Comprehensive NYS Chemistry Regents Study Guide

Atomic Structure and the Development of Atomic Theory

• The atom is the basic unit of matter and is composed of three primary subatomic particles: protons, neutrons, and electrons.

• Subatomic Particles:

  • Protons: Positively charged particles located in the nucleus. They have a mass of approximately $1.0\,amu$. The number of protons is the atomic number and determines the identity of the element.

  • Neutrons: Neutrally charged particles located in the nucleus. They have a mass of approximately $1.0\,amu$. Carbon-14 and Carbon-12 are isotopes because they have different numbers of neutrons.

  • Electrons: Negatively charged particles located in the orbitals (electron cloud) surrounding the nucleus. They have a negligible mass (approx. $\frac{1}{1836}\,amu$).

• Historical Models of the Atom:

  • Dalton Model: Atoms are solid, indivisible spheres.

  • J.J. Thomson (Plum Pudding Model): Discovered the electron using a cathode ray tube; suggested the atom is a positive sphere with embedded negative electrons.

  • Ernest Rutherford (Gold Foil Experiment): Bombarded gold foil with alpha particles. Most passed through, but some deflected. Conclusion: The atom is mostly empty space with a small, dense, positively charged nucleus.

  • Niels Bohr (Bohr Model): Electrons travel in fixed, circular orbits (energy levels or shells) around the nucleus.

  • Wave-Mechanical Model (Modern Model): Electrons do not travel in fixed paths but are found in regions of high probability called orbitals.

• Ground State vs. Excited State:

  • Ground State: Electrons occupy the lowest available energy levels.

  • Excited State: When an atom absorbs energy, electrons jump to higher energy levels. This state is unstable.

  • Bright-Line Spectrum: When an electron returns from an excited state to a ground state, it releases energy in the form of light. This spectrum serves as a "fingerprint" for identifying elements.

• Isotopes and Atomic Mass:

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

  • Average Atomic Mass: The weighted average of all naturally occurring isotopes. Formula: $\sum (\text{mass of isotope} \times \text{natural abundance})$.

The Periodic Table of Elements

• Organization: The Periodic Table is arranged by increasing atomic number.

• Groups (Vertical Columns):

  • Elements in the same group have the same number of valence electrons and similar chemical properties.

  • Group 1: Alkali Metals (highly reactive).

  • Group 2: Alkaline Earth Metals.

  • Groups 3-12: Transition Metals (form colored ions in solution).

  • Group 17: Halogens (highly reactive nonmetals).

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

• Periods (Horizontal Rows):

  • Elements in the same period have the same number of occupied energy levels (shells).

• Periodic Trends:

  • Atomic Radius: Increases down a group (more shells) and decreases across a period (greater nuclear charge pulling electrons in).

  • Electronegativity: A measure of the attraction for electrons in a bond. Decreases down a group and increases across a period (Fluorine is the highest at $4.0$).

  • First Ionization Energy: The energy required to remove the most loosely bound electron. Decreases down a group and increases across a period.

  • Metallic Character: Increases down a group and decreases across a period. Francium is the most metallic; Fluorine is the most nonmetallic.

• Metals, Nonmetals, and Metalloids:

  • Metals: Malleable, ductile, lustrous, good conductors of heat and electricity, low ionization energy, and low electronegativity.

  • Nonmetals: Brittle, dull, poor conductors, high ionization energy, and high electronegativity.

  • Metalloids (Semimetals): Located along the "staircase"; possess properties of both metals and nonmetals (B, Si, Ge, As, Sb, Te).

Chemical Bonding

• The Octet Rule: Atoms lose, gain, or share electrons to achieve a stable electron configuration (usually 8 valence electrons).

• Energy and Bonding:

  • Breaking a bond is an endothermic process (energy is absorbed).

  • Forming a bond is an exothermic process (energy is released).

• Types of Chemical Bonds:

  • Ionic Bond: Formed by the transfer of electrons from a metal to a nonmetal. Results from electrostatic attraction between ions. High melting points, hard, and conduct electricity in liquid ($l$) or aqueous ($aq$) states.

  • Covalent Bond: Formed by the sharing of electrons between two nonmetals. Molecular substances have low melting points and do not conduct electricity.

    • Nonpolar Covalent: Equal sharing of electrons (Electronegativity difference $\Delta EN = 0$ to $0.4$).

    • Polar Covalent: Unequal sharing of electrons (Electronegativity difference \Delta EN > 0.4).

  • Metallic Bond: Defined as a "sea of mobile electrons." Explains conductivity and malleability in metals.

• Molecular Polarity:

  • Symmetrical molecules (like $CH_4$ or $CO_2$) are nonpolar.

  • Asymmetrical molecules (like $H_2O$ or $NH_3$) are polar (T.I.C.S.: Transferred Ionic, Covalent Shared; or SNAP: Symmetrical Nonpolar, Asymmetrical Polar).

• Intermolecular Forces (IMF):

  • Hydrogen Bonding: A strong IMF occurring when Hydrogen is bonded to Nitrogen, Oxygen, or Fluorine (N, O, F). Explains the high boiling point of water.

  • Dipole-Dipole: Attraction between polar molecules.

  • London Dispersion Forces: Weakest forces, occurring in nonpolar molecules; strength increases with molecular size.

Matter, Energy, and Gas Laws

• Phases of Matter:

  • Solids ($s$): Definite shape and volume; particles vibrate in fixed positions.

  • Liquids ($l$): Definite volume, no definite shape; particles are close but can flow.

  • Gases ($g$): No definite shape or volume; particles move in random, straight-line motion.

• Heating and Cooling Curves:

  • Kinetic Energy ($KE$): Average kinetic energy is measured as temperature. $\text{Temperature} \propto KE$.

  • Potential Energy ($PE$): Changes during phase changes (plateaus on the graph), such as melting/fusion or boiling/vaporization.

• Thermodynamics Formulas:

  • Heat absorbed/released: $q = m \times c \times \Delta T$

  • Heat of Fusion (melting): $q = m \times H_f$

  • Heat of Vaporization (boiling): $q = m \times H_v$

• Kinetic Molecular Theory (KMT) for Ideal Gases:

  • Gas particles are in constant, random, straight-line motion.

  • Gas particles have negligible volume compared to the container.

  • There are no attractive or repulsive forces between particles.

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

• Real Gases vs. Ideal Gases:

  • Real gases behave most like ideal gases under conditions of High Temperature and Low Pressure.

  • Hydrogen and Helium are the most "ideal" gases because they are small.

• Gas Law Formulas:

  • Combined Gas Law: $\frac{P_1 V_1}{T_1} = \frac{P_2 V_2}{T_2}$, where temperature must always be in Kelvin ($K = ^\circ C + 273$).

  • Avogadro’s Hypothesis: Equal volumes of gases at the same temperature and pressure contain an equal number of molecules.

Stoichiometry and the Mole

• The Mole Concept:

  • $1\,\text{mole} = 6.02 \times 10^{23}\,\text{particles}$ (Avogadro's number).

  • Molar Mass (Gram Formula Mass): The sum of the atomic masses of all atoms in a formula ($g/mol$).

• Types of Chemical Formulas:

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

  • Molecular Formula: The actual number of atoms in a compound.

• Chemical Reactions:

  • Synthesis: $A + B \rightarrow AB$

  • Decomposition: $AB \rightarrow A + B$

  • Single Replacement: $A + BC \rightarrow AC + B$

  • Double Replacement: $AB + CD \rightarrow AD + CB$

• Conservation of Mass: Mass, charge, and energy must be conserved in all chemical reactions.

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

Kinetics and Equilibrium

• Collision Theory: For a reaction to occur, reactant particles must collide with effective orientation and sufficient energy (Activation Energy).

• Factors Affecting Reaction Rate:

  • Increasing Temperature: Increases speed and frequency of collisions.

  • Increasing Concentration: More particles result in more collisions.

  • Increasing Surface Area: More contact area for collisions.

  • Adding a Catalyst: Lowers the activation energy by providing an alternative pathway.

• Potential Energy Diagrams:

  • Exothermic Reaction: Products have less energy than reactants ($\Delta H$ is negative).

  • Endothermic Reaction: Products have more energy than reactants ($\Delta H$ is positive).

  • $\Delta H = H_{\text{products}} - H_{\text{reactants}}$

• Chemical Equilibrium:

  • Occurs when the rates of the forward and reverse reactions are equal.

  • The concentrations of reactants and products remain constant.

• Le Chatelier’s Principle: If a stress (change in concentration, temperature, or pressure) is applied to a system at equilibrium, the system shifts to relieve that stress.

  • Pressure: Increasing pressure shifts the equilibrium toward the side with fewer moles of gas.

• Entropy ($S$): A measure of randomness or disorder. Nature favors High Entropy and Low Energy (Enthalpy).

  • Phase order of increasing entropy: \text{solid} < \text{liquid} < \text{aqueous} < \text{gas}.

Oxidation-Reduction (Redox)

• Rules for Oxidation States:

  • Uncombined elements have an oxidation state of $0$.

  • Group 1 metals are always $+1$ in compounds.

  • Oxygen is usually $-2$ (except in peroxides where it is $-1$).

  • Hydrogen is usually $+1$ (except in metal hydrides where it is $-1$).

• LEO says GER (or OIL RIG):

  • Loss of Electrons is Oxidation (Oxidation number increases).

  • Gain of Electrons is Reduction (Oxidation number decreases).

• Electrochemical Cells:

  • Voltaic Cell: Spontaneous chemical reaction produces electricity (Chemical to Electrical). Anode is negative, Cathode is positive.

  • Electrolytic Cell: Electricity is used to force a non-spontaneous reaction (Electrical to Chemical). Anode is positive, Cathode is negative.

  • "AN OX and RED CAT": Oxidation occurs at the Anode; Reduction occurs at the Cathode.

Acids, Bases, and Salts

• Arrhenius Theory:

  • Acids produce $H^+$ (or $H_3O^+$ hydronium ions) as the only positive ion in aqueous solution.

  • Bases produce $OH^-$ (hydroxide ions) as the only negative ion in aqueous solution.

• Bronsted-Lowry Theory:

  • Acids are proton ($H^+$) donors.

  • Bases are proton ($H^+$) acceptors.

• pH Scale:

  • pH 0 to 6.9: Acidic.

  • pH 7: Neutral.

  • pH 7.1 to 14: Basic.

  • Each change of 1 pH unit represents a 10-fold change in $H^+$ concentration. Example: A change from pH 2 to pH 4 means the $H^+$ concentration decreased by 100 times ($10^2$).

• Neutralization: $\text{Acid} + \text{Base} \rightarrow \text{Salt} + \text{Water}$

• Titration: A process used to determine the concentration of an unknown acid or base.

  • Formula: $M_A V_A = M_B V_B$ (where $M$ is molarity and $V$ is volume).

Organic Chemistry

• Carbon Bonding: Carbon atoms always form 4 covalent bonds.

• Hydrocarbons: Compounds containing only Carbon and Hydrogen.

  • Alkanes: Single bonds ($C_n H_{2n+2}$), saturated.

  • Alkenes: One double bond ($C_n H_{2n}$), unsaturated.

  • Alkynes: One triple bond ($C_n H_{2n-2}$), unsaturated.

• Isomers: Compounds with the same molecular formula but different structural arrangements and different properties.

• Functional Groups:

  • Alcohols: Contain $-OH$ (hydroxyl group).

  • Organic Acids: Contain $-COOH$.

  • Halocarbons: Contain a halogen (F, Cl, Br, or I).

  • Esters: Formed by the reaction of an alcohol and an organic acid (Esterification), often smelling sweet.

• Organic Reactions:

  • Substitution: An atom is replaced in a saturated hydrocarbon.

  • Addition: Atoms added to an unsaturated hydrocarbon (breaking the double/triple bond).

  • Saponification: Reaction of a fat with a base to produce soap.

  • Fermentation: Production of ethanol and $CO_2$ from sugar using yeast.

  • Combustion: Reaction with Oxygen to produce $CO_2$ and $H_2O$.

  • Polymerization: Joining small units (monomers) to form large chains (polymers).

Nuclear Chemistry

• Stability: The stability of a nucleus depends on the ratio of neutrons to protons. Unstable nuclei are radioactive.

• Radiations/Decay Modes (Table N and O):

  • Alpha Particle ($\alpha$): ${}^4_2He$. Large mass, low penetrating power.

  • Beta Particle ($\beta^-$): ${}^0_{-1}e$. High-speed electron.

  • Positron ($\beta^+$): ${}^0_{+1}e$.

  • Gamma Radiation ($\gamma$): ${}^0_0\gamma$. Pure energy, highest penetrating power.

• Nuclear Reactions:

  • Natural Transmutation: Spontaneous decay of one element into another ($1$ reactant).

  • Artificial Transmutation: Bombardment of a nucleus with a particle to force change ($2$ reactants).

  • Fission: Splitting a heavy nucleus into lighter nuclei (produces massive energy).

  • Fusion: Combining light nuclei (like Hydrogen) to form a heavier nucleus (requires extreme heat and pressure; occurs in stars).

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

• Important Radioisotopes:

  • Carbon-14: Dating organic remains.

  • Uranium-238 / Lead-206: Dating geological formations (rocks).

  • Iodine-131: Diagnosing and treating thyroid disorders.

  • Cobalt-60: Cancer treatment through radiation therapy.

  • Technetium-99: Brain tumor imaging.