General Chemistry I: Exam V Study Guide – Lewis Structures, VSEPR, and Foundational Concepts

Overview of Exam Topics

This study guide covers key topics from General Chemistry I, focusing on Lewis structures, VSEPR theory, and foundational chemical concepts such as stoichiometry, molecular geometry, and types of chemical reactions. The guide is organized by major themes and subtopics to facilitate exam preparation.

Lewis Structures and VSEPR Theory

Valence Electrons

• Definition: Valence electrons are the outermost electrons of an atom and are involved in chemical bonding.

• Determination: The number of valence electrons is typically equal to the group number for main group elements.

• Example: Carbon (C) has 4 valence electrons; Oxygen (O) has 6.

Lewis Structures

• Definition: Lewis structures are diagrams that represent the bonding between atoms and the lone pairs of electrons in a molecule.

• Steps to Draw:

  1. Count total valence electrons.

  2. Arrange atoms (central atom is usually least electronegative).

  3. Connect atoms with single bonds.

  4. Distribute remaining electrons to complete octets (or duets for H).

  5. Form double/triple bonds if necessary to satisfy octet rule.

• Example: The Lewis structure for water (H2O) shows two single bonds between O and H, with two lone pairs on O.

Resonance Structures

• Definition: Resonance structures are two or more valid Lewis structures for the same molecule that differ only in the placement of electrons.

• Example: The nitrate ion (NO3-) has three resonance structures.

VSEPR Theory (Valence Shell Electron Pair Repulsion)

• Definition: VSEPR theory predicts the geometry of molecules based on the repulsion between electron pairs around a central atom.

• Common Geometries:

  • Linear: 2 electron groups, 180°

  • Trigonal planar: 3 electron groups, 120°

  • Tetrahedral: 4 electron groups, 109.5°

  • Trigonal bipyramidal: 5 electron groups, 90°/120°

  • Octahedral: 6 electron groups, 90°

• Example: Methane (CH4) is tetrahedral; carbon dioxide (CO2) is linear.

Molecular Geometry

• Definition: The three-dimensional arrangement of atoms in a molecule.

• Determined by: Number of bonding pairs and lone pairs on the central atom.

• Example: Ammonia (NH3) has a trigonal pyramidal geometry due to one lone pair on nitrogen.

Stoichiometry and Chemical Calculations

Precipitation Reactions

• Definition: Reactions in which two aqueous solutions combine to form an insoluble solid (precipitate).

• Example: $\text{AgNO}_3(aq) + \text{NaCl}(aq) \rightarrow \text{AgCl}(s) + \text{NaNO}_3(aq)$

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.

• Calculation: $\text{Molecular formula} = n \times \text{Empirical formula}$, where $n = \frac{\text{Molar mass}}{\text{Empirical formula mass}}$

Stoichiometric Calculations

• Definition: Calculations based on balanced chemical equations to determine the amounts of reactants and products.

• Key Steps:

  1. Write a balanced equation.

  2. Convert given quantities to moles.

  3. Use mole ratios to find moles of desired substance.

  4. Convert moles to grams or other units as needed.

• Example: How many grams of CO2 are produced from 10.0 g of C6H12O6 in cellular respiration?

Limiting Reactant and Percent Yield

• Limiting Reactant: The reactant that is completely consumed first, limiting the amount of product formed.

• Percent Yield: $\text{Percent yield} = \frac{\text{Actual yield}}{\text{Theoretical yield}} \times 100\%$

Mole Calculations

• Definition: The mole is the SI unit for amount of substance; 1 mole = $6.022 \times 10^{23}$ particles (Avogadro's number).

• Conversions:

  • Moles to grams: $\text{mass (g)} = \text{moles} \times \text{molar mass (g/mol)}$

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

Atomic Structure and Subatomic Particles

Subatomic Particles

• Protons: Positively charged particles in the nucleus; determine atomic number.

• Neutrons: Neutral particles in the nucleus; contribute to atomic mass.

• Electrons: Negatively charged particles in orbitals around the nucleus.

• Example: Carbon-12 has 6 protons, 6 neutrons, and 6 electrons.

Chemical Bonding and Nomenclature

Ionic vs. Covalent Compounds

• Ionic Compounds: Formed from the transfer of electrons between metals and nonmetals; composed of ions.

• Covalent Compounds: Formed from the sharing of electrons between nonmetals.

• Example: NaCl is ionic; H2O is covalent.

Acid Naming

• Binary Acids: Use the prefix "hydro-" and the suffix "-ic acid" (e.g., HCl = hydrochloric acid).

• Oxyacids: Based on the polyatomic ion:

  • "-ate" becomes "-ic acid" (e.g., HNO3 = nitric acid)

  • "-ite" becomes "-ous acid" (e.g., HNO2 = nitrous acid)

Summary Table: Key Concepts and Calculations

Concept	Definition	Key Formula/Example
Valence Electrons	Electrons in the outermost shell	Group number for main group elements
Lewis Structure	Diagram of bonds and lone pairs	H2O: O with two single bonds to H, two lone pairs
VSEPR Geometry	3D arrangement of atoms	CH4: Tetrahedral, 109.5°
Empirical Formula	Simplest ratio of atoms	CH2O for glucose
Percent Yield	Efficiency of reaction	$\frac{\text{Actual yield}}{\text{Theoretical yield}} \times 100\%$
Subatomic Particles	Protons, neutrons, electrons	Carbon-12: 6p, 6n, 6e
Ionic vs. Covalent	Transfer vs. sharing of electrons	NaCl (ionic), H2O (covalent)

Additional info:

• Some topics (e.g., molecular geometry, resonance, stoichiometric calculations) are expanded with standard academic context to ensure completeness.

• Specific calculation steps and examples are provided for clarity, even if not detailed in the original list.