General Chemistry Core Concepts: Structured Study Notes

Atoms & Elements

Atomic Structure and Electron Configuration

Atoms are the basic units of matter, composed of protons, neutrons, and electrons. The arrangement of electrons in shells and subshells determines chemical properties.

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

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

• Electron Configuration: Distribution of electrons among atomic orbitals, e.g., $1s^2 2s^2 2p^6$.

• Valence Electrons: Electrons in the outermost shell; crucial for chemical bonding.

• Example: Sodium (Na): 1 valence electron.

Periodic Properties

Elements are arranged in the periodic table by increasing atomic number, revealing periodic trends.

• Electronegativity: Tendency to attract electrons; increases across a period, decreases down a group.

• Atomic Radius: Distance from nucleus to outermost electron; increases down a group, decreases across a period.

• Ionization Energy: Energy required to remove an electron; increases across a period.

• Example: Fluorine has the highest electronegativity.

Chemical Reactions

Types of Chemical Reactions

Chemical reactions involve the rearrangement of atoms to form new substances.

• Synthesis: Two or more substances combine to form one product.

• Decomposition: One substance breaks into two or more products.

• Single Replacement: One element replaces another in a compound.

• Double Replacement: Exchange of ions between two compounds.

Balancing Chemical Equations

Balancing a chemical equation requires that the number of atoms of each element is the same on both sides of the equation.

• Steps: List elements, count atoms, adjust coefficients.

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

Chemical Quantities & Aqueous Reactions

Mole Concept and Avogadro's Number

The mole is a fundamental unit for counting particles in chemistry.

• Avogadro's Number: $6.022 \times 10^{23}$ particles/mole.

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

• Example: 1 mole of NaCl = 58.44 g.

Stoichiometry

Stoichiometry involves calculations based on balanced chemical equations.

• Steps: Convert moles to mass, use mole ratios, connect reactant to product via volume.

• Example: How many moles of NaCl are produced from 34.4 g of HCl?

Solutions and Concentration

Solutions are homogeneous mixtures; concentration is often expressed as molarity.

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

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

Gases

Gas Laws

Gas behavior is described by several laws relating pressure, volume, temperature, and amount.

• Ideal Gas Law: $PV = nRT$

• Boyle's Law: $P_1V_1 = P_2V_2$ (at constant T)

• Charles's Law: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$ (at constant P)

• Example: Calculate volume of 1 mol gas at STP.

Thermochemistry

Energy Changes in Reactions

Thermochemistry studies heat changes during chemical reactions.

• Enthalpy (H): Heat content of a system.

• Specific Heat Capacity: Amount of heat required to raise temperature of 1 g substance by $1^\circ$C.

• Example: Calculate heat required to raise temperature of 17 g substance by $12^\circ$C.

Quantum Mechanics

Electronic Structure and Light

Quantum mechanics explains atomic structure and the behavior of electrons.

• Energy Levels: Electrons occupy quantized energy levels.

• Photon Energy: $E = h\nu$

• Wavelength-Frequency Relationship: $c = \lambda \nu$

• Example: Calculate wavelength for frequency $\nu$.

Bonding & Molecular Structure

Covalent and Ionic Bonds

Bonds form when atoms share or transfer electrons to achieve stability.

• Ionic Bonds: Transfer of electrons between metals and nonmetals.

• Covalent Bonds: Sharing of electrons between nonmetals.

• Polar Covalent Bonds: Unequal sharing of electrons, leading to partial charges.

• Nonpolar Covalent Bonds: Equal sharing of electrons; no polar or nonpolar bonds.

Molecular Geometry and VSEPR Theory

The shape of molecules is predicted by the Valence Shell Electron Pair Repulsion (VSEPR) theory.

• Linear: 180° bond angle, e.g., CO2

• Trigonal Planar: 120° bond angle, e.g., BF3

• Tetrahedral: 109.5° bond angle, e.g., CH4

Solutions

Types of Mixtures

Mixtures can be homogeneous (solutions) or heterogeneous.

• Homogeneous Mixture: Uniform composition, e.g., salt water.

• Heterogeneous Mixture: Non-uniform, e.g., sand in water.

Solubility and Precipitation

Solubility is the ability of a substance to dissolve in a solvent. Precipitation occurs when solubility limits are exceeded.

• Example: Mixing AgNO3 and NaCl forms a precipitate of AgCl.

Chemical Kinetics & Equilibrium

Reaction Rates

Chemical kinetics studies the speed of reactions and factors affecting it.

• Rate Law: $\text{Rate} = k[A]^m[B]^n$

• Factors: Concentration, temperature, catalysts.

Chemical Equilibrium

At equilibrium, the rates of forward and reverse reactions are equal.

• Equilibrium Constant (K): $K = \frac{[\text{products}]}{[\text{reactants}]}$

• Le Châtelier's Principle: System shifts to counteract changes.

Acids, Bases, and Aqueous Equilibrium

Acid-Base Definitions

Acids donate protons (H+), bases accept protons.

• Strong Acid/Base: Completely ionize in water.

• Weak Acid/Base: Partially ionize.

• Example: Acetic acid is a weak acid; CH3COOH is a weak acid.

Thermodynamics

First Law of Thermodynamics

Energy cannot be created or destroyed, only transformed.

• Internal Energy (U): Total energy of a system.

• State Function: Depends only on current state.

• Enthalpy: Spontaneity determined by sign of $\Delta G$.

Electrochemistry

Redox Reactions

Electrochemistry studies electron transfer in chemical reactions.

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Oxidizing Agent: Causes reduction, is reduced.

• Reducing Agent: Causes oxidation, is oxidized.

• Example: Cu2+ reduced, Zn oxidized.

Lab Techniques and Procedures

Measurement and Significant Figures

Accurate measurement and reporting are essential in chemistry.

• Significant Figures: Digits that reflect precision, rules for counting and rounding.

• Example: 0.001867 has 4 significant figures.

Density Calculations

Density is mass per unit volume.

• Formula: $\text{Density} = \frac{\text{mass}}{\text{volume}}$

• Example: Find density of a gold bar weighing 8.0 kg with volume 0.41 L.

Mathematical Operations and Functions

Unit Conversions and Calculations

Chemistry often requires converting between units and using dimensional analysis.

• Example: Convert grams to moles using molar mass.

• Example: Calculate number of molecules in a given mass.

HTML Table Example: Periodic Trends

The following table compares periodic trends for selected elements.

Element	Electronegativity	Ionization Energy	Atomic Radius
F	Highest	High	Small
Na	Low	Low	Large
Cl	High	High	Medium

Additional info:

• Questions cover all major topics in General Chemistry, including atoms, bonding, reactions, stoichiometry, gases, thermochemistry, quantum mechanics, periodic properties, solutions, kinetics, equilibrium, acids/bases, thermodynamics, electrochemistry, and lab techniques.

• Some questions require calculation and application of formulas; others test conceptual understanding.

• Bonus section includes short answer and lab-based questions.