General Chemistry Final Exam Practice: Key Concepts and Applications

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Understanding the tools and units used in chemistry is essential for accurate experimentation and data analysis.

Pressure: Defined as force divided by unit area.
SI Units: The International System of Units includes kilogram (mass), meter (length), second (time), and mole (amount of substance).
Significant Figures: The number of meaningful digits in a measured or calculated quantity. Used to express precision.
Temperature Scales: Kelvin is the SI unit for temperature; absolute zero (0 K) is the lowest possible temperature.
Density: Mass per unit volume.
Measurement Accuracy: For precise measurements, use instruments with appropriate sensitivity and calibration.

Atoms are the fundamental units of matter, composed of protons, neutrons, and electrons.

Atomic Number (Z): Number of protons in the nucleus; determines the element.
Mass Number (A): Sum of protons and neutrons.
Isotopes: Atoms of the same element with different numbers of neutrons.
Electron Configuration: Distribution of electrons among atomic orbitals. Example: For Hg, the ground-state configuration is .
Valence Electrons: Electrons in the outermost shell, important for chemical bonding.
Oxidation Number: Indicates the charge of an atom in a compound; sulfur in is +6.

The periodic table organizes elements by increasing atomic number and recurring chemical properties.

Groups and Periods: Vertical columns are groups (families); horizontal rows are periods.
Alkali Metals: Group 1 elements, e.g., lithium.
Semimetals: Elements with properties intermediate between metals and nonmetals, e.g., germanium.
Atomic Radius: Generally increases down a group and decreases across a period.
Electron Configuration and Periodicity: Determines chemical reactivity and bonding.

Ionic bonds form between metals and nonmetals through the transfer of electrons.

Cations: Positively charged ions formed by loss of electrons (e.g., Na).
Anions: Negatively charged ions formed by gain of electrons (e.g., Cl).
Binary Ionic Compounds: Formed from cations and anions; typically involve elements from groups 1, 2, and 17.
Oxidation States: Used to balance chemical equations and determine electron transfer.

Gases are characterized by their ability to expand and fill containers, and their behavior is described by gas laws.

Ideal Gas Law:
Boyle's Law: (at constant temperature)
Charles's Law: (at constant pressure)
Avogadro's Law: Equal volumes of gases at the same temperature and pressure contain equal numbers of molecules.
Partial Pressure: The pressure exerted by a single gas in a mixture.
Mole Concept: 1 mole = particles.
Gas Stoichiometry: Used to relate volumes, masses, and moles of gases in reactions.

Covalent bonds involve the sharing of electron pairs between atoms, leading to the formation of molecules.

Lewis Structures: Diagrams showing bonding and lone pairs in molecules.
Hybridization: Mixing of atomic orbitals to form new hybrid orbitals. Example: hybridization in tetrahedral geometry.
Molecular Geometry: Determined by VSEPR theory; e.g., is tetrahedral.
Bond Lengths and Energies: Triple bonds are shorter and stronger than double or single bonds.
Polarity: Determined by differences in electronegativity and molecular shape.

Stoichiometry involves quantitative relationships in chemical reactions.

Balancing Equations: Ensures conservation of mass and charge.
Stoichiometric Coefficients: Indicate the relative number of moles of reactants and products.
Molarity: Concentration of a solution,
Limiting Reactant: The reactant that is completely consumed first, limiting the amount of product formed.

Thermochemistry studies the energy changes during chemical reactions.

Enthalpy (): Heat change at constant pressure.
Bond Dissociation Energy: Energy required to break a bond in a molecule.
Calorimetry: Measurement of heat changes using a calorimeter.
Specific Heat (): Amount of heat required to raise the temperature of 1 g of a substance by 1°C.
Example: Calculating the final temperature when a hot metal is placed in water using .

Solutions are homogeneous mixtures of solute and solvent.

Molarity ():
Solubility: The maximum amount of solute that can dissolve in a solvent at a given temperature.
Effect of Temperature: Increasing temperature generally increases solubility of solids and decreases density of solutions.

Accurate laboratory techniques are essential for reliable results.

Measurement: Use appropriate glassware and instruments for volume, mass, and temperature.
Calibration: Ensures accuracy of instruments.
Safety: Follow proper procedures to avoid accidents and contamination.

The periodic table groups elements by similar properties.

Group	Example Elements	Properties
Alkali Metals	Li, Na, K	Highly reactive, soft, form +1 cations
Alkaline Earth Metals	Mg, Ca	Reactive, form +2 cations
Halogens	F, Cl, Br	Reactive nonmetals, form -1 anions
Noble Gases	He, Ne, Ar	Inert, very low reactivity
Semimetals	B, Si, Ge	Intermediate properties

Some questions reference specific chemical reactions, periodic table trends, and molecular structures, which are core topics in general chemistry.
Practice questions cover a broad range of foundational concepts, suitable for exam preparation.