June 2026 Comprehensive Chemistry Regents Study Guide

Atomic Structure and the Fundamental Components of Matter

• Subatomic Particles: Atoms represent the smallest units of elements and are composed of three distinct subatomic components:

  • Protons: Carry a positive charge ($1+$) and have a mass of approximately $1\,\text{atomic mass unit (amu)}$. They are located in the nucleus.

  • Neutrons: Carry a neutral charge ($0$) and have a mass of approximately $1\,amu$. They are also located in the nucleus.

  • Electrons: Carry a negative charge ($1-$). Their mass is considered negligible ($\approx 0\,amu$ ). They orbit the nucleus in specific regions called shells or orbitals.

• Nuclear Composition and Atomic Volume:

  • The nucleus is the center of the atom, containing protons and neutrons. Because it consists of protons, the nucleus itself is positively charged. It is extremely small and dense, providing the vast majority of an atom's mass.

  • The electron cloud surrounds the nucleus. While electrons contribute nearly zero mass, they account for almost all of the atom's volume.

  • Atoms are described as being mostly empty space.

• Identity and Charge of Atoms:

  • Neutral Atoms: In a neutral atom, the number of protons is equal to the number of electrons (e.g., Oxygen has an atomic number of $8$, meaning it has $8$ protons and $8$ electrons if neutral).

  • Atomic Number: This value is defined solely by the number of protons in the nucleus. Changing the number of protons converts the atom into a new element.

  • Mass Number: The total sum of protons and neutrons ($\text{Mass Number} = \text{Protons} + \text{Neutrons}$).

Ions, Isotopes, and Average Atomic Mass

• Ions: Formed when there is a change in the number of electrons.

  • Cation: A positively charged ion formed by losing electrons.

  • Anion: A negatively charged ion formed by gaining electrons.

• Isotopes: Atoms of the same element that have the same number of protons but a different number of neutrons. This leads to atoms with different mass numbers.

  • Example: Carbon-13 ($6$ protons, $7$ neutrons) versus Carbon-14 ($6$ protons, $8$ neutrons).

  • Isotopic Notation: Written as $\text{C-14}$ or $\text{}_{6}^{14}\text{C}$, where the upper number ($14$) is the mass number and the lower number ($6$) is the atomic number.

• Average Atomic Mass Calculation: This is the weighted average of all naturally occurring isotopes of an element based on their abundance.

  • Example Scenario: If Carbon-13 ($20\,g/mol$) is $3\%$ abundant and Carbon-14 ($22\,g/mol$) is $97\%$ abundant, the calculation is:     $\text{Average Atomic Mass} = (0.03 \times 20\,g/mol) + (0.97 \times 22\,g/mol)$

Electron Energy Levels and the Electromagnetic Spectrum

• Energy States:

  • Ground State: The stable, relaxed state where all electrons occupy the lowest possible energy levels.

  • Excited State: Formed when an electron absorbs energy (often in the form of a photon) and is promoted to a higher energy orbital further from the nucleus.

  • Energy Transition: Moving from a lower to higher shell requires absorbing energy. Moving from a higher shell back to a lower shell (relaxation) results in the release of energy, often as visible light.

• Electromagnetic Spectrum Wave Properties:

  • Wavelength ($\lambda$): The distance between two consecutive peaks or troughs.

  • Frequency ($f$): The number of cycles passing a point per second.

  • Mathematical Relationships:

    • As wavelength decreases, energy and frequency increase ($\lambda \propto \frac{1}{E}$).

    • Energy and frequency are directly proportional ($E \propto f$).

  • Visible Light: Only a very small portion of the electromagnetic spectrum is visible to the human eye. Most waves (like Gamma rays) are invisible.

• Valence Electrons: Electrons located in the outermost shell of an atom.

  • The group number on the periodic table (dropping the leading '1' for groups 13-18) indicates the number of valence electrons.

  • Valence electrons determine the chemical reactivity of the element.

Periodic Table Organization and Trends

• Layout:

  • Arranged by increasing atomic number.

  • Metals: Located to the left of the bolded "staircase" line. They typically lose electrons to form cations. Properties include luster and conductivity.

  • Non-metals: Located to the right of the staircase. They typically gain electrons to form anions.

  • Metalloids: Elements touching the staircase line (e.g., Boron, Silicon) that possess properties of both metals and non-metals. Aluminum is an exception and is considered a metal.

• Families (Groups):

  • Group 1: Alkali Metals (highly reactive, $1$ valence electron).

  • Group 2: Alkaline Earth Metals ($2$ valence electrons).

  • Groups 3-12: Transition Metals (form colored ions, participate in metallic bonding with a "sea of electrons").

  • Group 17: Halogens (highly reactive non-metals, $7$ valence electrons).

  • Group 18: Noble Gases (unreactive, complete octet of $8$ valence electrons, except Helium which has $2$).

• Periodic Trends:

  • Across a Period (Left to Right):

    • Atomic Radius Decreases: Increased nuclear charge (more protons) pulls electrons closer.

    • Electronegativity Increases: Stronger attraction for electrons in a bond.

    • Ionization Energy Increases: More energy is required to remove an electron from a shell held tightly by a positive nucleus.

  • Down a Group (Top to Bottom):

    • Atomic Radius Increases: Addition of new principal energy levels (shells).

    • Electronegativity Decreases: Outer shells are further from the nucleus, weakening the pull.

    • Ionization Energy Decreases: Valence electrons are shielded and further away, making them easier to remove.

• Metallic and Non-metallic Character:

  • Metallic Character: The tendency to lose electrons. Increases as you go down a group (Francium is the most metallic).

  • Non-metallic Character: The tendency to gain electrons in reactions.

Chemical Bonding and Molecular Geometry

• Ionic Bonding: Occurs between a metal and a non-metal via the complete transfer of electrons. This creates oppositely charged ions ($\text{Cation}+$ and $\text{Anion}-$) that attract each other.

  • Example: $\text{NaCl}$. Sodium ($\text{Na}$) loses its $1$ valence electron ($\text{Na}^{+}$) and Chlorine ($\text{Cl}$) gains it ($\text{Cl}^{-}$).

• Covalent Bonding: Occurs between two non-metals via the sharing of electrons. These form molecular compounds.

  • Single Bond: $2$ electrons shared.

  • Double Bond: $4$ electrons shared.

  • Triple Bond: $6$ electrons shared (e.g., Diatomic Nitrogen, $\text{N}_2$).

  • S = N - A Rule: Shared electrons ($S$) equals Needed electrons ($N$) minus Available electrons ($A$).

• Bond Polarity:

  • Polar Covalent Bond: Unequal sharing of electrons due to elective negativity differences (DEN > 1.5 often indicates ionic, but polar bonds exist between $0.5$ and $1.5$). One atom is partially negative ($\delta-$) and one is partially positive ($\delta+$).

  • Non-polar Covalent Bond: Equal sharing of electrons (e.g., $\text{C-H}$ bonds or $\text{O}_2$).

• Molecular Polarity and Symmetry:

  • A molecule can have polar bonds but be non-polar overall if it is symmetrical (e.g., $\text{CO}_2$). The dipoles cancel out like a tug-of-war with equal strength on both sides.

  • Asymmetrical molecules with polar bonds are polar overall (e.g., $\text{H}_2\text{O}$).

Intermolecular Forces (IMF)

• Intermolecular Forces: Forces that act between molecules, not within them.

  • Hydrogen Bonding: The strongest IMF. Occurs when Hydrogen is bound to highly electronegative Nitrogen, Oxygen, or Fluorine ($\text{N}$, $\text{O}$, $\text{F}$).

  • Dipole-Dipole Attractions: Attractions between the partially positive end of one polar molecule and the partially negative end of another.

  • London Dispersion Forces: The weakest IMF. Temporary dipoles occurring in non-polar molecules (e.g., $\text{CH}_4$).

• Impact: Stronger IMFs result in higher boiling points because more energy is required to pull the molecules apart.

Chemical Nomenclature and Formula Writing

• Ionic Compounds:

  • Name the metal first, followed by the non-metal with the suffix "-ide" (e.g., Sodium Chloride).

  • Multivalent Metals: Transition metals use Roman Numerals to indicate the charge of the metal ion (e.g., Palladium (II) Oxide vs. Palladium (IV) Oxide).

  • Polyatomic Ions: Groups of atoms that act as a single ion with a special name (e.g., Ammonium $\text{NH}_4^{+}$ or Chlorate $\text{ClO}_3^{-}$). Found on the reference table.

• Covalent Compounds:

  • Use Greek prefixes to indicate the number of atoms (mono-, di-, tri-, tetra-, etc.).

  • Example: $\text{CO}_2$ is Carbon Dioxide; $\text{CO}$ is Carbon Monoxide.

Chemical Reactions and Balancing Equations

• Types of Reactions:

  1. Synthesis: Multiple reactants combine into one ($A + B \rightarrow C$).

  2. Decomposition: One reactant breaks into multiple products ($C \rightarrow A + B$).

  3. Single Replacement: One element replaces another in a compound ($\text{AB} + \text{C} \rightarrow \text{AC} + \text{B}$).

  4. Double Replacement: Two compounds exchange ions ($\text{AB} + \text{CD} \rightarrow \text{AC} + \text{BD}$).

  5. Combustion: A hydrocarbon reacts with Oxygen ($\text{O}_2$) to produce Carbon Dioxide ($\text{CO}_2$) and Water ($\text{H}_2\text{O}$).

• Conservation of Mass: In any reaction, mass must be conserved. The number of atoms on the reactant side (left) must equal the number of atoms on the product side (right).

  • Balancing: Use coefficients to adjust the number of molecules. Never change subscripts in a chemical formula.

Stoichiometry and the Mole Concept

• The Mole: A unit representing $6.02 \times 10^{23}$ particles (Avogadro's Number). It allows scientists to bridge the gap between microscopic particles and macroscopic grams.

• Molar Mass: The total mass of one mole of a compound, calculated by summing the atomic masses of its components.

• Conversions:     $\text{Moles to Grams}: \text{Multiply by Molar Mass}$     $\text{Grams to Moles}: \text{Divide by Molar Mass}$

• Stoichiometric Ratios: The coefficients in a balanced equation represent the mole ratio between the reactants and products.

  • Example: In $2\text{H}_2 + \text{O}_2 \rightarrow 2\text{H}_2\text{O}$, the ratio of oxygen to water is $1:2$. If you use $3\,moles$ of Oxygen, you produce $6\,moles$ of Water.

• Limiting Reagent: The reactant that is completely consumed first, thereby stopping the reaction and limiting the amount of product formed.

• Percent Composition: $\frac{\text{Mass of specific atom}}{\text{Total Mass of molecule}} \times 100$. This value is fixed for any given compound.

• Empirical Formula: The simplest whole-number ratio of atoms in a compound (e.g., $\text{C}_4\text{H}_8$ has an empirical formula of $\text{CH}_2$).

States of Matter and Kinetic Molecular Theory

• Solids: Low entropy (disorder), fixed volume and shape, high organization in a crystal lattice structure.

• Liquids: Medium entropy, fixed volume but takes the shape of the container. Properties include viscosity (thickness) and surface tension.

• Gases: Highest entropy, random movement, no fixed volume or shape. Compressible and expandable.

• Phase Changes:

  • Solid to Liquid: Melting (Heat of Fusion, $H_f$).

  • Liquid to Gas: Vaporization (Heat of Vaporization, $H_v$).

  • Gas to Liquid: Condensation.

  • Liquid to Solid: Freezing.

  • Solid to Gas: Sublimation.

  • Gas to Solid: Deposition.

• The Combined Gas Law:     $\frac{P_1 V_1}{T_1} = \frac{P_2 V_2}{T_2}$

  • Higher temperature leads to higher pressure or volume (proportional).

  • Higher pressure leads to lower volume (inverse relationship).

• Heating Curves:

  • Legs on a curve show an increase in kinetic energy (increased temperature).

  • Plateaus show a phase change where potential energy changes while kinetic energy stays constant.

  • Formula: $q = m C \Delta T$ calculates energy during heating; $q = m H_f$ and $q = m H_v$ calculate energy during phase changes.

Solutions and Solubility

• Vocabulary:

  • Solute: The substance that is dissolved (e.g., Salt).

  • Solvent: The substance doing the dissolving (e.g., Water is the "universal solvent").

• Saturation Levels:

  • Unsaturated: Can hold more solute.

  • Saturated: Holds the maximum amount of solute at a given temperature.

  • Super Saturated: Holds more than technically possible; unstable.

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

  • Dilution Formula: $M_1 V_1 = M_2 V_2$.

Acids and Bases

• pH Scale: Ranges from $0$ to $14$.

  • pH 7: Neutral.

  • pH < 7: Acidic (High concentration of $H_3O^+$ or $H^+$ions).

  • pH > 7: Basic (High concentration of $OH^-$ ions).

  • Mathematical Definition: $\text{pH} = -\log[H_3O^+]$.

• Theories:

  • Acids: $H^+$ (proton) donors.

  • Bases: $H^+$ (proton) acceptors.

• Neutralization: An acid reacts with a base to produce water and a salt (e.g., $\text{HCl} + \text{NaOH} \rightarrow \text{NaCl} + \text{H}_2\text{O}$).

Thermochemistry and Kinetics

• Energy Changes:

  • Endothermic: Heat is absorbed ($\Delta H$ is positive). Bonds are generally broken.

  • Exothermic: Heat is released ($\Delta H$ is negative). Bonds are generally formed.

• Reaction Rates: Increased by high temperature, high concentration, high surface area (using powder instead of a bar), and the addition of a catalyst.

  • Catalysts: Lower the activation energy by providing an alternate reaction pathway.

• Equilibrium:

  • Reached when the rate of the forward reaction equals the rate of the reverse reaction.

  • The concentrations of products and reactants remain constant (not necessarily equal).

• Le Chatelier's Principle: If a system at equilibrium is stressed (change in concentration, pressure, or temp), the system shifts to counteract the stress.

  • Adding reactant ($A$) causes a shift to the right to produce more product ($C$).

Redox and Electrochemistry

• Oxidation and Reduction (OIL RIG):

  • Oxidation Is Losing electrons ($\text{charge increases/more positive}$).

  • Reduction Is Gaining electrons ($\text{charge decreases/more negative}$).

• Electrochemical Cells:

  • Anode: Where oxidation occurs ("An Ox"); it is negative in voltaic cells and generally shrinks in size.

  • Cathode: Where reduction occurs ("Red Cat"); it is positive and generally grows in size.

  • Voltaic/Galvanic Cell: Spontaneous; generates electricity (a battery).

  • Electrolytic Cell: Non-spontaneous; requires an external battery/power source.

  • Salt Bridge: Allows for the flow of ions to keep the solutions neutral.

Nuclear Chemistry

• Radioactivity: Spontaneous breakdown of an unstable nucleus.

• Types of Emissions:

  1. Alpha Decay: Emissions of a Helium nucleus ($\text{}_{2}^{4}\text{He}$). Least penetrating; stopped by skin/paper.

  2. Beta Decay: Emission of an electron ($\text{}_{-1}^{0}\text{e}$). Stopped by aluminum foil.

  3. Positron Emission: Emission of a positively charged electron ($\text{}_{+1}^{0}\text{e}$).

  4. Gamma Radiation: Pure energy ($\gamma$). Most penetrating and dangerous; stopped by lead/concrete.

• Nuclear Processes:

  • Fission: Splitting a heavy nucleus into lighter nuclei (used in power plants).

  • Fusion: Combining light nuclei into a heavier nucleus (happens in the sun; releases more energy than fission).

• Half-Life: The time required for half of a radioactive sample to decay.     $\text{Remaining amount} = \text{Initial amount} \times \left(\frac{1}{2}\right)^{\frac{t}{h}}$     where $t$ is time elapsed and $h$ is the half-life period.