General Chemistry Study Guide: Atomic Structure, Chemical Reactions, and Properties of Matter

Electron Configuration

Basics of Electron Configuration

Electron configuration describes the arrangement of electrons in an atom's orbitals. It follows the Aufbau principle, Pauli exclusion principle, and Hund's rule.

• Aufbau Principle: Electrons fill orbitals starting with the lowest energy level first.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of quantum numbers.

• Hund's Rule: Electrons occupy degenerate orbitals singly before pairing.

• Example: Calcium (Ca, Z=20): $1s^2 2s^2 2p^6 3s^2 3p^6 4s^2$

Electron Configuration for Sodium

• Sodium (Na, Z=11): $1s^2 2s^2 2p^6 3s^1$

Atomic Structure

Atomic Notation

Atoms are represented as $^{A}_{Z}X$, where A is the mass number (protons + neutrons), Z is the atomic number (protons), and X is the element symbol.

• Example: An atom with 15 protons, 12 electrons, and 16 neutrons: $^{31}_{15}P^{3+}$

• Charge Calculation: Charge = Number of protons - Number of electrons

Valence Electrons

Determining Valence Electrons

Valence electrons are the electrons in the outermost shell, crucial for chemical bonding.

• Phosphorus (P): 5 valence electrons

• Chlorine (Cl): 7 valence electrons

Dipole Moment

Bond Polarity and Dipole Moment

The dipole moment measures the separation of charge in a bond, indicating bond polarity.

• O-H bond: Dipole moment ≈ 1.4 D

• H-F bond: Dipole moment ≈ 1.9 D

Oxidation State

Assigning Oxidation Numbers

Oxidation state indicates the degree of oxidation of an atom in a compound.

• Fe in Fe2O3: +3

• Cr in K2Cr2O7: +6

Chemical Formula

Writing Chemical Formulas

Chemical formulas represent the composition of compounds using element symbols and subscripts.

• Iron(III) sulfate: Fe2(SO4)3

• Copper(II) nitrate: Cu(NO3)2

Oxidation-Reduction (Redox)

Identifying Redox Processes

Redox reactions involve the transfer of electrons between species.

• Reduction: Gain of electrons

• Oxidation: Loss of electrons

• Example: Zn + Cl2 → ZnCl2: Cl2 is reduced, Zn is oxidized.

Types of Chemical Reactions

Classification of Reactions

Chemical reactions are classified based on the changes in reactants and products.

• Combination: Two or more substances 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.

• Combustion: Reaction with oxygen producing heat and light.

Balancing Chemical Equations

Principles of Balancing

Balancing equations ensures the conservation of mass and atoms.

• Example: $KClO_3 ightarrow KCl + O_2$; Balanced: $2KClO_3 ightarrow 2KCl + 3O_2$

Stoichiometry

Stoichiometric Calculations

Stoichiometry involves quantitative relationships between reactants and products.

• Example: $2Al + 3Cl_2 ightarrow 2AlCl_3$

• Mass-mole conversions: Use molar mass and balanced equations.

Gas Laws

Ideal Gas Law and Related Calculations

Gas laws describe the behavior of gases under varying conditions.

• Ideal Gas Law: $PV = nRT$

• Combined Gas Law: $\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$

Partial Pressure

Dalton's Law of Partial Pressures

The total pressure of a gas mixture is the sum of the partial pressures of each component.

• Partial Pressure: $P_X = X_X \times P_{total}$

• Mole Fraction: $X_X = \frac{n_X}{n_{total}}$

Molar Volume

Volume of Gases at STP

At standard temperature and pressure (STP), 1 mole of an ideal gas occupies 22.4 L.

• Example: 100.0 L of gas at STP contains $\frac{100.0}{22.4} = 4.46$ mol

Properties of Matter

Compressibility and States of Matter

Compressibility varies among solids, liquids, gases, and plasma.

• Gases: Most compressible

• Solids: Least compressible

Intermolecular Forces

Types and Effects

Intermolecular forces determine physical properties like boiling point and cohesion.

• Hydrogen Bonding: Strongest among neutral molecules, responsible for water's high boiling point

• London Dispersion: Weak, present in noble gases

• Dipole-Dipole: Between polar molecules

Electrolytes

Classification of Electrolytes

Electrolytes conduct electricity in solution; classified as strong, weak, or nonelectrolytes.

• Strong Electrolyte: NaCl

• Weak Electrolyte: CH3COOH

• Nonelectrolyte: Most covalent compounds

Solution Stoichiometry

Calculating Solution Concentrations

Solution stoichiometry involves molarity, volume, and mass relationships.

• Molarity (M): $M = \frac{mol\ solute}{L\ solution}$

• Example: Mass of NaOH needed for 0.400 M, 250 mL: $0.400 \times 0.250 \times 40.0 = 4.00$ g

Spectator Ions

Identifying Spectator Ions

Spectator ions do not participate in the net ionic equation.

• Example: In AgNO3 + NaCl → AgCl + NaNO3, Na+ and NO3- are spectator ions.

Thermodynamics

Exothermic and Endothermic Processes

Thermodynamics studies energy changes in chemical reactions.

• Exothermic: Releases energy to surroundings

• Endothermic: Absorbs energy from surroundings

Thermochemistry

Enthalpy Changes

Thermochemistry focuses on heat changes during chemical reactions.

• Standard Enthalpy Change ($\Delta H^\circ$): Calculated from reactants and products

• Example: $2NO(g) + O_2(g) \rightarrow 2NO_2(g)$; $\Delta H^\circ = -114.2$ kJ/mol

Heat Capacity

Calculating Heat Changes

Heat capacity is the amount of heat required to change temperature by 1°C.

• Formula: $q = mC\Delta T$

• Example: $q = 200.0 \times 0.385 \times (75-25)$

Empirical Formulas

Determining Empirical Formulas

Empirical formulas show the simplest whole-number ratio of atoms in a compound.

• Highest mass percentage of oxygen: H2O

Ionic Radii

Trends in Ionic Radii

Ionic radius depends on nuclear charge and electron configuration.

• Smallest ionic radius: Mg2+

• Trend: $Na^+ < Mg^{2+} < Al^{3+}$

Molecular Geometry

Shapes of Molecules

Molecular geometry is determined by the VSEPR theory.

• NO3-: Trigonal planar

• PCl5: Trigonal bipyramidal

Formal Charge

Calculating Formal Charge

Formal charge helps determine the most stable Lewis structure.

• Formula: $FC = V - (L + \frac{1}{2}B)$

• Example: Nitrogen in HN=NH: $FC = 5 - (2 + 3) = 0$

Electrostatic Potential Map

Visualizing Charge Distribution

Electrostatic potential maps show regions of positive and negative charge in molecules.

• Red regions: Most negative potential

• Blue regions: Most positive potential

Electron Transitions

Energy Changes in Atoms

Electron transitions between energy levels involve absorption or emission of energy.

• Absorption: Electron moves to higher energy level

• Emission: Electron drops to lower energy level, photon emitted

Hydrogen Electron Energy Levels

Calculating Wavelengths of Emitted Light

When an electron transitions in hydrogen, the wavelength of emitted light is given by the Rydberg formula:

• Formula: $\frac{1}{\lambda} = R_H \left( \frac{1}{n_1^2} - \frac{1}{n_2^2} \right)$

• Where: $R_H = 1.097 \times 10^7 \text{ m}^{-1}$

• Example: $n_2 = 4$ to $n_1 = 2$