General Chemistry IB Final Exam Study Guide: Key Concepts and Applications

Atoms, Elements, and Isotopes

Atomic Structure and Isotopes

Atoms are the fundamental units of matter, composed 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; defines the element.

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

• Isotopes: Atoms with the same atomic number but different mass numbers.

• Example: Carbon-12 and Carbon-14 are isotopes of carbon.

Formula:

$A = Z + N$

Additional info: Isotopic abundance is used to calculate average atomic mass.

Molecules, Compounds, and Chemical Bonding

Chemical Formulas and Bond Types

Chemical compounds are formed by atoms bonded together. The type of bond (ionic, covalent, metallic) determines the properties of the compound.

• Molecular Formula: Shows the exact number of atoms of each element in a molecule.

• Ionic Compounds: Formed from metals and nonmetals; involve transfer of electrons.

• Covalent Compounds: Formed from nonmetals; involve sharing of electrons.

• Example: $NaCl$ (ionic), $H_2O$ (covalent)

Chemical Reactions and Equations

Balancing and Types of Reactions

Chemical reactions involve the rearrangement of atoms to form new substances. Equations must be balanced to obey the law of conservation of mass.

• Types of Reactions: Synthesis, decomposition, single replacement, double replacement, combustion.

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

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

Stoichiometry and Chemical Quantities

Mole Concept and Calculations

Stoichiometry uses the mole concept to relate quantities of reactants and products in chemical reactions.

• Mole: $6.022 \times 10^{23}$ particles (Avogadro's number).

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

• Stoichiometric Calculations: Use balanced equations to convert between mass, moles, and number of particles.

• Example: How many moles of $K_2SO_4$ are produced from a given amount of $KMnO_4$?

Formula:

$\text{Moles} = \frac{\text{Mass (g)}}{\text{Molar Mass (g/mol)}}$

Thermochemistry

Enthalpy and Calorimetry

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

• Enthalpy ($\Delta H$): Heat content of a system at constant pressure.

• Calorimetry: Measurement of heat changes using a calorimeter.

• Standard Enthalpy of Formation ($\Delta H_f^\circ$): Enthalpy change for forming 1 mole of a compound from its elements.

• Example: Calculating $\Delta H_{rxn}$ using standard enthalpies.

Formula:

$\Delta H_{rxn} = \sum \Delta H_f^\circ (\text{products}) - \sum \Delta H_f^\circ (\text{reactants})$

Quantum Mechanics and Atomic Theory

Quantum Numbers and Electron Configuration

Quantum mechanics describes the behavior of electrons in atoms using quantum numbers.

• Principal Quantum Number (n): Energy level.

• Angular Momentum Quantum Number (l): Subshell (s, p, d, f).

• Magnetic Quantum Number ($m_l$): Orientation of orbital.

• Spin Quantum Number ($m_s$): Electron spin (+1/2 or -1/2).

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

• Example: Electron configuration of $O$: $1s^2 2s^2 2p^4$

Periodic Properties of the Elements

Trends in the Periodic Table

The periodic table organizes elements by increasing atomic number and reveals periodic trends.

• Atomic Radius: Decreases across a period, increases down a group.

• Ionization Energy: Increases across a period, decreases down a group.

• Electronegativity: Tendency to attract electrons; increases across a period.

• Example: Order elements by increasing atomic radius or ionization energy.

Chemical Bonding and Molecular Structure

Lewis Structures and VSEPR Theory

Lewis structures represent the arrangement of electrons in molecules. VSEPR theory predicts molecular shapes based on electron pair repulsion.

• Lewis Structure: Shows bonding and lone pairs.

• VSEPR: Valence Shell Electron Pair Repulsion; determines geometry (linear, trigonal planar, tetrahedral, etc.).

• Example: Draw the Lewis structure for $SO_2$ and determine its shape.

Liquids, Solids, and Intermolecular Forces

Types of Intermolecular Forces

Intermolecular forces determine the physical properties of substances.

• London Dispersion Forces: Present in all molecules; weakest.

• Dipole-Dipole Interactions: Between polar molecules.

• Hydrogen Bonding: Strong dipole interaction involving H and N, O, or F.

• Example: Identify which compounds exhibit hydrogen bonding.

Solutions and Concentrations

Calculating Molarity and Dilutions

Solutions are homogeneous mixtures; concentration is often expressed as molarity (mol/L).

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

• Dilution: $M_1V_1 = M_2V_2$

• Example: Calculate the volume of solution needed to dilute a stock solution.

Gases and Gas Laws

Ideal Gas Law and Applications

Gases are described by pressure, volume, temperature, and amount. The ideal gas law relates these variables.

• Ideal Gas Law: $PV = nRT$

• Variables: P = pressure, V = volume, n = moles, R = gas constant, T = temperature (K).

• Example: Calculate the volume of a gas at given pressure and temperature.

Chemical Kinetics and Equilibrium

Reaction Rates and Equilibrium

Chemical kinetics studies the speed of reactions; equilibrium describes the balance between forward and reverse reactions.

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

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

• Example: Calculate equilibrium concentrations.

Acids, Bases, and Aqueous Equilibria

pH, pOH, and Acid-Base Reactions

Acids donate protons; bases accept protons. pH measures the acidity of a solution.

• pH: $pH = -\log[H^+]$

• pOH: $pOH = -\log[OH^-]$

• Neutralization: Acid + Base $\rightarrow$ Salt + Water

• Example: Calculate pH from $[H^+]$ or $[OH^-]$.

Thermodynamics and Free Energy

Spontaneity and Gibbs Free Energy

Thermodynamics predicts whether reactions are spontaneous using Gibbs free energy.

• Gibbs Free Energy ($\Delta G$): $\Delta G = \Delta H - T\Delta S$

• Spontaneous Reaction: $\Delta G < 0$

• Example: Determine spontaneity from $\Delta H$ and $\Delta S$ values.

Electrochemistry

Redox Reactions and Cell Potentials

Electrochemistry studies electron transfer in redox reactions and the generation of electrical energy.

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Cell Potential ($E_{cell}$): $E_{cell} = E_{cathode} - E_{anode}$

• Example: Calculate cell potential for a galvanic cell.

Radioactivity and Nuclear Chemistry

Types of Radioactive Decay

Nuclear chemistry involves changes in the nucleus, including radioactive decay.

• Alpha Decay: Emission of $^4_2He$ nucleus.

• Beta Decay: Emission of electron ($\beta^-$) or positron ($\beta^+$).

• Gamma Decay: Emission of high-energy photons.

• Example: Write a nuclear equation for the alpha decay of $^{241}Am$.

Organic Chemistry and Functional Groups

Naming and Identifying Organic Compounds

Organic chemistry focuses on carbon-containing compounds and their functional groups.

• Functional Groups: Alcohol, amine, carboxyl, aldehyde, etc.

• IUPAC Naming: Systematic method for naming organic molecules.

• Example: Identify the functional group in acetaminophen; name branched alkanes.

Lab Techniques and Calculations

Density, Molarity, and Solution Preparation

Laboratory techniques involve accurate measurement and calculation of chemical quantities.

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

• Preparing Solutions: Use molarity and volume to prepare desired concentrations.

• Example: Calculate the density of a solution or prepare a dilution.

Sample Table: Bond Energies

Bond energies are used to estimate the enthalpy change of reactions.

Bond	Energy (kJ/mol)
H-H	436
O=O	498
H-O	463
C-H	414
C-C	348
C=O	799
Additional info: Values may vary slightly depending on source.

Additional info:

• Some questions reference specific chemical reactions, calculations, and molecular structures. For detailed practice, refer to textbook chapters on stoichiometry, thermochemistry, atomic theory, and organic chemistry nomenclature.

• Quantum mechanics and periodic trends are foundational for understanding chemical properties and reactivity.

• Lab calculations (density, molarity, solution preparation) are essential for practical chemistry skills.