H}_2\text{O}$) is 1 mole.

Solute vs. Solvent

A solution consists of a solute (substance dissolved) and a solvent (substance doing the dissolving).

• Solute: Present in lesser amount.

• Solvent: Present in greater amount.

• Example: Salt water—salt is solute, water is solvent.

Mass Percent Calculation

Mass percent expresses the concentration of a component in a solution.

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

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

Molarity Calculation

Molarity (M) is the number of moles of solute per liter of solution.

• Formula: $\text{Molarity} = \frac{\text{moles of solute}}{\text{volume of solution (L)}}$

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

Moles of Solute from Volume and Concentration

• Formula: $\text{moles} = \text{molarity} \times \text{volume (L)}$

• Example: 2.0 L of 0.1 M solution: $0.1 \times 2.0 = 0.2$ moles

Dilution Calculations

To dilute a solution, use the relationship between initial and final concentrations and volumes.

• Formula: $M_1 V_1 = M_2 V_2$

• Example: To make 250 mL of 0.2 M from 1.0 M stock: $V_1 = \frac{M_2 V_2}{M_1} = \frac{0.2 \times 0.25}{1.0} = 0.05$ L = 50 mL

Acids and Bases

Definitions: Arrhenius and Bronsted-Lowry

Acids and bases are defined by their behavior in aqueous solution.

• Arrhenius Acid: Produces H+ ions in water.

• Arrhenius Base: Produces OH- ions in water.

• Bronsted-Lowry Acid: Proton donor.

• Bronsted-Lowry Base: Proton acceptor.

• Example: HCl is an Arrhenius and Bronsted-Lowry acid; NH3 is a Bronsted-Lowry base.

pH Calculations

pH measures the acidity of a solution, calculated from the concentration of H+ or OH- ions.

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

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

• Relationship: $\text{pH} + \text{pOH} = 14$

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

Titration Calculations

Titration is used to determine the concentration of an acid or base by reacting it with a solution of known concentration.

• Formula: $M_1 V_1 = M_2 V_2$ (for monoprotic acids and bases)

• Example: 25 mL of 0.1 M NaOH neutralizes 50 mL of HCl. Find HCl concentration: $M_1 V_1 = M_2 V_2 \rightarrow 0.1 \times 0.025 = M_2 \times 0.05 \rightarrow M_2 = 0.05$ M

Summary Table: Functional Groups in Organic Chemistry

Functional Group	Structure	Example
Alcohol	-OH	Ethanol
Ether	R-O-R'	Diethyl ether
Amine	-NH2	Methylamine
Ketone	C=O (within chain)	Acetone
Carboxylic Acid	-COOH	Acetic acid
Ester	-COOR	Ethyl acetate
Aldehyde	C=O (end of chain)	Formaldehyde