General Chemistry II Midterm Study Guide: Key Concepts and Equations

Chapter 4: Chemical Composition and Nomenclature

Mass Percent Composition

Mass percent composition expresses the percentage by mass of each element in a compound. It is useful for determining the relative amounts of elements present.

• Definition: The mass percent of an element is the mass of the element divided by the total mass of the compound, multiplied by 100%.

• Formula: $\text{Mass percent} = \left( \frac{\text{mass of element in 1 mol of compound}}{\text{molar mass of compound}} \right) \times 100\%$

• Example: In water (H2O), calculate the mass percent of hydrogen and oxygen.

Empirical vs. Molecular Formulas

Empirical formulas show the simplest whole-number ratio of atoms in a compound, while molecular formulas show the actual number of atoms of each element in a molecule.

• Empirical Formula: Simplest ratio (e.g., CH2O for glucose).

• Molecular Formula: Actual number (e.g., C6H12O6 for glucose).

• Relationship: $\text{Molecular formula} = n \times \text{Empirical formula}$, where n is an integer.

Combustion Analysis

Combustion analysis is a method used to determine the empirical formula of a compound, typically containing C, H, and O, by burning it and measuring the amounts of CO2 and H2O produced.

• Key Steps: Measure mass of CO2 and H2O, calculate moles of C and H, and determine O by difference.

Chemical Nomenclature

Chemical nomenclature is the system for naming chemical compounds. It includes rules for naming ionic, covalent, transition metal, and polyatomic compounds, as well as hydrates.

• Ionic Compounds: Metal + nonmetal, use Roman numerals for transition metals.

• Polyatomic Ions: Ions composed of multiple atoms (e.g., SO42−, NO3−).

• Hydrates: Compounds with water molecules attached (e.g., CuSO4·5H2O).

• Example: Na2SO4 is sodium sulfate.

Chapter 5: Bonding and Molecular Structure

Bond Polarity and Electronegativity

Bond polarity arises from differences in electronegativity between atoms. Electronegativity is a measure of an atom's ability to attract electrons in a bond.

• Nonpolar Covalent Bonds: Electrons shared equally (e.g., H2).

• Polar Covalent Bonds: Electrons shared unequally (e.g., HCl).

• Ionic Bonds: Electrons transferred (e.g., NaCl).

• Electronegativity Chart: Provided on the exam; use to determine bond polarity.

Lewis Structures

Lewis structures represent the arrangement of electrons in molecules, showing bonds and lone pairs.

• Octet Rule: Atoms tend to have eight electrons in their valence shell.

• Octet Violators: Radicals (odd electrons), less than 8 (e.g., H, B), more than 8 (e.g., S, P).

• Formal Charge: Used to determine the most stable structure. $\text{Formal charge} = \text{valence electrons} - (\text{lone pair electrons} + \frac{1}{2} \times \text{bonding electrons})$

• Resonance: Multiple valid Lewis structures for a molecule.

Organic Molecules

Organic molecules are compounds containing carbon atoms, often forming chains or rings.

• Recognizing Line Structures: Each vertex or end represents a carbon atom unless otherwise specified.

VSEPR Theory

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the shapes of molecules based on electron domain repulsion.

• Electron Pair and Molecular Geometry: Shapes depend on the number of electron domains (bonding and lone pairs).

• 5 or 6 Domain Shapes: Trigonal bipyramidal (5), octahedral (6).

• Bond Angles: Ideal angles may deviate due to lone pairs.

• Molecular Shape and Polarity: Shape affects whether a molecule is polar or nonpolar.

Valence Bond Theory

Valence bond theory explains bonding using atomic orbital overlap and hybridization.

• Hybridization: Mixing of atomic orbitals (sp, sp2, sp3).

• No sp3d or sp3d2 hybridization for this exam.

• Sigma (σ) and Pi (π) Bonds: Sigma bonds are single bonds; pi bonds are found in double/triple bonds.

• Bond Strength: Sigma bonds are generally stronger than pi bonds.

• Rotation and Isomerism: Pi bonds restrict rotation, leading to cis/trans isomers.

• Hybridization in Organic Molecules: Used to explain bonding in carbon compounds.

Chapter 7: Chemical Reactions and Stoichiometry

Balancing Chemical Equations

Balancing equations ensures the same number of atoms of each element on both sides of a reaction.

• Steps: Identify reactants and products, count atoms, adjust coefficients.

• Example: $2H_2 + O_2 \rightarrow 2H_2O$

Reaction Stoichiometry

Stoichiometry involves calculating the amounts of reactants and products in chemical reactions.

• Mole Ratios: Use coefficients from balanced equations.

• Conversions: Grams ↔ Moles ↔ Molecules.

Limiting and Excess Reagents

The limiting reagent is the reactant that is completely consumed first, limiting the amount of product formed.

• Identify: Compare mole ratios of reactants.

• Excess Reagent: Reactant left over after reaction.

Theoretical, Actual, and Percent Yield

Yield calculations compare the amount of product obtained to the maximum possible.

• Theoretical Yield: Maximum possible product.

• Actual Yield: Amount actually obtained.

• Percent Yield Formula: $\text{Percent yield} = \left( \frac{\text{actual yield}}{\text{theoretical yield}} \right) \times 100\%$

Chapter 8: Solutions and Chemical Reactions

Solubility Rules

Solubility rules determine whether an ionic compound dissolves in water. A table of rules will be provided on the exam.

• Common Soluble Ions: Group 1 metals, NH4+, NO3−, etc.

• Common Insoluble Ions: Ag+, Pb2+, CO32−, etc.

Electrolytes

Electrolytes are substances that conduct electricity when dissolved in water. They are classified as strong or weak.

• Strong Electrolytes: Completely dissociate (e.g., NaCl).

• Weak Electrolytes: Partially dissociate (e.g., CH3COOH).

• Non-Electrolytes: Do not dissociate (e.g., sugar).

Molarity

Molarity (M) is the concentration of a solution, defined as moles of solute per liter of solution.

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

• Example: 0.5 mol NaCl in 1 L water = 0.5 M NaCl solution.

Dilution

Dilution involves adding solvent to decrease the concentration of a solution.

• Formula: $M_1V_1 = M_2V_2$

• Example: To dilute 1.0 M solution to 0.5 M, use the formula to find required volumes.

Types of Chemical Reactions

Chemical reactions are classified by the changes that occur.

• Double Displacement Reactions: Exchange of ions between compounds.

• Precipitation: Formation of an insoluble product.

• Molecular, Complete Ionic, and Net Ionic Equations: Show different levels of detail for reactions in solution.

Acid-Base Neutralization

Acid-base reactions involve the transfer of protons (H+) from acids to bases.

• Strong Acids: HCl, HBr, HI, HNO3, H2SO4, HClO4, HClO3

• Strong Bases: Alkali and Ba/Sr/Ca hydroxides (e.g., NaOH, Ca(OH)2).

• Example: $\text{HCl} + \text{NaOH} \rightarrow \text{NaCl} + \text{H}_2\text{O}$

Redox Reactions

Redox (reduction-oxidation) reactions involve the transfer of electrons between species.

• Oxidation States: Numbers assigned to atoms to track electron transfer.

• Oxidizing Agent: Causes oxidation, is reduced.

• Reducing Agent: Causes reduction, is oxidized.

• Combustion Reactions: Rapid reaction with oxygen, producing heat and light.

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

Additional info:

• Polyatomic ions and their charges should be memorized for nomenclature and reaction prediction.

• Tables for solubility rules and electronegativity will be provided on the exam; students should be familiar with their use.

• For VSEPR, students should know the shapes and bond angles for up to 6 electron domains.