General Chemistry Practice Final Exam Study Guide

General Chemistry Concepts

Units, Conversions, and Significant Figures

Understanding units and conversions is fundamental in chemistry for expressing measurements and performing calculations.

• Unit Conversion: Use dimensional analysis to convert between units (e.g., cm to km, atm to pascals).

• Significant Figures: The number of meaningful digits in a measurement. Rules determine which digits are significant.

• Example: Convert $2.18 \times 10^5$ cm to km: $2.18 \times 10^5 \text{ cm} \times \frac{1 \text{ m}}{100 \text{ cm}} \times \frac{1 \text{ km}}{1000 \text{ m}} = 2.18$ km.

Classification of Matter

Matter can be classified as elements, compounds, or mixtures based on composition.

• Element: Pure substance consisting of one type of atom (e.g., Carbon).

• Compound: Substance composed of two or more elements chemically bonded (e.g., Water).

• Mixture: Physical combination of substances (e.g., Air).

• Example: Classify carbon dioxide (compound), air (mixture), water (compound).

Chemical Formulas and Nomenclature

Chemical compounds are represented by formulas and named according to systematic rules.

• Writing Formulas: Use element symbols and subscripts to indicate the number of atoms (e.g., $\text{CO}_2$ for carbon dioxide).

• Naming Compounds: Follow IUPAC rules for ionic and molecular compounds.

• Example Table:

Compound Name	Formula
Vanadium (IV) Oxide	VO2
Potassium Sulfide	K2S
Barium Carbonate	BaCO3
Diphosphorus Pentoxide	P2O5

Atoms, Elements, and the Periodic Table

Atomic Structure

Atoms consist of protons, neutrons, and electrons. The arrangement of these particles determines the element's properties.

• Proton: Positively charged particle in the nucleus.

• Neutron: Neutral particle in the nucleus.

• Electron: Negatively charged particle in orbitals around the nucleus.

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

• Example: $^{63}\text{Cu}$ and $^{65}\text{Cu}$ are isotopes of copper.

Electron Configuration and Quantum Numbers

Electron configuration describes the arrangement of electrons in an atom. Quantum numbers specify electron properties.

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

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

• Magnetic Quantum Number (ml): Specifies orbital orientation.

• Spin Quantum Number (ms): Specifies electron spin.

• Example: For phosphorus (P), one set of quantum numbers for a valence electron: $n=3$, $l=1$, $m_l=0$, $m_s=+1/2$.

Periodic Trends

Elements show trends in properties across the periodic table.

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

• Ionization Energy: Energy required to remove an electron; increases across a period, decreases down a group.

• Electron Affinity: Energy change when an atom gains an electron.

• Example Table:

Element	Atomic Radius	First Ionization Energy
Na	Large	Low
Cl	Small	High

Chemical Bonding and Molecular Structure

Ionic and Covalent Compounds

Chemical bonds form between atoms to create compounds.

• Ionic Bond: Transfer of electrons from metal to nonmetal, forming ions (e.g., NaCl).

• Covalent Bond: Sharing of electrons between nonmetals (e.g., H2O).

• Lewis Structures: Diagrams showing valence electrons and bonding.

• Example: Draw Lewis structure for SO3 with octet around sulfur.

Intermolecular Forces (IMF)

IMFs are forces between molecules affecting physical properties.

• Types: London dispersion, dipole-dipole, hydrogen bonding.

• Strength Order: Hydrogen bonding > dipole-dipole > London dispersion.

• Example: Water exhibits hydrogen bonding, leading to high boiling point.

Chemical Reactions and Stoichiometry

Balancing Chemical Equations

Chemical equations must be balanced to obey the law of conservation of mass.

• Steps: Count atoms of each element on both sides; adjust coefficients as needed.

• Example: $2\text{HCl} + \text{Ba(OH)}_2 \rightarrow \text{BaCl}_2 + 2\text{H}_2\text{O}$

Stoichiometry

Stoichiometry involves calculations based on balanced equations to determine amounts of reactants and products.

• Mole Concept: 1 mole = $6.022 \times 10^{23}$ particles.

• Mass-Mole Conversions: Use molar mass to convert between grams and moles.

• Example: Calculate moles of HCl needed to neutralize Ba(OH)2.

Gases and Gas Laws

Properties of Gases

Gases have unique properties described by gas laws.

• Pressure (P): Force per unit area exerted by gas particles.

• Volume (V): Space occupied by gas.

• Temperature (T): Average kinetic energy of particles.

• Ideal Gas Law: $PV = nRT$

• Example: Calculate volume of gas at STP: $V = \frac{nRT}{P}$

Gas Law Applications

Gas laws relate pressure, volume, and temperature for gases.

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

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

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

• Example: If pressure is doubled and temperature is doubled, volume remains unchanged.

Solutions and Concentrations

Types of Solutions and Concentration Units

Solutions are homogeneous mixtures; concentration expresses the amount of solute in solvent.

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

• Normality (N): $N = \frac{\text{equivalents of solute}}{\text{liters of solution}}$

• Example: Calculate molarity of NaOH solution prepared by dissolving 17.4 g in 500 mL water.

Thermochemistry and Quantum Mechanics

Energy, Light, and Electromagnetic Radiation

Energy changes accompany chemical reactions and are related to light and quantum mechanics.

• Frequency ($\nu$) and Wavelength ($\lambda$): $c = \lambda \nu$

• Energy of Photon: $E = h\nu$

• Example: Calculate wavelength for frequency $3.22 \times 10^{15}$ s$^{-1}$: $\lambda = \frac{c}{\nu}$

Quantum Numbers and Atomic Orbitals

Quantum numbers describe electron positions and energies in atoms.

• Valid Set: $n, l, m_l, m_s$ must follow quantum rules (e.g., $n=2$, $l=1$, $m_l=0$, $m_s=+1/2$).

• Example: Which set of quantum numbers is valid for a 3p orbital?

Additional Info

• Many questions also cover chemical kinetics, equilibrium, acids and bases, and nuclear chemistry, but these are less emphasized in this particular exam.

• Practice problems include both conceptual and calculation-based questions, reflecting the breadth of a general chemistry curriculum.