1.01.01

Know how to collect, record, and analyze measurements appropriately.

1.01.02

Know and apply mathematical concepts and operations to evaluate and interpret experimental data including ratios, proportions, percentages, ranges, standard deviation, mean, median, mode, logarithms, and graphical analyses.

1.01.03

Demonstrate knowledge of the concepts of precision, accuracy, and error with regard to interpreting and recording numerical data acquired using a scientific instrument.

1.01.04

Know the appropriate use of tools and technology to make measurements including computer linked probes, pH meters, multimeters, spectrometers.

1.01.05

Plan and evaluate a specific scientific investigation on the basis of:

a. the formulation of testable questions and hypotheses

b. the accuracy and reproducibility of the data acquired

c. the distinction between variable and controlled parameters

d. distinctions between hypothesis, theory, and prediction as these terms are
used in science

e. the identification of discrepant results and the possible sources of error,
including uncontrolled conditions

g. the connections between the hypothesis proposed, tests conducted, data
collected and conclusions drawn

h. the conclusions and their fit into the underlying theory

i. the new questions leading to further investigations

1.01.06

Interpret numerical and non-numerical experimental results calculated and displayed in charts, maps, tables, models, graphs, and labeled diagrams.

1.01.07

Understand procedures for the appropriate and safe storage, handling, use, disposal, care, and maintenance of chemicals, materials, and equipment.

2.01.01

Describe the difference between mass and weight, compounds and mixtures, heterogeneous and homogeneous mixtures, elements and atoms, atoms and molecules, intensive and extensive, atoms and ions, and ionic and molecular compounds.

2.01.02

Know that atoms and molecules in the liquid state move randomly relative to one another, whereas in the solid state they vibrate about fixed positions and in the gaseous state they are mostly too far apart to interact significantly with one another.

2.01.03

Analyze a phase diagram for a pure substance.

2.01.04

Describe the motion of a body of mass m when the net force acting on it is zero (Newton's first law) and when the net force acting on it is a nonzero force (Newton's second law).

2.01.05

Describe the spectrum of electromagnetic waves. Understand the properties of each of the general wavelength ranges including the ways they can be generated, detected, and used for practical purposes. Explain that electromagnetic radiation of various wavelength ranges interacts with living tissue, and know how this can be hazardous.

2.01.06

Explain the basic functions and structure of DNA and RNA.

3.01.01

Know the contributions of the following individuals to the development of the atomic theory:

• John Dalton's atomic theory and its connection with the laws of combining volumes, definite proportions, conservation of matter, and observed chemical reactions

• J.J. Thomson's measurement of the charge/mass ratio of the electron

• The Rutherford-Geiger-Marsden experiment and the nuclear atom

• Robert A. Millikan's oil-drop experiment and the charge of the electron

• Niels Bohr's model of the atom, its electron arrangement, and the correlation with the hydrogen spectrum.

• Albert Einstein's explanation of the photoelectric effect

• Louis de Broglie's and Thomson's insights into the wave-particle duality

3.01.02

Know the key contributions of Aristotle, Democritus, Priestley, Lavoisier, Pasteur, Marie Curie, Schrödinger, and Pauling.

3.01.03

Use data to calculate average atomic masses from a given sample.

3.01.04

State the principles of the mass spectrometer and outline the main stages in its operation to determine relative isotopic and atomic masses.

3.01.05

Describe the structure and properties of the atom, including protons, neutrons, and electrons; atomic mass, mass number, and atomic number; energy levels; and isotopes.

3.02.01

Know the contributions of Mendeleyev, Moseley, Meyer, and Dobereiner to the development of the Periodic Table.

3.02.02

Identify major regions of the periodic table (metals, nonmetals, metalloids, lanthanides, actinides, alkali metals, alkaline earth metals, transition metals, gases, noble gases).

3.02.03

Explain the properties of metals, nonmetals, and metalloids/semi-metals.

3.02.04

Identify the connection between electron configuration, valence electrons, and the positioning of elements in the periodic table (s, p, d, f block elements).

3.02.05

Explain and identify periodic and group trends for ionization energy, electron affinity, electronegativity, atomic radius, and ionic radius and explain how these quantities relate to bond formation and reactivity.

3.02.06

Understand the difference between the trends down a group and across a period with specific reference to Group I metals, Group II metals, and halogens.

3.02.07

Identify and describe the general properties of the hydrides.

3.03.01

Describe the Bohr hydrogen atom and how its quantified structure accounts for the line spectrum of hydrogen (Balmer, Lyman, and Paschen series).

3.03.02

Understand the Heisenberg Uncertainty Principle with respect to the location of an electron and its momentum.

3.03.03

Use a set of quantum numbers to describe a particular electron in an orbital and to determine whether a given set of quantum numbers is possible.

3.03.04

Use the de Broglie relation and the Bohr model of the atom to show that electrons can only occupy a specific orbital with discrete energy levels.

3.03.05

Perform calculations based on the line spectrum of hydrogen.

3.03.06

Define the Aufbau principle, the Pauli Exclusion Principle, and Hund's rule, and apply them to the prediction of ground state electron configurations.

3.03.07

Write electron configurations for both atoms and simple ions (having noble gas and pseudo-noble gas configurations) and use orbital diagrams to predict the number of unpaired electrons.

3.03.08

Determine multiple oxidation states for transition elements using electron configuration.

3.03.09

Describe the general size, shape, energy, quantity, and spatial orientation of s, p, d, and f orbitals.

3.04.01

Explain how the Coulomb repulsion of the protons in an atomic nucleus implies that a strong attractive force must exist to hold the nucleus together and that decay implies that the strong force has a very short range.

3.04.02

Compare and contrast the properties of α, β, and γ emission in radioactive decay. Explain how radioactive isotopes are used.

3.04.03

Use mass and charge to balance nuclear reaction equations.

3.04.04

Use Einstein's equation E = mc² to explain the connection between nuclear mass defect and the energy observed in nuclear reactions.

3.04.05

Perform calculations involving half-lives of radioactive isotopes.

3.04.06

Explain the principles, uses, and limitations of various types of radiometric dating.

3.04.07

Distinguish between nuclear fission and fusion reactions and calculate the energy released by these reactions.

3.04.08

Explain the operation of fission reactors including the role of the moderator.

3.04.09

Describe the process used to convert nuclear energy into electrical energy in commercial reactors and compare to chemical energy processes (e.g. combustion of coal).

4.01.01

Draw electron-dot Lewis structures of molecules and polyatomic ions employing multiple bonding resonance structures, single bonds, double bonds, triple bonds, coordinate bonds, and lone pairs of electrons. Know when to expect the octet rule to be valid and when exceptions occur.

4.01.02

Use the periodic table and the octet rule to predict whether a given bond is ionic, polar covalent or non-polar covalent (including Ligands).

4.01.03

Describe the arrangement of atoms in molecules, ionic crystals, polymers, and metallic substances.

4.01.04

Compare and contrast the general physical and chemical properties of ionic and covalent compounds.

4.01.05

Relate the vapor pressure and boiling or melting point of a substance to its principle bonding mode: ionic, covalent, hydrogen or van der Waals (London dispersion forces) bonding.

4.01.06

Apply molecular orbital theory to chemical bonding.

4.01.07

Understand the use of hybridizations for molecular geometry of organic and inorganic compounds.

4.01.08

Use VSEPR theory to predict molecular shapes and the arrangement of atoms in a molecule.

4.02.01

Assign oxidation numbers to each atom in a chemical species to determine a formula.

4.02.02

Give formulas and names of common monatomic and polyatomic ions, ionic compounds, binary molecular compounds, hydrates and common acids and bases.

4.02.03

Know the IUPAC names of the alkanes C₁ to C₁₀.

4.02.04

Calculate:

• the mass percent composition of a compound

• the empirical formula of a compound

• the molecular formula of a compound

4.02.05

Identify compounds as basic, acidic, or amphoteric oxides, peroxides, or superoxides.

5.01.01

Identify types and predict products for various chemical reactions including composition (synthesis), decomposition, combustion, single and double displacement, and acid/base neutralization.

5.01.02

Understand that chemical reactions are processes in which atoms are rearranged into different combinations of molecules and, regardless of arrangement, the total mass and number of each kind of atom remains constant (i.e., conservation of mass).

5.01.03

Explain that chemical processes either release or absorb energy (are exothermic or endothermic).

5.01.04

Write balanced molecular, ionic, and net ionic equations for chemical reactions according to the law of conservation of matter (including but not limited to acid-base, redox, and water with alkali and alkaline earth metals).

5.01.05

Understand the mole concept with relationship to Avogadro's constants and laws as well as mole-mole ratio. Use this understanding to calculate molar mass from a chemical formula and the mass numbers of the constituent elements, moles to numbers of molecules, mass to number of moles, and the volume of gas at STP to moles.

5.01.06

Calculate quantitative relationships in chemical reactions involving solids, liquids or gases. Quantities may be specified from data given as mass, volume, density, concentration, and moles.

5.01.07

Calculate percent yield, limiting reactants, and excess reactants in chemical reactions.

5.02.01

Assess chemical reaction rates and the factors that affect reaction rates, including the roles of concentration, pressure, surface area, temperature, catalysts, light, and activation energy.

5.02.02

Express the relative rates of consumption of reactants and formation of products using the coefficients of a balanced chemical equation.

5.02.03

Use experimental data or plots to identify the order of a reaction and to determine the rate constant and average rate of a reaction.

5.02.04

Relate half-life to the rate constant for a reaction.

5.02.05

Given a reaction mechanism and an experimental rate law, identify the reaction intermediates and determine if the mechanism is consistent with the experimental rate law.

5.02.06

Understand the use of the Arrhenius equation and .

5.02.07

Evaluate a potential energy diagram showing activation energies for the forward and reverse reactions, and how they are affected by the addition of a catalyst.

5.03.01

Using half-reactions, identify the species oxidized and reduced as well as the oxidizing and the reducing agent.

5.03.02

Balance chemical equations for oxidation-reduction (redox) reactions in acidic, basic, and/or neutral solutions.

5.03.03

Understand the Daniell cell.

5.03.04

Identify where in an electrochemical cell oxidation and reduction occur and reference the direction of electron flow between the anode and cathode.

5.03.05

Use a table of standard reduction potentials to set up a voltaic cell and calculate its emf.

5.03.06

Using a standard hydrogen electrode, explain the measurements of standard reduction potentials to produce electrochemical series.

5.03.07

Use the Nernst equation to calculate the emf of a cell.

5.03.08

Understand the relationship between cell potential, electric work, and free energy in terms of the equations w = -qE and \'c4G = -nFE^\'e8.

5.03.09

Compare the factors involved in electrolysis of molten salts and diluted solutions.

5.03.10

Understand the application of electrolysis for purification of copper, production of aluminum, and electroplating.

5.04.01

Understand the concept of conservation of energy (qualitative and quantitative).

5.04.02

Understand the first and second laws of thermodynamics and the different forms of energy

5.04.03

Understand the differences between closed and open systems, surroundings, and the universe in state functions such as enthalpy, entropy, and/or free energy.

5.04.04

Know the relationship that exists between internal energy, heat, and work (PΔV) to and from the system and surroundings.

5.04.05

Calculate internal energy, heat, and work (PΔV) to and from the system and surroundings.

5.04.06

Perform calorimetry calculations using Q = mcΔt.

5.04.07

Explain the effects of the gain or loss of thermal energy on the temperature and state of solids, liquids and gases.

5.04.08

Understand that latent heat involves a change in the molecular potential energy of a substance as heat energy is released or absorbed with no temperature change.

5.04.09

Calculate the energy involved in a phase change given the latent heats. Given a heating curve, determine the enthalpy of an endothermic and/or exothermic process.

5.04.10

Use the Clausius-Clapeyron equation to calculate vapor pressure or heat of vaporization.

5.04.11

Use Hess's law to determine enthalpy, entropy and/or free energy.

5.04.12

Use the equation ΔG = ΔH - TΔS to:

• solve for the missing value under specific conditions and determine whether the reaction is spontaneous or non-spontaneous

• determine whether a forward reaction or its reverse reaction is favored or if the system is at equilibrium

• explain the relationship between energy transfer and entropy in a chemical reaction

• calculate the entropy change for a phase change, or the temperature at which the phase change (boiling, melting, or sublimation) occurs

5.04.13

Explain dynamic equilibrium. Calculate equilibrium constants given the concentrations of products and reactants at equilibrium.

5.04.14

Apply Le Châtelier's principle to predict changes in systems at equilibrium.

5.04.15

Use to find the equilibrium constant (K_C or K_P) for solutions or gases.

6.01.01

Analyze the basic assumptions of kinetic molecular theory and its applications, including how the ideal-gas model works, and understand deviations in the behavior of real gases.

6.01.02

Describe temperature in terms of the motion of molecules and describe the motion of molecules in terms of temperature.

6.01.03

Use Graham's law to calculate the molecular mass and/or relative rates of effusion of two different gases.

6.01.04

Apply Dalton's law of partial pressures to a mixture of gases.

6.01.05

Use Charles's, Boyle's, Gay-Lussac's law or the combined gas law for an ideal gas.

6.01.06

Use and apply the ideal gas law (pV = nRT).

6.02.01

Define compare, and contrast solutions, colloids, and suspensions.

6.02.02

Define saturated, unsaturated and supersaturated solutions and their relationship to precipitation.

6.02.03

Interpret the solution process in terms of solute-solute, solute-solvent, and solvent-solvent interactions.

6.02.04

With respect to solutions, understand the use of the following: ppm, v:v%, wt:wt%, M, m, and M₁V₁ = M₂V₂.

6.02.05

Understand the structure and special properties of water, including its cohesion, expansion on freezing, geometry of H-bonds, and solvent properties.

6.02.06

Describe the interparticle interactions (ion-ion, ion-dipole, dipole-dipole, hydrogen bonds, and London dispersion forces) in solutions.

6.02.07

Use Henry's law to predict the concentrations of gases in solution.

6.02.08

Use the concepts of colligative properties to calculate the theoretical freezing point and boiling point of the solution and/or molar mass of the solute.

6.02.09

Use Raoult's law to calculate the vapor pressure over a solution containing a nonvolatile solute and over a solution containing two volatile liquids.

6.02.10

Explain the circumstances under which diffusion and osmosis occur and explain osmotic pressure.

6.02.11

Determine whether a precipitate will form on mixing solutions of ionic compounds using relevant Ksp values, the solution volume, and concentration data.

6.03.01

Compare and contrast the Arrhenius, Brønsted-Lowry, and Lewis definitions of acids and bases.

6.03.02

Explain the origin of the pH scale and its relationship to H₃O⁺ concentration. Calculate H₃0⁺ concentration from OH- concentration and vice versa using K_w = [H₃O⁺][OH^-].

6.03.03

Write chemical equations showing how a given oxide or hydroxide exhibits amphoteric behavior.

6.03.04

Compare the relative strength of acids or bases in terms of their dissociation. Define strong and weak acids and know examples of each.

6.03.05

Identify the conjugate acid-base pairs in a chemical equation for a proton-transfer reaction.

6.03.06

Understand the use and relationship of pKa, pKb, Ka, Kb, Keq, Kp, Kc and Kw.

6.03.07

Understand buffer solutions in terms of composition, behavior, and capacity.

6.03.08

Given the concentrations of the conjugate acid and base at equilibrium and pKa, use the Henderson-Hasselbalch equation to calculate the pH of a buffer.

6.03.09

Using titration curves or experimental data, determine pH values for acid-base titrations (including diprotic acid-strong base titration). Differentiate between equivalence-point and end-point.

6.03.10

Know how an acid-base indicator works and be able to choose an appropriate indicator for a given titration.

7.01.01

Write condensed and line structures of organic molecules.

7.01.02

Apply the IUPAC naming system for alkanes, cycloalkanes, alkenes, alkynes, and aromatics and their structures.

7.01.03

Understand how organic molecules can exist as various isomers (structural, geometric, and stereoisomer).

7.01.04

Classify molecules based on their functional group and know their chemical structures and properties (with specific reference to alkanes, alkenes, alkynes, alcohols, aldehydes, ketones, carboxylic acids, amines, esters, and ethers).

7.01.05

Predict the products and write balanced chemical equations for the dehydrogenation and cracking (splitting) of alkanes.

7.01.06

Predict the products of and write balanced chemical equations for ethene with hydrogen, bromine, hydrogen-halide, and water.

7.01.07

Understand the structure of benzene using its chemical and physical properties.

7.01.08

Predict the products and write balanced chemical equations for electrophilic substitution reactions of benzene.

7.01.09

Understand the general structures of polymers such as polysaccharides, polypeptides, and nucleic acids and how they are built by condensation of monomers and degraded through the process of hydrolysis.

7.01.10

Outline the formation and mechanism of the production of dipeptides and esters.

7.01.11

Describe the use of chromatography and fractional distillation as analytical tools and give examples of chemical procedures in which they are especially useful.