General Chemistry I: Comprehensive Study Notes and Learning Outcomes

Safety, Matter, and Measurement

Elements, Compounds, and Mixtures

Understanding the classification of matter is fundamental in chemistry. Matter can be categorized as elements, compounds, or mixtures based on composition and properties.

• Element: A pure substance consisting of only one type of atom (e.g., O2, Fe).

• Compound: A substance formed from two or more elements chemically bonded in fixed proportions (e.g., H2O, NaCl).

• Mixture: A physical combination of two or more substances where each retains its own properties (e.g., air, saltwater).

Example: Air is a mixture, water is a compound, and gold is an element.

Density Calculations

Density is a physical property defined as mass per unit volume.

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

• Units: g/cm3 or kg/L

Example: If a block has a mass of 16.3 g and a volume of 2.5 cm3, its density is $\frac{16.3\,g}{2.5\,cm^3} = 6.52\,g/cm^3$.

Significant Figures

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

• All nonzero digits are significant.

• Zeros between nonzero digits are significant.

• Leading zeros are not significant; trailing zeros are significant only if there is a decimal point.

Example: 0.00340 has three significant figures.

Dimensional Analysis

Dimensional analysis is a method for converting units using conversion factors.

• Set up conversion factors so units cancel appropriately.

• Always check that the final units are correct.

Example: To convert 5.0 inches to centimeters: $5.0\,in \times \frac{2.54\,cm}{1\,in} = 12.7\,cm$.

Atoms, Molecules, and Ions

Atomic Structure and Isotopes

Atoms consist of protons, neutrons, and electrons. Isotopes are atoms of the same element with different numbers of neutrons.

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

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

• Isotope Notation: $^{A}_{Z}\text{Element Symbol}$

Example: $^{14}_{6}\text{C}$ is a carbon isotope with 6 protons and 8 neutrons.

Average Atomic Mass

The average atomic mass is the weighted average of all naturally occurring isotopes of an element.

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

Formulas and Nomenclature

Chemical formulas represent the composition of compounds. Nomenclature rules are used to name ionic and covalent compounds.

• Ionic Compounds: Name the cation first, then the anion (e.g., NaCl: sodium chloride).

• Covalent Compounds: Use prefixes to indicate the number of atoms (e.g., CO2: carbon dioxide).

Example: KNO3 is potassium nitrate; N2O3 is dinitrogen trioxide.

Common Polyatomic Ions

Name	Formula and Charge
Ammonium	NH4+
Sulfite	SO32-
Dichromate	Cr2O72-
Hydroxide	OH-
Phosphate	PO43-

Chemical Reactions and Stoichiometry

Types of Chemical Reactions

Chemical reactions can be classified into several types based on the rearrangement of atoms and molecules.

• Synthesis: $A + B \rightarrow AB$

• Decomposition: $AB \rightarrow A + B$

• Single Replacement: $A + BC \rightarrow AC + B$

• Double Replacement: $AB + CD \rightarrow AD + CB$

• Combustion: $C_xH_y + O_2 \rightarrow CO_2 + H_2O$

• Neutralization: $HA + BOH \rightarrow BA + H_2O$

Balancing Chemical Equations

Balancing ensures the law of conservation of mass is obeyed.

• Adjust coefficients to have the same number of each atom on both sides.

• Never change subscripts in chemical formulas.

Example: $CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O$

Stoichiometry and Mole Conversions

Stoichiometry involves quantitative relationships between reactants and products in a chemical reaction.

• Use balanced equations to relate moles of reactants to moles of products.

• Convert between grams, moles, and molecules using molar mass and Avogadro's number ($6.022 \times 10^{23}$ mol-1).

Example: For $2AlCl_3 + 3Na_2CO_3 \rightarrow Al_2(CO_3)_3 + 6NaCl$, 0.25 mol AlCl3 produces $0.25 \times \frac{6}{2} = 0.75$ mol NaCl.

Acids, Bases, and Solutions

Auto-Ionization of Water

Water can self-ionize to form hydronium (H3O+) and hydroxide (OH-) ions.

• Equation: $2H_2O \rightleftharpoons H_3O^+ + OH^-$

• The equilibrium constant for water at 25°C is $K_w = 1.0 \times 10^{-14}$.

Acid-Base Neutralization

Acids donate protons (H+), bases accept protons. Neutralization forms water and a salt.

• Example: $HCl + NaOH \rightarrow NaCl + H_2O$

Calculating Concentrations

Concentration is often expressed as molarity (M), the number of moles of solute per liter of solution.

• Formula: $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

Example: 2.0 g glucose (molar mass = 180 g/mol) in 0.5 L solution: $\frac{2.0\,g}{180\,g/mol} = 0.0111\,mol$, $M = \frac{0.0111}{0.5} = 0.0222\,M$.

Thermochemistry

Heat, Work, and Energy

Thermochemistry studies energy changes in chemical reactions, especially heat transfer.

• First Law of Thermodynamics: $\Delta E = q + w$ (change in internal energy equals heat plus work)

• Enthalpy Change ($\Delta H$): Heat absorbed or released at constant pressure.

• Endothermic: Absorbs heat ($\Delta H > 0$); Exothermic: Releases heat ($\Delta H < 0$).

Example: Combustion reactions are exothermic.

Calorimetry

Calorimetry measures heat flow in chemical reactions.

• Formula: $q = m c \Delta T$

• Where $q$ = heat, $m$ = mass, $c$ = specific heat, $\Delta T$ = temperature change.

Molecular Models and Bonding

Lewis Structures

Lewis structures represent valence electrons as dots and show bonding between atoms.

• Follow the octet rule for main group elements.

• Count total valence electrons, arrange atoms, and distribute electrons to satisfy octets.

Example: Lewis structure for ammonia (NH3):

H | H-N-H

Resonance Structures

Some molecules have multiple valid Lewis structures, called resonance forms.

• Resonance stabilizes the molecule by delocalizing electrons.

Example: The carbonate ion (CO32-) has three resonance structures.

VSEPR Theory and Molecular Geometry

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shapes based on electron pair repulsion.

• Linear: 2 electron groups (e.g., CO2)

• Trigonal Planar: 3 electron groups (e.g., BF3)

• Tetrahedral: 4 electron groups (e.g., CH4)

Periodic Properties

Trends in the Periodic Table

Periodic trends help predict element properties.

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

• Ionization Energy: Decreases down a group, increases across a period.

• Electronegativity: Decreases down a group, increases across a period.

Gases and Gas Laws

Properties of Gases

Gases have unique properties due to the large distances between particles.

• Compressible, expand to fill containers, low density.

Gas Laws

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

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

• Avogadro's Law: $\frac{V_1}{n_1} = \frac{V_2}{n_2}$ (at constant P and T)

• Ideal Gas Law: $PV = nRT$

Where P = pressure, V = volume, n = moles, R = gas constant (0.08206 L·atm·mol-1·K-1), T = temperature in Kelvin.

Kinetic Molecular Theory

Explains gas behavior based on particle motion.

• Gas particles are in constant, random motion.

• Collisions are elastic; average kinetic energy is proportional to temperature.

Formula: $KE_{avg} = \frac{3}{2}kT$

Solutions and Concentrations

Molarity and Titrations

Molarity (M) is used to express solution concentration. Titration is a technique to determine unknown concentrations.

• Formula: $M_1V_1 = M_2V_2$ (for dilution and titration calculations)

Example: To prepare 250 mL of 0.1 M NaCl from 1.0 M stock: $0.1 \times 0.250 = 1.0 \times V_2 \Rightarrow V_2 = 0.025$ L = 25 mL stock solution.

Additional info:

• Some content inferred and expanded for completeness, such as detailed explanations of periodic trends, VSEPR theory, and calorimetry.

• Tables and equations have been formatted for clarity and study purposes.