General Chemistry 100: Comprehensive Study Guide

General Chemistry Concepts and Problem-Solving

Atomic Structure and Electron Configuration

Understanding the structure of atoms is fundamental in chemistry. Atoms consist of protons, neutrons, and electrons, and their arrangement determines chemical properties.

• Atomic Number (Z): Number of protons in the nucleus; defines the element.

• Mass Number (A): Total number of protons and neutrons.

• Electron Configuration: Describes the arrangement of electrons in shells and subshells. For example, the configuration 1s2 2s2 2p6 3s2 3p6 4s2 3d6 corresponds to iron (Fe).

• Isotopes: Atoms of the same element with different numbers of neutrons.

Example: The element with configuration 1s2 2s2 2p6 3s2 3p6 4s2 3d6 is Fe (Iron).

Chemical Bonds and Molecular Geometry

Chemical bonds form when atoms share or transfer electrons. The geometry of molecules is determined by the number of bonding and lone pairs around the central atom.

• Ionic Bonds: Formed by transfer of electrons (e.g., NaCl).

• Covalent Bonds: Formed by sharing electrons (e.g., H2O).

• Nonpolar Molecules: Electrons are shared equally; no permanent dipole (e.g., CO2).

• Polar Molecules: Unequal sharing of electrons; permanent dipole (e.g., H2O).

• VSEPR Theory: Predicts molecular shapes based on electron pair repulsion.

Example: CO2 is a nonpolar molecule due to its linear geometry and symmetrical charge distribution.

Measurements, Units, and Significant Figures

Accurate measurements are essential in chemistry. Significant figures reflect the precision of a measurement.

• Density: $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$

• Significant Figures: All nonzero digits are significant; zeros between significant digits are significant; trailing zeros in a decimal are significant.

Example: A solution with 8.0 g of NaOH in 100 mL has a density of $0.08\ \text{g/mL}$.

Periodic Table and Chemical Properties

The periodic table organizes elements by increasing atomic number and similar chemical properties.

• Groups: Vertical columns; elements have similar valence electron configurations.

• Periods: Horizontal rows; properties change progressively across a period.

• Metals, Nonmetals, Metalloids: Classified based on physical and chemical properties.

Example: Halogens (Group 17) are diatomic and highly reactive nonmetals.

Chemical Formulas and Nomenclature

Chemical formulas represent the composition of compounds. Empirical formulas show the simplest ratio, while molecular formulas show the actual number of atoms.

• Empirical Formula: Simplest whole-number ratio of elements.

• Molecular Formula: Actual number of atoms in a molecule.

• Naming Compounds: Use systematic rules for ionic and covalent compounds.

Example: The empirical formula for a compound with 78% Boron and 22% Nitrogen is BN.

Stoichiometry and Chemical Reactions

Stoichiometry involves quantitative relationships in chemical reactions, including balancing equations and calculating reactant/product amounts.

• Balancing Equations: Ensure the same number of each atom on both sides.

• Mole Concept: $1\ \text{mol} = 6.022 \times 10^{23}$ particles (Avogadro's number).

• Molar Mass: Mass of one mole of a substance (g/mol).

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

• Theoretical Yield: Maximum amount of product possible from given reactants.

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

Example: If 5.0 g of H2 reacts with excess O2, the mass of water produced can be calculated using stoichiometry.

Solutions and Concentrations

Solutions are homogeneous mixtures of solute and solvent. Concentration is often expressed as molarity (M).

• Molarity (M): $\text{Molarity} = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Preparation of Solutions: Dissolving a known mass of solute in a specific volume of solvent.

Example: The molarity of a solution with 7.45 g NaCl in 100 mL is $1.28\ \text{M}$.

Gases and Gas Laws

Gases have unique properties described by several laws relating pressure, volume, temperature, and amount.

• Ideal Gas Law: $PV = nRT$

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

• Partial Pressure: The pressure exerted by a single gas in a mixture.

Example: At STP, 1 mole of any ideal gas occupies 22.4 L.

Acids, Bases, and pH

Acids and bases are classified by their ability to donate or accept protons. The pH scale measures acidity or basicity.

• Acid: Proton (H+) donor.

• Base: Proton (H+) acceptor.

• pH: $\text{pH} = -\log[\text{H}^+]$

Example: HCl is a strong acid; NaOH is a strong base.

Redox Reactions

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

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Oxidizing Agent: Accepts electrons.

• Reducing Agent: Donates electrons.

Example: In the reaction between Zn and CuSO4, Zn is oxidized and Cu2+ is reduced.

Thermochemistry

Thermochemistry studies the energy changes during chemical reactions, especially heat exchange.

• Endothermic: Absorbs heat (ΔH > 0).

• Exothermic: Releases heat (ΔH < 0).

• Calorimetry: Measurement of heat flow.

Example: Combustion reactions are typically exothermic.

Sample Table: Common Molecular Geometries (VSEPR Theory)

Electron Domains	Bonding Pairs	Lone Pairs	Geometry	Example
2	2	0	Linear	CO2
3	3	0	Trigonal planar	BF3
4	4	0	Tetrahedral	CH4
4	3	1	Trigonal pyramidal	NH3

Additional info:

• Some questions reference specific experiments (e.g., Millikan oil drop) and historical figures (e.g., Rutherford, Thomson) relevant to atomic theory.

• Calculations involving molar mass, empirical formulas, and limiting reactants are foundational for stoichiometry problems.

• Gas law problems often require unit conversions and understanding of STP conditions.

• Acid-base reactions and redox processes are central to chemical reactivity and laboratory analysis.