Exhaustive Chemistry Regents Study Guide

Unit 1 - Unit Conversion, Significant Figures, and Percent Error

• Unit Conversion and Setup

  • Uses the graph/conversion factors found on the reference sheet.

  • Proper setup is required for all conversions to ensure units cancel correctly.

• Percent Error

  • The formula for Percent Error is located on the Chemistry Reference Table.

  • Formula: $\text{Percent Error} = \frac{|\text{Measured Value} - \text{Accepted Value}|}{\text{Accepted Value}} \times 100\%$.

• Significant Figures (Sig Figs) Rules

  • Non-zero numbers: All non-zero digits are always significant.

  • Decimal Point Rule:

    • If a decimal point is ABSENT: Trailing zeros (zeros at the end) are NOT significant.

    • If a decimal point is PRESENT: Trailing zeros are significant, regardless of which side of the decimal point they occupy.

  • Leading Zeros: Leading zeros (zeros at the beginning of a number) are NEVER significant on either side of a decimal point.

• Mathematics with Significant Figures

  • Addition and Subtraction: The final answer should be rounded to the same number of decimal places as the measurement with the least number of decimal places.

  • Multiplication and Division: The final answer should have the same number of total significant figures as the measurement with the fewest significant figures.

  • Exact Numbers: Unit conversions are considered exact values; therefore, significant figure rules do not apply to them.

Unit 2 - Properties of Matter, Energy, and Temperature

• Properties of Matter

  • Physical Properties: Characteristics measured or observed through look, texture, smell, etc. Examples include mass, volume, and color.

  • Chemical Properties: Measures how chemically reactive a substance is when interacting with other substances.

• Physical and Chemical Changes

  • Physical Changes: Changes that do not alter the chemical composition of a substance.

    • Includes the creation or separation of a mixture.

    • Includes physical deformation (changing shape).

    • Includes physical relocation (moving the substance).

  • Chemical Changes: Changes that alter the chemical composition and transform a substance into something new.

  • Indications of Chemical Change:

    • Appearance of bubbles of gas.

    • Formation of a precipitate (solid).

    • Occurrence of an unexpected color change.

    • Gain or release of energy (heat or light).

    • Change in volume.

    • Change in electrical conductivity.

    • Change in melting point or boiling point.

    • Change in odor or taste.

    • Change in a distinctive chemical or physical property that is not easily reversed.

• Laws of Conservation

  • Conservation of Mass: Mass cannot be created nor destroyed in a chemical reaction.

  • Conservation of Energy: Energy is the ability to do work; it cannot be created or destroyed, only transformed from one form to another.

• Chemical Reactions and Thermodynamics

  • Chemical Reaction: A chemical change where one substance is transformed into another.

  • Exothermic: A process that releases heat or energy into the surroundings.

  • Endothermic: A process that absorbs heat or energy from the surroundings.

• Temperature

  • Definition: A measure of the average kinetic energy of the particles in a sample.

  • Scales used: Celsius and Kelvin. The conversion formula ($K = ^\circ C + 273$) is provided on the reference sheet.

Unit 3 - Chronological Order of Atomic Models

• 1. Dalton's Theory (Cannonball Model)

  • Matter is composed of atoms.

  • Atoms of the same element are identical in mass and other properties.

  • Chemical reactions involve the reorganization of atoms, but the atoms themselves do not change.

  • Refinement (What was wrong): Atoms can actually be divided, changed into other elements, and isotopes (atoms of the same element with different masses) exist.

• Cathode Ray Experiment

  • A cathode ray was repelled by a negative plate, indicating matter contains negatively charged components called electrons. This discovery led to the next model.

• 2. Thomson’s Model (Plum Pudding Model)

  • Electrons are distributed throughout a "sea" or sphere of positive charge.

  • Refinement (What was wrong): Electrons exist in specific orbitals, and there is a dense positive core.

• Gold Foil Experiment

  • Alpha particles were fired at a thin sheet of gold foil. While most passed through (as Thomson predicted), some were deflected at sharp angles.

  • Indications: This proved that atoms have a dense, positively charged core (the nucleus) and that most of the atom is empty space. This led to the Bohr model.

• 3. Bohr Model (Planetary Model)

  • Electrons travel in fixed orbits around the nucleus and cannot exist in the spaces between orbits.

  • Electrons can jump to higher orbits (absorption) and fall back to lower orbits (emission).

  • Refinement (What was wrong): This model only accurately predicts the behavior of atoms with one electron (like Hydrogen).

• 4. Wave-Mechanical Model (Modern Atomic Model)

  • Electrons do not move in fixed orbits but are found in orbitals.

  • Orbital: A region of space where there is a high probability of finding an electron.

• Subatomic Particles

  • Protons: Positively charged ($+1$). Located in the nucleus. Mass is $1\,amu$. The number of protons (Atomic Number, represented as $Z$) identifies the element and remains constant for all isotopes of that element.

  • Neutrons: No charge (neutral). Located in the nucleus. Mass is $1\,amu$. The sum of Protons + Neutrons = Atomic Mass (represented as $A$).

  • Electrons: Negatively charged ($-1$). Located outside the nucleus. Mass is approximately $0\,amu$. In a neutral atom, the number of electrons equals the number of protons.

• Ions, Isotopes, and Species

  • Ions: Charged particles resulting from a different number of protons and electrons.

    • Cation: Positively charged (+#); has more protons than electrons.

    • Anion: Negatively charged (-#); has more electrons than protons.

  • Isotopes: Atoms of the same element (same number of protons) but with different numbers of neutrons (different mass numbers).

  • Isoelectronic Species: Atoms, ions, or molecules that share the same total number of electrons and the same electron configuration (e.g., $Ne$, $Na^+$, and $F^-$), despite having different proton counts.

• Light and Spectra

  • Absorption: Electrons move from a ground state to an excited state by absorbing a photon.

  • Emission: Electrons move from an excited state back to a ground state, emitting a photon.

  • Bright Line Spectra: Emission from one energy level to another produces light with characteristic energy and specific wavelengths. This acts as a "fingerprint" for the element.

Unit 4 - Organization and Trends of the Periodic Table

• Organization

  • Groups: Vertical columns. Elements in the same group have similar chemical properties because they have the same number of valence electrons.

  • Periods: Horizontal rows. Valence electrons increase as you move from left to right across a period.

• Types of Elements

  • Metals: Located on the left side of the table. Usually solid (except Mercury), denser than water, malleable, ductile, shiny, and good conductors of heat and electricity. They have low ionization energy and low electronegativity. Metals tend to lose electrons to form cations.

  • Nonmetals: Located on the right side of the table. Mostly gases at room temperature; solids are brittle and dull. Poor conductors. They have high ionization energy and high electronegativity. They tend to gain electrons.

  • Metalloids: Found along the "staircase." They possess properties of both metals and nonmetals.

  • Diatomic Elements: Elements that exist naturally as molecules of two atoms bonded together. The mnemonic is "Have No Fear Of Ice Cold Beer": $H_2$, $N_2$, $O_2$, $F_2$, $Cl_2$, $Br_2$, and $I_2$.

• Periodic Trends

  • Atomic Radius: A measure of the size of an atom.

  • Shielding: Inner-shell electrons reduce the electrostatic attraction between the positive nucleus and the outer valence electrons.

  • Ionic Radius: Size of an ion.

    • Cations are smaller than their parent neutral atoms.

    • Anions are larger than their parent neutral atoms.

  • Ionization Energy: The energy required to remove an electron from the valence shell. Smaller atomic radii lead to higher ionization energy because the nucleus has a stronger pull on valence electrons.

  • Electronegativity (Electron Affinity): The tendency of an atom to attract shared electrons in a bond. Higher electronegativity means a stronger attraction. Same general trend as ionization energy: Metals are low, nonmetals are high.

• Specific Groups

  • Alkali Metals: Highly reactive metals that form $+1$ cations.

  • Halogens: Highly reactive nonmetals that form $-1$ anions.

  • Noble Gases: Traditionally unreactive gases with a full octet of electrons.

Unit 5 - Bonding and Intermolecular Forces

• BARF: Energy and Bonds

  • Break Absorb: Breaking a bond absorbs energy.

  • Release Form: Forming a bond releases energy.

• Types of Bonding

  • Ionic Bonding: Formed between positively charged metals and negatively charged nonmetals through electron transfer. Held together by electrostatic attraction. Includes Polyatomic Ions (found in Table E).

  • Covalent Bonding: Formed between two nonmetal atoms through electron sharing. Each bond contains 2 shared electrons.

    • Polar: Asymmetric molecule with an overall dipole (charge separation).

    • Non-Polar: Symmetric molecule with no overall dipole.

  • Metallic Bonding: Occurs in metals through Electron Delocalization (often called a "sea of mobile electrons").

• Lewis Structures and Formulas

  • Lewis Dot Structures: The element symbol (the kernel) represents the nucleus and non-valence electrons; dots represent valence electrons.

  • Criss-Cross Method: Determining chemical formulas for ionic compounds by crossing the charge numbers to become subscripts of the opposite ion.

• Properties of Solutions

  • Metallic solutions: Conductive due to mobile electrons.

  • Ionic solutions: Conductive only when dissolved ($aq$) or molten ($l$) because ions are free to move; solids do not conduct.

  • Covalent solutions: Generally not conductive (stay neutral). Exceptions include acids ($HCl$) and substances that ionize in water.

• Intermolecular Forces (IMF)

  • Dispersion: Weakest attractive forces present between all molecules.

  • Dipole-Dipole: Medium strength forces between polar molecules.

  • Ion-Dipole: Forces between an ionic compound and a polar compound (common in aqueous solutions).

  • Hydrogen Bonding: The strongest IMF. Occurs when a hydrogen atom bonded to $N$, $O$, or $F$ is attracted to a lone pair on a neighboring molecule's $N$, $O$, or $F$.

Unit 6 - Stoichiometry and Reactions

• The Mole and Formulas

  • Mole Calculations: Converting between Moles and Number of particles, Moles and Mass (using molar mass), and Moles and Volume (for a gas, $1\,mol = 22.4\,L$ at STP).

  • Empirical Formula: The simplest whole-number ratio of atoms in a compound.

  • Molecular Formula: The actual/exact number of atoms of each element in a compound.

• Chemical Equations

  • Conservation of Mass and Charge: Equations must be balanced so the number of atoms for each element and the total charge are identical on both sides.

  • Stoichiometric Ratios: Using coefficients to determine how many moles of a substance are produced or consumed.

• Types of Chemical Reactions

  • Synthesis (Combination): $A + B \rightarrow AB$.

  • Decomposition (Analysis): $AB \rightarrow A + B$.

  • Single Replacement (Displacement/Exchange): $AB + C \rightarrow AC + B$.

  • Double Replacement (Salt Metathesis): $AB + CD \rightarrow AD + CB$.

  • Combustion: An organic compound (containing $C$ and $H$) reacts with $O_2$ to produce $CO_2$ and $H_2O$.

Unit 7 - States of Matter, Gas Laws, and Solutions

• States of Matter

  • Solids: Closely bound, definite shape, definite volume.

  • Liquids: Less closely bound, no definite shape, definite volume.

  • Gases: Loosely bound, no definite shape, no definite volume.

• Kinetic Molecular Theory (KMT) for Ideal Gases

  • Particles are in constant, random, straight-line motion.

  • Particles collide with each other and container walls.

  • Collisions are perfectly elastic (no net loss of energy).

  • Volume of gas particles is negligible compared to the distance between them.

  • Particles have no attractive or repulsive forces.

  • Real Gases vs. Ideal Gases: Real gases behave most like ideal gases at high temperature and low pressure. They deviate most at high pressure (volume becomes non-negligible) and low temperature (IMFs become significant).

• Gas Laws

  • Boyle’s Law: Pressure and Volume are inversely proportional at constant Temperature.

  • Charles Law: Volume and Temperature ($K$) are directly proportional at constant Pressure.

  • Avogadro’s Law: Volume and Moles of gas are directly proportional at constant Temperature and Pressure.

  • Gay-Lussac’s Law: Pressure and Temperature ($K$) are directly proportional at constant Volume.

• Properties of Liquids and Solids

  • Vapor Pressure: Pressure exerted by a vapor in equilibrium with its liquid. Increases with temperature.

  • Boiling Point: The temperature where vapor pressure equals atmospheric pressure. High Vapor Pressure = Low Boiling Point = Weak IMFs.

  • Water Density: Unique because it becomes less dense upon freezing due to hydrogen bonding structure.

  • Solids: Crystalline (long-range repeating unit cells) vs. Amorphous (local structure only).

  • Allotropes: Different structural forms of the same element (e.g., graphite and diamond for carbon).

• Phase Changes

  • Sublimation: Solid to Gas (direct transition).

  • Deposition: Gas to Solid (direct transition).

  • Evaporation: Vaporization at the surface below the boiling point; energy is taken from the liquid, cooling it down.

• Solutions and Solubility

  • Unsaturated: Can dissolve more solute.

  • Saturated: Maximum solute dissolved at equilibrium.

  • Supersaturated: More than the maximum is dissolved (unstable).

  • Affecting Solubility: Increasing temperature increases solid solubility but decreases gas solubility. Increasing pressure increases gas solubility.

  • Colligative Properties: Adding solute increases the boiling point (elevation) and decreases the freezing point (depression).

• Separation of Mixtures

  • Filtration: Separates heterogeneous solid/liquid mixtures.

  • Distillation: Separates miscible liquids based on different boiling points.

  • Chromatography: Separates based on molecular weight and polarity.

• Thermochemistry

  • Enthalpy ($H$): Heat of reaction at constant pressure ($\Delta H$).

  • Exothermic: $-\Delta H$; Endothermic: $+\Delta H$.

Unit 8 - Kinetics, Equilibrium, and Thermodynamics

• Collision Theory

  • Reactions occur only if molecules collide with sufficient kinetic energy and correct orientation.

  • Rate Factors: Increased temperature, concentration, and pressure (for gases) increase reaction rates. Catalysts speed up reactions by lowering activation energy ($E_a$).

• Equilibrium and Le Chatelier’s Principle

  • Equilibrium: Rates of forward and reverse reactions are equal; concentrations are stable.

  • Le Chatelier's: A system at equilibrium will shift to oppose a stress (change in concentration, pressure, or temperature).

  • Common Ion Effect: Adding a common ion shifts equilibrium to the opposite side.

• Entropy ($S$) and Free Energy ($G$)

  • Entropy: Measure of disorder. Increases with phase changes toward gas, reducing particle size, increasing temperature, or dissolving solids.

  • Gibbs Free Energy Formula: $\Delta G = \Delta H - T\Delta S$.

  • If $\Delta G$ is negative, the reaction is spontaneous.

Unit 9 - Acids, Bases, and Salts

• General Properties

  • Acids: Sour taste, electrolytes in solution.

  • Bases: Bitter taste, soapy feel, electrolytes.

• Theories

  • Arrhenius: Acids produce $H^+$ ($H_3O^+$) in water; Bases produce $OH^-$ in water.

  • Bronsted-Lowry: Acids are $H^+$ donors; Bases are $H^+$ acceptors.

  • Lewis: Acids are electron pair acceptors; Bases are electron pair donors.

• pH and pOH

  • $pH = -\log[H^+]$. A change of 1 pH unit is a $10\times$ change in concentration.

  • $pH + pOH = 14$.

  • Neutral: $pH = 7$. Acidic: pH < 7. Basic: pH > 7.

• Titration

  • Process used to find the concentration of an unknown analyte using a titrant of known concentration and an indicator (like phenolphthalein).

Unit 10 - Redox and Electrochemistry

• Oxidation-Reduction

  • Oxidation: Loss of electrons (Oil); increase in oxidation number.

  • Reduction: Gain of electrons (Rig); decrease in oxidation number.

  • Oxidation Number Rules: Pure elements are 0. Group 17 is usually $-1$. Hydrogen is $+1$ (except with metals where it is $-1$). Oxygen is $-2$ (except in peroxides where it is $-1$).

• Electrochemical Cells

  • Voltaic Cells: Spontaneous; produce electricity. Anode is negative, Cathode is positive.

  • Electrolytic Cells: Non-spontaneous; require external voltage (Electroplating). Anode is positive, Cathode is negative.

  • Mnemonics: "An Ox" (Anode = Oxidation); "Red Cat" (Reduction = Cathode).

Unit 11 - Organic Chemistry

• Carbon Features

  • Carbon forms 4 bonds. Saturated chains have single bonds (tetrahedral). Unsaturated have double (planar) or triple (linear) bonds.

• Homologous Series

  • Alkanes: Single bonds, suffix "-ane". Boiling point increases with carbon count.

  • Alkenes: One double bond, suffix "-ene".

  • Alkynes: One triple bond, suffix "-yne".

• Organic Reactions

  • Substitution: Swapping $H$ for a halogen in an alkane.

  • Addition: Adding atoms across a multiple bond.

  • Esterification: Acid + Alcohol $\rightarrow$ Ester + Water.

  • Saponification: Triglyceride + Base $\rightarrow$ Soap + Glycerol.

  • Polymerization: Joining monomers (Addition or Condensation).

  • Fermentation: Sugar + Enzymes $\rightarrow$ Alcohol + $CO_2$.

Unit 12 - Nuclear Chemistry

• Radioactivity and Stability

  • Nuclei are held by the Nuclear Strong Force. Large nuclei (82+ protons) are unstable because repulsion overcomes the strong force.

  • Radioactive Particles: Alpha ($^4_2He$), Beta ($^0_{-1}e$), Positron ($^0_{+1}e$), and Gamma ($\gamma$) radiation.

• Nuclear Reactions

  • Transmutation: One element turning into another through decay or bombardment.

  • Fission: A heavy nucleus splits into smaller nuclei, releasing energy and creating a chain reaction.

  • Fusion: Light nuclei combine to form a heavier nucleus; occurs in the Sun ($H \rightarrow He$).

• Applications

  • Carbon-14: Dating organic matter.

  • Uranium-238: Geological dating.

  • Iodine-131: Thyroid cancer treatment.

  • Cobalt-60: Killing tumors and Anthrax Bacilli.

  • Technetium-99: Detecting tumors.

Note: There is no penalization for significant figure errors on this exam; rounding is permitted to a reasonable degree.