Chemistry Regents Exam - Comprehensive Review Notes

Chemistry Regents Exam Review

Topics Covered

• The Atom

• Moles and Stoichiometry

• Nuclear Chemistry

• Solutions

• Bonding

• Kinetics and Equilibrium

• Matter

• Acids, Bases, and Salts

• Energy

• Oxidation-Reduction (Redox)

• The Periodic Table

• Organic Chemistry

• Math and graphing skills are also required.

Packet Purpose

• The review packet is assembled from the NY State Core Curriculum.

• It is not a thorough review of the entire course but should be used with review sheets, past Regents exams, and Reference Tables.

• Emphasis on key ideas stressed by the Core Curriculum.

• Additional space is provided for personal notes.

• Active preparation is essential: study, ask questions, analyze problems, and attend review sessions.

Topic One: The Atom

1. Evolution of the Atomic Model:

   • Dalton's Model:

     • Elements are made of atoms.

     • Atoms of an element are the same.

     • Compounds are formed from combinations of atoms.

   • Rutherford Experiment:

     • Bombarded gold foil with alpha particles.

     • Showed atoms are mostly empty space with a small, dense, positively charged nucleus.

   • Bohr Model:

     • Small, dense, positively charged nucleus surrounded by electrons in circular orbits.

   • Wave-Mechanical Model (Modern Atomic Theory):

     • Small, dense, positively charged nucleus surrounded by electrons moving in an "electron cloud."

     • "Orbitals" are areas where an electron with a certain amount of energy is most likely to be found.

2. Atomic Structure:

   • Each atom has a positively charged nucleus with one or more orbiting, negatively charged electrons.

3. Nuclear Composition:

   • Protons and neutrons are found in the nucleus.

4. Charge of Subatomic Particles:

   • Protons: positive charge

   • Neutrons: no charge

   • Electrons: negative charge

5. Neutral Atoms:

   • The number of protons in an atom equals the number of electrons.

   • Positive charges of protons cancel negative charges of electrons, giving a neutral overall charge.

6. Mass of Subatomic Particles:

   • Mass of a proton: 1 amu

   • Mass of a neutron: 1 amu

   • Mass of an electron: almost 0 amu

   • 1 amu is defined as 1/12 the mass of a Carbon-12 atom.

   • The atomic mass of an atom is equal to the total number of protons and neutrons.

7. Electron Energy Levels:

   • Each electron in an atom has its own distinct amount of energy.

   • Ground State: When all electrons are at their lowest possible energy.

   • Electrons fill energy levels and orbitals starting with the one that requires the least energy (1s) and progressively move to those levels and orbitals that require increasing amounts of energy.

8. Excited State:

   • When an electron gains a specific amount of energy, it moves to a higher orbital.

9. Bright Line Spectrum:

   • When an electron returns from a higher energy state to a lower energy state, it emits a specific amount of energy, usually as light.

   • This emission can be used to identify an element (bright line spectrum).

   • A spectroscope is used to see the bright line spectrum.

10. Valence Electrons:

    • Outermost electrons are called valence electrons and affect the chemical properties of the element.

    • Atoms with a filled valence level are stable.

    • Most elements can have up to 8 electrons in their valence level, except for H and He, which can have only 2.

    • Atoms form bonds to fill their valence levels.

    • Orbital notation or Lewis structures can show the configuration of valence electrons.

11. Atomic Number:

    • Atoms of the same element contain the same number of protons.

    • Changing the number of protons changes the atom into a different element.

    • Atomic number is the number of protons in an atom of an element.

12. Isotopes:

    • Atoms with equal numbers of protons but different numbers of neutrons.

    • Isotopes of an element have the same atomic number (protons only) but different atomic masses (protons + neutrons).

13. Average Atomic Mass:

    • The average atomic mass of an element is the weighted average of its naturally occurring isotopes.

Topic Two: Nuclear Chemistry

1. Nuclear Stability:

   • The stability of an isotope depends on the ratio of protons to neutrons in the nucleus.

   • Most nuclei are stable, but some are unstable and spontaneously decay, emitting radiation.

   • Stable isotopes have a 1:1 ratio of protons and neutrons. Most radioactive isotopes have twice as many neutrons as protons.

   • All elements with an atomic number higher than 83 are radioactive.

2. Radioactive Decay:

   • Each isotope has a specific mode and rate of decay (see Table N).

   • The rate of decay is called half-life.

   • Half-life is a constant that can never be changed.

   • Half-life is the measure of the time it takes exactly one half of an amount of isotope to decay.

   • The amount of substance will never decay to zero.

3. Transmutation:

   • A change in the nucleus of an atom changes it to a new type of atom (i.e., a new element).

   • Transmutation can occur naturally or artificially.

   • Artificial transmutation requires the bombardment of a nucleus by high-energy particles.

4. Spontaneous Decay:

   • Spontaneous decay can involve the release of different particles from the nucleus.

   • The types of particles, as well as their masses and charges, can be found on Table O.

5. Nuclear Reactions:

   • Nuclear reactions include natural and artificial decay, nuclear fission, and nuclear fusion.

   • Nuclear fission occurs when the nucleus of an atom is split, artificially caused by bombarding the nucleus with a neutron.

   • Nuclear fusion combines two light nuclei to form heavier nuclei and is the process that powers the sun.

   • Nuclear fusion requires very high temperatures and is not yet ready for practical use. The main advantage is that the products are not radioactive wastes (as with fission).

6. Nuclear Equations:

   • Nuclear reactions can be represented by equations that include symbols representing atomic nuclei (with mass number and atomic number), subatomic particles (with mass and charge), and emitted particles.

7. Mass-Energy Conversion:

   • Energy from nuclear reactions comes from a very small fraction of mass that is lost – the reaction converts matter into energy.

   • Einstein's $E=mc^2$ describes the relationship between energy and matter.

8. Energy Release:

   • The energy released from nuclear reactions is much greater than that released from chemical reactions.

9. Risks of Radioactive Isotopes:

   • Risks include biological exposure (radiation poisoning and cancer), long-term storage and disposal, and nuclear accidents.

10. Uses of Radioactive Isotopes:

    • Radioactive isotopes may be used in medicine (tracing chemical and biological processes), radioactive dating, industrial measurement, nuclear power, and detection and treatment of disease.

Topic Three: Bonding

1. Formation of Chemical Compounds:

   • Chemical compounds are formed when atoms are bonded together.

   • Breaking a chemical bond is an endothermic process.

   • Forming a chemical bond is an exothermic process.

   • Compounds have less potential energy than the individual atoms they are formed from.

2. Categories of Compounds:

   • Two major categories: ionic and molecular (covalent) compounds.

3. Properties of Compounds:

   • Ionic substances:

     • High melting and boiling points.

     • Form crystals.

     • Dissolve in water (dissociation).

     • Conduct electricity in solution and as a liquid.

   • Covalent or molecular substances:

     • Lower melting and boiling points.

     • Do not conduct electricity.

   • Polar substances are dissolved only by another polar substance. Non-polar substances are dissolved only by other non-polar substances.

4. Types of Chemical Bonds:

   • Ionic: valence electrons are transferred from one atom to another.

   • Covalent: valence electrons are shared between atoms.

   • Metallic: valence electrons are mobile in a free-moving "sea" of electrons.

5. Multiple Covalent Bonds:

   • In multiple (double or triple) covalent bonds, more than 1 pair of electrons are shared between two atoms.

6. Molecular Polarity:

   • Polarity determined by shape and charge distribution.

   • Polar molecules must have polar bonds.

   • Polar molecules are asymmetrical.

   • Nonpolar molecules are symmetrical and/or have no polar bonds.

7. Ion Formation and Radius: Gaining Electrons

   • When an atom gains an electron, it becomes a negative ion, and its radius increases.

8. Ion Formation and Radius: Losing Electrons

   • When an atom loses an electron, it becomes a positive ion, and its radius decreases.

9. Stable Electron Configuration:

   • Atoms gain a stable electron configuration by bonding with other atoms.

   • Atoms are stable when they have a full valence level.

   • Most atoms need 8 electrons to fill their valence level.

   • H and He only need 2 electrons to fill their valence level.

   • The noble gasses (group 18) have filled valence levels and do not normally bond with other atoms.

10. Electron-Dot Diagrams (Lewis Structures):

    • Represent the valence electron arrangement in elements, compounds, and ions.

    • Electrons are arranged by their orbitals.

    • The first two electrons are placed together in the “s” orbital.

    • The remaining electrons are spread among the 3 “p” orbitals.

    • The “s” orbital must be filled first. Then each “p” orbital must have one electron before another “p” orbital gains a second.

11. Electronegativity:

    • Indicates how strongly an atom of an element attracts electrons in a chemical bond. Values are based on an arbitrary scale.

12. Bond Type and Polarity:

    • The electronegativity difference between two bonded atoms determines the type of bond and its polarity.

      • $0.0 - 0.4$ = non-polar covalent

      • $0.4 - 1.7$ = polar covalent

      • $1.7+$ = ionic

13. Bonding Guidelines:

    • Metals react with nonmetals to form ionic compounds.

    • Nonmetals bond with nonmetals to form covalent compounds (molecules).

    • Ionic compounds with polyatomic ions have both ionic and covalent bonds.

14. Intermolecular Forces (IMF):

    • Allow different particles to be attracted to each other to form solids and liquids.

    • Hydrogen bonds:

      • An example of a strong IMF between atoms.

      • Exist between atoms of hydrogen and oxygen, fluorine, or nitrogen.

      • Substances with hydrogen bonds tend to have much higher melting and boiling points than those without hydrogen bonds.

15. Physical Properties:

    • Explained in terms of chemical bonds and intermolecular forces, including conductivity, malleability, solubility, ductility, hardness, melting point, and boiling point.

Topic Four: Matter, Phases, and Gas Laws

1. Classification of Matter:

   • Classified as a pure substance or a mixture of substances.

   • Substance: fixed composition and uniform properties throughout the sample (elements and compounds).

2. Mixtures:

   • Composed of two or more different substances that may be physically separated.

   • Homogeneous (uniform – a solution) or heterogeneous (uneven).

   • Substances in a mixture retain their original properties.

   • Substances in a mixture may be separated by size, polarity, density, boiling and freezing points, and solubility.

   • Filtration and distillation are examples of processes used to separate mixtures.

3. Elements:

   • A substance composed of atoms with the same atomic number and cannot be broken down by chemical change.

4. Compounds:

   • Two or more elements bonded together, broken down by chemical changes.

   • Substances that form a compound gain new properties.

   • The ratio of substances in a compound is constant (e.g., water has a fixed $2:1$ ratio of hydrogen to oxygen).

5. Physical and Chemical Changes:

   • Physical change: rearrangement of existing particles (e.g., freezing, boiling).

   • Chemical change: formation of different substances with different properties.

   • Chemical and physical changes may be endothermic or exothermic.

6. Phases of Matter:

   • Solid, liquid, and gas, each with its own properties.

   • Solids:

     • Constant volume and shape.

     • Particles held in a rigid, crystalline structure.

   • Liquids:

     • Constant volume but a changing shape.

     • Particles are mobile but still held together by strong attraction.

   • Gases:

     • No set volume or shape.

     • Completely fill any closed container.

     • Particles have largely broken free of the forces holding them together.

7. Heating and Cooling Curves:

   • Trace the changes in temperature of a substance as it changes from solid to liquid to gas (or gas to liquid to solid).

   • During a phase change, there is no change in temperature; the line "flattens" until the phase change is complete.

   • During a phase change, the potential energy changes, while kinetic energy remains the same.

   • As temperature increases, kinetic energy increases.

8. Heat of Fusion ($H_f$):

   • Energy needed to convert one gram of a substance from solid to liquid.

9. Heat of Vaporization ($H_v$):

   • Energy needed to convert one gram of a substance from liquid to gas.

10. Specific Heat (C):

    • Energy required to raise one gram of a substance 1 degree (Celsius or Kelvin).

    • The specific heat of liquid water is $1 cal/g<em>°C$ or $4.2 J/g</em>K$.

11. Combined Gas Law:

    • States the relationship between pressure, temperature, and volume in a gas sample.

    • Increasing pressure causes a decrease in volume (inverse relationship).

    • Increasing temperature causes an increase in volume (direct relationship).

    • Increasing temperature causes an increase in pressure (direct relationship).

12. Ideal Gas Model:

    • Used to explain the behavior of gasses. A real gas is most like an ideal gas when it is at low pressure and high temperature.

13. Kinetic Molecular Theory (KMT):

    • All gas particles:

      • are in random motion.

      • have no forces of attraction between them.

      • have a negligible volume compared to the distances between them.

      • have collisions that result in the transfer of energy from one particle to another, but there is no net loss of energy from the collision.

14. Equal Volumes of Gasses:

    • Equal volumes of gasses at the same temperature and pressure have an equal number of particles.

Topic Five: Energy

1. Forms of Energy:

   • Chemical, electrical, electromagnetic, thermal, mechanical, nuclear.

   • Stored energy is potential energy.

   • Energy of motion is kinetic energy.

2. Law of Conservation of Energy:

   • Energy cannot be lost or destroyed, only changed from one form to another.

3. Heat:

   • A transfer of energy (often thermal) from a body of higher temperature to a body of lower temperature.

4. Temperature:

   • A measure of the average kinetic energy of the particles in a sample. NOT a form of energy; do not confuse with heat.

5. Kinetic and Potential Energy in Phase Changes:

   • Concepts can be used to explain physical processes such as fusion (melting), solidification (freezing), vaporization (boiling, evaporation), condensation, sublimation, and deposition.

6. Exothermic Processes:

   • Give off heat energy, typically causing the surrounding environment to become warmer.

7. Endothermic Processes:

   • Absorb energy, typically causing the surrounding environment to become colder.

Topic Six: The Periodic Table

1. Arrangement and Properties:

   • The placement of an element on the Periodic Table indicates its chemical and physical properties.

2. Increasing Atomic Number:

   • Elements are arranged in order of increasing atomic number.

3. Atomic Number and Element Identity:

   • The number of protons in an atom (atomic number) identifies the element.

   • Proton number changes only through nuclear reactions.

4. Atomic Mass:

   • The sum of protons and neutrons in the nucleus.

   • The mass number on the periodic table is a weighted average of the different isotopes of that element.

   • Electrons do not significantly add to the atomic mass.

5. Isotopic Notation:

   • Isotopes are identified by the sum of protons and neutrons.

   • Isotopes of the same element have the same number of protons and a different number of neutrons.

   • Examples: $^{14}_{6}C$, $^{14}C$, carbon-14, C-14

6. Classification of Elements:

   • Metals, non-metals, metalloids, and noble gasses, based on properties and location on the Periodic Table.

7. Differentiation by Physical Properties:

   • Examples include density, conductivity, malleability, hardness, ductility, solubility.

8. Differentiation by Chemical Properties:

   • Chemical properties describe how an element behaves in a chemical reaction.

9. Periods and Groups:

   • Elements arranged into periods (rows) and groups (columns).

10. Periods and Energy Levels:

    • Elements of the same period have the same number of occupied energy levels.

11. Groups and Valence Configuration:

    • Elements of the same group have the same valence configuration and similar chemical properties.

    • Group 1 (excluding H): alkali metals

    • Group 2: alkali earth metals

    • Group 17: halogens

    • Alkali metals, alkali earth metals, and halogens are highly reactive and not found as free elements in nature; they exist in compounds.

    • Group 18: noble or inert gasses with filled valence levels; they do not normally react.

12. Trends Down a Group:

    • Atomic radius increases.

    • Electronegativity decreases.

    • First ionization energy decreases.

    • Metallic character increases.

13. Trends Across a Period (Left to Right):

    • Atomic radius decreases.

    • Electronegativity increases.

    • First ionization energy increases.

    • Metallic character decreases.

14. Allotropes:

    • Some elements exist in two or more forms in the same phase with different molecular or crystal structures and properties.

    • Example: Carbon exists as graphite and diamond (a network solid).

Topic Seven: Moles and Stoichiometry

1. Compounds:

   • Substances composed of two or more different elements chemically combined in a fixed proportion.

   • Can only be broken down by chemical means.

2. Chemical Formulas:

   • Represented by specific formulas and assigned a name based on the IUPAC system.

3. Types of Chemical Formulas:

   • Empirical: elements in their simplest whole number ratios.

   • Molecular: the actual number of atoms per element in a single molecule.

   • Structural: shows the number of each type of atom as well as their physical arrangement.

4. Conservation Laws:

   • Chemical reactions show a conservation of mass, energy, and charge.

5. Balanced Chemical Equations:

   • Represent conservation of atoms.

6. Mole Ratios:

   • Coefficients in a balanced chemical equation determine mole ratios in the reaction.

7. Formula and Molar Mass:

   • Formula mass: sum of the atomic masses of its atoms.

   • Molar mass (gram formula mass): equals the mass of one mole of that substance.

8. Percent Composition:

   • Percent composition by mass of each element in a compound can be calculated mathematically.

9. Types of Chemical Reactions:

   • Synthesis, decomposition, single replacement, and double replacement.

Topic Eight: Solutions

1. Solutions Defined:

   • A homogeneous mixture of a solute dissolved in a solvent.

   • Solubility depends on temperature, pressure, and the nature of the solute and solvent.

   • "Like dissolves like": polar substances dissolve polar substances, and non-polar substances dissolve non-polar substances. Polar and non-polar do not mix.

2. Ionic Substances in Polar Solvents:

   • Ionic substances dissolve in polar solvents.

   • The positive ion is attracted to the negative end of the polar molecule, and the negative ion is attracted to its positive end.

3. Solution Concentration:

   • Expressed as molarity (M), percent by volume, percent by mass, or parts per million (ppm).

4. Boiling and Freezing Points:

   • Adding a solute to a solvent causes the boiling point of the solvent to increase and the freezing point to decrease.

5. Saturated Solutions:

   • Exist in equilibrium: the rate of crystallization equals the rate of dissolving.

Topic Nine: Kinetics and Equilibrium

1. Collision Theory:

   • A reaction is most likely to occur if reactant particles collide with the proper energy and orientation.

2. Reaction Rate Factors:

   • The rate of a chemical reaction depends on temperature, concentration, nature of the reactants, surface area, and the presence of a catalyst.

3. Equilibrium:

   • Some chemical and physical changes can reach equilibrium.

4. Equilibrium Rates:

   • At equilibrium, the rate of the forward reaction equals the rate of the reverse reaction.

5. Equilibrium Quantities:

   • The measurable quantities of reactants and products remain constant at equilibrium.

6. LeChatelier's Principle:

   • Predicts the effect of stress on a system in equilibrium.

   • Stresses include a change in pressure, volume, concentration, and temperature.

7. Potential Energy Diagrams:

   • Represent energy absorbed or released by a chemical reaction.

8. Heat of Reaction:

   • The amount of energy released or absorbed during a chemical reaction.

   • Heat of reaction equals the PE of the products – PE of reactants.

   • A positive heat of reaction implies an endothermic reaction.

   • A negative heat of reaction implies an exothermic reaction.

9. Catalysts:

   • Provide an alternative pathway for a chemical reaction.

   • The catalyzed reaction requires a lower activation energy than the uncatalyzed reaction.

   • Catalysts increase the rate of the forward and reverse reactions equally, so there is no shift in equilibrium.

10. Entropy:

    • A measure of the randomness or disorder in a system. A system with greater disorder has greater entropy.

11. Natural Tendencies:

    • Systems in nature tend to undergo changes towards lower energy and higher entropy.

12. Spontaneous Reactions:

    • Exothermic reactions that result in increased entropy are spontaneous.

Topic Ten: Acids, Bases, and Salts

1. Arrhenius Theory:

   • Explains the behavior of many acids and bases.

   • Arrhenius acids and bases are electrolytes.

2. Electrolytes:

   • Substances which, when dissolved in water, form a solution capable of conducting electricity.

   • The ability to conduct electricity depends on the concentration of ions.

3. Arrhenius Acids:

   • Yield $H^+(aq)$ ions as the only positive ion in solution.

   • $H^+(aq)$ ions may also be written as $H_3O^+(aq)$ ions (hydronium ions).

4. Arrhenius Bases:

   • Yield $OH^-(aq)$ ions as the only negative ion in solution.

   • Organic compounds with $OH^-$ are not bases.

   • Ammonia ($NH_3$) is a base.

5. Neutralization Reactions:

   • An Arrhenius acid and an Arrhenius base react to form salt and water.

   • The net ionic equation for all neutralization reactions is the same: $H^+(aq) + OH^-(aq) \rightarrow H_2O(l)$

6. Titration:

   • A lab process in which a volume of a solution of known concentration is used to determine the concentration of another solution.

   • A practical application of a neutralization reaction.

7. Other Acid-Base Theories:

   • One states that an acid is an $H^+$ donor and a base an $H^+$ acceptor.

8. pH Measurement:

   • The acidity or alkalinity of a solution can be measured by pH.

     • Low pH: higher concentration of $H^+$ ions than $OH^-$ ions.

     • High pH: lower concentration of $H^+$ ions than $OH^-$ ions.

     • Neutral pH (7): equal concentration of $H^+$ ions and $OH^-$ ions.

     • Pure water has a neutral pH.

9. pH Scale:

   • Each decrease of one pH unit represents a tenfold increase in $H^+$ ion concentration.

Topic Eleven: Oxidation-Reduction (Redox)

1. Electron Transfer:

   • A redox reaction involves the transfer of electrons ($e^-$).

2. Reduction:

   • The gain of electrons and decrease of oxidation number.

   • A half-reaction can be written to represent reduction.

3. Oxidation:

   • The loss of electrons and increase of oxidation number.

   • A half-reaction can be written to represent oxidation.

4. Electron Balance:

   • In redox, the number of electrons lost is equal to the number of electrons gained.

5. Oxidation Numbers:

   • Can be assigned to atoms and ions.

   • Changes in oxidation numbers indicate a redox reaction has occurred.

   • Double replacement reactions are not redox reactions.

   • A reaction where an element is alone on one side and part of a compound on the other is always a redox reaction.

6. Electrochemical Cells:

   • Can be either voltaic or electrolytic.

7. Electrodes:

   • In an electrochemical cell, oxidation occurs at the anode and reduction at the cathode.

8. Voltaic Cells:

   • Spontaneously convert chemical energy to electrical energy.

9. Electrolytic Cells:

   • Require energy to produce a chemical change; this is called electrolysis.

Topic Twelve: Organic Chemistry

1. Carbon Bonding:

   • Organic compounds consist of carbon atoms that bond to each other in chains, rings, and networks to form a variety of structures.

2. IUPAC Nomenclature:

   • Organic compounds can be named with the IUPAC system.

3. Hydrocarbons:

   • Compounds that contain only carbon and hydrogen.

     • Saturated hydrocarbons contain only single carbon-carbon bonds.

     • Unsaturated hydrocarbons contain at least one multiple carbon-carbon bond (double or triple bond).

4. Functional Groups:

   • Organic acids, alcohols, esters, aldehydes, ketones, ethers, halides, amines, amides, and amino acids are categories of organic molecules that differ in their structures.

5. Properties and Functional Groups:

   • Functional groups give organic molecules distinct physical and chemical properties.

6. Isomers:

   • Organic compounds that have the same molecular formula but different structures and properties.

7. Multiple Covalent Bonds:

   • In a multiple covalent bond, more than one pair of electrons are shared between two atoms.

   • Unsaturated organic compounds contain at least one double or triple bond.

8. Types of Organic Reactions:

   • Addition, substitution, polymerization, esterification, fermentation, saponification, and combustion.

Topic Thirteen: Lab Skills

1. General Lab Skills:

   • Any standard chemistry lab procedure is fair game for the Regents. Specific skills that may be tested include:

     • Using the scientific method for a controlled experiment.

     • Constructing a graph.

     • Using proper units of measurement.

     • Making accurate and precise measurements.

     • Using rules for significant figures.

     • Identifying and using lab equipment.

     • Lab safety.