Introduction to General Chemistry

Scientific Method and Key Concepts

• Hypothesis: A tentative explanation for an observation, which can be tested by experiments.

• Theory: A well-substantiated explanation of some aspect of the natural world that can incorporate facts, laws, inferences, and tested hypotheses.

• Law: A statement based on repeated experimental observations that describes some aspect of the universe (e.g., Law of Conservation of Mass).

• Physical Properties: Characteristics that can be observed without changing the substance's identity (e.g., melting point, density).

• Chemical Properties: Characteristics that describe a substance's ability to undergo chemical changes (e.g., flammability, reactivity).

• Physical Change: A change that does not alter the chemical composition (e.g., melting, boiling).

• Chemical Change: A change that results in the formation of new substances (e.g., rusting, combustion).

Classification of Matter

• Pure Substances: Matter with a fixed composition (elements and compounds).

• Mixtures: Physical combinations of two or more substances (homogeneous and heterogeneous).

• States of Matter: Solid (definite shape and volume), Liquid (definite volume, indefinite shape), Gas (indefinite shape and volume).

• Macroscopic vs. Microscopic: Macroscopic refers to what can be seen with the naked eye; microscopic refers to atomic or molecular level observations.

Mathematical Operations and Measurements

SI Units and Conversions

• SI Units: Standard units of measurement (meter, kilogram, second, mole, kelvin, etc.).

• Prefixes: Used to indicate multiples or fractions of units (e.g., kilo-, centi-, milli-).

• Temperature Conversions: Celsius to Kelvin:

• Extensive Properties: Depend on the amount of matter (e.g., mass, volume).

• Intensive Properties: Independent of the amount of matter (e.g., density, temperature).

• Density:

• Percent Composition:

Precision, Accuracy, and Significant Figures

• Precision: How close repeated measurements are to each other.

• Accuracy: How close a measurement is to the true value.

• Significant Figures: Digits in a measurement that are known with certainty plus one estimated digit.

• Exact Numbers: Values known with complete certainty (e.g., counting numbers, defined quantities).

Atoms, Elements, and the Periodic Table

Atomic Structure and Isotopes

• Subatomic Particles: Protons (positive, mass ≈ 1 amu), Neutrons (neutral, mass ≈ 1 amu), Electrons (negative, mass ≈ 1/1836 amu).

• Nuclear Atom Model: Atoms consist of a dense nucleus (protons and neutrons) surrounded by electrons.

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

• Ions: Atoms or molecules with a net charge due to loss or gain of electrons.

• Determining Subatomic Particles: For a neutral atom:

  • Protons = Atomic number

  • Neutrons = Mass number - Atomic number

  • Electrons = Atomic number (for neutral atoms)

• Average Atomic Mass:

Periodic Table and Properties

• Element Identification: Each element has a unique chemical symbol and atomic number.

• Periodic Trends: Properties such as atomic radius, ionization energy, and electronegativity vary predictably across the table.

Chemical Compounds and Nomenclature

Types of Compounds

• Ionic Compounds: Composed of metals and nonmetals; transfer of electrons.

• Molecular (Covalent) Compounds: Composed of nonmetals; sharing of electrons.

• Organic vs. Inorganic Compounds: Organic compounds contain carbon-hydrogen bonds; inorganic do not.

Naming and Writing Formulas

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

• Naming Rules: Systematic methods for naming ionic, molecular, and acid compounds.

• Hydrates: Ionic compounds containing water molecules in their structure.

• Functional Groups: Specific groups of atoms in organic molecules (e.g., alkanes, alcohols, carboxylic acids, amines).

Formulas and Calculations

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

• Avogadro’s Number: particles/mol.

• Conversions: Mass ↔ Moles ↔ Particles.

• Percent Composition:

• Empirical vs. Molecular Formula: Empirical is the simplest ratio; molecular is the actual number of atoms.

Chemical Reactions and Stoichiometry

Types of Reactions

• Writing Chemical Equations: Representing reactions with symbols, including states of matter.

• Combustion Reactions: Hydrocarbon reacts with to produce and .

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

Electrolytes and Solutions

• Electrolytes: Substances that conduct electricity in solution (strong, weak, nonelectrolytes).

• Solubility Rules: Guidelines for predicting if an ionic compound dissolves in water.

• Precipitation Reactions: Formation of an insoluble product from soluble reactants.

• Net Ionic Equations: Show only the species that change during the reaction.

Acid-Base and Redox Reactions

• Acids and Bases: Arrhenius (produce or in water), Bronsted-Lowry (proton donors/acceptors).

• Strong vs. Weak Acids/Bases: Strong dissociate completely; weak only partially.

• Oxidation-Reduction (Redox): Involves transfer of electrons; oxidation is loss, reduction is gain.

• Oxidizing/Reducing Agents: Oxidizing agent gains electrons; reducing agent loses electrons.

• Oxidation States: Assigned using rules to track electron transfer.

Stoichiometry and Yields

• Limiting Reactant: The reactant that determines the maximum amount of product.

• Theoretical Yield: Maximum possible amount of product.

• Actual Yield: Amount actually obtained from experiment.

• Percent Yield:

Solutions and Concentrations

Molarity and Dilutions

• Molarity (M):

• Dilution Equation:

• Solution Stoichiometry: Calculating amounts in reactions occurring in solution.

Thermochemistry

Energy and Heat

• System vs. Surroundings: System is the part studied; surroundings are everything else.

• Kinetic vs. Potential Energy: Kinetic is energy of motion; potential is stored energy.

• Heat (q) and Work (w): Forms of energy transfer.

• First Law of Thermodynamics:

• Specific Heat:

• Coffee-Cup Calorimetry: Used to measure heat changes at constant pressure.

• Enthalpy Change (): Heat change at constant pressure.

• Hess’s Law:

• Standard Enthalpy of Formation: for forming 1 mole from elements in standard states.

Quantum Mechanics and Atomic Structure

Light and Atomic Models

• Wave-Particle Duality: Light and electrons exhibit both wave and particle properties.

• Bohr Model: Electrons orbit nucleus in quantized energy levels.

• Energy of Light:

• de Broglie Wavelength:

• Uncertainty Principle:

Electronic Structure and Periodic Properties

Quantum Numbers and Orbitals

• Quantum Numbers: n (principal), l (angular momentum), ml (magnetic), ms (spin).

• Electron Configurations: Distribution of electrons in orbitals using Aufbau principle, Hund’s rule, and Pauli Exclusion Principle.

• Ground vs. Excited States: Ground is lowest energy; excited has higher energy electrons.

• Paramagnetic vs. Diamagnetic: Paramagnetic has unpaired electrons; diamagnetic has all electrons paired.

Periodic Trends

• Effective Nuclear Charge: Net positive charge experienced by valence electrons.

• Atomic and Ionic Radii: Size of atoms and ions; trends across periods and groups.

• Ionization Energy: Energy required to remove an electron.

• Electron Affinity: Energy change when an electron is added.

• Lattice Energy: Energy required to separate one mole of an ionic solid into gaseous ions.

Chemical Bonding and Molecular Structure

Lewis Structures and Bonding

• Lewis Symbols: Represent valence electrons as dots.

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

• Formal Charge:

• Resonance: Multiple valid Lewis structures for a molecule.

• Expanded Octet: Atoms in period 3 or higher can have more than 8 valence electrons.

• Bond Polarity: Difference in electronegativity between atoms.

• Bond Order: Number of shared electron pairs between two atoms.

• Bond Energy: Energy required to break a bond.

Molecular Shapes and Bond Theories

VSEPR and Molecular Geometry

• VSEPR Theory: Predicts molecular shapes based on electron group repulsions.

• Bond Angles: Determined by electron group geometry.

• Polarity: Determined by shape and bond polarity.

Bond Theories

• Valence Bond Theory: Bonds form from overlap of atomic orbitals; hybridization explains geometry.

• Molecular Orbital Theory: Atomic orbitals combine to form molecular orbitals; explains bond order, magnetism.

• Sigma and Pi Bonds: Sigma () bonds are end-to-end; pi () bonds are side-to-side overlaps.

Gases and Gas Laws

Gas Properties and Laws

• Pressure Units: atm, mmHg, torr, Pa.

• Boyle’s Law: (constant T, n)

• Charles’s Law: (constant P, n)

• Avogadro’s Law: (constant P, T)

• Ideal Gas Law:

• Combined Gas Law:

• STP Conditions: 1 atm, 273.15 K

• Dalton’s Law of Partial Pressures:

• Graham’s Law of Effusion:

• van der Waals Equation:

Additional Topics

Lab Techniques and Procedures

• Understanding and applying laboratory safety and measurement techniques.

• Using density, percent composition, and other properties as conversion factors in calculations.

Mathematical Operations and Functions

• Applying significant figures and unit conversions in all calculations.

• Using dimensional analysis for problem-solving.

Sample Table: Comparison of Physical and Chemical Properties

Sample Table: SI Prefixes

Example Calculation: Molarity

• To prepare 0.5 L of 1.0 M NaCl solution: mol

• Mass NaCl needed:

Additional info: This guide covers all major learning objectives for a two-semester general chemistry sequence, including foundational concepts, mathematical operations, atomic structure, periodic trends, chemical bonding, reactions, stoichiometry, thermochemistry, quantum mechanics, and gas laws. For each topic, students should practice with example problems and refer to their textbook for detailed explanations and additional examples.