MCAT General Chemistry Review

Mass Number and Atomic Number

• Mass number refers to the total number of protons and neutrons in an atom's nucleus.
• Atomic number denotes the number of protons in an atom's nucleus, determining the element's identity.

Scientist Contributions

• Different scientists including Dalton, Thomson, Rutherford, and Bohr made foundational discoveries in atomic theory and structure.
• Contributions include the development of the periodic table and the understanding of chemical bonds.

Light Energy and Constants

• Light exhibits wave-particle duality, demonstrating both wave-like and particle-like properties.
• Key constants in light energy include the speed of light (c) and Planck's constant (h).

AHED

• Refers to Acid, Hydrogen, Electron donor, which are key concepts in understanding acids and bases.

Diamagnetic vs. Paramagnetic

• Diamagnetic substances have no unpaired electrons and are not attracted to magnetic fields.
• Paramagnetic substances contain unpaired electrons and are attracted to magnetic fields.

Quantum Numbers

• Four quantum numbers describe electron positions: principal (n), azimuthal (l), magnetic (m_l), and spin (m_s).
• They determine electron configurations and energy levels in atomic orbitals.

3D Shapes of s, p, d, f Orbitals

• s orbitals are spherical, p orbitals are dumbbell-shaped, while d and f orbitals have more complex shapes.
• Each orbital shape influences how electrons interact in chemical reactions.

Atomic Orbitals on the Periodic Table

• Atomic orbitals correspond to the different blocks of the periodic table: s, p, d, and f blocks indicate the types of orbitals being filled.

The Aufbau Principle

• The Aufbau Principle states that electrons fill orbitals starting from the lowest energy level to the highest.

Periodic Table

• Organizes elements based on increasing atomic number and recurring properties.
• Groups signify elements with similar chemical behaviors, while periods indicate energy levels.

Periodic Table Trends

• Key trends include atomic radius, ionization energy, electron affinity, and electronegativity, influencing how elements interact chemically.

Covalent Bonds

• Formed when two atoms share electrons, typically between nonmetals.
• Strength and length of covalent bonds vary with the number of shared electron pairs.

Bond Type According to Electronegativity (EN)

• Differences in electronegativity determine bond type:
  • Nonpolar covalent: EN difference less than 0.5
  • Polar covalent: EN difference 0.5 to 1.6
  • Ionic: EN difference greater than 1.6.

Ionic Bonds

• Formed through the transfer of electrons from one atom to another, resulting in charged ions.
• Typically occur between metals and nonmetals.

Intermolecular Forces

• Include dipole-dipole interactions, hydrogen bonds, and London dispersion forces, impacting boiling points and melting points.

Sigma and Pi Bonds

• Sigma bonds are formed by head-on overlapping of orbitals, found in all single bonds.
• Pi bonds result from lateral overlapping, present in double and triple bonds.

Formal Charge

• Calculation method to determine charge distribution in molecules.
• The goal is to have formal charges as close to zero as possible for stability.

H-Bond Acceptor and Donor

• Hydrogen bonds form between an H atom covalently bonded to a highly electronegative atom and another electronegative atom.
• H-bond donors provide the hydrogen, while H-bond acceptors have lone pairs to form the bond.

Valence Shell Electron Pair Repulsion Theory (VSEPR)

• Predicts the 3D shapes of molecules based on repulsion between electron pairs surrounding a central atom.

Electron and Molecular Geometry

• Electron geometry considers all electron pairs, while molecular geometry focuses only on atoms.
• Shapes influence molecular properties and reactivity.

Equivalents and Normality

• Equivalents measure the reactive capacity of a substance, particularly in acid-base reactions.
• Normality is a concentration expression based on the equivalent, useful in titrations.

Types of Reactions

• Common reaction types include synthesis, decomposition, single replacement, double replacement, and combustion.

Chemical Kinetics

• The study of reaction rates and factors affecting them, such as temperature, concentration, and catalysts.

Zero, First, and Second Order Reactions

• Zero-order: Rate is constant regardless of reactant concentration.
• First-order: Rate is directly proportional to one reactant's concentration.
• Second-order: Rate is proportional to either the square of one concentration or the product of two concentrations.

Rate Law Equations

• Express the relationship between the rate of a chemical reaction and the concentration of its reactants.

Gibbs Free Energy

• A thermodynamic quantity that indicates the spontaneity of a reaction; negative ΔG indicates a spontaneous process.

Reaction Order and Michaelis-Menten Curve

• Reaction order shows how the rate is affected by the concentration of reactants.
• The Michaelis-Menten equation describes enzyme kinetics in relation to substrate concentration.

Reaction Mechanisms

• Detailed steps of how a reaction occurs, often involving elementary steps and intermediates.

Arrhenius Equation

• Relates the rate constant of a reaction to temperature, activation energy, and the frequency factor.

Equilibrium Constant

• Represents the ratio of products to reactants at equilibrium, reflecting the position of equilibrium in a reversible reaction.

Reaction Quotient

• A measure of the relative concentrations of products and reactants at any point in a reaction, used to predict directionality.

Kinetic (Ea) and Thermodynamic (ΔG) Control

• Kinetic control relates to activation energy's influence on reaction pathways, while thermodynamic control focuses on stable product formation.

Le-Chatelier's Principle

• States that if a system at equilibrium experiences a change in conditions, the system shifts to counteract that change.

Systems and Processes

• A system refers to the part of the universe being studied, while the process describes the transformations occurring within the system.

State and State Functions

• State functions, such as temperature, pressure, and volume, depend only on the current state of the system, not how it reached that state.

Gibbs Free Energy (G)

• Indicates the maximum reversible work performed by a thermodynamic system at constant temperature and pressure.

Gibbs Free Energy Equation

• ΔG = ΔH - TΔS relates changes in enthalpy (ΔH), temperature (T), and entropy (ΔS) to spontaneity.

Temperature and Heat (q)

• Temperature measures the average kinetic energy, while heat (q) is energy transferred between systems due to temperature difference.

Enthalpy (H)

• A measure of total energy in a system, including internal energy and the energy required to make space for it by displacing its environment.

Entropy (S)

• A measure of disorder or randomness in a system; higher entropy values indicate greater disorder and spontaneity.

Ideal Gases

• Follow the ideal gas laws, where gases behave predictably under certain conditions of low pressure and high temperature.

Ideal Gas Law

• PV = nRT relates pressure (P), volume (V), number of moles (n), the ideal gas constant (R), and temperature (T).

Other Gas Laws

• Include Boyle's law, Charles's law, and Avogadro's law, describing relationships between pressure, volume, and temperature.

Diatomic Gases

• Gases composed of two atoms, such as H₂, O₂, N₂, and Cl₂, following specific properties as gases.

Real Gases

• Deviate from ideal behavior under high pressures and low temperatures; interactions between molecules become significant.

Kinetic Molecular Theory

• Describes the behavior of gas molecules in terms of motion, kinetic energy, and collision theory.

Solutions Terminology

• Important terms include solute (substance dissolved) and solvent (substance doing the dissolving).

Concentration

• Measures the amount of solute in a given volume of solution; expressed in molarity, molality, or percent composition.

Solutions Equilibria

• Concerns the dynamic balance in a saturated solution where the rate of dissolving equals the rate of precipitation.

Solubility Rules

• Guidelines predicting the solubility of compounds in water; essential for understanding reactions and equilibria.

Colligative Properties

• Properties that depend on the number of solute particles in a solution, such as boiling point elevation and freezing point depression.

Acids and Bases Definitions

• Acids donate protons (H⁺), while bases accept protons, leading to various chemical reactions and properties.

Acids and Bases Properties

• Acids usually taste sour and turn blue litmus red, while bases taste bitter and turn red litmus blue.

Buffers

• Solutions that resist changes in pH when small amounts of acid or base are added; important for maintaining physiological pH levels.

Polyvalence and Normality

• Polyvalent acids and bases can donate or accept more than one proton; normality accounts for this reactivity in solution.

Titrations

• Analytical technique to determine concentration of a solute by reacting it with a solution of known concentration.

Titration Setup

• Includes burette, pipette, and flask to carry out the titration with precise measurement.

Titration Curve

• Graphical representation of pH changes during a titration, revealing information about acid-base strength and equivalence points.

Oxidation Reduction Definitions

• Redox reactions involve the transfer of electrons; oxidation refers to loss, while reduction

Prepare for the MCAT with our review sheets focused on general chemistry topics. This quiz covers essential concepts such as mass number, atomic number, and light energy, as well as key scientific contributions and quantum numbers. Hone your understanding of magnetic properties and electron configurations with this comprehensive tool.