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Study Guide: Regents Examination in Physical Setting / Chemistry: Chemistry Core Topic Outline
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Regents Examination in Physical Setting / Chemistry: Chemistry Core Topic Outline

By Fatskills Exam Guides Team — the exam nerds behind 28,500+ quizzes and 2.1M practice questions across 500+ global exams.

⏱️ ~26 min read

The Topic Outline is adapted from Appendix B of the New York State Physical Setting/Chemistry Core. All Regents Chemistry—Physical Setting Examinations are based on this core.

The topic outline is divided into 12 sections:
 

M. Math Skills
R. Reading Skills
I. Atomic Concepts
II. Periodic Table
III. Moles/Stoichiometry
IV. Chemical Bonding
V. Physical Behavior of Matter
VI. Kinetics/Equilibrium
VII. Organic Chemistry
VIII. Oxidation–Reduction
IX. Acids, Bases, and Salts
X. Nuclear Chemistry


Each section contains one or more of the following items:
The Major Understandings that you must have mastered for the examination
The Skills that you need to be able to demonstrate during the examination
The Key Points to Remember that remind you of key chemistry concepts that you should be familiar with to succeed on this examination.

M. Mathematics Skills Needed for Chemistry
M.1    Organize, graph, and analyze data gathered from laboratory activities or other sources.
Identify independent and dependent variables.
Create appropriate axes with label and scale.
Identify graph points clearly.
M.2    Interpret a graph constructed from experimentally determined data.
Identify direct and inverse relationships.
Apply data showing trends to predict information.
M.3    Measure and record experimental data and use the data in calculations.
Choose appropriate measurement scales and use units in recording.
Show mathematical work stating formula and steps for solution.
Estimate answers.
Use appropriate equations and significant digits.
Identify relationships within variables from data tables.
Calculate percent error.
M.4    Recognize and convert various scales of measurement.
Convert between Celsius (°C) and Kelvin (K).
Convert among kilometers (km), meters (m), centimeters (cm), and millimeters (mm).
Convert between grams (g) and kilograms (kg).
Convert between kilopascals (kPa) and atmospheres (atm).
M.5    Employ critical thinking skills in solving problems.
Apply algebraic or geometric concepts in the solution of mathematical problems.
Use knowledge of geometric arrangements to predict particle properties or behavior.
State the assumptions on which a particular mathematical equation is based.
Evaluate the appropriateness of an answer to a solved problem.

R. Reading Skills Needed for Chemistry
R.1    Reading Comprehension: Literal
Extract an answer from the text when the answer is “right there” in the information provided.
Read and interpret information provided in the Reference Tables for Chemistry.
R.2    Reading Comprehension: Inferential/Interpretive
Understand and make connections between prior content knowledge and information from the text when these connections are not explicitly stated.
Understand and apply new information provided in the text and solve problems using this information.
R.3    Reading Comprehension: Lexical
Understand and apply key vocabulary in a reading passage.

 

I. Atomic Concepts
  MAJOR UNDERSTANDINGS SKILLS

The student should be able to:		 KEY POINTS TO REMEMBER
I.1
The modern model of the atom has evolved over a long period of time through the work of many scientists.
relate experimental evidence to models of the atom
Examples:
J.J. Thomson and cathode ray experiment
Ernest Rutherford and gold foil experiment
I.2
Each atom has a nucleus, with an overall positive charge, surrounded by negatively charged electrons.
use models to describe the structure of an atom		  
I.3
Subatomic particles contained in the nucleus include protons and neutrons.		    
I.4
The proton is positively charged, and the neutron has no charge. The electron is negatively charged.		    
I.5
Protons and electrons have equal but opposite charges. The number of protons is equal to the number of electrons in an atom.
determine the number of protons or electrons in an atom or ion when given one of these values
Positive ions form when atoms lose electrons.

Negative ions form when atoms gain electrons.
I.6
The mass of each proton and each neutron is approximately equal to one atomic mass unit. An electron is much less massive than a proton or neutron.
calculate the mass of an atom, the number of neutrons, or the number of protons, given the other two values		  
I.7
In the wave-mechanical model (electron cloud), the electrons are in orbitals, which are defined as regions of most probable electron location (ground state).		 interpret electron configurations listed on the Periodic Table  
I.8
Each electron in an atom has its own distinct amount of energy.		   Electron configurations begin with the lowest energy shell near the nucleus and continue in order of increasing energy and distance from the nucleus.
I.9
When an electron in an atom gains a specific amount of energy, the electron is at a higher energy state (excited state).
distinguish between ground state and excited state electron configurations, e.g., 2–8–2 vs. 2–7–3
Excited state configurations contain the same number of electrons as ground state configurations.
I.10
When an electron returns from a higher energy state to a lower energy state, a specific amount of energy is emitted. This emitted energy can be used to identify an element.
identify an element by comparing its bright-line spectrum to given spectra		  
I.11
The outermost electrons in an atom are called the valence electrons. In general, the number of valence electrons affects the chemical properties of an element.
draw a Lewis electron-dot structure of an atom

distinguish between valence and non-valence electrons, given an electron configuration, e.g., 2–8–2		  
I.12
Atoms of an element that contain the same number of protons but a different number of neutrons are called isotopes of that element.		    
I.13
The average atomic mass of an element is the weighted average of the masses of its naturally occurring isotopes.
given an atomic mass, determine the most abundant isotope

calculate the atomic mass of an element, given the masses and ratios of naturally occurring isotopes		  

 

 

 

 

 

II. Periodic Table
  MAJOR UNDERSTANDINGS SKILLS

The student should be able to:		 KEY POINTS TO REMEMBER
II.1
The placement or location of an element on the Periodic Table gives an indication of physical and chemical properties of that element. The elements on the Periodic Table are arranged in order of increasing atomic number.
explain the placement of an unknown element in the Periodic Table based on its properties		  
II.2
The number of protons in an atom (atomic number) identifies the element. The sum of the protons and neutrons in an atom (mass number) identifies an isotope. Common notations that represent isotopes include: 14C, 146C, carbon-14, C-14.
interpret and write isotopic notation
Subtract the atomic number from the mass number to find the number of neutrons in an isotope!
The term “nuclear charge” refers to the number of protons in an atom’s nucleus. For example, a carbon atom with 6 protons would have a nuclear charge of 6+.
II.3
Elements can be classified by their properties and located on the Periodic Table as metals, nonmetals, metalloids (B, Si, Ge, As, Sb, Te), and noble gases.
classify elements as metals, nonmetals, metalloids, or noble gases by their properties		  
II.4
Elements can be differentiated by their physical properties. Physical properties of substances, such as density, conductivity, malleability, solubility, and hardness, differ among elements.
describe the states of the elements at STP suggest a simple separation strategy to separate two or more elements based on physical properties		  
II.5
Elements can be differentiated by chemical properties. Chemical properties describe how an element behaves during a chemical reaction.		  
Metal atoms lose electrons when they bond.
II.6
Some elements exist as two or more forms in the same phase. These forms differ in their molecular or crystal structure and hence in their properties.		  
Allotropes of carbon include graphite, diamond, and fullerene.
II.7
For Groups 1, 2, and 13–18 on the Periodic Table, elements within the same group have the same number of valence electrons (helium is an exception) and, therefore, similar chemical properties.
determine the group of an element, given the chemical formula of a compound, e.g., XCl or XCl2		  
II.8
The succession of elements within the same group demonstrates characteristic trends: differences in atomic radius, ionic radius, electronegativity, first ionization energy, and metallic/nonmetallic properties.
compare and contrast properties of elements within a group or a period for Groups 1, 2, 13–18 on the Periodic Table
Increasing the distance between the nucleus and an atom’s valence electrons causes the ionization energy and the electronegativity to decrease down each group of the Periodic Table.
II.9
The succession of elements across the same period demonstrates characteristic trends: differences in atomic radius, ionic radius, electronegativity, first ionization energy, and metallic/nonmetallic properties.		  
Increasing nuclear charge causes electronegativity and ionization energy to increase across each period of the Periodic Table.

 

 

 

 

 

III. Moles/Stoichiometry
  MAJOR UNDERSTANDINGS SKILLS

The student should be able to:		 KEY POINTS TO REMEMBER
III.1
A compound is a substance composed of two or more different elements that are chemically combined in a fixed proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by a specific chemical formula and assigned a name based on the IUPAC system.
determine how many total atoms are in a compound
determine how many atoms of each type are in a compound
use IUPAC nomenclature rules to name binary ionic, binary molecular, and polyatomic ionic compounds		 Compounds, along with elements, are pure substances.
III.2
Types of chemical formulas include empirical, molecular, and structural.		 identify examples of empirical, structural, or molecular formulas



convert a structural formula into a molecular formula or an empirical formula		  
III.3
The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that compound.
determine the molecular formula, given the empirical formula and molecular mass
determine the empirical formula from a molecular formula		  
III.4
In all chemical reactions there is a conservation of mass, energy, and charge.
interpret balanced chemical equations in terms of conservation of matter and energy
Atoms In = Atoms Out!
III.5
A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction.
balance equations, given the formulas for reactants and products

interpret balanced chemical equations in terms of conservation of matter and energy

create and use models of particles to demonstrate balanced equations

calculate simple mole–mole stoichiometry problems, given a balanced equation
Coefficients in a balanced chemical equation represent moles or particles, not grams!
III.6
The formula mass of a substance is the sum of the atomic masses of its atoms. The molar mass (gram-

formula mass) of a substance equals one mole of that substance.
calculate the formula mass and the gram-formula mass

use the molar mass to convert from grams to moles or from moles to grams
Formula masses are reported in atomic mass units (amu).
Molar masses are reported in grams/mol.
III.7
The percent composition by mass of each element in a compound can be calculated mathematically.
determine the number of moles of a substance, given its mass

determine the mass of a given number of moles of a substance




use lab data to determine the percent of water in a hydrate		  
III.8
Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement.
identify types of chemical reactions
Synthesis reactions create just one product:




Decomposition reactions have just a single reactant:




Single replacement reactions have uncombined elements on both sides of the reaction:




Double replacement reactions show “partners switching places”:

 

 

 

 

 

 

IV. Chemical Bonding
  MAJOR UNDERSTANDINGS SKILLS

The student should be able to:		 KEY POINTS TO REMEMBER
IV.1
Compounds can be differentiated by their chemical and physical properties.
distinguish among ionic, molecular, and metallic substances, given their properties		  
IV.2
Two major categories of compounds are ionic and molecular (covalent) compounds.		    
IV.3
Chemical bonds are formed when valence electrons are: transferred from one atom to another (ionic); shared between atoms (covalent); mobile within a metal (metallic).
demonstrate bonding concepts using Lewis dot structures representing valence electrons: transferred (ionic bonding); shared (covalent bonding); in a stable octet
Metallic bonds hold samples of pure metals together.

Ionic bonds form between metal ions and nonmetal ions.

Covalent bonds form between nonmetal or metalloid atoms.
IV.4
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.		 evaluate the number of electrons shared, or the number of electron pairs shared, in single, double, or triple covalent bonds  
IV.5
Molecular polarity can be determined by the shape and distribution of the charge. Symmetrical (nonpolar) molecules include CO2, CH4, and diatomic elements. Asymmetrical (polar) molecules include HCl, NH3, H2O.		  
Molecular symmetry cancels polarity!
Draw Lewis electron-dot diagrams to determine molecular shape.
IV.6
When an atom gains one or more electrons, it becomes a negative ion and its radius increases. When an atom loses one or more electrons, it becomes a positive ion and its radius decreases.		    
IV.7
When a bond is broken, energy is absorbed. When a bond is formed, energy is released.		    
IV.8
Atoms attain a stable valence electron configuration by bonding with other atoms. Noble gases have stable valence electron configurations and tend not to bond.
determine the noble gas configuration an atom will achieve when bonding
Remember the octet rule! Atoms (except hydrogen) tend to bond so as to achieve the stable octet of eight valence electrons.
IV.9
Physical properties of substances can be explained in terms of chemical bonds and intermolecular forces. These properties include conductivity, malleability, solubility, hardness, melting point, and boiling point.		 compare and contrast the physical properties of different compounds based on the types of chemical bonds and/or intermolecular forces they have  
IV.10
Electron-dot diagrams (Lewis structures) can represent the valence electron arrangement in elements, compounds, and ions.
demonstrate bonding concepts, using Lewis dot structures representing valence electrons: transferred (ionic bonding); shared (covalent bonding); in a stable octet
Check! Does your drawing show all of the valence electrons? Does each atom (except H) have 8 valence electrons?
IV.11
Electronegativity indicates how strongly an atom of an element attracts electrons in a chemical bond. Electronegativity values are assigned according to arbitrary scales.		    
IV.12
The electronegativity difference between two bonded atoms is used to assess the degree of polarity in the bond.
distinguish between nonpolar covalent bonds (two of the same nonmetals) and polar covalent bonds
Bond polarity and ionic character increase with an increasing difference in electronegativity.
Ionic character refers to an atom’s tendency to transfer electrons when forming a chemical bond. Ionic character, like molecular polarity, is determined by finding the electronegativity difference between the atoms in the bond.
IV.13
Metals tend to react with nonmetals to form ionic compounds. Nonmetals tend to react with other nonmetals to form molecular (covalent) compounds. Ionic compounds containing polyatomic ions have both ionic and covalent bonding.		    
IV.14
Ionic compounds will be soluble in water when the attraction between the ions and water molecules is greater than the attractions between ions in the solid compound.
predict solubility/insolubility of ionic compounds using Reference Table F

identify precipitates in double replacement reactions		  

 

 

 

 

V. Physical Behavior of Matter
  MAJOR UNDERSTANDINGS SKILLS

The student should be able to:		 KEY POINTS TO REMEMBER
V.1
Matter is classified as a pure substance or as a mixture of substances.		   A pure substance can either be an element or a compound.
V.2
The three phases of matter (solids, liquids, and gases) have different properties.
use a simple particle model to differentiate properties of a solid, a liquid, and a gas		  
V.3
A pure substance (element or compound) has a constant composition and has constant properties throughout a given sample and from sample to sample.
use particle models/diagrams to differentiate elements, compounds, and mixtures		  
V.4
Elements are substances that are composed of atoms that have the same atomic number. Elements cannot be broken down by chemical change.		    
V.5
Mixtures are composed of two or more different substances that can be separated by physical means. When different substances are mixed together, a homogeneous or heterogeneous mixture is formed.		    
V.6
The proportions of components in a mixture can be varied. Each component in a mixture retains its original properties.		    
V.7
Differences in properties such as density, particle size, molecular polarity, boiling point and freezing point, and solubility permit physical separation of the components of the mixture.
describe the process and use of filtration, distillation, and chromatography in the separation of a mixture		  
V.8
A solution is a homogeneous mixture of a solute dissolved in a solvent. The solubility of a solute in a given amount of solvent is dependent on the temperature, the pressure, and the chemical natures of the solute and solvent.
interpret and construct solubility curves

use solubility curves to distinguish saturated, supersaturated, and unsaturated solutions

apply the adage “like dissolves like” to real-world situations
Solid solutes increase in solubility and gas solutes decrease in solubility as the temperature increases.
V.9
The concentration of a solution may be expressed as: molarity (M), percent by volume, percent by mass, or parts per million (ppm).
describe the preparation of a solution, given the molarity

interpret solution concentration data

calculate solution concentrations in molarity (M), percent mass, and parts per million (ppm)		  
V.10
The addition of a nonvolatile solute to a solvent causes the boiling point of the solvent to increase and the freezing point of the solvent to decrease. The greater the concentration of solute particles, the greater the effect.		   Electrolytes have a greater effect on colligative properties than nonelectrolytes do because they dissociate into ions in aqueous solution.
V.11
Energy can exist in different forms, such as chemical, electrical, electromagnetic, heat, mechanical, and nuclear.		    
V.12
Heat is a transfer of energy (usually thermal energy) from a body of higher temperature to a body of lower temperature. Thermal energy is associated with the random motion of atoms and molecules.
distinguish between heat energy and temperature in terms of molecular motion and amount of matter

qualitatively interpret heating and cooling curves in terms of changes in kinetic and potential energy, heat of vaporization, heat of fusion, and phase changes		  
V.13
Temperature is a measure of the average kinetic energy of the particles in a sample of matter. Temperature is not a form of energy.
distinguish between heat energy and temperature in terms of molecular motion and amount of matter

explain phase changes in terms of the changes in energy and intermolecular distance

convert between Kelvin and Celsius temperature scales using the formula on Reference Table T		  
V.14
The concept of an ideal gas is a model to explain behavior of gases. A real gas is most like an ideal gas when the real gas is at low pressure and high temperature.		    
V.15
Kinetic molecular theory (KMT) for an ideal gas states all gas particles:

are in random, constant, straight-line motion
are separated by great distances relative to their size; the volume of gas particles is considered negligible
have no attractive forces between them
have collisions that may result in a transfer of energy between particles, but the total energy of the system remains constant.		  
Pressure is caused by collisions of gas particles against the inner walls of a container.
V.16
Kinetic molecular theory describes the relationships of pressure, volume, temperature, velocity, and frequency and force of collisions among gas molecules.
explain the gas laws in terms of KMT

solve problems, using the combined gas law		  
V.17
Equal volumes of gases at the same temperature and pressure contain an equal number of particles.
convert temperatures in Celsius degrees (°C) to kelvins (K), and kelvins to Celsius degrees
The molar volume of a gas at STP is equal to 22.4 L/mol.
V.18
The concepts of kinetic and potential energy can be used to explain physical processes that include: fusion (melting); solidification (freezing); vaporization (boiling, evaporation), condensation, sublimation, and deposition.
qualitatively interpret heating and cooling curves in terms of changes in kinetic and potential energy, heat of vaporization, heat of fusion, and phase changes

calculate the heat involved in a phase or temperature change for a given sample of matter

explain phase change in terms of the changes in energy and intermolecular distances
At a substance’s melting point, the solid and liquid phases are in equilibrium. At a substance’s boiling point, the liquid and gas phases are in equilibrium.
V.19
A physical change results in the rearrangement of existing particles in a substance. A chemical change results in the formation of different substances with changed properties.		    
V.20
Chemical and physical changes can be exothermic or endothermic.
distinguish between endothermic and exothermic reactions, using energy terms in a reaction equation, ΔH, potential energy diagrams or experimental data
Exothermic reactions have a –ΔH. Endothermic reactions have a +ΔH.
V.21
The structure and arrangement of particles and their interactions determine the physical state of a substance at a given temperature and pressure.
use a simple particle model to differentiate properties of solids, liquids, and gases		  
V.22
Intermolecular forces created by the unequal distribution of change result in varying degrees of attraction between molecules. Hydrogen bonding is an example of a strong intermolecular force.
explain vapor pressure, evaporation rate, and phase changes in terms of intermolecular forces

use vapor pressure curves (Reference Table H) to determine the boiling point of a substance		  
V.23
Physical properties of substances can be explained in terms of chemical bonds and intermolecular forces. These properties include conductivity, malleability, solubility, hardness, melting point, and boiling point.
compare the physical properties of substances based upon chemical bonds and intermolecular forces		  

 

 

 

 

 

VI. Kinetics/Equilibrium
  MAJOR UNDERSTANDINGS SKILLS

The student should be able to:		 KEY POINTS TO REMEMBER
VI.1
Collision theory states that a reaction is most likely to occur if reactant particles collide with the proper energy and orientation.
use collision theory to explain how various factors, such as temperature, surface area, and concentration, influence the rate of reaction
The reaction rate depends upon the number of effective collisions between reacting particles.
VI.2
The rate of a chemical reaction depends on several factors: temperature, concentration, nature of reactants, surface area, and the presence of a catalyst.		  
Reactions between ionic substances are faster than reactions between molecular substances.
VI.3
Some chemical and physical changes can reach equilibrium.
identify examples of physical equilibria as solution equilibrium and phase equilibrium, including the concept that a saturated solution is at equilibrium		  
VI.4
At equilibrium, the rate of the forward reaction equals the rate of the reverse reaction. The measurable quantities of reactants and products remain constant at equilibrium.
describe the concentration of particles and rates of opposing reactions in an equilibrium system
Equilibrium means equal and opposite rates, not equal amounts!
VI.5
Le Châtelier’s principle can be used to predict the effect of stress (change in pressure, volume, concentration, and temperature) on a system at equilibrium.
qualitatively describe the effect of stress on equilibrium, using Le Châtelier’s principle
A “shift to the right” means that product is being produced and reactants are being used up.
VI.6
Energy released or absorbed by a chemical reaction can be represented by a potential energy diagram.
read and interpret potential energy diagrams: PE of reactants and products, activation energy (with or without a catalyst), heat of reaction		  
VI.7
Energy released or absorbed by a chemical reaction (heat of reaction) is equal to the difference between the potential energy of the products and the potential energy of the reactants.
use Reference Table I to find the heat of reaction for specified moles produced from a particular reaction		  
VI.8
A catalyst provides an alternate reaction pathway that has a lower activation energy than an uncatalyzed reaction.		 draw a line (to show the effect of adding a catalyst) on a potential energy diagram  
VI.9
Entropy is a measure of the randomness or disorder of a system. A system with greater disorder has greater entropy.
compare the entropy of phases of matter		  
VI.10
Systems in nature tend to undergo changes toward lower energy and higher entropy.		    

 

 

 

 

 

VII. Organic Chemistry
  MAJOR UNDERSTANDINGS SKILLS

The student should be able to:		 KEY POINTS TO REMEMBER
VII.1
Organic compounds contain carbon atoms that bond to one another in chains, rings, and networks to form a variety of structures. Organic compounds can be named using the IUPAC system.
classify an organic compound based on its structural or condensed structural formula		  
VII.2
Hydrocarbons are 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.
draw structural formulas for alkanes, alkenes, and alkynes containing a maximum of ten carbon atoms

use Reference Table Q to classify hydrocarbons as alkanes, alkenes, or alkynes		  
VII.3
Organic acids, alcohols, esters, aldehydes, ketones, ethers, halides, amines, amides, and amino acids are types of organic compounds that differ in their structures. Functional groups impart distinctive physical and chemical properties to organic compounds.
classify an organic compound based on its structural or condensed structural formula

draw a structural formula with the functional group(s) on a straight chain hydrocarbon backbone, when given the correct IUPAC name for the compound
Use Reference Table R for help in naming and drawing organic compounds.
VII.4
Isomers of organic compounds have the same molecular formula, but different structures and properties.		    
VII.5
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.		    
VII.6
Types of organic reactions include: addition, substitution, polymerization, esterification, fermentation, saponification, and combustion.
identify types of organic reactions

determine a missing reactant or product in a balanced equation		  

 

 

 

 

 

VIII. Oxidation–Reduction
  MAJOR UNDERSTANDINGS SKILLS

The student should be able to:		 KEY POINTS TO REMEMBER
VIII.1
An oxidation–reduction (redox) reaction involves transfer of electrons (e–).
determine a missing reactant or product in a balanced equation		  
VIII.2
Reduction is the gain of electrons.		  
If an element’s oxidation number is getting more negative (or less positive), it is reducing (gaining electrons).
VIII.3
A half-reaction can be written to represent reduction.
write and balance half-reactions for oxidation and reduction of free elements and their monatomic ions
Reduction half-reactions show how many electrons are gained.
VIII.4
Oxidation is the loss of electrons.		  
If an element’s oxidation number is getting less negative (or more positive), it is oxidizing (losing electrons).
VIII.5
A half-reaction can be written to represent oxidation.		 write and balance half-reactions for the oxidation and reduction of free elements and their monatomic ions
Oxidation half-reactions show how many electrons are lost.
VIII.6
In a redox reaction, the number of electrons lost is equal to the number of electrons gained.		  
Electrons lost  =  electrons gained means that charge is conserved.
VIII.7
Oxidation numbers (states) can be assigned to atoms and ions. Changes in oxidation numbers indicate that oxidation and reduction have occurred.		 identify which reactions are redox reactions  
VIII.8
An electrochemical cell can be either voltaic or electrolytic. In an electrochemical cell, oxidation occurs at the anode and reduction at the cathode.
compare and contrast voltaic and electrolytic cells		  
VIII.9
A voltaic cell spontaneously converts chemical energy to electrical energy.
identify and label the parts of a voltaic cell (cathode, anode, salt bridge) and direction of electron flow, given the reaction equation

use an activity series to determine whether a redox reaction is spontaneous
Remember: Ions flow through the salt bridge. Electrons flow through the wire.
VIII.10
An electrolytic cell requires electrical energy to produce chemical change. This process is known as electrolysis.
identify and label the parts of an electrolytic cell (anode, cathode) and direction of electron flow, given the reaction equation		  

 

 

 

 

 

 

IX. Acids, Bases, and Salts
  MAJOR UNDERSTANDINGS SKILLS

The student should be able to:		 KEY POINTS TO REMEMBER
IX.1
Behavior of many acids and bases can be explained by the Arrhenius theory. Arrhenius acids and bases are electrolytes.
given properties, identify substances as Arrhenius acids or Arrhenius bases		  
IX.2
An electrolyte is a substance that, when dissolved in water, forms a solution capable of conducting an electric current. The ability of a solution to conduct an electric current depends on the concentration of ions.		  
The three categories of compounds that are electrolytes are acids, bases, and salts (ionic compounds).
IX.3
Arrhenius acids yield H+ (hydrogen ion) as the only positive ion in aqueous solution. The hydrogen ion may also be written as H3O+, hydronium ion.		    
IX.4
Arrhenius bases yield OH– (hydroxide ion) as the only negative ion in an aqueous solution.		    
IX.5
In the process of neutralization, an Arrhenius acid and an Arrhenius base react to form salt and water.
write simple neutralization reactions when given the reactants

at the equivalence point.
IX.6
Titration is a laboratory process in which a volume of solution of known concentration is used to determine the concentration of another solution.
calculate the concentration or volume of a solution, using titration data		  
IX.7
There are alternate acid–base theories. One such theory states that an acid is an H+ donor and a base is an H+ acceptor.		 identify acids and bases in a reversible reaction using the alternate theory  
IX.8
The acidity and alkalinity of an aqueous solution can be measured by its pH value. The relative level of acidity or alkalinity of a solution can be shown by using indicators.
interpret changes in acid–base indicator color

identify solutions as acid, base, or neutral based upon the pH
The pH of pure water is 7. Acids have pH values less than 7, and bases have pH values greater than 7.
IX.9
On the pH scale, each decrease of one unit of pH represents a tenfold increase in hydronium ion concentration.		    

 

 

 

 

 

 

X. Nuclear Chemistry
  MAJOR UNDERSTANDINGS SKILLS

The student should be able to:		 KEY POINTS TO REMEMBER
X.1
Stability of isotopes is based on the ratio of the neutrons and protons in its nucleus. Although most nuclei are stable, some are unstable and spontaneously decay emitting radiation.		    
X.2
Each radioactive isotope has a specific mode and rate of decay (half-life).
calculate the initial amount, the fraction remaining, or the half-life of a radioactive isotope, given two of the three variables

use Reference Table N to find the decay mode for a particular radioisotope		  
X.3
A change in the nucleus of an atom that converts it from one element to another is called transmutation. This can occur naturally or can be induced by the bombardment of the nucleus of high-energy particles.		  
Natural transmutation is also called radioactive decay.
X.4
Spontaneous decay can involve the release of alpha particles, beta particles, positrons, and/or gamma radiation from the nucleus of an unstable isotope. These emissions differ in mass, charge, ionizing power, and penetrating power.
determine decay mode and write nuclear equations showing alpha and beta decay
Radioactive decay continues until a stable isotope is reached.
X.5
Nuclear reactions include natural and artificial transmutation, fission, and fusion.
compare and contrast fission and fusion reactions		  
X.6
There are benefits and risks associated with fission and fusion reactions.		  
Fusion generates much more energy than fission. Fission produces far more dangerous radioactive waste than fusion.
X.7
Nuclear reactions can be represented by equations that include symbols that represent atomic nuclei (with the mass number and atomic number), subatomic particles (with mass number and charge), and/or emissions such as gamma radiation.
complete nuclear equations; predict missing particles from nuclear equations		  
X.8
Energy released in a nuclear reaction (fission or fusion) comes from the fractional amount of mass converted into energy. Nuclear changes convert matter into energy.		    
X.9
Energy released during nuclear reactions is much greater than the energy released during chemical reactions.		    
X.10
There are inherent risks associated with radioactivity and the use of radioactive isotopes. Risks can include biological exposure, long-term storage and disposal, and nuclear accidents.		    
X.11
Radioactive isotopes have many beneficial uses. Radioactive isotopes are used in medicine and industrial chemistry, e.g., radioactive dating, tracing chemical and biological processes, industrial measurement, nuclear power, and detection and treatment of diseases.
identify specific uses of some common radioisotopes, such as: I-131 in diagnosing and treating thyroid disorders; C-14 to C-12 ratio in dating living organisms; U-238 to Pb-206 ratio in dating geological formations; Co-60 in treating cancer		  

 

 

 

 


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Global History & Geography Transition Regents Review Regents Physics Review: Applications of Waves / Diffraction Regents Physics Review: DC Circuits / Electrical Power Regents Physics Review: Work, Power, and Energy / Conservation of Energy Regents Physics Review: Work, Power, and Energy / Work/Power Regents Physics Review: Impulse and Momentum / Impulse 10000+ Science Questions Regents Physics Review: Electrostatics / Coulomb's Law English Regents Review Regents Physics Review: Math Skills / Graph Interpretation

 
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