Chemistry Regents Comprehensive Study Guide

Periodic Table and Trends

• Elements are organized into specific structures:

  • Periods: Horizontal rows across the table.

  • Groups/Families: Vertical columns. Elements in the same group possess the same number of valence electrons and exhibit similar chemical properties.

• Important Groups and Characteristics:

  • Group 1: Alkali Metals; $1$ valence electron; Oxidation number: $+1$.

  • Group 2: Alkali Earth Metals; $2$ valence electrons; Oxidation number: $+2$.

  • Group 13: Boron family; $3$ valence electrons.

  • Group 14: Carbon family; $4$ valence electrons.

  • Group 15: Nitrogen family; $5$ valence electrons.

  • Group 16: Oxygen family; $6$ valence electrons; Oxidation number of Oxygen: $-2$.

  • Group 17: Halogens; $7$ valence electrons; Oxidation number of Fluorine: $-1$.

  • Group 18: Noble gases; $8$ valence electrons (except Helium, which has $2$).

• Definitions of Trends and Forces:

  • Oxidation #: The charge an atom possesses when it becomes an ion.

  • Shielding effect: Occurs when inner electrons block or shield the attraction between the nucleus and the valence electrons.

  • Nuclear charge: The total positive charge located in an atom’s nucleus, which is determined solely by the number of protons.

• Periodic Trends:

  • 1. Atomic Radius (Size of the atom):

    • Across a period (left to right): Decreases because more protons are added, which pull the electrons closer to the nucleus.

    • Down a group: Increases because more energy shells are added and the shielding effect occurs.

  • 2. Ionization Energy (Energy required to remove an electron):

    • Across a period: Increases because a higher number of protons hold onto electrons more tightly.

    • Down a group: Decreases because valence electrons are positioned farther from the nucleus.

  • 3. Electronegativity (Ability of an atom to attract electrons within a bond):

    • Across a period: Increases because more protons hold onto electrons more tightly.

    • Down a group: Decreases because the nucleus is effectively smaller/farther relative to the bond.

  • 4. Metallic Character (Strength of metallic behavior):

    • Across a period: Decreases.

    • Down a group: Increases.

Characteristics of Elements

• Metals:

  • Location: Left and center of the periodic table.

  • Properties: Shiny (lustrous), conductive of heat and electricity, and opaque.

  • Behavior: They lose electrons to fill orbitals, becoming cations (positive ions). They are typically solid at room temperature.

• Nonmetals:

  • Location: Right side of the periodic table.

  • Properties: Dull, nonconductive, and brittle.

  • Behavior: They gain electrons to become anions (negative ions). Most nonmetals are gases at room temperature.

• Metalloids:

  • Location: The zigzag staircase between metals and nonmetals.

  • Properties: Possess a mix of metal and nonmetal characteristics. They can act as metals or nonmetals depending on the situation. They act as semiconductors when heated.

Unit 1: States of Matter and Energy

• Comparison of States of Matter:

  • Solid: Definite shape and definite volume; particles are tightly packed and vibrate in place; lowest kinetic energy; strongest intermolecular forces.

  • Liquid: Takes the shape of its container and has a definite volume; particles are close together but can move/slide past each other; medium kinetic energy; moderate intermolecular forces.

  • Gas: No definite shape and no definite volume; particles are far apart and move freely/rapidly; highest kinetic energy; weakest intermolecular forces.

  • Plasma: A superheated state of matter consisting of charged particles (ions and free electrons). Particles are very far apart and move extremely fast/randomly. It conducts electricity.

• Phase Changes:

  • Melting: Solid to Liquid.

  • Freezing: Liquid to Solid.

  • Vaporization: Liquid to Gas.

  • Condensation: Gas to Liquid.

  • Sublimation: Solid to Gas.

• Energy Types:

  • Kinetic Energy: Energy of motion; increases as temperature increases.

  • Potential Energy: Stored energy due to position or arrangement.

• Electrostatic Potential Energy and Distance:

  • Like Charges ($+$/$+$ or $-$/$-$):

    • As distance increases: Electrostatic potential energy decreases because charges want to move away naturally.

    • As distance decreases: Electrostatic potential energy increases because work is required to force them together against repulsion.

  • Opposite Charges ($+$/$-$ or $-$/$+$):

    • As distance increases: Electrostatic potential energy increases because work is required to pull the attracted charges apart.

    • As distance decreases: Electrostatic potential energy decreases because they are naturally attracted to each other.

Atomic Structure and Electrostatics

• Subatomic Particles:

  • Atoms are the basic units of matter. The nucleus contains Protons and Neutrons. The electron cloud outside the nucleus contains Electrons and makes up the atom's volume.

  • Protons: $1+$ charge, $1\,u$ mass. Effect: Increases charge; changes the element if added or removed.

  • Electrons: $1-$ charge, approximately $0$ ($0.0005\,u$) mass. Effect: Decreases charge.

  • Neutrons: $0$ charge (neutral), $1\,u$ mass. Effect: Increases mass.

• Quantifying the Atom:

  • Atomic number = number of protons.

  • Mass number = Protons + Neutrons.

  • Measured in unified atomic mass units ($u$).

• Electrostatics and Polarization:

  • Electrostatic Fields: Invisible spaces surrounding charged objects influencing interactions.

  • Coulomb’s Law:

    • $F \propto \frac{1}{d^2}$ (Force is inversely squared proportional to distance; as distance decreases, force increases).

    • $F \propto C$ (Force is directly proportional to charge; as charge increases, force increases).

  • Polarization: Separation of positive and negative charges.

  • Induced Polarization: A neutral object develops partial charges when close to a charged object.

  • Example (Balloon, Sweater, Wall):

    1. All start neutral.

    2. Balloon rubs sweater; electrons transfer to balloon (static electricity).

    3. Balloon becomes negatively charged with a strong field.

    4. Electrons in the wall repel the balloon's electrons.

    5. The side of the wall closest to the balloon becomes partially positive; the far side becomes partially negative.

    6. The charge imbalance in the wall is created through induced polarization.

• Definitions:

  • Electrostatic force: Attraction or repulsion between charged objects.

  • Electricity: Movement of electrons.

  • Static electricity: Buildup of charges on a surface via electron transfer.

  • Partial charges: Uneven charge distribution from polarization.

  • Empirical: Evidence based on observations/experiments.

  • Theoretical: Evidence based on calculations/models.

Properties of Water and Electrolytes

• Electrolytes: Substances that help conduct electricity. Salts made of cations and anions conduct when dissolved and dispersed in water.

• Water Properties:

  • Cohesion: Attraction between molecules of the same substance; allows droplets to form. Water is polar, attracting other water molecules.

  • Adhesion: Attraction between different substances. Water rises in tubes (capillary action) when adhesion forces are greater than cohesion.

  • Surface tension: A thin "film" on water caused by cohesion.

  • Specific heat: Water absorbs and releases energy slowly, meaning it takes significant energy to change its temperature.

  • Hydrogen bonds: Attraction between a hydrogen atom (partially positive) and an electronegative atom on another molecule.

  • Attraction to Salt ($NaCl$): Oxygen in $H_2O$ is partially negative (attracted to Cations like $Na^+$) and Hydrogen is partially positive (attracted to Anions like $Cl^-$).

• Intermolecular vs. Intramolecular Forces:

  • Intramolecular: Strong forces holding atoms inside a molecule together.

  • Intermolecular (IMF): Weak attractions between molecules affecting boiling/melting points and evaporation.

  • Strength and Evaporation:

    • Strong IMFs: Require more thermal energy to evaporate; result in a slower evaporation rate.

    • Weak IMFs: Require less thermal energy; result in a faster evaporation rate.

  • Vapor pressure: Pressure from vapor above a liquid; indicates how easily molecules escape the system.

Unit 2: Chemical Bonding

• Stability and the Octet Rule:

  • Atoms bond to become stable by achieving a full outer shell of $8$ valence electrons.

• Valence Electron Behavior:

  • Metals: Have $1$-$3$ valence electrons. It is easier to lose them than to gain more. Losing electrons reveals the full energy shell underneath and turns the metal into a Cation.

    • Example: $Na \rightarrow Na^+ + e^-$.

  • Nonmetals: Have $5$-$7$ valence electrons. It is easier to gain electrons to reach an octet, turning them into Anions.

    • Example: $Cl + e^- \rightarrow Cl^-$.

• Ionic Bonds and Compounds:

  • Formation: Metal loses electrons (becomes positive) + Nonmetal gains electrons (becomes negative). Opposite charges attract to form the bond.

  • Properties: Hard solids, tightly packed, crystalline/geometric form, high melting/boiling points, conduct electricity when in water.

  • Charges: Ionic compounds are always neutrally charged (total charge = $0$).

  • Naming: Metal comes first, then the nonmetal with the suffix changed to "-ide."

• Electron Configuration Shell Limits:

  • 1st shell: Max $2$ electrons.

  • 2nd shell: Max $8$ electrons.

  • 3rd shell: Max $8$-$18$ electrons.

• Models:

  • Bohr Model: Shows core (inner) and valence (outer) electrons.

  • Lewis Dot Structure: Shows only valence electrons and unbonded lone pairs.

• Metallic Bonding:

  • Characterized by delocalized/free electrons (mobile valence electrons not attached to a specific atom).

  • Durability: Metals deform (malleability) instead of shattering; electrons act as a "flexible glue" between cations.

  • Conductivity: Delocalized electrons easily carry electrical charge.

  • Ductility: Mobile electrons hold ions together while they are drawn into wires.

  • Luster: Shining and reflecting light.

• Covalent Bonding:

  • Occurs between two nonmetal atoms sharing electrons. The shared electrons are called a "bonding pair."

  • Bond Polarity:

    • Non-polar: Equal sharing; identical/similar electronegativities; no partial charges.

    • Polar: Unequal sharing; different electronegativities; creates dipoles (partial charges).

  • Molecular Polarity:

    • Non-Polar molecule: Symmetrical shape; no partial charges.

    • Polar molecule: Asymmetrical shape; electrons pulled toward one side, creating slightly positive and negative ends.

Unit 3: Stoichiometry

• Law of Conservation of Mass: Mass cannot be created or destroyed. Total mass of reactants must equal total mass of products.

• The Mole Concept:

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

  • $1$ mole = Gram Formula Mass ($GFM$) in grams.

  • $1$ mole = $22.4\,L$ of gas at STP.

• Conversion Equations:

  • Moles to Particles: \text{# of particles} = 6.02 \times 10^{23} \times \text{# of moles}.

  • Moles to Grams: \text{# of moles} = \frac{\text{Given Mass}}{GFM}.

  • Particles to Moles: \text{# of moles} = \frac{\text{particles}}{6.02 \times 10^{23}}.

  • Mole to Volume: \text{# of moles} = \frac{\text{Given Volume}}{22.4\,L}.

  • Mole to Mole Ratio: $\frac{\text{Mole of Unknown (from equation)}}{\text{Mole of Known (from equation)}}$.

Unit 4: Kinetics and Equilibrium

• Enthalpy (Delta H):

  • $\Delta H = (\text{potential energy of products}) - (\text{potential energy of reactants})$.

  • Exothermic: Delta H is negative; energy is released (appears on product side); products have lower potential energy; environment gets hotter.

    • General Form: $\text{Reactants} \rightarrow \text{Products} + \text{Energy}$.

  • Endothermic: Delta H is positive; energy is absorbed (appears on reactant side); products have higher potential energy; environment gets cooler.

    • General Form: $\text{Reactants} + \text{Energy} \rightarrow \text{Products}$.

• Reaction Requirements:

  • Activation Energy: Minimum energy required to start a reaction. A catalyst lowers this energy but does not change enthalpy.

  • Activated Complex: Transition stage at impact where bonds are broken and reformed.

  • Collision Theory: Particles must collide with Proper Orientation and Enough Energy (activation energy).

• Equilibrium and Le Chatelier’s Principle:

  • Dynamic Equilibrium: Rate of forward reaction = Rate of reverse reaction. Concentrations remain constant.

  • Shifts:

    • Add reactant: Shift right.

    • Remove reactant: Shift left.

    • Add product: Shift left.

    • Remove product: Shift right.

    • Temperature (Exothermic): Increase T shifts left; Decrease T shifts right.

    • Temperature (Endothermic): Increase T shifts right; Decrease T shifts left.

    • Pressure: Only affects gases. Increase pressure shifts toward the side with fewer moles of gas; decrease pressure shifts toward the side with more moles of gas.

• Types of Reactions:

  • Synthesis: $A + B \rightarrow AB$.

  • Decomposition: $AB \rightarrow A + B$.

  • Single Replacement: $A + BC \rightarrow AC + B$ (More active metal replaces less active metal).

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

  • Combustion: $\text{C}_X\text{H}_Y + \text{O}_2 \rightarrow \text{CO}_2 + \text{H}_2\text{O}$.

• Oxidation Number Rules:

  • 1. Free element = $0$.

  • 2. Monatomic ions = their charge.

  • 3. Group 1 = $+1$; Group 2 = $+2$.

  • 4. Hydrogen = $+1$; Oxygen = $-2$; Fluorine = $-1$.

  • 5. Sum in a neutral compound = $0$.

  • 6. Sum in a polyatomic ion = ion charge.

Unit 5: Titration and Redox

• Titration:

  • Laboratory technique to find unknown concentration. Analyte (unknown concentration) reacts with Titrant (known concentration).

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

  • Titration Formula: $M_A V_A = M_B V_B$.

  • Molarity Formula: $M = \frac{\text{moles of solute}}{\text{liters of solution}}$.

• Redox Reactions (Reduction + Oxidation):

  • Oxidation (OIL): Loss of electrons; oxidation number increases.

    • Example: $Na \rightarrow Na^+ + e^-$ ($0$ to $+1$).

  • Reduction (RIG): Gaining of electrons; oxidation number decreases.

    • Example: $Cl_2 + 2e^- \rightarrow 2Cl^-$ ($0$ to $-1$).

• Cells:

  • An Ox / Red Cat: Anode is Oxidation; Cathode is Reduction.

  • Voltaic Cell: Spontaneous; Chemical to Electrical energy; contains salt bridge and 2 cells; Anode is negative (-), Cathode is positive (+).

  • Electrolytic Cell: Non-spontaneous (requires battery); Electrical to Chemical energy; 1 cell; Anode is positive (+), Cathode is negative (-).

  • Electroplating: Process of coating metal. Anode (metal coating) loses mass; Cathode (object being coated) gains mass.

Unit 6: Gases and Solutions

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

  • Temperature must be in Kelvin ($K = \text{Celsius} + 273.15$).

  • Gas Relationships:

    • Pressure vs. Volume: Indirect (P up, V down).

    • Volume vs. Temperature: Direct (V up, T up).

    • Pressure vs. Temperature: Direct (P up, T up).

• Solubility:

  • Unsaturated: Below saturation point; more can dissolve.

  • Saturated: At saturation point; no more can dissolve.

  • Supersaturated: Above saturation point; extra solute settles at bottom.

• Colligative Properties (Depend on # of particles, not type):

  • Boiling point elevation: Adding solute raises boiling point.

  • Freezing point depression: Adding solute lowers freezing point.

  • Vapor pressure lowering: Adding solute lowers vapor pressure.

Units 7-10: Acids, Bases, and Nuclear Chemistry

• Acids vs. Bases:

  • Acids: pH < 7; taste sour; produce $H^+$/$H_3O^+$ ions; react with metals to make $H_2$ gas; electrolyes.

  • Bases: pH > 7; taste bitter; slippery; produce $OH^-$ ions; electrolyes.

  • Arrhenius Theory: Acid makes $H_3O^+$; Base makes $OH^-$ in water.

  • Bronsted-Lowry Theory: Acid is a proton ($H^+$) donor; Base is a proton ($H^+$) acceptor.

• pH Indicators:

  • Range Logic: If pH is below lower limit, it shows the left color. If above the higher limit, it shows the right color. Between limits, it shows a mix.

  • Methyl Orange Range ($3.1$ - $4.4$): Red if < 3.1; Yellow if > 4.4.

• Nuclear Chemistry:

  • Unstable nucleus has an unstable ratio of protons to neutrons.

  • Radioactive Decay: $\text{Parent atom} \rightarrow \text{Daughter atom} + \text{Radioactive energy}$.

  • Isotopes: Same protons, different neutrons (different mass).

  • Notation: $^A_Z X$ where $A$ is mass and $Z$ is atomic number.

• Electromagnetic Spectrum (Low energy to High energy):

  • 1. Radio, 2. Microwave, 3. Infrared, 4. Visible Light, 5. Ultraviolet, 6. X-rays, 7. Gamma.

  • Relationships:

    • Frequency vs. Energy: Direct.

    • Wavelength vs. Frequency: Inverse.

    • Wavelength vs. Energy: Inverse.

  • Ionizing Energy: High frequency (UV, X-rays, Gamma); can remove electrons and damage DNA.

• Photosynthesis:

  • Transforms visible light (electromagnetic energy) into chemical energy (glucose).

  • Uses $CO_2$ and $H_2O$; releases oxygen.

  • Chlorophyll absorbs blue and red light; reflects green. Photons absorbed are high frequency and emitted as lower frequency.