High School Chemistry Science Curriculum Assessment Instructional Guide

High School Chemistry

Science Curriculum

Assessment Instructional Guide

Salinas Union High School District

Updated June 19, 2009

Salinas Union High School District High School Chemistry Instructional Guide

Table of Contents

Course Description: High School Chemistry page 3 SUHSD Student Performance Objectives: Chemistry page 4 Investigation and Experimentation page 15 Terminology Guide page 17 Instructional Guide: High School Chemistry page 20 Atomic Structure and the Periodic Table page 20 Chemical Bonding and Balancing Equations page 27 Moles and Stoichiometry page 35 Gas Laws page 42 Introduction to Solutions page 51 Chemical Bonds and Structure page 57 Thermodynamics page 66 Solutions page 72 Equilibrium page 77 Reaction Rates page 83 Acids and Bases page 88 Safety page 93

Joyce Wilkinson, Science Resource 2


DEPARTMENT: SCIENCE COURSE TITLE: Chemistry GRADE LEVEL: 9-12 LENGTH: 1 year NUMBER OF CREDITS: 10 credits PREREQUISITE: NONE TEXTBOOK: Chemistry Prentice-Hall 2002 SUPPLEMENTARY MATERIALS: Materials provided by the publisher including study guide and lab manual, CD accompaniment, Videodisc Programs, Web Links

COURSE DESCRIPTION: A lab-based chemistry course which encompasses concepts set forth in the California State Science Content Standards for California Public Schools and the Science Framework for California Public Schools. Laboratory experiences are a major component of the course in addition to traditional instruction. This course meets the California and Salinas Union High School District graduation requirements for physical science. It qualifies as college lab course, meeting the University of California “A-G” requirements. SOURCES OF CURRICULUM CONTENT: The curriculum incorporates all High School Life Science Blueprint content standards adopted by the California State Board of Education 2004, as required. Clarification of topics and unifying themes have been drawn from Benchmarks for Scientific Literacy (American Association for the Advancement of Science) and from National Science Education Standards (National Research Council). Much of the clarification of state standards and development of concept units is a result of meetings with the planning committee composed of high school chemistry teachers in the Salinas Union High School District who met during 2007-08.



Chemistry Course Outline Grades: 9-12

Major Concepts

1. Introduction to Chemistry: (Safety and Emergency Response, Laboratory Techniques and Measurement,

Significant Figures and Dimensional Analysis) 2. Introduction to Atoms, the Periodic Table, and Nomenclature 3. Conservation of Matter and Stoichiometry 4. Gas Laws 5. Chemical Bonds 6. Thermodynamics 7. Solutions 8. Chemical Equilibrium 9. Rates of Reaction 10. Acids and Bases 11. Organic Chemistry and Biochemistry (after CST) 12. Nuclear Processes (after CST)

MAJOR CONCEPTS/EXPECTED STUDENT OUTCOMES SUGGESTED TEACHER ACTIVITIES AND STRATEGIES

Unit 1: An Introduction to Chemistry

Safety and Emergency Response

State Standards Investigation and Experimentation 1. Scientific progress is made by asking meaningful questions and conducting careful investigations. Students

will: a. Select and use appropriate tools and technology (such as computer-linked probes, spreadsheets, and

graphing calculators) to perform tests, collect data, analyze relationships, and display data.

Student Outcomes 1. Follow all safety rules and precautions as highlighted in the laboratory manual and by the teacher.

Students will be expected to pass a safety quiz before the first lab with safety concerns. A grade of 80% o or better must be obtained before students may participate in laboratory activities.

2. Follow good housekeeping procedures in their laboratory area. This includes good facilities management, equipment and materials storage, chemical handling techniques, and glassware use and storage.

3. Know the location of and how to operate all safety equipment available in the classroom. Protect eyes, face, hands, and body while conducting class and laboratory activities.

4. Dispose of dangerous waste chemicals and materials as prescribed by appropriate standards and laws. 5. Know the appropriate emergency steps to take in the event of an accident or emergency including

evacuation procedures.



Possible Strategies and Activities 1. Review the appropriate safety procedures before activities and briefly explain specific precautions and

hazards associated with the activity. 2. Demonstrate how to store and handle laboratory equipment. Ensure that chemicals are kept in a safe

place with limited access to prevent the possibility of hazardous accidents. 3. Check smoke detectors, safety showers, and eyewash facilities on a regular basis. Know how to access

all master shutoffs where applicable. 4. Clarify all waste disposal procedures before every activity. Provide separate labeled waste receptacles

for waste as needed. Also provide labeled receptacles for broken glass, waste paper, and used matches if applicable.

5. Post your school's accident policy and procedures; make accident reports promptly, accurately, and completely.

Laboratory Techniques and Measurement

State Standards Investigation and Experimentation 1. Scientific progress is made by asking meaningful questions and conducting careful investigations.

Students will: a. Select and use appropriate tools and technology (such as computer-linked probes, spreadsheets, and

graphing calculators) to perform tests, collect data, analyze relationships, and display data.

Student Outcomes 1. In advance of each laboratory experiment, write a prelab or flow chart which expresses a hypothesis

and describes the experimental procedure in the student's own words. 2. Learn the names and functions of the laboratory equipment used in experiments. 3. Understand and use common laboratory terminology. 4. Make all quantitative measurements to the proper level of precision, and state the absolute error for

each measurement. 5. Demonstrate competence in performing laboratory procedures. These tasks may include, but are not

limited to: transfer of chemicals from stock containers, use of balance (both digital and mechanical, if possible), reading a thermometer, filtration, decantation, operating a Bunsen burner, reading volumes in graduated cylinders and gas measuring tubes, use of buret, cleaning and rinsing glassware, handling hot equipment and/or solutions, operation of specialized equipment such as colorimeters, pH meters, or spectrophotometers.

6. Perform appropriate analysis of experimental data, draw logical conclusions, and report findings in a clearly written laboratory report.

Possible Strategies and Activities 1. Provide guidelines for proper prelab procedures. 2. Identify the equipment used in each laboratory experiment, clarify the function of each item, and insist

that students use the proper terms when referring to the equipment during class discussion and in written laboratory reports.

3. Model the use of proper terminology when discussing laboratory procedures, and require students to employ the correct terms when discussing their experiments.

4. Demonstrate how to make quantitative measurements to the highest possible level of precision with each instrument that is used in the laboratory. Follow-up may include putting emphasis on precision



statements in written laboratory reports as well as consistent modeling when, discussing treatment of quantitative laboratory data during class discussions.

5. Demonstrate the correct procedure for each specialized lab task that is introduced in student experiments, and correct students when necessary while observing their work in, the laboratory. Supporting strategies include requiring students to clearly explain each lab technique in their written flow sheets, and giving a "lab practical" exam in which students' lab skills are formally evaluated.

6. Provide students with a format for writing laboratory reports, evaluate students' reports in a timely manner, and give students suggestions for improving their work as the course progresses.

Significant Figures and Dimensional Analysis (Factor Label Method)

State Standards Investigation and Experimentation 1. Scientific progress is made by asking meaningful questions and conducting careful investigations.

Students will: b. Select and use appropriate tools and technology (such as computer-linked probes, spreadsheets, and

graphing calculators) to perform tests, collect data, analyze relationships, and display data. c. Identify and communicate sources of unavoidable experimental error. d. Identify possible reasons for inconsistent results, such as sources of error or uncontrolled conditions. e. Formulate explanations by using logic and evidence. f. Solve scientific problems by using quadratic equations and simple trigonometric, exponential, and

logarithmic functions. k. Recognize the issues of statistical variability and the need for controlled tests. m. Analyze situations and solve problems that require combining and applying concepts from more than

one area of science.

Student Outcomes 1. Demonstrate a consistent use of units in measurements and calculations. 2. Use the factor label method in the solution of all problems and laboratory calculations assigned

throughout the year. 3. Use estimation techniques incorporating dimensional analysis to arrive at approximate values when

assigned to do so by the teacher. 4. Demonstrate skills in the use of fractions, scientific notation, exponents, graphing, graph

interpretation, and the correct use of significant digits to indicate the appropriate degree of uncertainty.

Possible Strategies and Activities 1. Provide direct instruction and consistently model the use of appropriate units associated with

measurements. 2. Teach factor labeling as a means of solving of problems both from laboratory experiences and from

hypothetical situations. 3. Model and provide opportunities for estimation using orders of magnitude; arriving at approximate

values; and indicating appropriate degrees of uncertainty. 4. Review basic mathematical skills for use in chemistry. 5. Monitor and evaluate student learning. Make adjustments as shown necessary in the monitoring

process, with prompt correction and review. 6. Select appropriate materials and media to enhance learning, including some of the following:

measurement lab with significant figures calculations; density lab; graphing lab



Unit 2: Introduction to Atoms, the Periodic Table, and Chemical Nomenclature

State Standards Atomic and Molecular Structure 1. The periodic table displays the elements in increasing atomic number and shows how periodicity of the

physical and chemical properties of the elements relates to atomic structure. As a basis for understanding this concept: a. Students know how to relate the position of an element in the periodic table to its atomic number and

atomic mass. b. Students know how to use the periodic table to identify metals, semimetals, non-metals, and halogens. c. Students know how to use the periodic table to identify alkali metals, alkaline earth metals and

transition metals, trends in ionization energy, electronegativity, and the relative sizes of ions and atoms. d. Students know how to use the periodic table to determine the number of electrons available for

bonding. e. Students know the nucleus of the atom is much smaller than the atom yet contains most of its mass. f. *Students know how to use the periodic table to identify the lanthanide, actinide, and transactinide

elements and know that the transuranium elements were synthesized and identified in laboratory experiments through the use of nuclear accelerators.

h. *Students know the experimental basis for Thomson’s discovery of the electron, Rutherford’s nuclear atom, and Millikan’s oil drop experiment.

Chemical Bonds 2. Biological, chemical, and physical properties of matter result from the ability of atoms to form bonds from

electrostatic forces between electrons and protons and between atoms and molecules. As a basis for understanding this concept: a. Students know atoms combine to form molecules by sharing electrons to form covalent or metallic

bonds or by exchanging electrons to form ionic bonds.

Student Outcomes 1. Students will be able to describe the basic structure of the atom and the experimental evidence that

lead to our understanding of atomic structure. 2. Students will know the basic organization of the periodic table. 3. Students will know that compounds are formed when ions are held together by electrostatic forces or

when atoms share electrons to form molecules. 4. Students will be able to name and write formulas for ionic and molecular compounds and distinguish

between the two by noting whether a metal and a non-metal or two non-metals are combining. In order to do this; students must be familiar with both the locations of metals and non-metals on the periodic table, and with the use and the naming of polyatomic ions.

Possible Strategies and Activities 1. Provide a historic overview of the history of the discovery of the atom, subatomic particles, and the

periodic table. This overview will be accompanied by lab, audiovisual, lecture, and research (including Internet assignments).

2. Explain all the parts of the periodic table including the distinction between groups and families, metals and non-metals, specific families, valence numbers, and how to determine ionic charges.

3. Have students memorize the formula and name for the common polyatomic ions. 4. Have students write IUPAC names for compounds based on their formulas and formulas based on

names.



Unit 3: Conservation of Matter and Stoichiometry

State Standards Conservation of Matter and Stoichiometry 3. The conservation of atoms in chemical reactions leads to the principle of conservation of matter and the

ability to calculate the mass of products and reactants. As a basis for understanding this concept: a. Students know how to describe chemical reactions by writing balanced equations. b. Students know the quantity one mole is set by defining one mole of carbon 12 atoms to have a mass of

exactly 12 grams. c. Students know one mole equals 6.02 x 1023 particles (atoms or molecules). d. Students know how to determine the molar mass of a molecule from its chemical formula and a table of

atomic masses and how to convert the mass of a molecular substance to moles, number of particles, or volume of gas at standard temperature and pressure.

e. Students know how to calculate the masses of reactants and products in a chemical reaction from the mass of one of the reactants or products and the relevant atomic masses.

f. *Students know how to calculate percent yield in a chemical reaction.

Student Outcomes 1. Balance chemical equations using the proper chemical symbols and coefficients. 3a 2. Utilize Avogadro's hypothesis to solve conversion problems involving atomic weight, molar mass, the

mole, and molar volume. 3b, 3c, 3d, 3e 3. Students will be able to calculate the percent yield in a chemical reaction.

Possible Strategies and Activities 1. Provide direct instruction that clarifies, defines, and provides a consistent pattern for solving problems

in the laboratory and from hypothetical situations. 2. Give students numerous opportunities to perform molar calculations, using stoichiometric ratios,

Avogadro's number, molar mass, and molarity. 3. Provide extensive practice with prompt correction and review of assigned problems. 4. Select appropriate materials in the form of films and software to supplement lecture. 5. Provide laboratory opportunities including calculations of molar mass, mass relationships

accompanying chemical changes, and requiring the use of stoichiometric calculations including the calculation of percent yield.

Unit 4: Gas Laws

State Standards Gas Laws 4. The kinetic molecular theory describes the motion of atoms and molecules and explains the properties of

gases. As a basis for understanding this concept: a. Students know the random motion of molecules and their collisions with a surface create the observable

pressure on that surface. b. Students know the random motion of molecules explains the diffusion of gases. c. Students know how to apply the gas laws to relations between the pressure, temperature, and volume of

any amount of an ideal gas or any mixture of ideal gases. d. Students know the values and meanings of standard temperature and pressure (STP). e. Students know how to convert between the Celsius and Kelvin temperature scales. f. Students know there is no temperature lower than 0 Kelvin.



Student Outcomes 1. Be able to explain the characteristics of gases. 2. Use the following gas laws: Dalton’s Law of Partial Pressures, Charles' Law, Boyle's Law, Avogadro’s

Principle, the Ideal Gas Law, the Combined Gas Law, and Graham's Law. 4c, 4d 3. Be able to graph and exchange the changes of state using heating and cooling curves and label: melting

point, boiling point, heat of fusion, heat of vaporization 4e. 4f

Possible Strategies and Activities 1. Use models and lecture to illustrate the differences between a solid, liquid and gas on the macro and

molecular level. 2. Define the gas laws and utilize them for problem solving, then provide guided practice and laboratory

activities for the students. 3. Select experiments that will facility the students use of laboratory data, and when possible computer

graphing software to illustrate heating and cooling curves. UNIT 5: CHEMICAL BONDS State Standards 2. Biological, chemical, and physical properties of matter result from the ability of atoms to form bonds from

electrostatic forces between electrons and protons and between atoms and molecules. As a basis for understanding this concept: a. Students know atoms combine to form molecules by sharing electrons to form covalent or metallic bonds

or by exchanging electrons to form ionic bonds. b. Students know chemical bonds between atoms in molecules such as H2, CH4, NH3, H2 CCH2, N2, Cl2,

and many large biological molecules are covalent. c. Students know salt crystals, such as NaCl, are repeating patterns of positive and negative ions held

together by electrostatic attraction. d. Students know the atoms and molecules in liquids move in a random pattern relative to one another

because the intermolecular forces are too weak to hold the atoms or molecules in a solid form. e. Students know how to draw Lewis dot structures.

Students Outcomes 1. Review the nature and rational of chemical bonding, ionic model, ionic character, oxidation states of

families, and covalent bonding. 2. Draw and explain Lewis dot structure, molecular structure, octet rule, valence electrons/pairs. 3. Describe molecular structure, octet rule and exception, resonance. 4. Explain VSEPR model, hybridization, and molecular orbital model. Possible Strategies and Activities 1. Students build or observe models of crystal lattices. 3. Build molecular shapes and structures of certain molecules. 4. Use of molecular model sets. 5. Use of Styrofoam balls and colored pipe cleaners to demonstrate valence electrons and molecular shapes.



UNIT 6: Thermodynamics State Standards 7. Energy is exchanged or transformed in all chemical reactions and physical changes of matter. As a basis for

understanding this concept: a. Students know how to describe temperature and heat flow in terms of the motion of molecules (or

atoms). b. Students know how chemical processes can either release (exothermic) or absorb (endothermic) thermal

energy. c. Students know energy is released when a material condenses or freezes and is absorbed when a material

evaporates or melts. d. Students know how to solve problems involving heat flow and temperature changes, using known values

of specific heat and latent heat of phase change. Students Outcomes 1. Explain why energy always flows from warnter to cooler. 2. Describe specific heat is the amount of energy needed to raise the temperature of a substance.. 3. Know that specific heat is an inherent property of substances. 4. Chemical and physical changes involve the flow of energy. 5. Be able to solve problems involving specific heat and temperature change. 6. Temperature does not change during a phase change. Possible Strategies and Activities 1. Demonstrate calorimeter, Enthalpy, thermo – chemical equations, heats of formation. 2. Solve heat flow problems 3. Calorimeter labs using Styrofoam cups and different metal pieces to identify specific heats of various

metals. 4. Demonstrate athletic trainer’s cold packs (endothermic) and hot packs (exothermic). Ziploc bag reactions

also work. 5. Show students models of molecular motion of gases, solids and liquids to show energy loss or gain

during phase change.

Unit 7: Solutions

State Standards Solutions 6. Solutions are homogenous mixtures of two or more substances. As a basis for understanding this concept:

a. Students know the definitions of solute and solvent. b. Students know how to describe the dissolving process at the molecular level by using the concept of

random molecular motion. c. Students know temperature, pressure, and surface area affect the dissolving process. d. Students know how to calculate the concentration of a solute in terms of grams per liter, molarity, parts

per million, and percent composition. 2. Biological, chemical, and physical properties of matter result from the ability of atoms to form bonds from

electrostatic forces between electrons and protons and between atoms and molecules. As a basis for understanding this concept: d. Students know the atoms and molecules in liquids move in a random pattern relative to one another

because the intermolecular forces are too weak to hold the atoms or molecules in a solid form.



Student Outcomes 1. Be able to state the properties and states of matter verbally and to diagram the states of mater at a

macro and micro level. 6b 2. Compare solids, liquids, and gases and be able to explain the distinguishing characteristics of each. 6b 3. Be able to explain the structure and properties of liquids, including surface tension, surface area, and

how the intermolecular forces in water make it a unique substance. 2d, 6c 4. Be able to explain what a solution is (distinguishing between solute and solvent) including the process

of solvation, factors affecting solubility and conductivity, and methods of purification. 6a, 6b 5. Students will be able to calculate the concentration of a solute in terms of grams per liter, molarity,

and percent composition. 6d Possible Strategies and Activities 1. Use models and lecture to illustrate the differences between a solid, liquid and gas on the macro and

molecular level. 2. Select experiments that will facility the students use of laboratory data, and when possible computer

graphing software to illustrate heating and cooling curves, and to calculate heats of phase changes. 3. Supplement lecture with laboratory investigations and demonstrations to illustrate the properties

of solutions and the colligative effects of various solutes.

Unit 8: Equilibrium Chemical Equilibrium

State Standards 9. Chemical equilibrium is a dynamic process at the molecular level. As a basis for understanding this

concept: a. Students know how to use LeChatelier’s principle to predict the effect of changes in concentration,

temperature, and pressure. b. Students know equilibrium is established when forward and reverse reaction rates are equal.

Student Outcomes 1. Describe the equilibrium condition. 2. Describe the microscopic processes that occur in a system when it is at equilibrium. 9b 3. Use Le Chatelier's principle to explain how addition of a catalyst or changes in concentration,

temperature, and/or pressure affect a system at equilibrium. 9a 4. Determine equilibrium exp0ressions for a given reaction. Possible Strategies and Activities 1. Present examples that clearly illustrate the similarities and differences between steady state and

equilibrium conditions. 2. Emphasize the dynamic processes that occur at the microscopic level within a system that is at

equilibrium. 4. Give students numerous opportunities to apply Le Chatelier's principle, including written problems,

demonstrations, and/or experiments. 5. Illustrate how to derive an equilibrium law expression; then provide practice opportunities with

monitoring of student progress. 6. Show several variations of how equilibrium law calculations are used to find unknown concentrations or

equilibrium constants.



Unit 9: Rates of Reaction

State Standards Reaction Rates 8. Chemical reaction rates depend on factors that influence the frequency of collision of reactant molecules.

As a basis for understanding this concept: a. Students know the rate of reaction is the decrease in concentration of reactants or the increase in

concentration of products with time. b. Students know how reaction rates depend on such factors as concentration, temperature, and pressure. c. Students know the role a catalyst plays in increasing the reaction rate.

Student Outcomes 1. Be able to explain how factors such as concentration, temperature, pressure, and the influence of

catalysts affect reaction rates. 8b, 8c 2. Be able to graph the change in amount of products versus reactants over time, including the affects of

catalyst on the activation energy. 8a Possible Strategies and Activities 1. Give students numerous opportunities to apply Le Chatelier's principle, including written problems,

demonstrations, and/or experiments. 2. Graph concentration verses time for iodine clock reaction. 3. Use graphs of activation energy to demonstrate catalyst affects on reactions. 4. Demonstrate catalyst using split tests for oxygen using manganese dioxide in the decomposition of

hydrogen peroxide.

Unit 10: Acids & Bases

State Standards Acids and Bases 5. Acids, bases, and salts are three classes of compounds that form ions in water solutions. As a basis for

understanding this concept: a. Students know the observable properties of acids, bases, and salt solutions. b. Students know acids are hydrogen-ion-donating and bases are hydrogen-ion-accepting substances. c. Students know strong acids and bases fully dissociate and weak acids and bases partially dissociate. d. Students know how to use the pH scale to characterize acid and base solutions.

Student Outcomes 1. Demonstrate. how to calculate pH from a hydrogen-ion concentration in a laboratory exercise utilizing

different aqueous concentrations. 5c 2. Use a pH meter and hydronium paper to identify solutions that are acidic and basic. 5d 3. Apply the Arrhenius, Brønsted-Lowery, and Lewis acid-base concepts to lecture problems and

homework practice exercises. 5c 4. Identify the observable properties of acids, bases, and salt solutions. 5a 5. Compare and contrast the differences between acids that are hydrogen-ion-donating and bases that are

hydrogen-ion-accepting substances. 5b Possible Strategies and Activities 1. Clarify how acids are hydrogen-ion-donating and bases are hydrogen-ion-accepting substances. 2. Provide classroom demonstrations that differentiate between acids, bases, and salt solutions.



3. Monitor student's use of pH meters during laboratory experiments with different strengths of aqueous

acidic and aqueous basic solutions. 4. Synthesize the concepts of Arrhenius, Brønsted-Lowery, and Lewis Acid-base solutions in simplifying

how each concept is involved in a solution. 5. Evaluate student understanding of how strong bases and bases fully dissociate and weak acids and

bases partially dissociate with practice problems, quizzes, and homework. 6. Teachers will facilitate learning with classroom demonstrations, Internet investigations of current

topics, and the use of relevant software.

Unit 11: Organic Chemistry

State Standards Organic Chemistry and Biochemistry 10. The bonding characteristics of carbon allow the formation of many different organic molecules of varied

sizes, shapes, and chemical properties and provide the biochemical basis of life. As a basis for understanding this concept: a. Students know large molecules (polymers), such as proteins, nucleic acids, and starch, are formed by

repetitive combinations of simple subunits. b. Students know the bonding characteristics of carbon that result in the formation of a large variety of

structures ranging from simple hydrocarbons to complex polymers and biological molecules. c. Students know amino acids are the building blocks of proteins.

Student Outcomes 1. Be able to explain the importance of carbon in the formation of many organic molecules essential to life.

10a 2. Define the role of proteins, nucleic acids, and starch as they relate to biological structure, movement,

catalysts, transport, storage, energy transformation, protection, and buffering. 10b 3. Identify the thirty essential elements and know the R-group structure of amino acids and how they

combine to form the polypeptide backbone structure of proteins. 10c Possible Strategies and Activities 1. Illustrate and provide an opportunity for students to differentiate between simple linear hydrocarbons,

hydrocarbons with double and, triple bonds, and molecules that contain benzene rings through demonstration, laboratory experiments, and lecture.

2. Supplement lecture with overheads, film, and models to illustrate the thirty essential elements used by amino acids as the building blocks of protein, and the basic structure of proteins.

3. Explain the role of proteins, nucleic acids, and starch using a systemic approach to the biochemical processes in the human body.

4. Supplement lecture with classroom demonstrations, Internet investigations, laboratory assignments and where appropriate relevant software will be incorporated into content area.

5. Provide opportunities for the students to complete laboratory experiments, analyze data, and write laboratory reports in the above content areas.

Unit 12: Nuclear Chemistry

State Standards Nuclear Processes 11. Nuclear processes are those in which an atomic nucleus changes, including radioactive decay of naturally

occurring and human-made isotopes, nuclear fission, and nuclear fusion. As a basis for understanding this concept:



a. Students know protons and neutrons in the nucleus are held together by nuclear forces that overcome the

electromagnetic repulsion between the protons. b. Students know the energy release per gram of material is much larger in nuclear fusion or fission

reactions than in chemical reactions. The change in mass (calculated by E=mc2) is small but significant in nuclear reactions.

c. Students know some naturally occurring isotopes of elements are radioactive, as are isotopes formed in nuclear reactions.

d. Students know the three most common forms of radioactive decay (alpha, beta, and gamma) and know how the nucleus changes in each type of decay.

e. Students know alpha, beta, and gamma radiation produce different amounts and kinds of damage in matter and have different penetrations.

Student Outcomes 1. Be able to explain how nuclear forces allow protons to overcome their natural repulsion and reside in

the nucleus. lla 2. Be able to state the difference between nuclear fusion and nuclear fission, and be able to relate these

differences to the amount of energy generated in each process. llb 3. Describe the existence of radioactive isotopes, both natural, and those formed as a byproduct of

nuclear reactions, and explain how a radiation counter is used to measure radioactive isotopes. llc 4. Compare the different amounts of radiation that alpha, beta, and gamma particles produce, know how

the nucleus changes in each type of decay, and be able to describe the penetration ability, and types of damage that each will produce in matter. 11 d, 11 e

5. Students will be able to explain how a nuclear reactor works, and the importance of chain reactions in this process (including the idea of critical mass).

Possible Strategies and Activities 1. Explain and model nuclear forces as they pertain to atomic structure and illustrate how energy released

per gram of material is much larger via nuclear fusion or fission than in previously studied chemical reactions.

2. Use a radiation counter, illustrations, and lecture to demonstrate the radiation emitted from alpha, beta, and gamma radiation and explain the differences in penetration and damage between the three types of radiation.

3. Use guided practice, and examples to show students how to utilize the half-life rule to calculate the amount of material that will decay during a specified amount of time.

4. Use classroom demonstrations, Internet investigations and computer technology where appropriate to compliment lecture and provide the opportunity for students to conduct laboratory experiments, analyze data, and complete laboratory reports to highlight these concepts.



INVESTIGATION AND EXPERIMENTATION Scientific progress is made by asking meaningful questions and conducting careful investigations. Students will be able to: • Plan and conduct a scientific investigation to test a hypothesis. • Evaluate the accuracy and reproducibility of data. • Distinguish between variables and controls in a test. • Design, construct and interpret graphs using data from investigations. • Apply simple mathematical relationships to determine a missing quantity in a mathematic

expression, given the two remaining terms. The teacher will integrate the investigation and experimentation topics into the content Focus Standards throughout the year, instead of an isolated unit, spiraling content until mastery is achieved by the end of the year. The teacher will devote a minimum of 20% of the instruction time to hands on experiences in the laboratory.

June, 2005, at the Summer Science Institute at Salinas Union School District, teachers adopted by consensus the following criteria and standards an investigation, activity, or experiment must have to be considered as part of the minimum 20 % required investigation and experimentation Focus Standard. Definition of lab: All labs will include: • A clear objective • A descriptive procedure • Data collection • Qualitative and/or quantitative observations • Analysis of data • A conclusion – communication of results

An experimental lab will include: • All of the above and a(n)

1. Hypothesis 2. Control 3. Variable(s)

Lab required elements: • Title • Background information (optional) • Objective with standard(s) • Questions/problem • Hypothesis (for experimental labs only) • Materials listed • Safety standards (worked into procedure) • Procedure



• Data • Analysis of Data • Conclusion

*Scaffolding of required elements can be changed according to student needs. Teacher information: • Grading rubric (to hand out to students) • Prelab activities • Expected results • Possible difficulties • Set-up time • Lab completion time • Resources/references • Differentiation • Vocabulary/skills needed • Answers to lab questions

Addition demonstrations, activities, manipulatives, and other type of lab investigations are encouraged if they enhance student learning, but are not to be considered as part of the minimum 20% of investigations and experimentation. SUGGESTED TEACHING STRATEGIES: • Whole group instructions and strategies. • Cooperative groups. • Individualized instruction. • Peer teaching. • Independent practice. • Audio-visual aided instruction. • Individual and or group projects. • Field trips. • Interactive media. • Computer assisted learning. • Internet activities. • Laboratory activities. • Demonstrations.

EVALUATION TECHNIQUES: • Teacher developed tests. • Peer Evaluation. • Laboratory write-ups. • Homework. • Directed study. • Class participation. • Independent research and/or experimental projects.



TERMINOLOGY GUIDE

Standard A goal statement that identifies the knowledge and skills to be learned. A content standard specifies what we want students to know and be able to do. The standards identified here are derived from the Science Content Standards for California Public Schools (1998) and serves as the basis of statewide student assessments, the science curriculum framework, and the evaluation of instructional materials. Attention to and familiarity with these standards identified should be the responsibility of the teacher.

Focus Standard (Integrated Focus Standards) A standard, or parts of a standard, identified by teachers locally as “essential” for students to learn. While teachers may teach more than the identified Focus Standards, they will teach these Focus Standards as targets for student learning and assessment. Cumulatively, the identified Focus Standards form the core of what should be taught, learned and assessed during the school year.

Related California State Standards Prior standards identified from the Science Content Standards for California Schools that are vertically aligned to the current Focus Standards for each science content area or grade level. These standards are introduced in prior science classes, but provide support for student learning of the current content standards. The No Child Left Behind framework includes identified Related California State Standards as a part of the targets for student learning and assessment. Teachers should integrate these related standards into the current Focus Standards to ensure student understanding.

Enduring Understanding The core concepts, principles, theories, and processes that serve as the focal point of curriculum, instruction, and assessment. Enduring Understandings are important, enduring, and transferable beyond the scope of a particular unit. Enduring Understandings identify the “key concepts” or overarching ideas derived from the California Science Content Standards and Framework (2004). Enduring Understandings support the content for that grade level, but may also be applicable at other grade levels. Enduring Understandings are written for student access and understanding at that grade level. Enduring Understandings maybe repeated across the grades to provide “bridges” for students and a sense of connectedness across place and time.

Essential Question A provocative question designed to engage student interest and guide inquiry into the important ideas in a field of study. Rather than yielding pat answers, Essential Questions are intended to stimulate discussion and rethinking over time. These may be seen as “doorways” to the Enduring Understandings.



Integrate Integration of curriculum takes on different meanings in different situations. For the purposes of these curriculum guides, integration should be understood as: • The blending of core concepts in teaching Science lessons/units (i.e., teaching science investigation and experimentation, life science, chemistry, earth science, physical science and/or other science areas simultaneously); • Teaching more than one subject within a lesson/unit (i.e., Science and Mathematics, Language Arts, History-Social Science, etc.). • Combining two or more parts of a standard (e.g., Life Science 1.c and 1.f) to teach one lesson/unit.

Benchmark Benchmarks are a series of assessments used to determine how students are progressing in their learning throughout the school year. Benchmark assessments may be administered at the end of each quarter, for example, to show “where students are” at that time. Benchmarks provide a snapshot of student achievement and may be indicative if teaching resources and strategies are effective in helping students achieve maximum learning. Benchmarks are forms of assessment used to inform students and teachers of the student’s on-going progress. The results of these assessments are used to continue/extend the learning process and/or modify teaching strategies and resources in order to benefit the student.

Blueprint Blueprints are guides developed by the California Department of Education to assist teachers in identifying which standards will be assessed (and to what degree each standard will be weighted in the total assessment) through the California Standards Test and/or No Child Left Behind test. The Science Framework for California Public Schools is the blueprint for reform of the science curriculum, instruction, professional preparation and development, and instructional materials in California. The blueprint also identifies the reporting clusters for each test.

Assessments Techniques to analyze student accomplishments against specific goals and criteria. Includes paper-pencil tests, exhibits, investigative and experimental labs, exhibits, interviews, surveys and observations. Good assessment requires a balance of techniques because each technique is limited and prone to error. Assessment is sometimes synonymous with evaluation, though a subtle difference exists. A teacher might assess the strengths and weaknesses of a student’s performance without placing a value or a grade on that performance.

Assessment Sample Sample performance evaluation and district test questions that are keyed to specific objectives and focus standards. Samples include enhanced multiple-choice questions, open-ended questions, and performance tasks. Assessments samples are to serve as a guide for instructors to make meaningful aligned assessments within the content and context of district, state, and national standards at the appropriate level to ensure student understanding and learning.

Template A guide or framework for designers. In its original usage, a template was a thin form, constructed of paper, wood, or sheet metal, whose edge provided a guide for cutting a particular shape. In




Understanding by Design, the unit-planning template provides a conceptual guide to apply the elements of backward design in developing or refining a unit of study.

Backward Planning A process used to design curriculum or units by beginning with the end in mind and designing plans for teaching, resources, learning and assessment toward that end. Why is such a view backward when it seems logical? Many teachers begin their unit design with textbooks, favored lessons, and time-honored activities rather than deriving ideas from targeted goals or standards. In Backward Planning, one starts with the end – the desired results- and then identifies the evidence necessary to determine that the results have been achieved – the assessments. With the results and assessments clearly specified, one can determine the necessary knowledge and skills, and then the teaching needed to equip students to perform.

Understanding by Design A theory developed by researchers and educational consultants Jay McTighe and Grant Wiggins with the intent tot teach for understanding using a method of Backward Planning.

Unit: Atomic Structure and the Periodic Table Assessment - BMK 1 Time: 4 Hours Concepts:

• The atom is composed of protons, neutrons and electrons • Bohr’s model of the atom consists of a small dense nucleus made up of protons and neutrons with electrons orbiting

around the nucleus • The periodic table is a graphic representation of the relationship between groups or families of elements

The teacher will integrate the investigation and experimentation topics into the content Focus Standards throughout the year instead of an isolated unit devoting 40% of the instructional time to hands on experiences in the laboratory. State Standard Goal: ch.1 - The periodic table displays the elements in increasing atomic number and shows how periodicity of the physical and chemical properties of the elements relates to atomic structure. As a basis for understanding this concept: . Focus Standards: ch.1a: how to relate the position of an element in the periodic table to its atomic number and atomic mass. ch.1b: use the periodic table to identify metals, semimetals, nonmetals, and halogens. ch.1c: use the periodic table to identify alkali metals, alkaline earth metals and transition metals ch.1d: use the periodic table to determine the number of electrons available for bonding. ch.1e: the nucleus of the atom is much smaller than the atom yet contains most of its mass Enduring Understandings:

1. The atom is the basic unit of matter. 2. Elements are placed on the periodic table in a specific arrangement.

Essential Questions:

1. Why are elements in families? 2. What makes a noble gas noble?



Key Vocabulary for Atomic Structure and the Periodic Table alkali metals group/family period alkaline earth metals halogen periodic law atom metals periodic table atomic mass neutron physical properties atomic number noble gas proton chemical properties nonmetals semimetals, metalloids electron nucleus transition metals

Supporting Vocabulary periodicity isotope atomic mass unit (amu) representative elements

Academic Vocabulary volume



State Standard Goal: Assessment – BMK 1 Time: ch.1 - The periodic table displays the elements in increasing atomic number and shows how periodicity of the physical and chemical properties of the elements relates to atomic structure I Standard: Focus

• Student Objectives II Evidence of Student Learning

Performance Indicators III Vocabulary

Periodic Table: ch.1a: how to relate the position of an element in the periodic table to its atomic number and atomic mass. Objective: Student will

• Know how to interpret information from the periodic table to find elements by their atomic number

• Demonstrate an understanding of the regions on the periodic table

• On a quiz of elements and their symbols will identify the correct symbol 80% of the time

• From a diagram of the periodic table be able to identify specific elements and give their atomic number and atomic mass

periodic table periodic law group/family period alkali metal halogen noble gas

noble gas alkaline earth metals atomic number atomic mass atomic mass unit (amu) isotope periodicity

IV Previous Concepts /Prior Knowledge

V Suggested Teaching Strategies, Activities, and Labs

VI Suggested Instructional Resources

Physical properties of elements: solids, liquids, gases

Introduction – Periodic table game – using “pun” sheet to become familiar with elements

Article – “Periodic Table” Small scale lab – “Atomic mass of Candium” p. 122

Textbook: Chemistry – Addision-Wesley, Prentice Hall 2002 (Chemistry – Holt 2004) pp. 113-121 pp. 123-126 (pp.84-88) (pp. 124-131) www.scilinks.org –Periodic Table HW 4094 Science News at www.phschool.com

VII Teacher Notes: Vocabulary terms - * used in previous objective/unit; Italic = supporting vocabulary

http://www.scilinks.org/






Periodic Table: ch.1b: use the periodic table to identify metals, semimetals, nonmetals, and

halogens. ch.1c: use the periodic table to identify alkali metals, alkaline earth metals and

transition metals Objective: Student will

• Identify the location of metals, semimetals, nonmetals and halogens on the periodic table and relate the properties of elements to their chemical families

• Identify elements as alkali metals, alkaline earth metals and transition metals from their position on the periodic table

• On a periodic table identify, using color coding, for example, the location of metals, semimetals, nonmetals, halogens, alkali metals, alkaline earth metals and transition metals

metals nonmetals semimetals (metalloids) halogens* alkali metals* alkaline earth metals* transition metals chemical properties representative elements physical properties noble gas*




Review: element symbols

Use sample periodic table as class activities, distinguishing areas by using different colors

Lab – flame test using metal samples

Textbook: Chemistry – Addision-Wesley, Prentice Hall 2002 (Chemistry – Holt 2004) pp.123-126 (pp.115-131)

VII Teacher Notes: Vocabulary terms - * used in previous objective/unit; Italic = supporting vocabulary semimetals = metalloids use both terms with students






Periodic Table: ch.1d: use the periodic table to determine the number of electrons available for

bonding. ch.1e: the nucleus of the atom is much smaller than the atom yet contains most

of its mass Objective: Student will • Be able to determine the number of valence electrons for given elements • Discuss the atomic structure and explain the location and size/mass of the

component parts

• Quick write on the atomic structure of an atom and the location and relative mass of its parts

• Determine for a list of elements the number of valence electrons for each element

valence electrons proton neutron electron nucleus




Review: • basic atomic structure of an

atom

Textbook: Chemistry – Addision-Wesley, Prentice Hall 2002 (Chemistry – Holt 2004) pp. 107-112 (pp. 79-83) Video - Rutherford

VII Teacher Notes:



Assessment Sample:

Multiple Choice Other 1.

Iodine is a halogen. It would have chemical properties most like (ch1b)

A manganese (Mn) B tellurium (Te) C chlorine ( Cl)* D xenon (Xe)

2. Which statement best describes the density of an atom’s nucleus? (ch1e)

A. The nucleus occupies most of the atom’s volume but contains little of its mass.

B. The nucleus occupies very little of the atom’s volume and contains little of its mass.*

C. The nucleus occupies most of the atom’s volume and contains most of its mass.

D. The nucleus occupies very little of the atom’s volume but contains most of its mass.



Videos:

Additional Websites: Research Possibilities:



Unit: Chemical Bonding and Balancing Equations Assessment - BMK 1 Time: 25 Days Concepts:

• Naming and writing of chemical formulas is essential for writing balanced equations • Chemical bonding is a result of electron sharing or transfer • The number of atoms for each element must be equal on both sides in a balanced equation.

The teacher will integrate the investigation and experimentation topics into the content Focus Standards throughout the year instead of an isolated unit devoting 40% of the instructional time to hands on experiences in the laboratory. State Standard Goal: ch.2: Chemical Bonding: Biological, chemical, and physical properties of matter result from the ability of atoms to form

bonds from electrostatic forces between electrons and protons and between atoms and molecules. ch.3: The conservation of atoms in chemical reactions leads to the principle of conservation of matter and the ability to

calculate the mass of products and reactants. Focus Standards: Naming and writing chemical formulas ch.2a: atoms combine to form molecules by sharing electrons to form covalent or metallic bonds or by exchanging electrons

to form ionic bonds. ch.3a: how to describe chemical reactions by writing balanced equations Determining types of chemical reactions Enduring Understandings:

1. A molecule is two or more atoms chemically combined. 2. In a chemical reaction the amount of reactants have to equal the amount of products. 3. Chemical bonds are all about electrons.


1. What is the difference between an atom and a molecule? 2. Why does the number and type of each atom have to be the same on both sides of a chemical equation? 3. Yo – what’s up with the electrons?



Key Vocabulary for Chemical Bonding and Balancing Equations anions chemical reaction ions polyatomic ions balanced equation coefficient metallic bond products cations covalent bond molecular compound reactants chemical equation ionic bond molecule valence electron chemical formula ionic compound monatomic ion

Supporting Vocabulary

binary compound molecular formula formula unit ternary compound Law of Definite Proportion

Academic Vocabulary combine sharing






Naming and Writing Formulas: Objective: Student will

• Be able to write and name chemical formulas

• Practice sheet of writing formulas for compounds

• Mini-assessment of writing formulas and naming compounds

chemical formula coefficient molecular formula




Review: • Symbols of elements • Valence electrons of

elements • Covalent and ionic bonding

• Practice - naming and writing chemical formulas • Graphic organizers – flow charts for naming

compounds

Chemistry – Addision-Wesley, Prentice Hall 2002(Chemistry – Holt 2004) Chapter 6 - Chemical Names and Formulas pp. 132-156 (Chapter 5- pp. 176-180) www.scilinks.org Naming compounds HW4081 (Chapter 6 – pp. 190-193, 206-207)

VII Teacher Notes: Vocabulary terms - * used in previous objective/unit; Italic = supporting vocabulary




State Standard Goal: Assessment – BMK 1 Time: ch.2: Chemical Bonding: Biological, chemical, and physical properties of matter result from the ability of atoms to form bonds

from electrostatic forces between electrons and protons and between atoms and molecules. I Standard: Focus



Chemical bonds: ch.2a: atoms combine to form molecules by sharing electrons to form

covalent or metallic bonds or by exchanging electrons to form ionic bonds.

Objective: Student will • Discuss the similarities and differences between covalent, metallic

and ionic bonds

• Quick write on the nature of covalent, metallic and ionic bonds

• Identify whether a compound has a covalent, or ionic bond

• Graphic organizer comparing and contrasting ionic and covalent bonds

electron chemical bond covalent bond metallic bond ionic bond




Review: • valence electrons • chemical nature of elements

from periodic table

Conductivity lab – Testing conductivity of salt water solution Lab – Names and formulas of ionic compounds (PH p. 157) Mini-lab – Making Ionic Compounds (PH p.163)

Textbook: Chemistry – Addision-Wesley, Prentice Hall 2002 (Chemistry – Holt 2004) Chapter 8 - pp. 203-211 (Chapter 8 – pp. 267-274, 286-288)

VII Teacher Notes:



State Standard Goal: Assessment – BMK 1 Time: ch.3: The conservation of atoms in chemical reactions leads to the principle of conservation of matter and the ability to calculate the

mass of products and reactants. I Standard: Focus



Chemical equations: ch.3a: how to describe chemical reactions by writing balanced equations Objective: Student will

• Draw a diagram of the particles in a chemical equation • Write a balanced equation

• draw a diagram of a balanced chemical reaction using atoms and molecule shapes

• Given a chemical equations without coefficients, balance the equation by adding the correct coefficients

chemical equation conservation of matter balanced equation reactants products atom molecule coefficient




Review: • Valence electrons • Covalent and ionic bonds

Activity: Using paper positive and negative ion models, work to form compounds. Practice sheets for writing balanced equations

Textbook: Textbook: Chemistry – Addision-Wesley, Prentice Hall 2002 (Chemistry – Holt 2004) Chapter 8 – pp. 212-228 (Chapter 8 – pp. 275-285)

VII Teacher Notes:







Determining types of chemical reactions: Objective: Student will

• Be able to identify the different types of chemical reactions, including: combination (synthesis), decomposition, single replacement, double-replacement, and combustion reactions

• From a given set of chemical equations, identify the type of chemical reactions

chemical reaction combination reaction (synthesis reaction) decomposition reaction single replacement reaction double replacement reaction combustion reaction




Review: • Compounds • Balanced chemical equations

• Precipitation reactions – formation of solids (PH p. 229)

Textbook: Chemistry – Addision-Wesley, Prentice Hall 2002 (Chemistry – Holt 2004) pp.212-226 (pp. 275-285)

VII Teacher Notes:



Assessment Sample:

Multiple Choice Other 1. C3H8 + O2 CO2 + H2O This chemical equation represents the combustion of propane. When correctly balanced, what will be the coefficient for water? (ch3a) A 2 B 4* C 8 D 16

2. Which of the following pairs of elements are most likely to form an ionic bond? (ch2a)

A. Br and Ca B. Br and N C. Ca and Mg D. Ca and Fe*



Videos: “World of Chemistry”

Additional Websites: www.gohrw.com/liveink Research Possibilities:

http://www.gohrw.com/liveink



Unit: Moles and Stoichiometry Assessment - BMK 2 Time: 14 Hours Concepts:

• The mole is the central unit in stoichiometry

The teacher will integrate the investigation and experimentation topics into the content Focus Standards throughout the year instead of an isolated unit devoting 40% of the instructional time to hands on experiences in the laboratory. State Standard Goal: ch.3: The conservation of atoms in chemical reactions leads to the principle of conservation of matter and the ability to

calculate the mass of products and reactants. Focus Standards: ch.3b: the quantity one mole is set by defining one mole of carbon 12 atoms to have a mass of exactly 12 grams. Ch.3c: one mole equals 6.02x1023particles (atoms or molecules). Ch.3d: how to determine the molar mass of a molecule from its chemical formula and a table of atomic masses and how to

convert the mass of a molecular substance to moles, number of particles, or volume of gas at standard temperature and pressure.

Ch.3e: how to calculate the masses of reactants and products in a chemical reaction from the mass of one of the reactants or products and the relevant atomic masses.

Enduring Understandings:

1. Matter cannot be created nor destroyed. 2. Moles count atoms.

. Essential Questions:

1. Who is Ms Mole?



Key Vocabulary for Moles and Stoichiometry Avogadro’s number molar mass Law of Conservation of Mass mole Law of Definite Proportions standard temperature & pressure (STP) Law of Multiple Proportions stoichiometry

Supporting Vocabulary

actual yield coefficients percent yield atomic mass empirical formula products balanced reaction limiting reactant reactants chemical reaction mass number theoretical yield

Academic Vocabulary conservation







Moles: ch.3b: the quantity one mole is set by defining one mole of carbon 12

atoms to have a mass of exactly 12 grams. Ch3c: one mole equals 6.02x1023particles (atoms or molecules) Objective: Student will be able to

• define a mole • define Avogadro’s number

• Quick write on the concept of a mole • Discuss an analogy to the concept of a

mole

Avogadro’s number mole




Review: • atoms, molecules, compounds • exponents and coefficients • numbers as quantities with

units

Analogies about the concept of the mole (i.e. 1 doz = 12 as 1 mol = Avogadro’s number)

Practice problems

Textbook: Chemistry – Addision-Wesley, Prentice Hall 2002 (Chemistry – Holt 2004) Chapter 7 – Chemical Quantities pp. 170-186 (Chapter 9 pp. 301-310)

VII Teacher Notes:



State Standard Goal: Assessment – BMK 2 Time: ch.3: The conservation of atoms in chemical reactions leads to the principle of conservation of matter and the ability to calculate the mass of products and reactants. I Standard: Focus



Molar Mass & Stoichiometry: ch.3d: how to determine the molar mass of a molecule from its chemical formula and a table of atomic masses and how to convert the mass of a molecular substance to moles, number of particles, or volume of gas at standard temperature and pressure. Objective: Student will

• Determine the molar mass of a molecule from its chemical formula and atomic masses.

• Convert the mass of molecular substances to moles, mber of particles, or volume of gas at S

• Students should know to calculate molar

mass from its chemical formula. • Balanced the equations.

Molar mass stoichiometry limiting reactants excress reactants mole ratio percent yield




Review: balancing chemical equations atomic/molecular mass algebraic manipulation of

formulas and unit conversion significant figures numbers as quantities with

units exponents and coefficients

Small Scale Lab – measuring mass as a means of counting (PH p. 187)

Practice problems

Textbook: Chemistry – Addision-Wesley, Prentice Hall 2002 (Chemistry – Holt 2004)

VII Teacher Notes:



State Standard Goal: Assessment – BMK 2 Time: ch.3: The conservation of atoms in chemical reactions leads to the principle of conservation of matter and the ability to calculate the mass of products and reactants. I Standard: Focus



Stoichiometry: ch.3e: how to calculate the masses of reactants and products in a chemical reaction from the mass of one of the reactants or products and the relevant atomic masses. Objective: Student will • Calculate the mass of reactants and/or products from a balanced

equation given the mass of one of the reactants or products • Predict the mass of products in a single replacement reaction or other

chemical reaction in a lab setting • Demonstrate use of dimensional analysis (factor labeling) to solve

stoichiometry problems

• Correctly calculate the mass of a product and/or reactant on practice problems

• Demonstrate correct use of dimensional analysis (factor labeling) on a set of stoichiometry problems

• Predict the mass of a product produced in a lab experiment

Law of Multiple proportions Law of Definite Proportions Law of Conservation of Mass Avogadro’s number stoichiometry standard temperature & pressure (STP) mole molar mass




Review: balancing chemical equations atomic/molecular mass algebraic manipulation of

formulas and unit conversion significant figures numbers as quantities with

units exponents and coefficients

Dimensional analysis and explanation of conversion factors and a mole ratio to introduce stoichiometry

Cooking analogies (i.e. two pieces of bread, one slice of bread)

Stoichiometry lab that shows that you can predict the mass of the products (such as labs with a single replacement reaction)

Practice problems Lab – Conservation of mass: Analysis of Baking Soda

(PH p. 251)


VII Teacher Notes:



Assessment Sample:

Multiple Choice Other 1. A mole is defined as (std 3b) A. the amount of mass in exactly 1.00 ml B. the number of atoms in 12.0 grams of any element C. a dozen D. the number of atoms in 12.0 grams of carbon-12 * 2. How many ions are in 0.5 moles of NaCl? (std 3d) A. 1.204 x 10 23 ions B. 3.011 x 10 23 ions C. 6.022 x 10 23 ions * D. 9.033 x 10 23 ions 3. How many moles of Calcium are in a serving of milk containing 290

mg of calcium? (std 3e)

A. 2.90 x 10-3 B. 7.23 x 10-3* C. 2.90 mole D. 7.23 mole

___ NH3(g) + __ O2(g) __ N2(g) + __ N2O(g) 4. When the reaction above is completely balanced, the coefficient for

NH3 will be (std. 3a) A. 2. B. 3. C. 4. D. 6.

1. How many grams of oxygen will be produced by the reaction of 10 moles of O3 in the reaction? (std 3d)

2O2 3O3

A. 48 g B. 96 g C. 480 g * D. 960 g



Videos: “The World of Chemistry – The Mole”




Unit: Gas Laws Assessment - BMK 2 – Time: 4 Hours Concepts:

• Gas laws explain the properties of gases The teacher will integrate the investigation and experimentation topics into the content Focus Standards throughout the year instead of an isolated unit devoting 40% of the instructional time to hands on experiences in the laboratory. State Standard Goal: ch.4: The kinetic molecular theory describes the motion of atoms and molecules and explains the properties of gases. Focus Standards: ch.4a: the random motion of molecules and their collisions with a surface create the observable pressure on that surface. ch.4b: the random motion of molecules explains the diffusion of gases. ch.4c: how to apply the gas laws to relations between the pressure, temperature, and volume of any amount of an ideal gas or

any mixture of ideal gases. ch.4d: the values and meanings of standard temperature and pressure (STP). ch.4e: how to convert between the Celsius and Kelvin temperature scales. ch.4f: there is no temperature lower than 0 Kelvin. Enduring Understandings:

1. All gases follow a group of predictable properties. 2. Temperature is a measurement of the kinetic energy of particles.


1. Does hot air really rise?



Key Vocabulary for Gas Laws atmospheric pressure (atm) compressible/compressibility Graham's law partial pressure Avogadro's Law Dalton's Law of Partial Pressure ideal gas constant ( R) Pascal (Pa) barometer diffusion ideal gas law standard pressure Boyle's Law effusion Kelvin temperature standard temperature Charles' Law elastic collision kinetic energy collision gas pressure Kinetic-Molecular Theory combined gas law Gay-Lussac's Law Newton (N)

Supporting Vocabulary monometer temperature pressure vacuum Standard Temperature & Pressure (STP) volume

Academic Vocabulary random theoretical, theoretically equivalent

Important Equations



State Standard Goal: Assessment – BMK 2 Time: ch.4: The kinetic molecular theory describes the motion of atoms and molecules and explains the properties of gases. I Standard: Focus


Performance Indicators

III Vocabulary

Gas Pressure: ch.4a: the random motion of molecules and their collisions with a surface create the observable pressure on that surface. Objective: Student will

• Explain how the motion of gases and their collisions create the observed pressure a gas exerts on a surface

• Quiz explaining how the kinetic molecular theory explains the pressure exerted by a gas on its containier

pressure collision elastic collision kinetic energy gas pressure compressible barometer monometer kinetic molecular theory atmospheric




Review: • Pressure, mass, velocity,

kinetic energy and area

Demonstrations of Boyle’s Law with a vacuum pump Demo using BBs in a clear dish (such as a petri dish)

on the overhead to show the motion of gas particles and how they create pressure

Minilab – Carbon dioxide from Antacid Tablets (PH p. 346)

Textbook: Chemistry – Addision-Wesley, Prentice Hall 2002 (Chemistry – Holt 2004) Chapter 10 pp. 267-272 Chapter 12 pp.353 (Chapter 12 pp. 414-442)

VII Teacher Notes:






Diffusion of gases: ch.4b: the random motion of molecules explains the diffusion of gases. Objective: Student will

• Explain how gases diffuse due to the random motion and elastic collisions of molecules

• Quiz explaining how the kinetic molecular theory explains the diffusion of gases

random collision* elastic collision* diffusion kinetic molecular theory* effusion




Review: • Kinetic molecular theory • Std 4a – random movement

gas particles and gas pressure

• velocity

Open perfume or ammonia on one side of the room and smell it on the other

Cotton wad dipped in ammonia and another dipped in HCl in a tube (or just opened near each other in the room) to see the white cloud that results from the diffusion of the gases. Can also incorporate Graham’s Law here.

Textbook: Chemistry – Addision-Wesley, Prentice Hall 2002 (Chemistry – Holt 2004) pp.326-354

VII Teacher Notes:






Gas Laws: ch.4c: how to apply the gas laws to relations between the pressure, temperature, and volume of any amount of an ideal gas or any mixture of ideal gases. Objective: Student will

• be able to use the combined gas law to solve for temperature, pressure or volume

• be able to use the ideal gas law to solve for one variable when the others are known

• be able to use Dalton’s law of partial pressures to calculate total pressure or partial pressure of one gas in a mixture

• successfully solve problems given initial conditions of temperature, pressure, and volume when one value has changed

• successfully solve problems using the ideal gas equation to find the fourth variable

• successfully solve problems using Dalton’s law of partial pressure to solve for total pressure or partial pressure

temperature * pressure volume ideal gas Kelvin temperature Newton ideal gas constant Pascal partial pressure Boyles’ law Graham’s law Gay-Lussac combined gas law




Review: • algebraic manipulation of

equations with multiple variables

• moles • units of measuring pressure,

volume, and temperature • converting degrees Celsius to

Kevins

Scuba diving ideas, such as why not to hold your breath as you come up to the surface

Butane lighter lab for PV=nRT (empirically calculating the constant R)

Practice problems and student presentations of practice problems

cooking times are longer at high altitude because lower atmospheric pressure causes lower boiling point of water

Small-scale lab: Reactions of Acids with Carbonates (PH p. 329)


VII Teacher Notes:






STP: ch.4d: the values and meanings of standard temperature and pressure (STP). Objective: Student will

• state or write the meaning of STP • state or write what values of pressure and temperature are at

STP

• when asked state or quick write the values of standard temperature and pressure

STP standard temperature standard pressure atmospheric pressure Pascal




Review: • pressure, temperature • units for measuring pressure

and temperature

Direct instruction of standard temperature and pressure

Use “STP” instead of giving numerical values for temperature and pressure in Charles’ Law problems.


VII Teacher Notes:






Kelvin: ch.4e: how to convert between the Celsius and Kelvin temperature

scales. Ch.4f: there is no temperature lower than 0 Kelvin. Objective: Student will

• convert between Celsius and Kelvin temperature scales • explain why there can be no temperature lower than 0 K

• Given a temperature, be able to convert the unit of measure in Kelvins or degrees Celsius

• Quiz – correctly use the kinetic molecular theory to explain why there is no temperature below absolute zero

Kevin degrees Celsius absolute zero kinetic molecular theory




Review: • Kinetic molecular theory

Direct instruction and practice problems to practice converting between Celsius and Kelvin.

Charles’ Law lab in which students extrapolate volume and temperature data to x-axis, which is absolute zero


VII Teacher Notes:



Assessment Sample:

Multiple Choice Multiple Choice 4. The volume of 400 mL of chlorine gas at 400 mm Hg is decreased to

200 mL at constant temperature. What is the new gas pressure? (std 4c)

A 400 mm Hg B 300 mm Hg C 800 mm Hg D 650 mm Hg

Other

1. If you open a bottle of perfume in the front of the room, you can

soon smell it on the other side of the room. How do the perfume molecules cross the room? (std 4b)

A. Air currents flow from one side of the room to the other. B. Molecules want to move to areas of low concentration. C. Random motion of gas particles lead to an even distribution throughout the room.*

D. The perfume molecules attach to O2 molecules by intermolecular forces and are carried across the room.

2. Standard temperature and pressure (STP) are defined as (std 4d)

A. 0° C and 1.0 atm pressure.* B. 0° C and 273 mm Hg pressure. C. 0 K and 1.0 atm pressure. D. 0 K and 760 mm Hg pressure.

3. What is the equivalent of 423 kelvin in degrees Celsius? (std 4e)

A. -223 °C B. -23 °C C. 150 °C D. 696 °C

1. A sample of gas at 20° C and 1 atm pressure is heated to 120° C at constant volume. What is the pressure at this new temperature?

2. Use kinetic molecular theory to explain how perfume will diffuse

across a room.



Videos:




Unit: Introduction to Solutions Assessment – BMK 2 Time: 7 Hours Concepts: The teacher will integrate the investigation and experimentation topics into the content Focus Standards throughout the year instead of an isolated unit devoting 40% of the instructional time to hands on experiences in the laboratory. State Standard Goal: ch.6: solutions are homogenous mixtures of two or more substances Focus Standards: Ch.6a: the definitions of solute and solvent Ch.6b: how to describe the dissolving process at the molecular level by using the concept of random molecular motion Enduring Understandings:

1. Solutions are homogenous mixtures. 2. Substances that dissolve become solutions


1. Why does soda go flat? 2. What is a solution? 3. How are solutions made?



Key Vocabulary for Solutions aqueous solution mass percent solute unsaturated composition molarity solution concentration moles solvent

dilute ppm standard solution

electrostatic forces saturated super saturated Supporting Vocabulary boiling point dissolve, dissolving ions vaporization boiling point elevation evaporation miscible water hydration colligative properties freezing point depression normal boiling point

dissociate immiscible vapor pressure Academic Vocabulary Important Equations



State Standard Goal: Assessment – BMK 2 Time: ch.6: solutions are homogenous mixtures of two or more substances I Standard: Focus



Solutions: Ch.6a: the definitions of solute and solvent Objective: Student will

• compare and contrast solute & solvent

• T-chart to compare and contrast solute & solvent • The definitions of solvent solute + solutions.

Solubility solution solvent aqueous solution solute dissolving dissolve




Review: • boiling point • melting point • OF, OC + K • pressure • vaporization • evaporation

• Direct instruction, demonstrations and analogies • Growing crystals with CuSO4 + Alum

Textbook: Chemistry – Addision-Wesley, Prentice Hall 2002 (Chemistry – Holt 2004) Chapter 10 – pp. 274-279 Chapter 17 – pp. 482-488 Chapter 18 – pp. 500-526 (Chapter 13 – pp.453-486) www.scilinks.org Solutions HW4118

VII Teacher Notes:




State Standard Goal: Assessment – BMK 2 Time: ch.6: solutions are homogenous mixtures of two or more substances I Standard: Focus



Solutions: Ch.6b: how to describe the dissolving process at the molecular level by using the concept of random molecular motion Objective: Student will

• Describe the dissolving process at the molecular level by using the concept of random molecular motion.

• Quiz and Formative assessments

solubility ions molecular motion solvent solute electrostatic forces




Review – ions

Cartoon videos and drawings to illustrate this concept. For example, Chem ASAP compact disc.

Small-Scale Lab – Electrolytes (PH p. 489)

Textbook: Chemistry – Addision-Wesley, Prentice Hall 2002 (Chemistry – Holt 2004) Chapter 10 – pp. 274-279 Chapter 17 – pp. 482-488 Chapter 18 – pp. 500-526 (Chapter 13 – pp.453-486)

VII Teacher Notes:



Assessment Sample:

Multiple Choice Other

1. Which of the substances in the table above acts as either the

solute or the solvent when mixed with 100 grams of water at 20° C? (std 6a)

A NH3 B C6H5COOH C MgCl2 D CH3CH2OH

2. If the attractive forces among solid particles are less than the attractive forces between the solid and a liquid, the solid will

(std 6b) A probably form a new precipiatate as its crystal lattice is

broken and reformed. B be unaffectede because attractive forces within the crystal

lattice are too strong for the dissolution to occur. C begin the process of melting to form a liquid. D dissolve as particles are pulled away from the crystal lattice

by the liquid molecules.

3. What type of bond do all of the molecules in the table above have in common? (std 2b) A covalent* B ionic C metallic D polar

Table of Common Molecules Name hydrogen ammonia methane Molecular formula

H2 NH3 CH4



Videos:




Unit: Chemical Bonds and Structure Assessment – BMK 3 Time: 15 hours Concepts: The teacher will integrate the investigation and experimentation topics into the content Focus Standards throughout the year instead of an isolated unit devoting 40% of the instructional time to hands on experiences in the laboratory. State Standard Goal: ch.2: Biological, chemical, and physical properties of matter result from the ability of atoms to form bonds from electrostatic

forces between electrons and protons and between atoms and molecules. Focus Standards: ch.2b: Students know chemical bonds between atoms in molecules such as H2 , CH4 , NH3 , H2 CCH2 , N2 , Cl2 , and many

large biological molecules are covalent. Ch.2c: salt crystals, such as NaCl, are repeating patterns of positive and negative ions held together by electrostatic

attraction. Ch.2d: the atoms and molecules in liquids move in a random pattern relative to one another because the intermolecular

forces are too weak to hold the atoms or molecules in a solid form. Ch.2e: how to draw Lewis dot structures ch.1c: use the periodic table to identify trends in ionization energy, electronegativity, and the relative sizes of ions and atoms. Enduring Understandings:

1. Electrons are responsible for the formation of chemical bonds. Essential Questions:

1. How do atoms join together to form molecules? 2. Why is a diamond so hard?



Key Vocabulary for Chemical Bonds and Structure

anion lattice energy

cation octet rule

ion polyatomic ion

isotope Supporting Vocabulary biological molecules electrostatic attraction (electrostatic forces) metallic bond covalent bond intermolecular forces organic crystal lattice ionic bond electro-negativity ionization energy

Academic Vocabulary attraction random maximize relative minimize



State Standard Goal: Assessment – BMK 3 Time: ch.2: Biological, chemical, and physical properties of matter result from the ability of atoms to form bonds from electrostatic forces

between electrons and protons and between atoms and molecules. I Standard: Focus



Covalent bonds: ch.2b: Students know chemical bonds between atoms in molecules such as H2 , CH4 , NH3 , H2 CCH2 , N2 , Cl2 , and many large biological molecules are covalent. Objective: Student will

• Describe covalent bonding in biological and organic molecules

• Identify whether a compound will have

ionic or covalent or covalent bonds

biological molecules organic anion cation ion




Review – Ionic and covalent bonds Definition of organic and

biological molecules


VII Teacher Notes:







Crystal lattice: ch.2c: salt crystals, such as NaCl, are repeating patterns of positive and negative ions held together by electrostatic attraction. Objective: Student will be able to

• Describe a crystal lattice structure in terms of electrostatic forces between the ions

• Draw and label a diagram of ions in a crystal

lattice structure, like salt • Quick write about the electrostatic forces

that hold a salt crystal together

cctet rule crystal lattice anion cation electrostatic attraction electrostatic forces




Review: crystal lattice ions electrostatic forces

Students read labels of food, cosmetics etc and write correct names and formulas of ionic compounds in ingredient (example: calcium chloride)

Students build or observe models of crystal lattices


VII Teacher Notes:







Intermolecular forces: ch.2d: the atoms and molecules in liquids move in a random pattern relative to one another because the intermolecular forces are too weak to hold the atoms or molecules in a solid form. Objective: Student will be able to

• Describe the intermolecular forces in solids, liquids and gases

• Describe the changes in internal energy necessary for phase changes to occur

• Draw and label diagrams of molecules during phase change, like liquid to gas or solid to liquid

• In a quick write, describe what is happening at the molecular level in solids, liquids and gases in terms of intermolecular forces

intermolecular forces




Review: • Particle properties in solids,

liquids and gases • Phase changes: melting,

freezing, evaporation, condensation, boiling

• Lattice structure

Direct instruction to explain this concept, along with a demonstration using BBs in petri dishes shown on the overhead to show and explain the concept of intermolecular forces

Use the concept of surface tension of water and mercury, and an activity involving water’s surface tension to demonstrate hydrogen bonding


VII Teacher Notes:







Lewis dot structures: ch.2e: how to draw Lewis dot structures Objective: Student will be able to

• Draw a Lewis dot structure for an element • Draw a Lewis dot structure of a simple covalent molecule,

such as N2, CO2, H2O • Identify trends families of elements and their Lewis dot

structures • Explain the octet rule as it applies to Lewis dot structures

• Draw Lewis dot structures of elements and simple covalent compounds

• In a quick write, explain the trends of valence electrons and Lewis dot structures in chemical families

• Explain the octet rule in Lewis dot structures of simple molecules and compounds

octet rule




Review: valence electrons covalent and ionic bonding

For Lewis structures, students need direct instruction and many practice problems

Students can draw in Lewis structures on the entire periodic table to see valence electron patterns

Molecule Building activity: students build, draw Lewis dot structures and name common covalent molecules such as ammonia


VII Teacher Notes:






Trends in atoms: ch.1c: use the periodic table to identify trends in ionization energy, electronegativity, and the relative sizes of ions and atoms. Objective: Student will

• Describe the trends in ionization energy, electronegativity and relative size of atoms and ions on a periodic table

• Draw arrows on a periodic table illustrating the trends in ionization energy, electronegativity and relative size of atoms

• Describe the reasons for the differences in relative size of atoms and their ions

• From a list of atoms, identify the largest or smallest atom, smallest radius, or highest ionization energy, using a periodic table

ionization energy electronegativity




Review:

periodic table – periods, families, metals, nonmetals

ions

• On a periodic table, students will draw arrows to show trends in ionization energy, electronegativity and relative size of atoms and ions


VII Teacher Notes:



Assessment Sample:

Multiple Choice Other 1. When cations and anions join, they form what kind of chemical

bond? (std 2c)

A ionic B hydrogen C metallic D covalent

2. Which of the following elements has the same Lewis dot structure as silicon? ( std 2e) A germanium (Ge)* B aluminum (Al) C arsenic (As) D gallium (Ga)

3. Which of the following atoms has the largest atomic radius? (std 1c)

A barium (Ba)* B chlorine (Cl) C iodine (I) D magnesium (Mg)



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Unit: Thermodynamics BMK 3 Time: 10 Days Concepts: The teacher will integrate the investigation and experimentation topics into the content Focus Standards throughout the year instead of an isolated unit devoting 40% of the instructional time to hands on experiences in the laboratory. State Standard Goal:

ch.7: Energy is exchanged or transformed in all chemical reactions and physical changes of matter.

Focus Standards: ch7a: how to describe temperature and heat flow in terms of the motion of molecules (or atoms). ch7b: chemical processes can either release (exothermic) or absorb (endothermic) thermal energy. ch7c: energy is released when a material condenses or freezes and is absorbed when a material evaporates or melts. ch7d: how to solve problems involving heat flow and temperature changes, using known values of specific heat and latent heat of phase change. Enduring Understandings:

1. Energy flows from warmer to cooler. 2. Specific heat is the amount of energy needed to raise the temperature of a substance. 3. Specific heat is an inherent property of substances. 4. Chemical and physical changes involve the flow of energy. 5. Energy flows during chemical and physical changes. 6. Temperature does not change during a phase change.


1. Why do you put ice in a warm drink? 2. What is the relationship between heat and temperature? 3. How does the release or absorption of heat affect substances?



Key Vocabulary for Thermodynamics endothermic latent heat exothermic specific heat heat capacity thermal energy

Supporting Vocabulary Calorie heat flow pressure calorimetry joules temperature enthalphy kilojoules heat phase change

Academic Vocabulary absorb random boundary, boundaries release processes transfer, transferred

Important Formulas Q = m(∆T)Cp q=mC∆T (from textbook)



State Standard Goal: Assessment – BMK 3 Time: ch.7: Energy is exchanged or transformed in all chemical reactions and physical changes of matter. I Standard: Focus



Temperature: ch7a: how to describe temperature and heat flow in terms of the motion of molecules (or atoms). Objective: Student will be able to • describe the differences between temperature and heat flow in

terms of the motion of atoms or molecules

• In a quick write, discuss the differences between temperature and heat flow

• Draw and label what happens to atoms or molecules during heat flow from a cold solution in a warm room

heat flow




Review: • temperature verses heat • kinetic energy

Calorimeter labs using Styrofoam cups and different metal pieces to identify specific heats of various metals.


VII Teacher Notes:






Thermal energy: ch7b: chemical processes can either release (exothermic) or absorb (endothermic) thermal energy. ch7c: energy is released when a material condenses or freezes and is absorbed when a material evaporates or melts. Objective: Student will be able to

• Describe what happens to the energy of a system during a chemical reaction

• Identify the stages of a reaction on a graph, including: reactants, transition states, and products and levels of potential energy

• Identify if a chemical reaction is exothermic or endothermic

• In a quick write, discuss what happens during an endothermic reaction in terms of net heat of the system

• label a graph of a endothermic or exothermic reaction including :reactants, transition states, products and levels of potential energy

exothermic endothermic




Review: • States of matter: solid,

liquid, gas • Exothermic and endothermic

reactions • Potential energy and kinetic

energy

7b Demonstrate athletic trainer’s cold packs (endothermic) and hot packs (exothermic). Ziploc bag reactions also work. 7c Show students models of molecular motion of gases, solids and liquids to show energy loss or gain during phase change.


VII Teacher Notes:






Heat Problems: ch7d: how to solve problems involving heat flow and temperature changes, using known values of specific heat and latent heat of phase change. Objective: Student will be able to

• Solve problems involving heat flow and temperature change

• Solve problems involving temperature change, specific heat

specific heat heat capacity joules kilojoules enthalpy




Review: • phase changes • heat and temperature • solving algebraic problems

Direct instruction contrasting touching metal or water exposed to same amount of heat

Practice problems using q= s x m x T(change) and q=m x H(fusion or vaporization)

Interpret and read graph of phase change Phase Change Lab: observe and graph no temperature

change as ice melts and water evaporates


VII Teacher Notes: Note - the formula on the CST formula sheet is different from the textbook Q = m(∆T)Cp (from CST) q=mC∆T (from textbook)



Assessment Sample:

Multiple Choice Other 1 When is the random molecular motion of a substance greatest? (std 7a) A in condensed form B in a liquid C when frozen D in a gas * 2. Which of these is an example of an exothermic chemical process? (std 7b) A evaporation of water B melting ice C photosynthesis of glucose D combustion of gasoline* 3. The specific heat of copper is about 0.4 joules/gram °C. How much heat is needed to change the temperature of a 30-gram sample of copper from 20.0 °C to 60.0 °C ? (std 7d) A 1000 J B 720 J C 480 J* D 240 J

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Unit: Solutions BMK 4 Time: 4 Days Concepts: The teacher will integrate the investigation and experimentation topics into the content Focus Standards throughout the year instead of an isolated unit devoting 40% of the instructional time to hands on experiences in the laboratory.

State Standard Goal: ch 6: Solutions are homogeneous mixtures of two or more substances.

Focus Standards:

ch 6c: temperature, pressure, and surface area affect the dissolving process.

ch 6d: how to calculate the concentration of a solute in terms of grams per liter, molarity, parts per million, and percent composition.

Enduring Understandings: 1. The dissolving of a solid into a solution will be affected by temperature, pressure and surface area.

Essential Questions: 1. Which will dissolve faster in a cup of coffee, a spoon of sugar or a sugar cube?



Key Vocabulary for Solutions aqueous solution molarity standard solution concentration moles supersaturated dilute parts per million (ppm) unsaturated solution mass percent saturated solution

Supporting Vocabulary dissolving surface area Academic Vocabulary composition generalization



State Standard Goal: Assessment – BMK 4 Time: ch 6: Solutions are homogeneous mixtures of two or more substances I Standard: Focus



Solutions: ch6c: temperature, pressure, and surface area affect the dissolving process. Objective: Student will

• Describe how temperature, pressure and surface area can affect the dissolving process

• In a quick write, discuss how temperature, pressure and surface area can affect dissolving of a solute

• Draw and label a diagram of different size particles dissolving in a solution

• Describe how changes in temperature or pressure can affect the dissolving process

surface area dissolving aqueous solution saturated solution super saturated solution unsaturated solution




Review: • solvent, solute, solution • temperature and heat • pressure • dissolving verses melting

Demonstration of cube of sugar vs. powdered sugar dissolving, gas pressure in a can of soda

Example: the Bends in SCUBA diving Crystal growing activity shows supersaturated

solutions


VII Teacher Notes:



State Standard Goal: Assessment – BMK 4 Time: ch 6: Solutions are homogeneous mixtures of two or more substances. I Standard: Focus



Concentration: ch6d: how to calculate the concentration of a solute in terms of grams per liter, molarity, parts per million, and percent composition. Objective: Student will

• correctly calculate molarity and percent composition

• Calculate the concentration of solute in terms of gram per liter, molarity, parts per millions and percent composition.

dilute concentration unsaturated mass percent dissolving dissolute standard solution saturated solution solution molarity moles ppm




Have students make solutions of copper sulfate or other colored solutions, assigning various different molarities so that the color of the solutions shows the different concentrations.

Homework assignment where students determine the molarity of a drink from home, such as a soft drink or Vitamin Water.


VII Teacher Notes:



Assessment Sample:

Multiple Choice Other 1. How many moles of HNO3 are needed to prepare 5.0 liters of a 2.0 M solution of HNO3? (std 6d) A 2.5 moles B 5 moles C 10 moles* D 20 moles 2. If the solubility of NaCl at 25C is 36.2 g/100g H2O, what mass

of NaCl can be dissolved in 50.0 g of H2O? (std 6d)

A 18.1 g* B 36.2 g C 72.4 g D 86.2 g

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Unit: Equilibrium BMK 4 Time: 10 Days

Concepts: The teacher will integrate the investigation and experimentation topics into the content Focus Standards throughout the year instead of an isolated unit devoting 40% of the instructional time to hands on experiences in the laboratory.

State Standard Goal: ch 9: Chemical equilibrium is a dynamic process at the molecular level.

Focus Standards: ch 9a: how to use Le Chatelier's principle to predict the effect of changes in concentration, temperature, and pressure. ch 9b: equilibrium is established when forward and reverse reaction rates are equal.

Enduring Understandings: 1. During chemical equilibrium forward and reverse reactions are occurring at the same rate. Essential Questions: 1. How do conditions “stress” chemical equilibrium?



Key Vocabulary for Equilibrium entropy inhibitor equilibrium law of disorder equilibrium constant Le Chatelier's Principle first order reaction reversible reaction

Supporting Vocabulary pressure concentration reaction rates Academic Vocabulary dynamic spontaneous non-spontaneous stress proportional



State Standard Goal: Assessment – BMK 4 Time: ch 9: Chemical equilibrium is a dynamic process at the molecular level. I Standard: Focus



Reaction rates: ch 9b: equilibrium is established when forward and reverse reaction rates

are equal. Objective: Student will be able

• Discuss the nature of an equilibrium reaction at a molecular level

• Draw and label a diagram of a solutions in equilibrium

• Write equilibrium equations

reaction rates equilibrium




Review: • solution, solvent, solute • balanced equations

Use the water cycle as an example of a system that has established equilibrium.

Students interpret graphs to identify when equilibrium has been established

Textbook: Chemistry – Addision-Wesley, Prentice Hall 2002 (Chemistry – Holt 2004) p.543, 545 (p. 513, 515)

VII Teacher Notes:



State Standard Goal: Assessment – BMK 4 Time: ch 9: Chemical equilibrium is a dynamic process at the molecular level. I Standard: Focus



LeChatelier’s principle: ch 9a: how to use Le Chatelier’s principle to predict the effect of changes in concentration, temperature, and pressure. Objective: Student will be able to

• Predict the effect of changes in concentration, temperature or pressure on a chemical reaction in equilibrium

• Draw and label diagrams of a solutions in equilibrium and with changes in pressure, temperature or pressure

• Write equilibrium equations and predict what will happen if there is a change in concentration, temperature or pressure

concentratiuon equilibrium Le Chatelier’s principle




Review: • reactant • solution, solute, solvent • pressure • temperature

Direct instruction of Le Chatelier’s principle and practice problems where students predict the shift of a reaction

Demonstration of changes in temperature and pressure in a can of soda

Visual demonstrations of equilibrium: p.543, p.545 Prentice Hall 2002 Edition and Holt 2007 Edition p. 513 and 515)

Equilibrium of iron thiocyanate (caution: cyanide gas produced if in contact with acid)

Textbook: Chemistry – Addision-Wesley, Prentice Hall 2002 (Chemistry – Holt 2004) p.543, 545 (p. 513, 515)

VII Teacher Notes:



Assessment Sample:

Multiple Choice Other 1. When a reaction is at equilibrium and more reactant is added,

which of the following changes is the immediate result? (std 9a)

A The reverse reaction rate remains the same. B The forward reaction rate increases.* C The reverse reaction rate decreases. D The forward reaction rate remains the same.

2. 4HCl(g) + O2(g) 2 H2O(l) + 2Cl2(g) + 113 kJ ⇔

Which action will drive the reaction to the right? (std 9a)

A heating the equilibrium mixture B adding water to the system C decreasing the oxygen concentration D increasing the system’s pressure*

3. In a sealed bottle that is half full of water, equilibrium will be

attained when water molecules (std 9b) A cease to evaporate. B begin to condense. C are equal in number for both the liquid and the gas phase. D evaporate and condense at equal rates.*



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Unit: Reaction Rate BMK 4 Suggested Time: 11 Days

Concepts: The teacher will integrate the investigation and experimentation topics into the content Focus Standards throughout the year instead of an isolate unit devoting 40% of the instructional time to hands on experiences in the laboratory.

State Standard Goal: ch 8: Chemical reaction rates depend on factors that influence the frequency of collision of reactant molecules.

Focus Standards: ch 8a: the rate of reaction is the decrease in concentration of reactants or the increase in concentration of products with time. ch 8b: how reaction rates depend on such factors as concentration, temperature, and pressure. ch 8c: the role a catalyst plays in increasing the reaction rate.

Enduring Understandings: 1. Concentration, temperature and pressure can affect reaction rates.


1. What is the function of a catalyst? 2. What is rate?



Key Vocabulary for Reaction Rate

reaction rate activation energy

catalyst reversible reaction

free energy Supporting Vocabulary

collision equilibrium

concentration magnetic field

enzyme Academic Vocabulary frequency influence



State Standard Goal: Assessment – BMK 4 Time: ch 8: Chemical reaction rates depend on factors that influence the frequency of collision of reactant molecules. I Standard: Focus



Reaction rate: ch 8a: the rate of reaction is the decrease in concentration of reactants or the increase in concentration of products with time. ch 8b: how reaction rates depend on such factors as concentration, temperature, and pressure. Objective: Student will be able to

• Define reaction rate in terms of what happens to the concentration reactants and products during a chemical reaction

• Describe how concentration of reactants, temperature or pressure affect reaction rates in chemical reactions

• In a quick write, discuss how changes in temperature or pressure can affect the reaction rate of a chemical reaction.

reaction rate




Review: • energy • Balanced equations,

reactants and products • concentration • pressure • temperature3 • reaction

Students graph concentration vs. time for iodine clock reaction or sulfur clock reaction


VII Teacher Notes: rate = change/time



State Standard Goal: Assessment – BMK 4 Time: ch 8: Chemical reaction rates depend on factors that influence the frequency of collision of reactant molecules. I Standard: Focus



Catalyst: ch 8c: the role a catalyst plays in increasing the reaction rate. Objective: Student will be able to

• Describe how a catalyst can affect reaction rate • Describe different types of catalysts • Recognize the affect of a catalyst on activation energy on a

reaction curve

• In a quick write, describe how a catalyst can affect a reaction

• Interpret a reaction curve showing the lowering of activation energy by a catalyst

catalyst enzyme




Review: • Examples of how enzymes act

as catalysts in biological systems

Lab: splint tests for oxygen using manganese dioxide as a catalyst for the decomposition of hydrogen peroxide

Show graph of activation energy Example: Enzymes are catalysts Demonstration of the decomposition of hydrogen

peroxide using catalysts of sodium potassium tartarate and cobalt chloride as a visual example of how catalysts are not consume in a reaction

Inquiry Lab: using baking soda and vinegar, how do you get balloons to fill up the most quickly?


VII Teacher Notes:



Assessment Sample:


3. A catalyst can speed up the rate of a give chemical reaction by (std 8c)

A increasing the equilibrium constant in favor of products. B lowering the activation energy required for the reaction to occur.* C raising the temperature at which the reaction occurs. D increasing the pressure of reactants, thus favoring products.

H2 + Cl2 2HCl

2. Which of these describes the rate of this chemical reaction? (std 8a) A an increase in the concentration of HCl and H2 with time B an increase in the concentration of HCl with time C an increase in H2 and Cl2 with time D a decrease in HCl and Cl2 with time catalyst C6H6 + Br2 C6H5Br + HBr 3. Which of the following changes will cause an increase in the rate of the

above reaction? (std 8b) A increasing the concentration of Br2* B decreasing the concentration of C6H6 C increasing the concentration of HBr D decreasing the temperature

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Unit: Acids and Bases BMK 4 Suggested Time: 14 Days Concepts: The teacher will integrate the investigation and experimentation topics into the content Focus Standards throughout the year instead of an isolated unit devoting 40% of the instructional time to hands on experiences in the laboratory. State Standard Goal: ch 5: Acids, bases, and salts are three classes of compounds that form ions in water solutions.

. Focus Standards: 5a: the observable properties of acids, bases, and salt solution 5b: acids are hydrogen-ion-donating and bases are hydrogen-ion-accepting substances. 5c: strong acids and bases fully dissociate and weak acids and bases partially dissociate. 5d: how to use the pH scale to characterize acid and base solutions. Enduring Understandings: 1. Acids donate hydrogen ions and bases accept hydrogen ions Essential Questions: 1. Why is lemon juice an acid?



Key Vocabulary for Acids and Bases

acid pH scale amphoteric substance

base weak acid acidic solution

salt strong acid basic solution

hydrogen ion dissociate (dissociated, dissociates) neutral solution Supporting Vocabulary BrØnsted-Lowry model conductivity properties Academic Vocabulary observable



State Standard Goal: Assessment – BMK 4 Time: ch 5: Acids, bases, and salts are three classes of compounds that form ions in water solutions.

I Standard: Focus • Student Objectives

II Evidence of Student Learning Performance Indicators

III Vocabulary

Properties of acids and bases: 5a: the observable properties of acids, bases, and salt solution 5d: how to use the pH scale to characterize acid and base solutions. Objective: Student will be able to

• Describe the properties of acids and bases and salts • Differentiate acids and bases from salts • Determine the pH of a solution on the pH scale • Recognize acids and bases using indicators and the pH

scale

• Complete a graphic organizer, such as a Venn diagram, about the properties of acids, bases and salts

• In a quick write, describe the pH scale and the values for acids and bases

• Draw and label a pH scale and correctly place various acids and bases on the scale

• Detemine if solutions are acids, bases or salts, depending on their pH

acid base salt weak acid strong acid pH scale neutral solution properties




Review: • pH scale • examples of acids and bases

Assignment: use pH paper to take the pH of common substances at home or at school

Graphic organizer/table for students to fill in about characteristics of acids, bases and salts (taste, feel, conductivity, pH, etc.), for example Venn diagram

Activity where students experiment with a wide variety of indicators with various common substances

Place common substances on the pH scale based on their approximate pH, or state whether a substance is an acid or a base based on a given pH


VII Teacher Notes:



State Standard Goal: Assessment – BMK 4 Time: ch 5: Acids, bases, and salts are three classes of compounds that form ions in water solutions

I Standard: Focus • Student Objectives

II Evidence of Student Learning Performance Indicators

III Vocabulary

Acids and bases: 5b: acids are hydrogen-ion-donating and bases are hydrogen-ion-accepting substances. 5c: strong acids and bases fully dissociate and weak acids and bases partially dissociate. Objective: Student will be able to • Describe acids and bases by their ability for hydrogen-ion

donation or acceptance • Predict products in reactions between acids and bases • Discuss the dissociation of acids and bases and their properties,

including strength of solution, electrical conductivity

• In quick write, discuss the relationship of the dissociation of acids and bases and their properties, such as electrical conductivity

• Discuss the strength of an acid or base according to their dissociation ability

hydrogen ion dissociate conductivity BrØnsted-Lowry model

Assessment Sample:




Review: • atoms, ions

Titration lab Direct instruction, “mapping” the H+ ions through

color coding the H+ ions to track H+ ions during chemical reactions

Predict products of reactions between acids and bases Use conductivity meter or light bulbs to gain an

understanding of full dissociation vs. partial dissociation in common household substances (can be done as a demo or a lab)


VII Teacher Notes:



Multiple Choice Other 1. Equal volumes of 1 molar hydrochloric acid (HCl) and 1 molar

sodium hydroxide base (NaOH) are mixed. After mixing, the solution will be (std 5a)

A strongly acidic. B weakly acidic. C nearly neutral.* D weakly basic.

2. Why is potassium hydroxide (KOH) a strong base? (std 5c) A It easily releases hydroxide ions. B It does not dissolve in water. C It reacts to form salt crystals in water. D It does not conduct an electric current. 3. There are four different laboratory solutions. What would be the

pH of the solution with the highest acidity? (std 5d) A 11 B 7 C 5 D 3

Videos:




Unit : Safety & Lab Procedures Time: 1 week Concepts:

• Safety practices • Safety symbols • Safety equipment • Legal obligations • Environmental health

State Standard Goal: Focus Standards: Enduring Understandings:


Key Vocabulary Safety contract Safety symbols



Safety and Lab Procedures

Standard Assessed

Objectives – Mastery Suggested Teaching Activities, Labs, and Strategies Resources: Textbook (pages and ancillary items)

Safety rules

Safety symbols Safety equipment Safety practices a. heat b. electricity c. chemicals d. animals e. plant f. glassware

g. student behavior

Conservation, recycling, and disposal of chemicals and various lab materials.

Safety contract

Have students make posters concerning different aspects of classroom and laboratory safety. Address students clothing, shoes, hair, and jewelry safety in the laboratory. Students identify and explain symbols. Demonstrate use of eye wash, fire blanket, goggles, and all equipment available to ensure safety of all students. Demonstrate and explain all safety expectations. Discuss procedures concerning all lab materials. Explain contract. Students and parents are to sign before students are allowed to participate in laboratory experiences



Assessment Sample:


1. When a solution is heated in an open test tube, the test tube should always be pointed –

A. so bubbles are visible B. at a 1800 angle from the flame C. toward a ventilated area D. away from nearby people

2. A science class is conducting an experiment that produces noxious fumes. Because of inadequate ventilation, some students begin to feel nauseated and dizzy. The first response should be to

A. neutralize the solution that is reacting to produce the

noxious fumes B. carry the reactants outside, away from other students C. leave the room and go to an area with fresh air D. spray the reaction with a fire extinguisher



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High School Chemistry Science Curriculum Assessment Instructional Guide

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Transcript of High School Chemistry Science Curriculum Assessment Instructional Guide