Comprehensive Regents Chemistry Encyclopedia

Matter and Its Classifications

• Pure Substance
  • Contains only one element or compound.
  • Exhibits Definite Composition.
• Elements
  • Located on the Periodic Table (PT).
  • Consists of only one type of atom.
  • The smallest form is the atom.
  • Elements cannot be decomposed or broken down.
  • Monatomic: Consists of 1 atom.
  • Diatomic: Consists of 2 atoms.
• Compounds
  • Consists of two or more atoms chemically combined.
  • Elements take on new properties once they are part of a compound.
  • The smallest form is the molecule.
  • Can be decomposed or broken down into constituent elements via chemical means such as Electrolysis, Light, or Heat.
• Mixtures
  • Contains two or more pure substances (elements or compounds) physically combined.
  • Exhibits Varying Composition.
  • Substances within the mixture retain their original physical properties.
  • Can be physically separated using various methods.
• Homogeneous Mixture
  • Uniform (evenly distributed) throughout.
  • Individual substances cannot be distinguished from one another.
  • Often referred to as Solutions.
• Heterogeneous Mixture
  • Not uniform (unevenly distributed) throughout.
  • Different substances can be told apart; often shows visible Layers.

Physical vs. Chemical Properties and Changes

• Physical Property/Physical Change
  • Properties and changes measured that do not change the substance.
  • The appearance is different, but the chemical composition stays the same.
• Chemical Property/Chemical Change
  • Properties and changes measured that change the nature of the substance.
  • The appearance changes as well as the composition; something new is formed.

Separation Techniques

• Filtration: Separates large particles via a membrane. It does not separate dissolved particles in a solution.
• Distillation: Separates liquids based on their different boiling points. Each liquid boils off and is collected one at a time.
• Chromatography: Separates substances by differences in polarity, Intermolecular Forces (IMF), or attraction.
• Evaporation: Used to separate a solid from a liquid in a solution.
• Centrifuge: Separates components based on density by spinning particles at high speeds.
• Magnet: Used to separate magnetic materials from non-magnetic materials.
• Electrolysis: A chemical method used to decompose compounds into elements.

Atomic Theory and Structure

• Historical Development of the Atom
  • Dalton: Solid sphere model.
  • Thomson: Plum pudding model; discovery of electrons using a cathode ray tube.
  • Rutherford: Gold foil experiment; discovery of the dense, positive nucleus; atom is mostly empty space.
  • Bohr: Planetary model; electrons orbit the nucleus in specific energy levels.
  • Today: Wave-mechanical model; electrons are found in orbitals (regions of high probability).
• Subatomic Particles
  • Proton ($p^+$): Charge of $+1$, Mass of $1\,amu$, located inside the nucleus. Atomic Number = number of protons ($p^+$).
  • Neutron ($n^0$): Charge of $0$, Mass of $1\,amu$, located inside the nucleus.
  • Electron ($e^-$): Charge of $-1$, Mass of $0\,amu$, located outside the nucleus (orbitals).
• Atomic Calculations
  • Mass Number: Sum of protons and neutrons ($p^+ + n^0$).
  • Neutral Atom: Number of protons equals the number of electrons.
  • Atomic Mass: The weighted average of all naturally occurring isotopes of an element.
    • Example: Calculation for Carbon
    • Isotope $C-12$: Mass $12.0000\,amu$, Abundance $98.90\%$
    • Isotope $C-13$: Mass $13.0035\,amu$, Abundance $1.10\%$
    • Set up: $(12.0000 \times 0.9890) + (13.0035 \times 0.0110)$
• Isotopes
  • Atoms of the same element that have the same number of protons but different numbers of neutrons. This results in different masses.
  • Examples: $C-12$, $C-13$, $C-14$ (also written as $^{12}C$, $^{13}C$, $^{14}C$).

Electron Arrangement and Behavior

• Valence Electrons: Electrons on the outermost level or shell of an atom. Example: Argon (Ar) with configuration $2-8-8$ has $8$ valence electrons.
• Lewis Dot Diagram: Shows valence electrons as dots around an element's symbol (e.g., $\text{:Ar:}$ with dots on top, bottom, and sides).
• Ground State: All electrons are in the lowest possible energy levels (as found on the Periodic Table). Example: Ar $2-8-8$.
• Excited State: One or more electrons move from the ground state to a higher energy level. Example: Ar $2-8-7-1$. Note: The total number of electrons must still equal the element's atomic number.
• Bright-Line Spectra and Flame Colors
  • Produced when electrons in the ground state absorb energy ($E$) and move to an excited state (higher energy level).
  • When they fall back to the ground state, they emit energy as light.
  • Every element has a unique spectrum based on its number of electrons; spectra are used for identification.
• Ions
  • Negative Ions (Anions): Gained electrons. Example: $N^{3-}$ has $10$ electrons.
  • Positive Ions (Cations): Lost electrons. Example: $Na^+$ has $10$ electrons.
  • Both above examples are isoelectronic with Neon ($Ne$).

The Periodic Table

• Structural Organization
  • Groups (Down): Also called Families. Elements have the same number of valence electrons and similar chemical properties.
  • Periods (Across): Indicate the number of energy levels (shells) in an atom.
  • Periodic Law: Properties of elements are functions of their Atomic Numbers ($At.\,\#$).
  • Elements are arranged in order of increasing Atomic Number. Mendeleev is the "Father of the PT" (though Moseley corrected it to Atomic Number).
• Major Categories
  • Metals: Located to the left of the "step" (zig-zag). They lose electrons to form positive cations. They have low Ionization Energy (IE) and Electronegativity (EN). Physical traits: Good conductors, high Melting Point (MP) and Boiling Point (BP), hard, malleable, and ductile.
  • Metalloids (Semimetals): Found along the step. Have properties of both metals and nonmetals. Examples: $B, Si, Ge, As, Sb, Te, Po, At$.
  • Nonmetals: Located to the right of the "step". They gain electrons to form negative anions. They have high IE and EN. Physical traits: Poor conductors, low MP and BP, soft, brittle.
• Specific Groups
  • Group 1: Alkali Metals. Most reactive metals; explosive in water; found only as compounds in nature; $1$ valence electron.
  • Group 2: Alkaline Earth Metals. Less reactive than Group 1; $2$ valence electrons.
  • Groups 3-12: Transition Metals. Have multiple oxidation states (charges) and form colored compounds and solutions.
  • Group 17: Halogens. Most reactive nonmetals; react readily with Group 1; all three states of matter (solid, liquid, gas) represented; $7$ valence electrons.
  • Group 18: Noble (Inert) Gases. Unreactive because they have a stable, complete octet of $8$ valence electrons.
• States of Matter at Room Temperature
  • Most elements are solids.
  • Liquids: $Br_2$ and $Hg$.
  • Gases: $N_2, Cl_2, H_2, O_2, F_2$ and the Noble Gases.
• Diatomic Elements
  • Elements that exist as two bonded atoms: $Br_2, I_2, N_2, Cl_2, H_2, O_2, F_2$.

Periodic Trends (Table S)

• Atomic Radii ($pm$)
  • The size of an atom ($\frac{1}{2}$ the distance between nuclei).
  • Decreases toward the top right.
  • Across a Period: Increased nuclear charge pulls the valence shell tighter.
  • Down a Group: More shells make the atom bigger.
• Ionization Energy (IE) ($kJ$ or $J$)
  • Energy required to remove the outermost valence electron.
  • Increases toward the top right (smaller atoms have a stronger pull on electrons).
• Electronegativity (Scale 0-4)
  • The ability to attract electrons in a bond. Fluorine has the highest EN ($4.0$).
  • Increases toward the top right.
• Ionic Radii
  • Metals: Lose electrons; ionic radius is smaller than atomic radius (more protons pulling on fewer electrons).
  • Non-Metals: Gain electrons; ionic radius is larger than atomic radius (same protons pulling on more electrons; weaker pull).

Chemical Bonding

• Stability: Atoms bond to become more stable by achieving lower potential energy and a "stable octet" ($8$ valence electrons). Hydrogen is an exception, stable with a "duplet" ($2$ valence electrons).
• Energy in Bonding (BARF)
  • Breaking bonds Absorbs energy (Endothermic).
  • Releasing energy happens when bonds Form (Exothermic).
• Ionic Bonds
  • Electrons are transferred from a Metal to a Nonmetal.
  • Ions are formed; attraction is electrostatic.
  • $\Delta EN > 1.7$.
  • Properties: Crystal lattice, high MP/BP, hard, brittle, conductive only as liquids (molten) or in solution (aq).
• Covalent Bonds
  • Electrons are shared ($2$ electrons per bond) between Nonmetals only.
  • $\Delta EN \le 1.7$.
  • Nonpolar Bond: Equal sharing (same EN).
  • Polar Bond: Unequal sharing (different EN).
  • Properties: Soft, low MP/BP, nonconductors, volatile (odorous).
• Metallic Bonds
  • Occurs in Metals only.
  • Electrons are delocalized in a "sea of mobile electrons".
  • Properties: Good conductors of heat and electricity, malleable, ductile.
• Special Cases
  • Polyatomic Ion compounds: Contain both ionic and covalent bonds. Example: $NaNO_3$.
  • Network Solids: Extended network of covalent bonds (e.g., Diamond, $SiO_2$). Very hard, very high MP/BP, nonconductors.

Lewis Dot Structures

• Ionic Compounds
  • Use brackets for ions.
  • Metal ion has no dots (lost valence shell) and shows a positive charge.
  • Nonmetal ion has $8$ dots and shows a negative charge.
• Covalent (Molecular) Compounds
  • Lines represent shared pairs ($1$ line = $2\,e^-$).
  • All atoms (except H) must have an octet ($8$ dots/lines).
  • If electrons are insufficient, form double or triple bonds.

Molecular Shape and Polarity

• SNAP Rule: Symmetrical = Nonpolar; Assymetrical = Polar.
• Shapes
  • Linear: Can be symmetrical or asymmetrical.
  • Tetrahedral: Typically symmetrical (Nonpolar if external atoms are same).
  • Pyramidal: Asymmetrical (Polar).
  • Bent: Asymmetrical (Polar).

Intermolecular Forces (IMF)

• Hydrogen Bonding: Not a real bond, but the strongest IMF. Occurs between polar molecules where H is bonded to F, O, or N (FON). Examples: $H_2O, NH_3, HF$.
• Dipole-Dipole: Between two polar molecules.
• Molecule-Ion (Dipole-Ion): Between a polar molecule (like water) and an ion. Occurs in aqueous solutions ($aq$).
• London Dispersion Forces (Van der Waals): Weakest IMF. Occurs between nonpolar molecules.
• Relationship: Stronger IMF leads to higher Melting/Boiling points and lower volatility.

Formula Writing and Nomenclature

• General Rules
  • Total charge of a compound is zero.
  • Use Table S for names and Table E for Polyatomic Ions (PAI).
  • State symbols: $(s)$ solid, $(l)$ liquid, $(g)$ gas, $(aq)$ aqueous solution.
• Writing Formulas
  • Binary Ionic: Metal + Nonmetal. Criss-cross oxidation numbers and drop signs.
  • Polyatomic Ionic: Use Table E. Use parentheses if more than one PAI is needed.
  • Stock System: Roman numerals indicate the charge of the metal.
  • Binary Molecular: Two nonmetals. Use prefixes (mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca-).
• Naming
  • Binary Ionic: Name metal, name nonmetal with "-ide" ending.
  • Molecular: Use prefixes. Example: $PF_3$ is Phosphorus trifluoride.

Chemical Reactions

• Synthesis: $A + B \rightarrow AB$. Two reactants form one product.
• Decomposition: $AB \rightarrow A + B$. One reactant breaks into multiple products.
• Single Replacement: $A + BC \rightarrow AC + B$. Use Table J (Activity Series); the single metal must be higher on the table to replace the metal in the compound.
• Double Replacement: $AB + CD \rightarrow AD + CB$. Use Table F to determine if an insoluble precipitate ($ppt$) forms.
• Combustion: Hydrocarbon + Oxygen ($O_2$) $\rightarrow$ Carbon Dioxide ($CO_2$) + Water ($H_2O$).

The Mole and Stoichiometry

• The Mole: A counting unit. $1\,mole = 6.02 \times 10^{23}$ representative units.
• Conservation of Mass: Mass of Reactants = Mass of Products. Atoms are conserved.
• Gram Formula Mass (GFM): Molar mass. Sum of the atomic masses of all elements in a formula ($g/mol$).
• Percent Composition (Table T): $\frac{\text{part mass}}{\text{total mass}} \times 100$.
• Hydrates: Compounds with water in their structure. Percent water = $\frac{\text{mass of } H_2O}{\text{total mass of hydrate}} \times 100$.
• Empirical Formula: The lowest whole-number ratio of atoms in a compound.
• Molecular Formula: The actual formula. To find it: $\frac{\text{Molecular Mass}}{\text{Empirical Mass}} = \text{multiplier}$. Multiply empirical subscripts by this whole number.

States of Matter and Thermodynamics

• Solids: Definite volume/shape; regular geometric or crystalline structure; particles vibrate in fixed positions; strong IMF.
• Liquids: Definite volume; indefinite shape; particles flow; moderately strong IMF.
• Gases: Indefinite volume/shape; constant random straight-line motion; particles far apart; weakest IMF.
• Phase Changes
  • Endothermic (Absorb energy, increase entropy): Melting (Fusion), Boiling (Vaporization), Sublimation (Solid to Gas, e.g., $CO_2, I_2$).
  • Exothermic (Release energy, decrease entropy): Freezing (Solidification), Condensation, Deposition (Gas to Solid).
• Heating/Cooling Curves
  • Plateaus: Phase changes. Temperature is constant. Average Kinetic Energy ($KE$) is constant. Potential Energy ($PE$) increases (heating) or decreases (cooling).
  • Slopes: Single phase. Temperature changes. $KE$ changes. $PE$ is constant.
• Heat Calculations (Table T)
  • $q = m \times C \times \Delta T$: Used for temperature changes.
  • $q = m \times H_f$: Used for melting/freezing (heat of fusion).
  • $q = m \times H_v$: Used for vaporization/condensation (heat of vaporization).
  • Temperature is the measure of Average Kinetic Energy.

Gas Laws

• Kinetic Molecular Theory (KMT) for Ideal Gases
  • Particles move in constant, random, straight-line motion.
  • Collisions are elastic (energy is transferred, not lost).
  • Particle volume is negligible compared to the distance between them.
  • There are no attractive (intermolecular) forces between particles.
• Real Gases vs. Ideal Gases
  • Real gases behave MOST like ideal gases at High Temperature and Low Pressure and when molecules are small (e.g., $H_2, He$).
• Gas Formulas (Table T)
  • Combined Gas Law: $\frac{P_1 V_1}{T_1} = \frac{P_2 V_2}{T_2}$. Note: Temperature must be in Kelvin ($K = {^\circ}C + 273$).
  • Boyle's Law: $PV = k$.
  • Charles's Law: $\frac{V}{T} = k$.
  • Gay-Lussac's Law: $\frac{P}{T} = k$.
• Avogadro's Hypothesis: Equal volumes of different gases at the same Temperature and Pressure contain an equal number of molecules.
• Standard Temperature and Pressure (STP)
  • Pressure: $1\,atm = 101.3\,kPa = 760\,torr = 760\,mmHg$.
  • Temperature: $0\,{^\circ}C = 273\,K$.
• Vapor Pressure: Pressure exerted by vapor over a liquid. Equilibrium boiling occurs when Vapor Pressure ($P_{vap}$) equals Atmospheric Pressure ($P_{atm}$). Weaker IMF results in higher vapor pressure and lower boiling points.

Solutions

• Terms:
  • Solute: Substance being dissolved.
  • Solvent: Substance doing the dissolving (e.g., $H_2O$).
  • Miscible: Liquids that mix.
  • Solubility: "Like Dissolves Like" (Polar dissolves polar/ionic; nonpolar dissolves nonpolar).
• Solubility Conditions
  • Solids: Solubility increases with higher temperature.
  • Gases: Solubility increases with lower temperature and higher pressure.
• Table G (Solubility Curves)
  • Unsaturated: Below the line; can hold more solute.
  • Saturated: On the line; maximum solute dissolved.
  • Supersaturated: Above the line; contains more than the maximum (unstable).
• Concentration Formulas
  • Molarity (M): $\frac{\text{moles of solute}}{\text{L of solution}}$.
  • Parts Per Million (ppm): $\frac{\text{grams of solute}}{\text{grams of solution}} \times 1,000,000$.
  • Percent by Mass/Volume: $\frac{\text{part}}{\text{whole}} \times 100$.
  • Dilution: $M_1 V_1 = M_2 V_2$.
• Colligative Properties:
  • As particle concentration increases, Boiling Point increases and Freezing Point decreases.
  • Ionic substances (electrolytes) have a greater effect than covalent substances because they dissociate into multiple particles (e.g., $CaCl_2 \rightarrow 3$ particles; $C_6H_{12}O_6 \rightarrow 1$ particle).

Kinetics and Equilibrium

• Collision Theory: For a reaction to occur, particles must collide with proper Energy and proper Orientation.
• Factors Affecting Rate
  • Temperature: Increases $\text{KE}$, more effective collisions.
  • Concentration: More particles, more collisions. For gases, increase pressure/decrease volume.
  • Surface Area: More exposed particles.
  • Nature of Reactants: Ionic reactions are faster than covalent.
  • Catalyst: Lowers Activation Energy ($E_a$) by providing an alternate pathway. Does not change $\Delta H$.
• Thermodynamics
  • Enthalpy (\Delta H): Heat of reaction. $\Delta H = PE_{products} - PE_{reactants}$.
    • Negative ($-$): Exothermic (heat released).
    • Positive ($+$): Endothermic (heat absorbed).
  • Entropy (\Delta S): Measure of disorder. Gases have highest entropy. Solids have lowest.
  • Spontaneity: Nature favors Lower Energy (-\Delta H) and Higher Entropy (+\Delta S).
• Equilibrium
  • Rate of Forward Reaction = Rate of Reverse Reaction.
  • Concentrations of reactants and products remain constant.
• Le Chatelier's Principle
  • Concentration: Shift away from additions; shift toward removals.
  • Temperature: Shift away from heat (Exo/Endo determines which side heat is on).
  • Pressure: Increase pressure shifts to the side with fewer moles of gas.

Acids, Bases, and Salts

• Properties
  • Acids (Table K): Sour, electrolytes ($H^+$ ions), pH $< 7$, turn litmus red, phenolphthalein is colorless, react with metals to produce $H_2(g)$.
  • Bases (Table L): Bitter, slippery, electrolytes ($OH^-$ ions), pH $> 7$, turn litmus blue, phenolphthalein is pink.
• Definitions
  • Arrhenius: Acids produce $H^+$; Bases produce $OH^-$ in solution.
  • Bronsted-Lowry: Acids are proton ($H^+$) donors; Bases are proton ($H^+$) acceptors.
• pH Scale: Each change of $1$ pH unit is a $10\text{-fold}$ change in $[H^+]$. A decrease in pH means an increase in acidity.
• Neutralization: $\text{Acid} + \text{Base} \rightarrow \text{Salt} + H_2O$.
• Titration: Used to find unknown concentration. Formula: $M_A V_A = M_B V_B$.

Oxidation-Reduction (Redox)

• LEO says GER
  • Loss of Electrons = Oxidation (Oxidation number increases).
  • Gain of Electrons = Reduction (Oxidation number decreases).
• Oxidation Numbers
  • Free elements = $0$.
  • Ions = their charge.
  • Group 1 atoms in compounds = $+1$.
  • Fluorine = always $-1$.
  • Hydrogen = $+1$ (with nonmetals) or $-1$ (with metals).
  • Oxygen = $-2$ (except $-1$ in peroxides or $+2$ with $F$).
• Electrochemical Cells
  • Voltaic Cell: Spontaneous; chemical energy to electrical energy; involves two half-cells and a salt bridge (to balance ion flow). An Ox, Red Cat (Anode = oxidation; Cathode = reduction). Electrons flow from Anode ($-$ ) to Cathode ($+$).
  • Electrolytic Cell: Nonspontaneous; requires a battery/power source; electrical energy to chemical energy. Anode is positive; Cathode is negative. Used for electroplating (object to be plated is the Cathode).

Organic Chemistry

• Hydrocarbons (Table P & Q)
  • Alkanes: Only single bonds; saturated; formula $C_n H_{2n+2}$.
  • Alkenes: One double bond; unsaturated; formula $C_n H_{2n}$.
  • Alkynes: One triple bond; unsaturated; formula $C_n H_{2n-2}$.
• Functional Groups (Table R)
  • Alcohols: $-OH$ (hydroxyl).
  • Ethers: $-O-$ bridge.
  • Aldehydes: $CHO$ at end.
  • Ketones: $C=O$ in middle.
  • Organic Acids: $-COOH$.
  • Esters: $-COOC-$ (smell/fragrance).
• Isomers: Same molecular formula, different structural formula (different properties).
• Organic Reactions
  • Substitution: Halogen replaces H on a saturated alkane.
  • Addition: Halogen added across a double/triple bond.
  • Fermentation: $\text{Sugar} + \text{enzyme} \rightarrow \text{Alcohol} + CO_2$.
  • Esterification: $\text{Acid} + \text{Alcohol} \rightarrow \text{Ester} + H_2O$.
  • Saponification: Hydrolysis of fats with base to make soap.
  • Polymerization: Joining small monomers into a polymer chain.

Nuclear Chemistry

• Stability: Depends on the ratio of protons to neutrons. All elements above Atomic Number $82$ are unstable.
• Transmutation: Changing the identity of a nucleus.
  • Natural: Spontaneous decay ($1$ reactant).
  • Artificial: Bombarding a nucleus with a particle ($2$ reactants).
• Decay Particles (Table N)
  • Alpha ($\alpha$): $^4_2 He$.
  • Beta ($\beta^-$): $^0_{-1} e$.
  • Positron ($\beta^+$): $^0_{+1} e$.
  • Gamma ($\gamma$): Pure energy; no mass/charge.
• Fission: Splitting a heavy nucleus ($U-235$) into lighter nuclei.
• Fusion: Combining light nuclei ($H + H$) into a heavier one ($He$). Releaces more energy than fission.
• Applications:
  • $C-14$ (Dating organic fossils).
  • $U-238$ (Dating rocks/geology).
  • $I-131$ (Thyroid issues).
  • $Co-60$ (Cancer treatment).
  • $Tc-99$ (Brain tumors).

Math and Laboratory Skills

• Significant Figures
  • Leading zeros never count ($0.0045$ is $2$ sig figs).
  • Trailing zeros only count if a decimal is present ($200.0$ is $4$ sig figs).
• Error Calculation: $\text{\% Error} = \frac{|accepted - measured|}{accepted} \times 100$.
• Measurement: Always estimate one final digit past the marking on the instrument scale.
• Precision vs. Accuracy: Precision is consistency; Accuracy is closeness to target value.
• Safety: Always wear goggles and exercise caution with burners and chemicals.