Introduction to Chemistry: Comprehensive Study Notes

Chapter 1: Foundations of Chemistry

Scientific Method

The scientific method is a systematic approach used in scientific investigation to acquire new knowledge and validate existing concepts.

• Observation: Gathering information through the senses or instruments.

• Hypothesis: A tentative explanation or prediction that can be tested.

• Experiment: A controlled procedure to test the hypothesis.

• Law: A statement that describes consistent natural phenomena (e.g., Law of Conservation of Mass).

• Theory: A well-substantiated explanation of some aspect of the natural world.

Data Analysis

Data analysis involves identifying patterns and interpreting results from experiments.

• Patterns in Data: Recognizing trends or regularities in collected data.

• Interpreting Graphs: Understanding graphical representations of data.

  • Y-intercept: The value where the graph crosses the y-axis.

  • Slope: Indicates the rate of change; calculated as rise over run.

  • Data Points: Individual measurements plotted on the graph.

  • Using Graphs to Predict Outcomes: Extrapolating or interpolating data to make predictions.

Chapter 2: Measurements and Calculations

Scientific Notation

Scientific notation is a method for expressing very large or very small numbers using powers of ten.

• Writing Numbers: $a \times 10^n$ where $1 \leq a < 10$ and $n$ is an integer.

• Converting Decimal to Scientific Notation: Move the decimal point to create a number between 1 and 10, count the places moved for the exponent.

• Converting Scientific Notation to Decimal: Expand the number by multiplying by the power of ten.

Significant Figures (SigFigs)

Significant figures indicate the precision of a measured or calculated quantity.

• Determining SigFigs: All nonzero digits are significant; zeros between nonzero digits and trailing zeros in decimals are significant.

• Calculations: For multiplication/division, use the least number of sigfigs; for addition/subtraction, use the least number of decimal places.

• Rounding: Round to the correct number of significant figures based on calculation rules.

Measurements

Measurements in chemistry require understanding units and prefixes.

• Units: Standard quantities (e.g., meter, gram, liter).

• Prefix Multipliers: MEMORIZE common prefixes (e.g., kilo-, centi-, milli-).

• Precision and Uncertainty: Precision refers to reproducibility; uncertainty is the estimated error in measurement.

Unit Conversions

Unit conversions use conversion factors to change from one unit to another.

• Example: To convert 10 cm to meters: $10 \text{ cm} \times \frac{1 \text{ m}}{100 \text{ cm}} = 0.1 \text{ m}$

Chapter 3: Matter and Energy

Molecules and States of Matter

Matter exists in different forms and can be classified by composition.

• Molecules: Groups of atoms bonded together.

• States of Matter: Solid, liquid, gas.

• Mixtures: Homogeneous (uniform composition) vs. Heterogeneous (non-uniform composition).

Physical and Chemical Properties

Properties describe how substances behave and interact.

• Physical Properties: Observable without changing composition (e.g., melting point).

• Chemical Properties: Describe reactivity (e.g., flammability).

Law of Conservation of Mass

In chemical reactions, mass is conserved.

• Statement: The mass of reactants equals the mass of products.

Energy

Energy is the capacity to do work or produce heat.

• Kinetic Energy: Energy of motion.

• Potential Energy: Stored energy due to position.

• Chemical Energy: Energy stored in chemical bonds.

• Electrical Energy: Energy from moving electrons.

• Thermal Energy: Energy due to temperature.

• Energy Units: Joule (J), calorie (cal).

• Converting Energy Units: $1 \text{ cal} = 4.184 \text{ J}$

Work and Conservation of Energy

• Work: Energy transferred by a force.

• Law of Conservation of Energy: Energy cannot be created or destroyed, only transformed.

Exothermic and Endothermic Processes

• Exothermic: Releases energy (heat).

• Endothermic: Absorbs energy (heat).

Temperature and Heat

• Temperature: Measure of average kinetic energy.

• Heat: Energy transferred due to temperature difference.

• Calculating Heat: $q = mC\Delta T$ Where $q$ = heat, $m$ = mass, $C$ = specific heat capacity, $\Delta T$ = change in temperature.

• Temperature Scales: Fahrenheit, Celsius, Kelvin.

• Conversions: $K = ^\circ C + 273.15$

Chapter 4: Atomic Structure and the Periodic Table

Atomic Theory

Atoms are the fundamental units of matter, composed of subatomic particles.

• Protons: Positively charged, located in nucleus.

• Neutrons: Neutral, located in nucleus.

• Electrons: Negatively charged, located outside nucleus.

• Charges and Locations: Know the charge and location of each particle.

Atomic Mass and Number

• Atomic Number: Number of protons in nucleus.

• Atomic Mass: Sum of protons and neutrons.

Chemical Symbols and Periodic Law

• Chemical Symbols: Shorthand notation for elements (e.g., H for hydrogen).

• Periodic Law: Properties of elements repeat periodically when arranged by atomic number.

Periodic Table

• Metals: Good conductors, malleable, shiny.

• Metalloids: Properties intermediate between metals and nonmetals.

• Nonmetals: Poor conductors, brittle.

• Main Group Elements: Groups 1, 2, 13-18.

• Transition Metals: Groups 3-12.

• Groups: Columns in the periodic table (e.g., noble gases, alkali metals, alkaline earth metals, halogens).

Ions and Isotopes

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

• Charge of Ions: Number of protons minus number of electrons.

• Chemical Symbols for Ions: E.g., Na+, Cl-.

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

• Mass Number: Total number of protons and neutrons.

• % Relative Abundance: Proportion of each isotope in nature.

• Weighted Average Atomic Mass: $\text{Average mass} = \sum (\text{isotope mass} \times \text{fractional abundance})$

Chapter 5: Chemical Formulas and Nomenclature

Chemical Formulas

• Order: Write the more metallic element first.

• Subscripts: Indicate the number of atoms of each element.

• Types: Empirical (simplest ratio), molecular (actual number), structural (shows arrangement).

Polyatomic Ions

• Definition: Atoms covalently bonded that act as a unit.

• Notation: Use parentheses and subscripts for multiples (e.g., Ca(NO3)2).

Diatomic Molecules

• Common Diatomics: H2, O2, F2, N2, Br2, Cl2, I2.

Nomenclature

• Type 1 Ionic Compounds: Name of cation + base name of anion + "ide".

• Type 2 Ionic Compounds: Name of cation (charge in Roman numerals) + base name of anion + "ide".

• Ionic Compounds with Polyatomic Ions: Name cation + name of polyatomic ion.

• Molecular/Covalent Compounds: Prefix name of 1st element + prefix base name of 2nd element + "ide".

• Prefixes: mono-, di-, tri-, tetra-, penta-, etc.

• Acids: Binary acids: hydro + base name of nonmetal + "ic acid"; Oxyacids: polyatomic ion ends in -ite: base name + "ous acid", ends in -ate: base name + "ic acid".

Formula Mass

• Calculation: $(\text{atoms in 1st element} \times \text{atomic mass of 1st element}) + (\text{atoms in 2nd element} \times \text{atomic mass of 2nd element}) + \ldots$

• Units: Atomic mass unit (amu).

Chapter 6: The Mole and Molar Mass

Moles

• Definition: 1 mole = $6.022 \times 10^{23}$ entities (Avogadro's number).

• Conversions: Between moles, atoms, molecules, and grams.

Molar Mass

• Definition: Mass of 1 mole of a substance (g/mol).

• Calculation: Add atomic masses of all atoms in the formula.

• Conversions: Use molar mass as a conversion factor between grams and moles.

Percent Composition and Empirical/Molecular Formulas

• Mass Percent Composition: $(\frac{\text{mass of element in 1 mole}}{\text{molar mass}}) \times 100\%$

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

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

Chapter 7: Chemical Reactions and Aqueous Solutions

Indications of Chemical Reactions

• Color change, gas formation, precipitate formation, energy change.

Chemical Equations

• Writing Equations: Use chemical formulas and indicate phases (s, l, g, aq).

• Balancing Equations: Ensure equal numbers of each atom on both sides.

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

• Spectator Ions: Ions that do not participate in the reaction.

Aqueous Solutions and Solubility

• Strong Electrolytes: Dissociate completely in water.

• Solubility: Soluble substances dissolve; insoluble do not.

• Precipitation Reactions: Formation of an insoluble product (precipitate).

• Solubility Chart: Used to predict precipitate formation.

Acid-Base Reactions

• Acids: Produce H+ in solution.

• Bases: Produce OH- in solution.

• Strong Acids/Bases: MEMORIZE common examples.

Gas-Evolution and Redox Reactions

• Gas-Evolution: Reaction produces a gas.

• Redox (Oxidation-Reduction): Electron transfer occurs; OIL RIG (Oxidation Is Loss, Reduction Is Gain of electrons).

Types of Chemical Reactions

• Synthesis: Two or more reactants form one product.

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

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

• Double-Displacement: Exchange of ions between two compounds.

Chapter 9: Electromagnetic Radiation and Atomic Structure

Electromagnetic Radiation

• Wavelength ($\lambda$): Distance between two peaks.

• Frequency ($\nu$): Number of waves per second.

• Photon: Quantum of light energy.

• Electromagnetic Spectrum: Includes gamma rays, X-rays, UV, visible light, infrared, microwaves, radio waves.

Orbitals and Electron Configuration

• Shapes: s, p, d, f orbitals.

• Principal Quantum Number ($n$): Indicates energy level.

• Electron Configuration: Arrangement of electrons in orbitals.

• Orbital Diagrams: Visual representation of electron arrangement.

• Energy Ordering: Aufbau principle, Pauli exclusion, Hund's rule.

• Ionization Energy: Energy required to remove an electron.

• Metallic Character: Tendency to lose electrons.

• Atomic Size: Increases down a group, decreases across a period.

Chapter 10: Chemical Bonding and Molecular Structure

Lewis Structures

• Valence Electrons: Electrons in the outermost shell.

• Lewis Structures: Diagrams showing bonding and lone pairs.

• For Elements, Ions, Ionic Molecules, Molecular/Covalent Molecules: Indicate single, double, triple bonds and resonance structures.

• Polyatomic Ions: Draw Lewis structures for ions with multiple atoms.

Predicting Molecular Shape (VSEPR Theory)

• VSEPR: Valence Shell Electron Pair Repulsion theory predicts molecular geometry.

• Linear: 2 electron groups, 2 bonding groups, 180° bond angle.

• Trigonal Planar: 3 electron groups, 3 bonding groups, 120° bond angle.

• Bent (Trigonal Planar): 3 electron groups, 2 bonding groups, 120° bond angle.

• Tetrahedral: 4 electron groups, 4 bonding groups, 109.5° bond angle.

• Trigonal Pyramidal: 4 electron groups, 3 bonding groups, 109.5° bond angle.

• Bent (Tetrahedral): 4 electron groups, 2 bonding groups, 109.5° bond angle.