Comprehensive Study Notes for June 2026 Chemistry Regents

Fundamental Atomic Structure and Subatomic Particles

• Definition of Atoms: Atoms are the fundamental building blocks of all elements. Every atom is composed of three primary subatomic components: electrons, protons, and neutrons.

• Subatomic Particle Properties:

  • Electrons: These carry a negative charge. Their mass is considered negligible, which is mathematically close to $0\,\text{amu}$. Electrons occupy the space around the nucleus, orbiting in shells or orbitals. They dictate the volume and size of the atom.

  • Protons: These carry a positive charge. Each proton has a mass of exactly $1\,\text{amu}$. They are located within the nucleus (the center) of the atom. The number of protons determines the identity of the element.

  • Neutrons: These carry a neutral charge (a charge of $0$). Like protons, they have a mass of $1\,\text{amu}$ and reside in the nucleus.

• The Nucleus:

  • The nucleus is described as a small, dense, and positively charged center.

  • The positive charge of the nucleus results specifically from the presence of protons.

  • The majority of an atom's mass is concentrated in the nucleus (protons + neutrons).

• Atomic Space: Atoms are described as being mostly empty space. While the nucleus contains the mass, the vast electron cloud/density around it determines the volume.

Ionization, Isotopes, and Average Atomic Mass

• Ions: These are formed when there is a change in the number of electrons. Atoms become charged particles via the loss or gain of electrons.

  • Cations: Positively charged ions formed by losing electrons ($\text{Losing a negative charge makes the atom positive}$).

  • Anions: Negatively charged ions formed by gaining electrons ($\text{Gaining an extra negative charge makes the atom negative}$).

• Isotopes: These occur when atoms of the same element have a different number of neutrons, resulting in different mass numbers.

  • Identity Retention: Changing the neutrons does not change the element itself (protons remain the same).

  • Example: Carbon-13 ($^{13}_6C$) and Carbon-14 ($^{14}_6C$). Both have $6$ protons. Carbon-13 has $7$ neutrons ($13 - 6 = 7$), while Carbon-14 has $8$ neutrons ($14 - 6 = 8$).

  • Notation: For an isotope symbol like $^{14}_6C$, the top number ($14$) is the Mass Number (Protons + Neutrons) and the bottom number ($6$) is the Atomic Number (Protons).

• Average Atomic Mass: This is calculated based on the mass and the relative abundance of all naturally occurring isotopes of an element.

  • Calculation Method: Multiply the abundance (in decimal form) by the mass of each isotope, then add the products together.

  • Transcript Scenario: If Carbon-13 ($20\,g/mol$) is $3\%\,\text{abundant}$ and Carbon-14 ($22\,g/mol$) is $97\%\,\text{abundant}$:

  • $\text{Average Atomic Mass} = (0.03 \times 20) + (0.97 \times 22)$

• Proton Changes: Unlike changes in electrons or neutrons, changing the number of protons results in the formation of a brand-new element with unique properties.

• Neutral Atoms: In a neutral molecule (no net charge), $\text{Number of Protons} = \text{Number of Electrons}$. Example: Oxygen has an atomic number of $8$, meaning it has $8$ protons and $8$ electrons if neutral.

Electron Orbitals, Excitation, and the Electromagnetic Spectrum

• Orbital Energy Levels: Electrons exist in shells. The distance from the nucleus correlates with energy:

  • Ground State: The relaxed state where all electrons are in their lowest possible energy levels/proper shells.

  • Excited State: Occurs when an electron absorbs energy and is promoted to a higher, normally unoccupied orbital.

• Energy Transfer:

  • Absorption: Moving to a higher shell requires the absorption of energy in the form of light or a photon.

  • Emission: When an excited electron relaxes back to the ground state, it releases energy, often as visible light.

• Electromagnetic Spectrum (EMS):

  • Light travels in waves with peaks and dips.

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

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

  • Relationships:

  • As Wavelength ($\lambda$) decreases, Energy Increases ($\text{Inverse relationship}$).

  • As Frequency ($f$) increases, Energy Increases ($\text{Direct relationship}$).

  • Gamma Rays: High-energy waves with wavelengths ranging from $10^{-3}\,m$ to $10^{-1}\,m$.

  • Visible Light: A very small fraction of the total spectrum that humans can see.

• Valence Electrons: These are the electrons in the outermost shell. They are determined by the group number on the periodic table (e.g., Group 13 has $3$ valence electrons by dropping the "1"). Valence electrons dictate chemical reactivity.

Periodic Table Organization and Trends

• Organization: The table is arranged in order of increasing atomic number (number of protons).

• Sectional Divisions:

  • Metals: Located on the left (represented in green in the visual aid).

  • Non-metals: Located on the right (represented in blue).

  • Metalloids: Elements that touch the bold zigzag line (red). They have both metallic and non-metallic properties. Note: Aluminum is a metal, not a metalloid.

• Key Groups (Families):

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

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

  • Groups 3-12: Transition Metals (metallic bonding, sea of electrons).

  • Group 17: Halogens (highly reactive non-metals, high electronegativity).

  • Group 18: Noble Gases (unreactive due to a complete octet of $8$ valence electrons).

• Periodic Trends (Across a Period - Left to Right):

  • Electronegativity Increases: The nucleus becomes more positive (more protons), attracting negative electrons more strongly.

  • Atomic Radius Decreases: The stronger positive charge pulls electron shells closer to the nucleus.

  • Ionization Energy Increases: It requires more energy to remove an electron because it is held tighter by the nucleus.

• Group Trends (Down a Group):

  • Atomic Radius Increases: New electron shells are added (e.g., Lithium has $2$, Sodium has $3$, Potassium has $4$).

  • Electronegativity Decreases: The distance between the positive nucleus and valence electrons increases, weakening the attraction.

  • Ionization Energy Decreases: It is easier to lose electrons that are far away from the nucleus.

• Metallic vs. Non-metallic Character:

  • Metallic Character: The tendency to lose electrons (highest on the bottom left, e.g., Francium).

  • Non-metallic Character: The tendency to gain electrons (highest on the top right, excluding Noble Gases).

Chemical Bonding: Ionic and Covalent Interactions

• The Octet Rule: Atoms react to achieve a stable configuration of $8$ valence electrons. Exception: Hydrogen only needs $2$ electrons.

• Ionic Bonding:

  • Occurs between a Metal and a Non-metal.

  • Involves a complete transfer of electrons from the metal (which becomes a cation) to the non-metal (which becomes an anion).

  • Example (NaCl): Sodium ($1$ valence electron) gives its electron to Chlorine ($7$ valence electrons). This results in $Na^+$ and $Cl^-$, which attract each other due to opposite charges.

  • Formula derivation: If Oxygen ($6$ valence electrons) needs $2$ to reach $8$, it requires two Sodium atoms to provide them ($Na_2O$).

• Covalent Bonding:

  • Occurs between two Non-metals.

  • Involves the sharing of electrons.

  • Single Bond: Shares $2$ electrons ($\text{notated as } -$).

  • Double Bond: Shares $4$ electrons ($\text{notated as } =$).

  • Triple Bond: Shares $6$ electrons ($\text{notated as } \equiv$).

• Shared Electrons Formula: $S = N - A$

  • $S$ = Number of shared electrons.

  • $N$ = Number of electrons needed ($8$ for most, $2$ for H).

  • $A$ = Number of electrons available (valence count).

  • Example ($N_2$): Two nitrogens need $16$ ($8 \times 2$). They have $10$ available ($5 \times 2$). $16 - 10 = 6$ shared electrons, indicating a triple bond.

Molecular Polarity, Lewis Structures, and VSEPR Concepts

• Bond Polarity: Determined by the difference in electronegativity.

  • Polar Covalent Bond: Electronegativity difference is significant (up to $1.5$); electrons are shared unequally. Example: $HF$ ($F$ is highly electronegative, pulling electrons toward itself, becoming partially negative $\delta^-$).

  • Non-polar Covalent Bond: Electrons are shared evenly (e.g., $CH_4$ or diatomic molecules like $N_2$).

• Molecular Polarity and Symmetry:

  • A molecule can have polar bonds but be non-polar overall if it is symmetrical.

  • Example ($CO_2$): Oxygen pulls electrons away from Carbon in opposite directions. Like a tug-of-war between two equally strong men, the forces cancel out, making the molecule non-polar.

  • Asymmetrical Molecules: Generally polar (e.g., $H_2O$ or $HCl$).

Intermolecular Forces and Physical Properties

• Intermolecular Forces (IMF): Forces of attraction between molecules.

• Ranking by Strength:

  1. Hydrogen Bonding: The strongest IMF. Occurs when Hydrogen is bonded to Fluorine, Oxygen, or Nitrogen (F-O-N). These create very strong dipoles.

  2. Dipole-Dipole Attractions: Attractions between polar molecules that do not involve F-O-N.

  3. London Dispersion Forces: The weakest IMF. Occurs in non-polar molecules due to temporary, brief shifts in electron density.

• Boiling Point Connection: Stronger IMFs lead to higher boiling points because more energy is required to break the attraction between molecules to phase change into a gas.

Chemical Nomenclature and Formula Writing

• Ionic Naming:

  • Format: Name of Metal + Name of Non-metal ending in "-ide".

  • Example: $NaCl$ is Sodium Chloride.

  • Multivalent Metals (Transition Metals): Metals that can have multiple oxidation states require Roman Numerals.

  • Example: Palladium can be $+2$ or $+4$. $PdO$ is Palladium (II) Oxide because Oxygen needs $2$ electrons.

• Covalent Naming:

  • Uses prefixes to indicate the number of atoms: Mono (1), Di (2), Tri (3), Tetra (4).

  • Format: Prefix + First Non-metal, Prefix + Second Non-metal ending in "-ide".

  • Example: $CO_2$ is Carbon Dioxide; $CO$ is Carbon Monoxide.

• Polyatomic Ions: Groups of covalently bonded atoms with a charge (e.g., Ammonium $NH_4^+$, Chlorate $ClO_3^-$). They have specific names found on reference tables and should not be renamed with prefixes.

Chemical Reactions and Balancing Equations

• Types of Reactions:

  • Synthesis: Two or more reactants combine to form one product ($A + B \rightarrow C$).

  • Decomposition: One reactant breaks down into two or more products ($C \rightarrow A + B$).

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

  • Double Replacement: Components of two compounds swap places ($AB + CD \rightarrow AC + BD$).

  • Combustion: A hydrocarbon reacts with Oxygen ($O_2$) to produce Carbon Dioxide ($CO_2$) and Water ($H_2O$).

• Conservation of Mass: In a chemical reaction, the number of atoms of each element on the reactant side (left) must equal the number of atoms on the product side (right).

• Balancing Technique:

  • Use Coefficients (numbers in front) to adjust atom counts.

  • Subscripts (small numbers within the formula) cannot be changed.

  • Example: $2H_2 + O_2 \rightarrow 2H_2O$

  • Reactants: $4\,H$, $2\,O$

  • Products: $4\,H$, $2\,O$

The Mole Concept, Stoichiometry, and Composition

• Avogadro's Number: One mole equals $6.02 \times 10^{23}$ particles. The mole allows scientists to bridge the gap between microscopic particles and macroscopic grams.

• Molar Mass Conversions:

  • Moles to Grams: $\text{Moles} \times \text{Molar Mass}$.

  • Grams to Moles: $\text{Grams} / \text{Molar Mass}$.

  • Example: Molar mass of $O_2$ is approximately $32\,g/mol$ ($16 \times 2$).

• Stoichiometry: Using balanced equation coefficients as molar ratios to predict product yield.

  • If the ratio of $H_2$ to $H_2O$ is $2:2$ ($1:1$), then $4$ moles of $H_2$ will produce $4$ moles of $H_2O$.

• Limiting and Excess Reagents:

  • Limiting Reagent: The reactant that is completely consumed first and limits the amount of product made.

  • Excess Reagent: The reactant that is left over after the reaction stops.

• Percent Composition: $\frac{\text{Mass of Part}}{\text{Mass of Whole}} \times 100$.

  • Example: Nitrogen in $NH_4$ ($\text{mass } 18.05\,g/mol$) is $\frac{14.01}{18.05} \approx 77\%$.

• Empirical vs. Molecular Formulas:

  • Molecular Formula: The actual number of atoms (e.g., $C_4H_8$).

  • Empirical Formula: The simplest whole-number ratio (e.g., dividing $C_4H_8$ by the common factor $4$ yield $CH_2$).

States of Matter and Kinetic Molecular Theory

• Solids:

  • Lowest entropy (disorder).

  • Atoms are locked in a Crystal Lattice structure with minimal vibration.

• Liquids:

  • Moderate entropy.

  • Particles are still close but can flow past one another.

  • Properties: Viscosity (thickness) and Surface Tension.

• Gases:

  • Highest entropy.

  • Particles move randomly and are far apart.

  • No attraction between particles (in an ideal sense).

• Phase Changes:

  • Melting: Solid to Liquid.

  • Vaporization: Liquid to Gas.

  • Condensation: Gas to Liquid.

  • Freezing: Liquid to Solid.

  • Sublimation: Solid straight to Gas.

  • Deposition: Gas straight to Solid.

Thermodynamics, Heating Curves, and Gas Laws

• Heating Curves:

  • Temperature (Kinetic Energy) is on the Y-axis vs. Heat added.

  • Plateaus indicate a phase change (potential energy increases, but temperature stays constant).

  • Heat of Fusion ($H_f$): Energy to melt/freeze.

  • Heat of Vaporization ($H_v$): Energy to boil/condense.

• Specific Heat Formula: $Q = mc\Delta T$

  • $Q$ = heat energy.

  • $m$ = mass.

  • $c$ = specific heat capacity.

  • $\Delta T$ = change in temperature ($T_{final} - T_{initial}$).

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

  • Pressure and Volume are inversely proportional (High P = Low V).

  • Pressure and Temperature are directly proportional (High T = High P).

Solutions, Solubility, and Concentration

• Terminology:

  • Solute: Substrate that is being dissolved (e.g., Salt).

  • Solvent: The dissolving agent (e.g., Water is the "Universal Solvent").

• Saturation States:

  • Saturated: Contains the maximum amount of solute for a given temperature.

  • Unsaturated: Can hold more solute.

  • Super Saturated: Contains more solute than should be possible (often achieved by heating and cooling).

• Solubility Rules: Use reference tables to determine if a compound will dissolve. Example: All Nitrates ($NO_3^-$) are soluble.

• Concentration (Molarity): $M = \frac{n}{V}$

  • $M$ = Molarity ($mol/dm^3$).

  • $n$ = number of moles.

  • $V$ = volume in Liters.

  • Dilution Formula: $M_1V_1 = M_2V_2$.

Acid-Base Chemistry and pH

• pH Scale: Ranges from $0$ to $14$.

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

  • Neutral: pH = $7$.

  • Basic (Alkaline): pH > $7$ (Low concentration of $H^+$).

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

• Theory:

  • Acids: $H^+$ (Proton) donors.

  • Bases: $H^+$ (Proton) acceptors.

  • Reaction Example: $HCl + NaOH \rightarrow NaCl + H_2O$. The $HCl$ donates a proton to the $OH$, forming water.

Oxidation-Reduction (Redox) and Electrochemistry

• OIL RIG Mnemonic:

  • Oxidation Is Losing electrons (Oxidation number increases).

  • Reduction Is Gaining electrons (Oxidation number decreases).

• Oxidation Numbers: Used to track electron flow. Standard rules: Oxygen is usually $-2$, Group 1 is $+1$. Sum of numbers must equal the charge of the molecule.

• Electrochemical Cells:

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

  • Electrolytic Cell: Non-spontaneous, requires an external power source.

  • Anode: Site of Oxidation. It usually shrinks because solid metal turns into ions.

  • Cathode: Site of Reduction. It usually grows as ions plate onto the solid metal. (Mnemonic: Cathode has a "T" like a plus sign; it is positive in galvanic cells).

  • Salt Bridge: Permits the flow of ions between beakers to maintain neutrality.

Nuclear Chemistry and Radioactive Decay

• Radioactivity: Spontaneous breakdown of unstable nuclei.

• Types of Decay:

  • Alpha Decay ($\alpha$): Emits a helium nucleus ($^{4}_{2}He$). Reduces mass by $4$ and atomic number by $2$. Least penetrating (stopped by paper).

  • Beta Decay ($\beta^-$): Emits an electron ($^{0}_{-1}e$). Neutron is converted to a proton. Stopped by aluminum foil.

  • Positron Emission: Emits a positively charged electron ($^{0}_{+1}e$).

  • Gamma Radiation ($\gamma$): Pure energy emission. No mass change. Most penetrating (requires lead or concrete).

• Half-Life: The time required for half of a radioactive sample to decay.

  • Formula: $\text{Amount remaining} = \text{Initial amount} \times (0.5)^{\frac{t}{h}}$, where $t$ is time elapsed and $h$ is half-life.

• Fission vs. Fusion:

  • Fission: A heavy nucleus splits into smaller nuclei (used in nuclear power plants).

  • Fusion: Small nuclei combine to form a heavier nucleus (occurs in the Sun; releases significantly more energy than fission).