Final Exam Review
Focus Guide Final Exam
Density
Definition: Density is defined as mass per unit volume.
Formula:
Units: Typically measured in grams per cubic centimeter (g/cm³) or kilograms per cubic meter (kg/m³).
General Naming Rules
Compounds: The systematic naming often follows the IUPAC rules based on the composition of elements.
Ionic Compounds: Name the metal first followed by the non-metal with an '-ide' ending.
Covalent Compounds: Use prefixes to indicate the number of atoms (e.g., CO = carbon monoxide, CO₂ = carbon dioxide).
Difference between Element, Compound, Heterogeneous, and Homogeneous Mixtures
Element: A pure substance that cannot be broken down into simpler substances (e.g., Iron (Fe)).
Compound: A substance formed when two or more elements are chemically bonded (e.g., Water (H₂O)).
Heterogeneous Mixtures: Mixtures that are not uniform and have distinct components (e.g., salad).
Homogeneous Mixtures: Mixtures that are uniform throughout (e.g., salt water).
Balancing Equations
Importance: Balancing chemical equations ensures that the law of conservation of mass is adhered to, where the quantity of each element is conserved.
Method: Adjust coefficients to balance each side of the equation.
Percent Mass
Definition: The percent composition of an element in a compound; calculated using the formula:
Simple Stoichiometry
Definition: Stoichiometry involves using balanced chemical equations to calculate the relationships between reactants and products.
Example: If 2H₂ + O₂ → 2H₂O, stoichiometry can be used to determine how much water is produced when a certain amount of hydrogen reacts.
Net Ionic Equations
Definition: A chemical equation that shows only the species that participate in the reaction, omitting spectator ions.
Importance: Helps simplify reactions in solutions and showcases the actual chemical change.
General Characteristics of Reactions
Synthesis Reaction: A + B → AB; two or more substances combine to form one product.
Decomposition Reaction: AB → A + B; one compound breaks down into two or more simpler substances.
Metathesis (Double Replacement) Reaction: AB + CD → AD + CB; involves the exchange of partners.
Neutralization Reaction: An acid reacts with a base to yield water and a salt.
Titration Formula: where M = molarity, V = volume.
Combustion Reaction: A substance combines with oxygen, releasing energy, primarily producing CO₂ and H₂O.
REDOX Reactions: Reactions involving the transfer of electrons; reduction involves gain of electrons, while oxidation involves loss.
Determining Protons, Neutrons, and Electrons in an Atom
Protons: The number of protons in an atom defines the atomic number.
Neutrons: The number of neutrons can be calculated by subtracting the atomic number from the atomic mass (rounded).
Electrons: In a neutral atom, the number of electrons equals the number of protons.
Groups and Trends on the Periodic Table
Groups: Vertical columns in the periodic table where elements share similar properties.
Periodic Trends: Includes atomic radius, ionization energy, electron affinity, and electronegativity trends across periods and down groups.
Pauli Exclusion Principle, Aufbau Principle, Hund’s Rule
Pauli Exclusion Principle: No two electrons in an atom can have the same set of quantum numbers.
Aufbau Principle: Electrons are added to the lowest energy orbitals first.
Hund’s Rule: Electrons will fill degenerate orbitals singly before pairing begins.
Radius and Isoelectronic Ions
Atomic Radius: The distance from the nucleus to the outer boundary of the electron cloud; it generally decreases across a period and increases down a group.
Isoelectronic Ions: Ions that have the same number of electrons (e.g., Na⁺ and Mg²⁺ both have 10 electrons).
Paramagnetic vs. Diamagnetic
Paramagnetic: A substance with unpaired electrons that will be attracted to a magnetic field.
Diamagnetic: A substance with all paired electrons that will be repelled by a magnetic field.
JJ Thomson’s Cathode Ray Tube, Millikan’s Oil Drop Experiment, Rutherford’s Experiment
JJ Thomson’s Cathode Ray Tube: Demonstrated the existence of electrons as negative charged particles.
Millikan’s Oil Drop Experiment: Measured the charge of the electron, leading to the determination of its mass.
Rutherford’s Experiment: Concluded the nucleus is small, dense, and positively charged based on the gold foil experiment results.
Endothermic and Exothermic Principles
Endothermic Reaction: A reaction that absorbs heat from the surroundings.
Exothermic Reaction: A reaction that releases heat to the surroundings.
Heat Formulas
Heat Transfer: Calculated using the equation:
where ( q ) is heat absorbed or released, ( m ) is mass, ( C ) is specific heat capacity, and ( \Delta T ) is the change in temperature.
Specific Heat and Mass Effects on Energy Transfer
Specific Heat: The amount of heat required to raise the temperature of one gram of a substance by one degree Celsius. Higher specific heat means slower temperature change.
Mass: Greater mass can absorb more heat without a significant rise in temperature.
Hess’s Law
Definition: States that the total enthalpy change in a reaction is the sum of the changes in the individual steps of the reaction.
Calculation: Can be applied by adding the enthalpy changes of each step to find the overall change.
Molecular Geometry
Definition: The three-dimensional arrangement of atoms in a molecule, which can affect the physical and chemical properties of the substance.
Common Shapes: Linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral.
Hybridization
Definition: The mixing of atomic orbitals to form new hybrid orbitals that are degenerate (of the same energy).
Types: sp, sp², sp³ and their corresponding geometries.
Quantum Numbers
Principal Quantum Number (n): Indicates the energy level and size of the orbital.
Angular Momentum Quantum Number (l): Indicates the shape of the orbital.
Magnetic Quantum Number (ml): Indicates the orientation of the orbital in space.
Spin Quantum Number (ms): Indicates the spin of the electron (either +1/2 or -1/2).
Shapes of Orbitals
s Orbital: Spherical shape, one orbital.
p Orbitals: Dumbbell-shaped, three orbitals (px, py, pz).
d Orbitals: More complex shapes, five orbitals.
f Orbitals: Even more complex shapes, seven orbitals.
Number of Orbitals in Each Sublevel
s Sublevel: 1 orbital
p Sublevel: 3 orbitals
d Sublevel: 5 orbitals
f Sublevel: 7 orbitals
Simple Electron Configurations with Orbital Notation
Example: The electron configuration of Carbon: 1s² 2s² 2p².
Electron Deficient and Electron Abundant Molecules
Electron Deficient: Molecules with fewer total electrons than required for octets (e.g., BF₃).
Electron Abundant: Molecules that have more than enough electrons available, often with lone pairs (e.g., NH₃).
Formal Charge
Definition: A concept used to determine the charge distribution in a molecule. Calculated with the formula:
Combined Gas Law
Formula: , where P = pressure, V = volume, and T = temperature.
Application: Used to calculate changes in state of a gas when temperature, volume, or pressure is altered.
Calculating Partial Pressure Using Mole Fraction
Definition: The partial pressure of a gas in a mixture is calculated using the equation:
where ( X{gas} ) is the mole fraction of the gas.
Kinetic Molecular Theory
Postulates:
Gases consist of a large number of tiny particles that are in constant motion.
The volume of gas particles is negligible compared to the volume of the container.
Gas particles collide elastically, meaning total kinetic energy is conserved.
The average kinetic energy of gas particles is proportional to the temperature of the gas in Kelvin.
Intermolecular Forces
London Dispersion Forces: Weak forces due to temporary dipoles in molecules.
Dipole-Dipole Forces: Attractive forces between polar molecules.
Hydrogen Bonding: Strong type of dipole-dipole interaction experienced when hydrogen is bonded to electronegative atoms like O, N, or F.
Ion-Dipole Forces: Forces between an ion and a polar molecule; crucial in solutions such as salt in water.