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TOWNSHIP OF UNION PUBLIC SCHOOLS

AP Chemistry

Curriculum Guide2012

Curriculum Guide Approved June 2011

Board MembersFrancis “Ray” Perkins, President

Versie McNeil, Vice President

Gary Abraham

David Arminio

Linda Gaglione

Richard Galante

Thomas Layden

Vito Nufrio

Judy Salazar

TOWNSHIP OF UNION PUBLIC SCHOOLSAdministration

District Superintendent …………………………………………………………………...…………………….... Dr. Patrick Martin

Assistant Superintendent …………………………………………………………..……………………….….…Mr. Gregory Tatum

Director of Elementary Curriculum ……………………………….………………………………..…………….Ms. Tiffany Moutis

Director of Secondary Curriculum ……………………………….………………………….…………………… Dr. Noreen Lishak

Director of Student Information/Technology ………………………………..………………………….…………. Ms. Ann M. Hart

Director of Athletics, Health, Physical Education and Nurses………………………………..……………………Ms. Linda Ionta

DEPARTMENT SUPERVISORS

Language Arts/Social Studies K-8 ……..………………………………….…………………………………….. Mr. Robert Ghiretti

Mathematics K-5/Science K-5 …………………………………………….………………………………………. Ms. Deborah Ford

Guidance K-12/SAC …..………………………………………………………………………………….……….Ms. Bridget Jackson

Language Arts/Library Services 8-12 ….………………………………….…………………………………….…Ms. Mary Malyska

Math 8-12…………………………………………………………………………………………………………..Mr. Jason Mauriello

Science 6-12…….............…………………………………………………….………………………………….Ms. Maureen Guilfoyle

Social Studies/Business………………………………………………………………………………………..…….Ms. Libby Galante

World Language/ESL/Career Education/G&T/Technology….…………………………………………….….Ms. Yvonne Lorenzo

Art/Music …………………………………………………………………………………………………………..….Mr. Ronald Rago

Curriculum Committee

John Kronis

Academic Area

AP Chemistry

Table of Contents

Title Page

Board Members

Administration

Department Supervisors

Curriculum Committee

Table of Content

District Mission/Philosophy Statement

District Goals

Course Description

Recommended Texts

Course Proficiencies

Curriculum Units

Appendix: New Jersey Core Curriculum Content Standards

Mission Statement

The Township of Union Board of Education believes that every child is entitled to an education designed to meet his or her individual needs in an environment that is conducive to learning. State standards, federal and state mandates, and local goals and objectives, along with community input, must be reviewed and evaluated on a regular basis to ensure that an atmosphere of learning is both encouraged and implemented. Furthermore, any disruption to or interference with a healthy and safe educational environment must be addressed, corrected, or when necessary, removed in order for the district to maintain the appropriate educational setting.

Philosophy Statement

The Township of Union Public School District, as a societal agency, reflects democratic ideals and concepts through its educational practices. It is the belief of the Board of Education that a primary function of the Township of Union Public School System is to formulate a learning climate conducive to the needs of all students in general, providing therein for individual differences. The school operates as a partner with the home and community.

Statement of District Goals

Develop reading, writing, speaking, listening, and mathematical skills. Develop a pride in work and a feeling of self-worth, self-reliance, and self

discipline. Acquire and use the skills and habits involved in critical and constructive

thinking. Develop a code of behavior based on moral and ethical principals. Work with others cooperatively. Acquire a knowledge and appreciation of the historical record of human

achievement and failures and current societal issues. Acquire a knowledge and understanding of the physical and biological

sciences. Participate effectively and efficiently in economic life and the development

of skills to enter a specific field of work. Appreciate and understand literature, art, music, and other cultural

activities. Develop an understanding of the historical and cultural heritage. Develop a concern for the proper use and/or preservation of natural

resources. Develop basic skills in sports and other forms of recreation.

Course Description

The AP Chemistry course is a college level course which meets the requirements as outlined in the AP Chemistry course description. This course is designed to help students develop a conceptual framework for modern Chemistry and an appreciation of science as a process. Primary emphasis in this course is placed on understanding, analyzing, and applying rather than memorization. Essential to conceptual understanding is recognition of unifying themes that integrate the major topics as well as problem solving skills and scientific inquiry. Students are prepared to be critical and independent thinkers who are able to function effectively in a scientific and technological society. The textbook for the course is the 9th edition of General Chemistryby Ebbing and Gammon. In addition to the textbook, internet sources are used to illuminate concepts that are discussed in class. Classes meet each day for forty-two minutes. Twice a week an additional forty-two minute period is designed specifically for labs. Hands on lab components are designed to challenge the students’ ability to understand the nature of the problems, formulate hypotheses, design and implement experiments, interpret data, and draw conclusions. Each lab is followed by a complete and formal report in a separate lab notebook ready for college review if required for lab credits. A post lab discussion of each lab exercise is done to emphasize science as a process and its relationship to the theoretical material and a test is administered as well. In addition to the book material, every chapter is summarized with problems and essays that we solve and discuss in class. Homework is required to be handed at the end of each chapter so corrections can be made. Students are assessed with quizzes, essays from old AP exams, take home assessments, and tests. The scoring is used to determine individual weaknesses so the student will know what to review and improve in those particular topics. Students are prepared so that they will be able to achieve an acceptable grade on the AP Exam in May. After the AP Exam, the remainder of the year is devoted to topics (including labs developed by students request) that were not covered before the exam and a research project into chemical, biochemical, and pharmaceutical companies. Students are to present to the class their own analysis of the company they chose to do research on.

Recommended Textbook

General Chemistry by Ebbing and Gammon, 9th Edition

Course Proficiencies

Students will be able to… Demonstrate the proper handling of chemicals and lab equipment Perform scientific measurements and apply mathematical operations Use the structure of the atom and electron arrangement to identify and explain the trends in the periodic table Distinguish between the following types of bonds: ionic, covalent and metallic, and their relationship to atomic structure and

develop molecular orbitals. Write chemical formulas and name substances Identify the relationship between molecular shape and polarity for small molecules. Identify the types of reactions and be able to predict products from reactants Uphold the Law of Conservation of Matter by identifying and writing balanced chemical equations Understand the mathematics of chemistry by using and understanding the mole concept examine mathematical relationships in reactions using stoichiometric calculations Understand kinetics of reactions and integrate it to reaction quotient and its applications Distinguish between the four states of matter by applying the kinetic molecular theory Perform gas law calculations and their applications to everyday life Examine the factors that affect solubility and determine the concentration of solutions by different methods. Understand chemical equilibrium and its relationship to Kp, Ke, ∆G, and Ksp

Distinguish between acids and bases and relate it to pH. Use calorimetry to understand enthalpy change and its affect on bond energy. Apply oxidation reduction reactions to the development of voltaic cell and calculations of their potential.

Curriculum UnitsUnit 1: Basics of Chemistry Unit 2:Atomic and Molecular Structure

1. Chemistry and Measurement 7. Quantum Theory of the Atom2. Atoms, Molecules and ions 8. Electronic Configurations and Periodicity3. Calculations with Chemical Formulas and Equations 9. Ions and covalent Bonding

4. Chemical Reactions 10. Molecular geometry and Bonding theory5. The Gaseous State6. Thermochemistry

Unit 3: States Of Matter and Solutions Unit 4: Chemical Reactions and Equilibrium 11. Liquids and Solids 13. Rates of Reactions 12. Solutions 14. Chemical Equilibrium

15. Acids and Bases16. Acid-Base Equilibria17. Solubility and Complex-ion Equilibria18. Thermodynamics19. Electrochemistry

Unit 5: Chemistry of the Elements 20. Nuclear Chemistry 23. Organic Chemistry

Pacing Guide- Course

Content Number of Days

Unit 1:Basics of Chemistry 371. Chemistry and Measurement2. Atoms, Molecules and ions3. Calculations with Chemical Formulas and Equations4. Chemical Reactions5. The Gaseous State6. Thermochemistry

Unit 2:Atomic and Molecular Structure 187. Quantum Theory of the Atom8. Electronic Configurations and Periodicity9. Ions and covalent Bonding10.Molecular geometry and Bonding Theory

Unit 3:States Of Matter and Solutions 1011.Liquids and Solids12.Solutions

Unit 4:Chemical Reactions and Equilibrium 8513. Rates of Reactions14. Chemical Equilibrium15. Acids and Bases16. Acid-Base Equilibria17. Solubility and Complex-ion Equilibria18. Thermodynamics19. Electrochemistry

Unit 5:Chemistry of the Elements 8

Unit 1: Basics of Chemistry

Essential Questions Instructional Objectives/ Skills and Benchmarks(CPIs)

Activities Assessments

1. Chemistry and Measurement

a) What is the Law of Conservation of Mass and how do you apply it?

b) What is the difference between precision and accuracy as applied to measured quantities.

c) What are the Rules for determining significant figures and how are they used in reporting calculated values?

d) What are the SI units of measurement and the prefixes?

e) How is density used to relate mass and volume?

f) What is dimensional analysis and how is it used to sovle numerical problems?

5.1.12.A.1

5.1.12.A.2

5.1.12.A.3

5.1.12.B.1

5.1.12.B.2

5.1.12.B.3

5.1.12.B.4

5.1.12.C.1

5.1.12.C.2

5.1.12.C.3

5.1.12.D.1

5.1.12.D.2

Apply the Law of Conservation of Mass to a chemical problem.

Make measurements using a variety of equipment to distinguish between precision and accuracy.

Perform mathematical operations rounding to the proper number of significant digits.

Make measurements and report them using SI units.

Determine the density of different materials and substances.

Perform calculations using dimensional analysis.

Textbook problems. Concept Exploration. Text Online Resource

Questions CRSIRM’s.

Lab – Equipment identification and Measurement.

Lab - Density



2. Atoms, Molecules and Ions

1. What are the subatomic particles and how do you distinguish between them?

2. What is a nuclide and how is it used to signify elements and their isotopes?

3. How is atomic mass determined?

4. What are the features of the Periodic Table?

5. What are the distinguishing characteristics between ionic and molecular compounds?

6. What are the names, formulas and charges of common ions, including polyatomic ions?

7. How are ionic compounds named?

8. How are molecular formulas written?

9. How are molecular compounds named?

10. How are acids identified and named?

11.How are chemical equations written and balanced?

Distinguish between the subatomic particles.

Distinguish between Groups and Periods and identify other features of the Periodic Table.

5.2.12.A.1: Use atomic models to predict the behaviors of atoms in interactions.

Write nuclides describing elements and isotopes.

Determine atomic masses from appropriate data.

Write chemical formulas from names and name chemical compounds from formulas – both ionic and molecular and acids (including oxoacids).

Write and balance chemical equations.

Text Problems. Exercises. Concept Exploration. Text Online Resource


Lab - Chemical Reactions



3. Calculations with Chemical Formulas and Equations.

1. What is a mole and how is it related to molecular and formula masses of substances, masses of substances in grams, numbers of particles of substances (using Avogadro’s Number) and Volumes of a gas at STP?

2. What is percent composition and how is it applied?

3. What is an elemental analysis and how is it calculated?

4. How are empirical and molecular formulas determined?

5. What is the molar interpretation of coefficients in a balanced equation and how are coefficients used to relate quantities of reactants and products in chemical reactions?

6. How do you identify limiting and excess reactants in a chemical reaction?

7. How do calculate percent yield?

5.2.12.B.3: Balance chemical equations by applying the Law of conservation of Mass.

Perform all types of Stoichiometric calculations between mass, particles and volumes of gases.

Calculate percent composition of elements in a compound.

Calculate the mass of an element in a given mass of a compound.

Calculate the percentage of C, H and O from combustion data.

Determine empirical formulas from mass data or percent data.

Determine molecular formula percent composition and molecular mass.

Use stoichiometric calculations to relate amounts of reactants and products in chemical reactions.

Calculate limiting reactant and excess reactant in a chemical reaction.

Calculate percent yield by first calculating theoretical yield.

Text Problems Concept Exploration Exercises Text Online Resource


Lab – Determining Percent Carbonate in An Unknown Carbonate By Precipitation.



4. Chemical Reactions.

1. What is the Ionic Theory of Solutions and how is it applied using Solubility Rules for ionic compounds.

2. How do you distinguish between complete and net ionic equations?

3. How do you identify a precipitation reaction and determine whether a precipitate will form?

4. What are the different definitions for acids and bases and how are the equations written and the species identified?

5. What are the different types of acid-base reactions?

6. How are oxidation numbers determined?

7. How are half-reactions used to balance Redox reactions?

8. What are the different types of Redox reactions?

9. What is molarity?10. How are solutions diluted?11. What are gravimetric and

volumetric analyses?

Distinguish between non-electrolytes, and strong and weak electrolytes and their solutions.

Learn and apply the Solubility Rules.

Define Arrhenius, Bronsted-Lowry and Lewis acids and bases, and write corresponding equations for each, appropriately identifying the species.

Distinguish between strong acids and bases.

Recognize decomposition, combination, displacement and combustion reactions as Redox reactions.

5.2.12.B.2: Describe oxidation and reduction reactions, and give examples of oxidation and reduction reactions that have an impact on the environment , such as corrosion and the burning of fuel.

Write complete and net ionic equations.

Write molecular, complete and net ionic equations for precipitation reactions.

Assign oxidation numbers according to the Rules.

Balance Redox reactions using Half reactions.

Calculate Molarity. Determine dilution

concentrations. Perform gravimetric and

volumetric calculations.



Lab – Ionic Reactions and Aqueous Solutions. (Precipitation).

Lab – How Much Acid Is in Vinegar (by titration, volumetric analysis).



5. The Gaseous State

1. What is gas pressure and how is it measured?

2. What are the empirical Gas Laws and how are they used?

3. What is the Ideal Gas Law

4. And how is it used?5. How do you solve

stoichiometric problems involving gas volumes?

6. What is Dalton’s Law of partial pressures and how is it related to mole fractions?

7. What are the postulates of the Kinetic-Molecular Theory and how are they used to describe the Gas laws?

8. How do you calculate molecular speeds of Diffusion and effusion?

9. How does Van der Waals equation help distinguish between real and ideal gases?

Define pressure and its units.

List the 5 postulates of the kinetic molecular theory.

Describe how the root-mean-square (rms) molecular speed of gas molecules varies with temperature.

Define effusion and diffusion.

5.2.12.A.2: Account for the differences in physical properties of solids, liquids and gases.

5.2.12.C.1: Use the kinetic molecular theory to describe and explain the properties of solids, liquids and gases.

Express and use Boyle’s Law, Charles’ Law, Avogadro’s Law and combined Gas Law.

Use the Ideal Gas Law to determine gas density, and molecular mass of a vapor.

Solve stoichiometric problems involving gases.

Calculate partial pressures and mole fractions of gas in a mixture.

Calculate the amount of gas collected over water.

Calculate the ratio of effusion rates.

Use van der Waals equation.



Lab – Determine Molar Volume of A Gas.



6. Thermochemistry

1. How are the different forms of energy defined and what are the units?

2. What is Heat of Reaction?

3. What is enthalpy and enthalpy of reaction, and how are they used?

4. What is a thermochemical equation, how is it written and how are they manipulated?

5. How is stoichiometry applied to Heats of Reaction?

6. How do you measure the Heat of Reaction?

7. What is Hess’ Law and how is it applied?

8. What are standard enthalpies of formation and how are they used in calculations?

Define kinetic, potential and internal energy, and the common units of energy.

Define Heat of Reaction. Distinguish between

exothermic and endothermic processes.

Explain how the terms enthalpy of reaction and heat of reaction are related.

Define heat capacity and specific heat, and perform calculations of heat absorbed or evolved.

5.2.12.D.2: Describe the potential commercial applications of exothermic and endothermic reactions.

Write thermochemical equations and manipulate them applying the appropriate Rules.

Calculate the heat absorbed or evolved in a reaction given enthalpy and mass data.

Perform calorimetric calculations.

Apply Hess’ Law to otain enthalpy changes.

Use standard enthalpies of formation in calculations.



Lab – Calorimetry

Unit 2: Atomic and Molecular Structure



7. Quantum Theory of the Atom

1. What are the features of light and the electromagnetic spectrum?

2. What are the Quantum Effects and photons, and how are they applied to Quantum Theory?

3. How is the Bohr Theory of the Hydrogen Atom related to the energy of a photon in energy levels of an atom?

4. What is quantum mechanics and how is it applied to electron motion?

5. How are quantum numbers used to describe atomic orbitals of atoms?

Define wavelength and frequency of a wave, and relate them to the speed of light for electromagnetic radiation.

Define Planck’s constant and calculate the energy of a photon.

Describe the photoelectric effect.

State Heisenberg’s Uncertainty Principle.

Relate wave function for an electron to the probability of its location.

Define the quantum numbers for atomic orbitals.

State the rules for the allowed values of each quantum number, and apply them.


Determine the wavelength or frequency of a hydrogen atom transition.

Calculate the wavelength of a moving particle according to DeBroglie.



Lab – Atomic spectrum of hydrogen.



8. Electron Configurations andPeriodicity.

1. What are electron configurations and orbital diagrams?

2. What is Pauli Exclusion Principle?

3. What is the building-upprinciple and how is it related to the Periodic Table?

4. How do you determine paramagnetism and diamagnetism.

5. What are the Periodic Trends of properties in the Periodic Table?

6. What are basic, acidic and amphoteric oxides?

Determine and write electron configurations and orbital diagrams for atoms using the building-up principle, Hund’s Rule and Pauli Exclusion Principle.

Define and determine paramagnetism and diamagnetism of atoms.


5.2.12.A.3: Predict the placement of unknown elements on the Periodic Table based on their physical and chemical properties.

Identify the Trends of atomic radius, ionization energy and electron affinity across Periods and down Groups in the Periodic Table.

Determine relative successive 1st, 2nd, 3rd, etc ionization energies of atoms of elements.

Define basic, acidic and amphoteric oxides.





9. Ionic and Covalent Bonding.

1. What distinguishes ionic and covalent bonds and how are they illustrated using electron dot diagrams?

2. What are the characteristics of ionic bonds, ionic compounds, covalent bonds and covalent compounds?

3. What are the charges of the transitions metals?

4. How are Lewis Electron-dot formulas determined and written?

5. What are some exceptions to the octet Rule?

6. What is Formal Charge and how is it applied to writing Lewis Formulas?

7. How are bond length and bond order related?

8. What is bond energy and how are they used to estimate enthalpies?

Describe the energetics of ionic bonding.

Define lattice energy. Define isoelectronic ions. Define coordinate

covalent bond. Distinguish between

polar and nonpolar covalent bonds.

5.2.12.B.1: Model how the outermost electrons determine the reactivity of elements and the nature of the chemical bonds they tend to form.

Use the Born-Haber cycle to calculate lattice energies.

Identify bonding pairs and lone pairs of electrons.

Identify single, double and triple covalent bonds.

Write Lewis formulas. Write resonance

formulas. Write Lewis formulas that

are exceptions to the Octet Rule.

Use formal charges to identify the best formulas.

Determine bond orders. Use bond energies to

calculate delta H’s.





10.Molecular Geometry and Bonding theory.

1. How is the Valence-Shell Electron Pair Repulsion Model used to determine molecular geometry?

2. How do you determine molecular polarity?

3. How is Valence bond Theory applied to determine bonding?

4. How is Molecular Orbital Theory applied to determine bonding in small molecules.

5. What is delocalized bonding?

Define the Valence-Shell Electron Pair Repulsion Model.

State the 5 steps in describing bonding following Valence Bond Theory.

Define sigma and pi bonds, and explain cis- and trans geometric isomers in terms of pi bonding.

Define Molecular Orbital Theory.

Define bonding and antibonding orbitals.

Define bond order.

5.2.12.B.1: Model how the outermost electrons determine the reactivity of elements and the nature of the chemical bonds they tend to form

Predict and illustrate molecular geometries of small molecules applying the principles of VSEPR Model, showing any lone pairs of electrons.

Determine molecular polarity.

Apply Valence Bond Theory.

Using MO Theory, determine the molecular orbital electron configurations for homo-and heteronucleardiatomic molecules.

Text Problems. Exercises. Concept Exploration Text Online Resource


Lab – Molecular model-building.

Unit 3: States of Matter and Solutions



11. States of Matter; Liquids and Solids.

1. How does Kinetic Molecular Theory explain the different behaviors and properties of solids, liquids and gases.

2. What are the characteristics of the different phase transitions?

3. What is a phase diagram and how is it used?

4. What is the significance of viscosity and surface tension of liquids?

5. What are the intermolecular forces and how do they explain liquid properties?

6. What is the classification of solids by type of attraction of units?

7. What are crystal lattices and Unit Cells?

8. What are the structures of some crystalline solids

Compare a gas, liquid and solid using Kinetic molecular Theory.

Define the six phase changes and the characteristics of each as dynamic equilibria.

Define vapor pressure. Describe the features of

a phase diagram. Define surface tension

and viscosity. Describe capillary action. Identify intermolecular

forces and characteristics of each.

Identify types of solids and predict relative properties based on type of solid.

Distinguish between crystalline and amorphous solids.

Distinguish between 3 types cubic unit cells.

Calculate the heat required for a phase change of a given mass of substance.

Apply the Clausius-Clapeyron equation.

Calculate vapor pressure and heat of vaporization.

Draw phase diagrams from given information.

Predict relative properties based on intermolecular forces.

Determine the number of atoms in a unit cell.

Calculate atomic mass from unit-cell dimension and density.

Calculate unit cell dimension from unit-cell type and density.



Lab – Determine density of the vapor of a volatile liquid?

5.2.12.A.2: Account for the differences in physical properties of solids, liquids and gases.

5.2.12.C.1: Use the kinetic molecular theory to describe and explain the properties of solids, liquids and gases.



12.Solutions

1. What are the different types of Solutions?

2. What determines solubility?

3. How is concentration expressed?

4. What is vapor pressure?5. What are colligative

properties?6. What is osmosis?7. What are colloids?

Define terms related to solubility.

Give examples of solutions.

Distinguish between saturated, unsaturated and supersaturated.

Describe the factors that determine solubility.

Distinguish between molarity, molality, mass percent and mole fraction as ways of expressing concentration.

State Raoult’s Law. Determine colligative

properties of ionic solutions.

Define colloid. Explain Tyndall Effect.

5.2.12.A.5: Describe the process by which solutes dissolve insolvents.

Apply Henry’s Law. Perform conversions

between the different methods of expressing concentration.

Calculate tvapor pressure lowering.

Calculate boiling point elevation and freezing point depression.

Calculate molecular mass of solute from molality.

Calculate molecular mass from freezing point depression.

Calculate osmotic pressure.



Lab – Determination of molecular mass by freezing point depression.

Unit 4: Chemical Reactions and Equilibrium



13. Rates of Reaction

1. What is Reaction Rate?2. How are reaction rates

determined experimentally?

3. How does a rate depend on concentration?

4. What is an intergrated rate law?

5. What is half-life?6. What is Collision

Theory?7. What is Transition State

Theory?8. What is the Arrhenius

Equation and how is it used?

9. What is an elementary reaction?

10.What is a Rate Law?11.What is a catalyst?

Define terms related to reaction rates.

Describe how reaction rates can be determined experimentally.

Provide examples of rate law, rate constant and reaction order.

Learn half-life equations for zero-, first- and second order reactions,

State the postulates of Collision Theory.

Define an activated complex?

Define elementary reaction, reaction mechanism and reaction intermediate.Explain the rate-determining step

5.2.12.D.5: Model the change in rate of a reaction by changing a factor.

Determine order from the rate law.

Determine rate law from initial rates.

Use an integrated rate law.

Plot kinetic data to determine a rate law.

Use the Arrhenius Equation.

Write the overall chemical equation from a mechanism.

Determine molecularity of an elementary reaction.

Write the rate equation for an elementary reaction.

Determine rates laws from mechanisms with an initial slow step, fast step, or equilibrium step.



Lab – Kinetics



14. Chemical Equilibrium

1. What is a dynamic equilibrium and how does it apply to a chemical equilibrium?

2. What is an equilibrium constant, and how are they interpreted?

3. What is a heterogeneous equilibrium?

4. How is the direction of a reaction predicted?

5. How are equilibrium concentrations calculated?

6. What happens when an equilibrium is altered?

7. What is the effect of a catalyst on an equilibrium?

Define:Dynamic equibriumEquilibrium constantEquilibrium expressionHomogeneous eq’m.Heterogeneous eq’mReaction quotientCatalyst

Describe Kp. State Le Chatelier’s

Principle.

5.2.12.D.5: Model the change in rate of a reaction by changing a factor.

Apply stoichiometry to an equilibrium mixture.

Write equilibrium cinstantexoressions.

Obtain an equilibrium constant from a reaction composition.

Write Kc for a reaction with pure solids and liquids.

Use the reaction quotient.

Obtain one equilibrium concentration given others.

Solve equilibrium problems involving a linear equation or quadratic equation.

Apply Le Chatelier’s Principle.



Lab – Keq of FeSCN.



15. Acids and Bases

1. What are acids and bases according to Arrhenius, Bronsted-Lowry and Lewis?

2. What is a conjugate acid-base pair?

3. How are the relative strengths of acids and bases determined?

4. What is the self-ionization of water?

5. What is pH and how is it used?

Define acids and bases according to Arrhenius, Bronsted-Lowry and Lewis.

Understand the rules for determining the relative strengths of acids (including polyprotic acids) and bases.

Define self-ionization of water and the ion-product constant for water.

Define pH. Describe the use of a pH

meter and indicators to determine pH.

5.2.12.A.6: Relate the pH scale to the concentrations of various acids and bases.

Identify acid and base species.

Identify acid-base conjugate pairs.

Decide whether reactants or products are favored in an acid-base s in solutions of a strong acid or base.

Calculate concentrations of hydronium ion and hydroxide ion.

Calculate pH from hydronium ion concentration and visa versa.



Lab – Standardization of a strong basic solution.



16.Acid-Base Equilibria

1. How are acid and base equilibria described and expressed?

2. How are equilibria of salt solutions expressed?

3. What is the Common Ion Effect?

4. What is a buffer?5. What is the Henderson-

Hasselbach Equation?6. What are the features of

the various types of acid-base titration curves?

5.2.12.A.6: Relate the pH scale to the concentrations of various acids and bases.

Define acid and base ionization constants and degree of ionization.

State the general trend in ionization constant of a polyprotic acid.

Write chemical equations for a weak acid or weak base undergoing ionization in aqueous solution.

Write the hydrolysis reactions of ions forming acid or base solutions.

Explain the common-ion effect.

Define buffer and buffer capacity.

State when the Henderson-Hasselbach equation can be applied.

Define equivalence point. Describe the features of

a titration curve and how they vary depending on their combination of weak or strong acids and bases.

Determine Ka from solution pH.

Calculate conc’n of species in a weak acid or base sol’n usingKa’s or Kb’s.

Calculate conc’n of species in a sol’n of a polyprotic acid.

Predict whether a salt solution is acid, basic or neutral.

Obtain Ka from Kb or Kb from Ka.

Calculate conc’n of species in a salt sol’n.

Calculate the common-ion effect on acid ionization (the effect of a strong acid, or a conjugate base).

Calculate the pH of a buffer from given volumes of solution.

Calculate the pH of a buffer when a strong acid or base is added.

Calculate the pH at the equivalence point.


Questions CRSIRM’s.Lab – Determine the molar mass (and therefore the identity) of an unknown weak acid.



17.Solubility and Complex-Ion Equilibria

1. What is a solubility product constant?

2. How is the solubility of a salt affected by the presence of a common-ion?

3. How can precipitation be predicted?

4. How does pH affect solubility?

5. What are complex ions and how do they form?

6. What is the strategy in qualitative analysis?

Define:KspMolar solubilityIon productFractional precipitationComplex ionLigandFormation, stability and dissociation constantsAmphoteric hydroxideQualitative analysis.

Explain the common-ion effect as applied to solubility.

State the criterion for precipitation.

Explain the qualitative effect of pH on solubility of a slightly soluble salt.

5.2.12.A.5: Describe the process by which solutes dissolve in solvents.

Write solubility product expressions.

Calculate Ksp from solubility and visa versa.

Calculate solubilities in solutions containing common ions.

Predict whether precipitation will occur, given ion concentrations.

Determine the qualitative effect of pH on solubility.

Calculate the concentrations of metal ion-complex ion equilibria.

Predict whether a precipitate will form in the presence of a complex ion.



Lab – Determine A Solubility Product Constant.



18. Thermodynamics and Equilibrium

1. What is the First Law of Thermodynamics: Enthalpy?

2. What is the Second Law of Thermodynamics; Entropy?

3. What is the Third Law of Thermodynamics?

4. How are enthalpy and entropy related to Free Energy and Spontaneity of a reaction?

5. What are the different interpretations of Free Energy?

6. How is ΔG related to an Equilibrium constant?

7. How is free energy affected by Temperature?

Define work. Define the First Law Of

Themodynamics. Define enthalpy and

relate it to heat, q. Define entropy and a

spontaneous process as related to chemical processes.

State the Second Law of Thermodynamics and distinguish between changes in the surroundings and the system.

Describe how ΔH-TΔS, and ΔG, function as criteria of a spontaneous process.

State the Third Law of Thermodynamics.

Define Free Energy and standard free energy change.

Define K (thermodynamic)

Describe how the free energy change in a reaction and the reaction quotient are related.

Calculate the entropy change for a phase transition.

Predict the sign of an entropy change for a reaction.

Calculate ΔS for a reaction.

Calculate ΔG from ΔH and ΔS.

Calculate ΔG from standard free energies of formation.

Interpret the sign of ΔG. Write the expression for

a thermodynamic equilibrium constant.

Calculate K from standard free energy change.

Calculate ΔG and K at various temperatures.

5.2.12.D.2: Describe the potential commercial applications of exothermic and endothermic reactions.



Lab – Spontaneity – Sodium nitrate Solution.



19. Electrochemistry

1. How are Redox reactions balanced using half-reactions depending on whether they are in acidic or basic solution?

2. What are the features and operability of a Voltaic cell?

3. What is cell potential and how is it calculated?

4. How are Standard Cell potentials and standard electrode potentials interpreted?

5. How are equilibrium constants related to cell potentials?

6. How are cell potentials related to concentration?

7. How are voltaic cells used commercially?

8. What is electrolysis?

Learn the steps for balancing Redox reactions.

Describe the features and functioning of a voltaic cell , including the salt bridge.

Distinguish between reducing agent and oxidizing agent.

Identify the reactions that occur at each electrode (anode and cathode).

Define cell potential and volt.

Interpret Tables of standard reduction potentials.

Define electrolysis.

Activites (continued) Predict the half-reaction

in aqueous electrolysis Calculate the amount of

charge from the amount of product in an electrolysis, and visa versa.

Balance reaction by the half-reaction method (acidic and basic solution)

Sketch and label a voltaic cell.

Write a cell reaction from cell notation and visa versa.

Determine the direction of spontaneity from electrode potentials.

Calculate cell potential from electrode potentials.

Calculate free energy change from electrode potentials.

Calculate cell potential from free energy change.

Calculate the equilibrium constant from cell potential.

Calculate cell potential for nonstandard conditions using Nernst Equation.



Lab – Electrochemical Cell

5.2.12.B.2: Describe oxidation and reduction reactions, and give examples of oxidation and reduction reactions that have an impact on the environment , such as corrosion and the burning of fuel.

Unit 5: Chemistry Of The Elements.



20. Nuclear Chemistry

1. What is radioactivity?2. What are the products of

nuclear bombardment reactions?

3. What are some sources of radioactivity?

4. How is the rate of radioactive decay measured and its significance?

5. What are some applications of Radioisotopes?

6. What is the difference between nuclear fission and fusion?

Define:Radioactive decayNuclear bombardment reaction.Radioactive decay series.Transmutation.CurieRad and remHalf-life

Write a nuclear equation and deduce the product or reactant.

List 6 types of radioactive decay.

Draw a half-life decay curve.

Describe how radioisotopes are used for medical therapy.

5.2.12.D.3: Describe the products and potential applications of fission and fusion reactions.

Predict relative stability of nuclides.

Predict the typ of radioactive decay.

Use the notation for a bombardment reaction.

Calculate half-lives.





23. Organic Chemistry1. What are the different

ways in which carbon typically bonds?

2. What are the structural features that distinguish alkanes, alkenes and alkynes?

3. What are cycloalkanes?4. What are geometric

isomers?5. What are aromatics

compounds?6. How are simple organic

compounds named?7. What are the major

functional groups?

Distinguish between saturated, unsaturated and aromatic hydrocarbons.

Distinguish between alkanes, alkenes and alkynes, with respect to the presence of a multiple bond.

Distinguish between cis-and trans-isomers.

Identify aromatic compounds.

Learn the IUPAC system for numbering and naming simple compounds.

Identify functional Groups:Alcohols, ethers, aldehydes, ketones, carboxylic acids, esters, amines, amides.

Write structural formulas for all compounds being studied.

Name compounds.


Questions CRSIRM’sLab – Molecular modeling of isomers of hexane. (optional)

New Jersey Core Curriculum Content StandardsAcademic AreaAP Chemistry

5.1 Science Practices: All students will understand that science is both a body of knowledge and an evidence-based, model-building enterprise that continually extends, refines, and revises knowledge. The four Science Practices strands encompass the knowledge and reasoning skills that students must acquire to be proficient in science.

A. Understand Scientific Explanations: Students understand core concepts and principles of science and use measurement and observation tools to assist in categorizing, representing, and interpreting the natural and designed world.B. Generate Scientific Evidence Through Active Investigations: Students master the conceptual, mathematical, physical, and computational tools that need to be applied when constructing and evaluating claims.C. Reflect on Scientific Knowledge: Scientific knowledge builds on itself over time.D. Participate Productively in Science: The growth of scientific knowledge involves critique and communication, which are social practices that are governed by a core set of values and norms.

5.2 Physical Science: All students will understand that physical science principles, including fundamental ideas about matter, energy, and motion, are powerful conceptual tools for making sense of phenomena in physical, living, and Earth systems science.

A. Properties of Matter: All objects and substances in the natural world are composed of matter. Matter has two fundamental properties: matter takes up space, and matter has inertia.B. Changes in Matter: Substances can undergo physical or chemical changes to form new substances. Each change involves energy.C. Forms of Energy: Knowing the characteristics of familiar forms of energy, including potential and kinetic energy, is useful in coming to the understanding that, for the most part, the natural world can be explained and is predictable.D. Energy Transfer and Conservation: The conservation of energy can be demonstrated by keeping track of familiar forms of energy as they are transferred from one object to another.E. Forces and Motion: It takes energy to change the motion of objects. The energy change is understood in terms of forces.

township of union public schools ... school... · web viewunderstand kinetics of reactions and...

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