Physical Setting/Chemistry Core Curriculum

Physical Setting/

Chemistry

Core Curriculum

THE UNIVERSITY OF THE STATE OF NEW YORK THE STATE EDUCATION DEPARTMENT

http://www.nysed.gov

THE UNIVERSITY OF THE STATE OF NEW YORK

Regents of The University

CARL T. HAYDEN, Chancellor, A.B., J.D. ............................................................................Elmira

DELAIDE L. SANFORD, Vice Chancellor, B.A., M.A., P.D. .................................................Hollis

IANE O’NEILL MCGIVERN, B.S.N., M.A., Ph.D. ...............................................................Staten Island

AUL B. COHEN, B.A., M.A., Ph.D......................................................................................New Rochelle

AMES C. DAWSON, A.A., B.A., M.S., Ph.D. .......................................................................Peru

OBERT M. BENNETT

, B.A., M.S. ........................................................................................Tonawanda

OBERT M. JOHNSON, B.S., J.D. .........................................................................................Huntington

NTHONY S. BOTTAR, B.A., J.D. .........................................................................................North

Syracuse

ERRYL H. T

ISCH, B.A., M.A. ............................................................................................New York

NA L. FARLEY, B.A., M.A., Ph.D. .....................................................................................Brockport

ERALDINE D. CHAPEY, B.A., M.A., Ed.D...........................................................................Belle Harbor

RNOLD B. GARDNER

, B.A., LL.B........................................................................................Buffalo

HARLOTTE K. FRANK, B.B.A., M.S.Ed., Ph.D. ..................................................................New York

ARRY PHILLIPS, 3

, B.A., M.S.F.S. ...................................................................................Hartsdale

OSEPH E. BOWMAN, JR., B.A., M.L.S., M.A., M.Ed., Ed.D ...............................................Albany

ORRAINE A. CORTÉS-VÁZQUEZ, B.A., M.P.A......................................................................Bronx

President of The University and Commissioner of Education

ICHARD

P. MILLS

Chief Operating Officer

ICHARD H. CATE

Deputy Commissioner for Elementary, Middle, Secondary, and Continuing Education

AMES A. KADAMUS

Assistant Commissioner for Curriculum, Instruction, and Assessment

OSEANNE DEFABIO

Assistant Director for Curriculum and Instruction

NNE SCHIANO

The State Education Department does not discriminate on the basis of age, color, religion, creed, dis-

ability, marital status, veteran status, national origin, race, gender, genetic predisposition or carrier sta-

tus, or sexual orientation in its educational programs, services, and activities. Portions of this publica-

tion can be made available in a variety of formats, including braille, large print or audio tape, upon

request. Inquiries concerning this policy of nondiscrimination should be directed to the Department’s

Office for Diversity, Ethics, and Access, Room 152, Education Building, Albany, NY 12234.

CONTENTS

Acknowledgments . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .iv

Core Curriculum . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .1

Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .3

Process Skills Based on Standards 1, 2, 6, and 7 . . .5

Process Skills Based on Standard 4 . . . . . . . . . . . .12

Standard 4:

The Physical Setting

. . . . . . . . . . . . . . . . . . . . .16

Appendix A:

Chemistry Core Topics . . . . . . . . . . . . . . . . . . . . . . . .26

Appendix B:

Chemistry Content Connections Table . . . . . . . . . . . .34

Chemistry iii

ACKNOWLEDGMENTS

The State Education Department acknowledges the assistance of teachers and school administrators from across

New York State and the New York State Chemistry Mentors. In particular, the State Education Department would

like to thank:

Robert Dayton Rush-Henrietta High School, Henrietta

Mary Dery Dutchess BOCES, Poughkeepsie

David Hanson SUNY at Stony Brook, Stony Brook

Linda Hobart Finger Lakes Community College, Canandaigua

Silvana Jaus Edgemont High School, Scarsdale

Carol Jemmott Bishop Loughlin Memorial High School, Brooklyn

Elaine Jetty Ravena-Coeymans-Selkirk Senior High School

Patrick Kavanah (retired) Monroe Woodbury Senior High School, Central Valley

David Kiefer Midwood High School, Brooklyn

Elise Hilf Levine Scarsdale High School, Scarsdale

Joan Laredo-Liddell St. Barnabas High School, Bronx

June Kasuga Miller Queens College, Flushing

Theresa Newkirk Saratoga Springs Sr. High School, Saratoga Springs

Linda Padwa Port Jefferson High School, Port Jefferson

Cynthia Partee Division High School, Levittown

Diane Pillersdorf Richmond Hill High School, Richmond Hill

Lee Roberts Wellsville High School, Wellsville

Lance W. Rudiger Potsdam Senior High School, Potsdam

David L. Shelc Portville Jr./Sr. High School, Portville

Thomas Shiland Saratoga Springs Senior High School, Saratoga Springs

Virginia M. Trombley AuSable Valley High School, Clintonville

Alice Veyvoda Half Hollow Hills High School West, Dix Hills

Beatrice G. Werden Bronx High School of Science, Bronx

Harvey Weiner John F. Kennedy High School, Bellmore

The project manager for the development of the Chemistry Core Curriculum was Diana Kefalas Harding, Associate in

Science Education, with content and assessment support provided by Sharon Miller, Associate in Educational

Testing, and Elise Russo, Associate in Science Education. Special thanks go to Jan Christman for technical expertise.

Chemistry iv

Physical Setting/

Chemistry

Core Curriculum

Chemistry 2

INTRODUCTION

The Physical Setting/Chemistry Core Curriculum has been

written to assist teachers and supervisors as they pre-

pare curriculum, instruction, and assessment for the

chemistry content and process skills in the New York

State Learning Standards for Mathematics, Science, and

Technology. This core curriculum is an elaboration of the

science content of that document and its key ideas and

performance indicators. Key ideas are broad, unifying,

general statements of what students need to know. The

performance indicators for each key idea are statements

of what students should be able to do to provide evi-

dence that they understand the key idea.

The Chemistry Core Curriculum presents major under-

standings that give more specific detail to the concepts

underlying the performance indicators in Standard 4.

In addition, portions of Standards 1, 2, 6, and 7 have

been elaborated to highlight skills necessary to allow

students to evaluate proposed explanations of natural

phenomena. The concepts and skills identified in the

introductions and the major understandings of each

key idea in the core curriculum will provide the mater-

ial from which Regents examination items will be

developed. Occasionally, examples are given in an

effort to clarify information. These examples are not

inclusive lists. Therefore, teachers should not feel lim-

ited by them.

This core is not a syllabus. This is a core for the prepara-

tion of high school curriculum, instruction, and assess-

ment. The lack of detail in this core is not to be seen as

a shortcoming. Rather, the focus on conceptual under-

standing in the core is consistent with the approaches

recommended in the National Science Education Standard

(National Research Council) and Benchmarks for Science

Literacy (American Association for the Advancement of

Science). The local courses designed using this core cur-

riculum are expected to prepare students to explain

both accurately and with appropriate depth concepts

and models relating to chemistry. The core addresses

only the content and skills to be assessed at the com-

mencement level by the Physical Setting/Chemistry

Regents examination. The core curriculum has been

prepared with the assumption that the content, skills,

and vocabulary as outlined in the Learning Standards for

Mathematics, Science, and Technology at the elementary

and intermediate levels have been taught previously.

Work in grades 9-12 must build on the knowledge,

understanding, and ability to do science that students

have acquired in their earlier grades.

It is essential that instruction focus on the understand-

ing of concepts, relationships, processes, mechanisms,

models, and applications. Less important is the memo-

rization of specialized terminology and technical

details. In attaining scientific literacy, students will be

able to demonstrate these understandings, generate

explanations, exhibit creative problem solving and rea-

soning, and make informed decisions. Future assess-

ments will test students’ ability to explain, analyze, and

interpret chemical processes and phenomena, and use

models and scientific inquiry. The major understand-

ings in this guide will also allow teachers more flexibil-

ity, making possible richer creativity in instruction and

greater variation in assessment. The general nature of

the major understandings in this core will encourage

the teaching of science for understanding, rather than

for memorization.

The order of presentation and numbering of all state-

ments in this guide are not meant to indicate any rec-

ommended sequence of instruction. Ideas have not

been prioritized, nor have they been organized to indi-

cate teaching time allotments or test weighting. Many

of the major understandings in this document are

stated in a general rather than specific manner. It is

expected that teachers will provide examples and

applications in their teaching/learning strategies to

bring about understanding of the major concepts

involved. Teachers are encouraged to help students

find and elaborate conceptual cross-linkages that inter-

connect many of the chemistry key ideas to each other,

and to other mathematics, science, and technology

learning standards.

Historical Content

The study of chemistry is rich in historical develop-

ment. The learning standards encourage the inclusion

not only of important concepts but also of the scientists

who were responsible for discovering them. Robert

Boyle, generally regarded as one of the fathers of mod-

ern chemistry, introduced systematic experimental

methods into the study of chemistry. John Dalton laid

down the tenets of the atomic theory at the beginning

of the 19th century. By mid-century Mendeleev had

completed most of his work organizing the Periodic

Chemistry 3

Table, and Amedeo Avogadro had provided keen

insights into the relationships of gaseous molecules.

Ernest Rutherford discovered the nucleus, and soon

afterward Henry Moseley identified the atomic number

as the identifying factor of the elements. Soon after,

Albert Einstein proposed the insight into the interrela-

tionship of matter and energy. Marie Curie worked

with radioactive substances showing natural transmu-

tations. Linus Pauling provided insights into the nature

of the chemical bond in the 1930s, and introduced elec-

tronegativity values, an important tool in understand-

ing bonding.

To successfully teach chemistry, teachers can inter-

weave both the concepts and the scientists who were

responsible for discovering them. Chemistry will be far

more interesting when the human element can be

incorporated into the lessons.

Scientific Thinking and a Scientific Method

Modern science began around the late 16th century with

a new way of thinking about the world. Few scientists

will disagree with Carl Sagan’s assertion that “science is a

way of thinking much more than it is a body of knowl-

edge” (Broca’s Brain, 1979). Thus science is a process of

inquiry and investigation. It is a way of thinking and act-

ing, not just a body of knowledge to be acquired by

memorizing facts and principles. This way of thinking,

the scientific method, is based on the idea that scientists

begin their investigations with observations. From these

observations they develop a hypothesis, which is

extended in the form of a predication, and challenge the

hypothesis through experimentation and thus further

observations. Science has progressed in its understanding

of nature through careful observation, a lively imagina-

tion, and increasing sophisticated instrumentation.

Science is distinguished from other fields of study in that

it provides guidelines or methods for conducting

research, and the research findings must be reproducible

by other scientists for those findings to be valid.

It is important to recognize that scientific practice is not

always this systematic. Discoveries have been made

that are serendipitous and others have not started with

the observation of data. Einstein’s theory of relativity

started not with the observation of data but with a kind

of intellectual puzzle.

Laboratory Requirements

Critical to understanding science concepts is the use of

scientific inquiry to develop explanations of natural

phenomena. Therefore, as a prerequisite for admission

to the Physical Setting/Chemistry Regents

Examination, students must have successfully com-

pleted 1200 minutes of laboratory experience with

satisfactory reports on file. Because of the strong

emphasis on student development of laboratory skills,

a minimum of 280 minutes per week of class and

laboratory time is recommended.

Prior to the written portion of the Regents examination,

students will be required to complete a laboratory per-

formance test during which concepts and skills from

Standards 1, 2, 4, 6, and 7 will be assessed.

The Laboratory Setting

Laboratory safety dictates that a minimum amount of

space be provided for each individual student.

According to the National Science Teachers

Association, recommended space considerations

include:

• A minimum of 60 ft

/pupil (5.6m

) which is

equivalent to 1440 ft

(134m

) to accommodate a

class of 24 safely in a combination

laboratory/classroom.

Or,

• A minimum of 45 ft

/pupil (4.2m

) which is

equivalent to 1080 ft

(101m

) to accommodate a

class of 24 safely in a stand-alone laboratory.

It is recommended that each school district comply

with local, State, and federal codes and regulations

regarding facilities and fire and safety issues.

Systems of Units

International System (SI) units are used in this core cur-

riculum. SI units that are required for the chemistry

core are listed in the Reference Tables. SI units are a log-

ical extension of the metric system. The SI system

begins with seven basic units, with all other units being

derived from them (see Reference Tables). While some

of the basic and derived units of the SI system are com-

monly used in chemistry (mole, kelvin, kilogram,

meter, joule, volt), there are other units that are used in

chemistry that are exceptions. Thus, in addition to the

SI units, you will find liters used in volume measure-

ments, atmospheres and torr used as pressure units,

and Celsius as a temperature indicator.

Uncertainty of Measurements and Significant Figures

It is an important concept in chemistry that all mea-

surements contain some uncertainty. Such data is

reported in significant figures to inform the reader of

the uncertainty of the measurement. When these values

are used in calculations, it is vital that the answers to

such calculations are not misleading, and hence, rules

for addition, subtraction, multiplication, and division

should be followed.

Chemistry 4

PROCESS SKILLS

BASED ON STANDARDS 1, 2, 6, AND 7

Science process skills should be based on a series of discoveries. Students learn most effectively when they have a central role

in the discovery process. To that end, Standards 1, 2, 6, and 7 incorporate in the Chemistry Core Curriculum a student-

centered, problem-solving approach to chemistry. This list is not intended to be an all-inclusive list of the content or skills

that teachers are expected to incorporate into their curriculum. It should be a goal of the instructor to encourage science

process skills that will provide students with background and curiosity to investigate important issues in the world around

them.

Note: The use of e.g. denotes examples which may be used for in-depth study. The terms for example and such as denote

material which is testable. Items in parentheses denote further definition of the word(s) preceding the item and are testable.

STANDARD 1—Analysis, Inquiry, and Design

Students will use mathematical analysis, scientific inquiry, and engineering design, as appropriate, to pose

questions, seek answers, and develop solutions.

STANDARD 1

Key Idea 1:

Analysis, Inquiry,

Abstraction and symbolic representation are used to communicate mathematically.

and Design

M1.1 Use algebraic and geometric representations to describe and compare data.

• organize, graph, and analyze data gathered from laboratory activities or other

MATHEMATICAL

sources

ANALYSIS:

 identify independent and dependent variables

 create appropriate axes with labels and scale

 identify graph points clearly

• measure and record experimental data and use data in calculations

 choose appropriate measurement scales and use units in recording

 show mathematical work, stating formula and steps for solution

 estimate answers

 use appropriate equations and significant digits

 show uncertainty in measurement by the use of significant figures

 identify relationships within variables from data tables

 calculate percent error

• recognize and convert various scales of measurement

 temperature

§ Celsius (°C)

§ Kelvin (K)

 length

§ kilometers (km)

§ meters (m)

§ centimeters (cm)

§ millimeters (mm)

 mass

§ grams (g)

§ kilograms (kg)

 pressure

§ kilopascal (kPa)

§ atmosphere (atm)

• use knowledge of geometric arrangements to predict particle properties or

behavior

Chemistry 5

STANDARD 1

Analysis, Inquiry,

and Design

MATHEMATICAL

ANALYSIS:

continued

Key Idea 2:

Deductive and inductive reasoning are used to reach mathematical conclusions.

M2.1 Use deductive reasoning to construct and evaluate conjectures and arguments, rec-

ognizing that patterns and relationships in mathematics assist them in arriving at

these conjectures and arguments.

• interpret a graph constructed from experimentally obtained data

 identify relationships

§ direct

§ inverse

 apply data showing trends to predict information

Key Idea 3:

Critical thinking skills are used in the solution of mathematical problems.

M3.1 Apply algebraic and geometric concepts and skills to the solution of problems.

• state assumptions which apply to the use of a particular mathematical equation

and evaluate these assumptions to see if they have been met

• evaluate the appropriateness of an answer, based on given data

STANDARD 1

Analysis, Inquiry,

and Design

SCIENTIFIC INQUIRY:

Key Idea 1:

The central purpose of scientific inquiry is to develop explanations of natural phenomena

in a continuing, creative process.

S1.1 Elaborate on basic scientific and personal explanations of natural phenomena, and

develop extended visual models and mathematical formulations to represent

thinking.

• use theories and/or models to represent and explain observations

• use theories and/or principles to make predictions about natural phenomena

• develop models to explain observations

S1.2 Hone ideas through reasoning, library research, and discussion with others,

including experts.

• locate data from published sources to support/defend/explain patterns

observed in natural phenomena

S1.3 Work towards reconciling competing explanations, clarifying points of agreement

and disagreement.

• evaluate the merits of various scientific theories and indicate why one theory

was accepted over another

Key Idea 2:

Beyond the use of reasoning and consensus, scientific inquiry involves the testing of pro-

posed explanations involving the use of conventional techniques and procedures and usu-

ally requiring considerable ingenuity.

S2.1 Devise ways of making observations to test proposed explanations.

• design and/or carry out experiments, using scientific methodology to test pro-

posed calculations

S2.2 Refine research ideas through library investigations, including information

retrieval and reviews of the literature, and through peer feedback obtained from

review and discussion.

• use library investigations, retrieved information, and literature reviews to

improve the experimental design of an experiment

Chemistry 6

STANDARD 1

Analysis, Inquiry,

and Design

SCIENTIFIC INQUIRY:

continued

S2.3 Develop and present proposals including formal hypotheses to test explanations,

i.e.; they predict what should be observed under specific conditions if their expla-

nation is true.

• develop research proposals in the form of “if X is true and a particular test Y is

done, then prediction Z will occur”

S2.4 Carry out a research plan for testing explanations, including selecting and devel-

oping techniques, acquiring and building apparatus, and recording observations

as necessary.

• determine safety procedures to accompany a research plan

Key Idea 3:

The observations made while testing proposed explanations, when analyzed using conven-

tional and invented methods, provide new insights into phenomena.

S3.1 Use various means of representing and organizing observations (e.g., diagrams,

tables, charts, graphs, equations, and matrices) and insightfully interpret the

organized data.

• organize observations in a data table, analyze the data for trends or patterns,

and interpret the trends or patterns, using scientific concepts

S3.2 Apply statistical analysis techniques when appropriate to test if chance alone

explains the result.

S3.3 Assess correspondence between the predicted result contained in the hypothesis

and the actual result, and reach a conclusion as to whether or not the explanation

on which the prediction is supported.

• evaluate experimental methodology for inherent sources of error and analyze

the possible effect on the result

• compare the experimental result to the expected result; calculate the percent

error as appropriate

S3.4 Using results of the test and through public discussion, revise the explanation and

contemplate additional research.

S3.5 Develop a written report for public scrutiny that describes the proposed explana-

tion, including a literature review, the research carried out, its results, and sugges-

tions for further research.

STANDARD 1

Analysis, Inquiry,

and Design:

ENGINEERING

DESIGN

Key Idea 1:

Engineering design is an iterative process involving modeling and optimization (finding

the best solution within given constraints); this process is used to develop technological

solutions to problems within given constraints.

If students are asked to do a design project, then:

• Initiate and carry out a thorough investigation of an unfamiliar situation and

identify needs and opportunities for technological invention or innovation.

• Identify, locate, and use a wide range of information resources, and document

through notes and sketches how findings relate to the problem.

• Generate creative solutions, break ideas into significant functional elements,

and explore possible refinements; predict possible outcomes, using mathemati-

cal and functional modeling techniques; choose the optimal solution to the

problem, clearly documenting ideas against design criteria and constraints; and

explain how human understandings, economics, ergonomics, and environmen-

tal considerations have influenced the solution.

• Develop work schedules and working plans which include optimal use and cost

of materials, processes, time, and expertise; construct a model of the solution,

incorporating developmental modifications while working to a high degree of

quality (craftsmanship).

Chemistry 7

• Devise a test of the solution according to the design criteria and perform the

test; record, portray, and logically evaluate performance test results through

quantitative, graphic, and verbal means. Use a variety of creative verbal and

graphic techniques effectively and persuasively to present conclusions, predict

impact and new problems, and suggest and pursue modifications.

STANDARD 1

Analysis, Inquiry,

and Design

ENGINEERING

DESIGN:

continued

STANDARD 2—Information Systems

Students will access, generate, process, and transfer information using appropriate technologies.

STANDARD 2

Key Idea 1:

Information technology is used to retrieve, process, and communicate information as a

INFORMATION

tool to enhance learning.

SYSTEMS:

Examples include:

• use the Internet as a source to retrieve information for classroom use, e.g.,

Periodic Table, acid rain

Key Idea 2:

Knowledge of the impacts and limitations of information systems is essential to its

effectiveness and ethical use.

Examples include:

• critically assess the value of information with or without benefit of scientific

backing and supporting data, and evaluate the effect such information could

have on public judgment or opinion, e.g., environmental issues

• discuss the use of the peer-review process in the scientific community and

explain its value in maintaining the integrity of scientific publication, e.g., “cold

fusion”

STANDARD 6—Interconnectedness: Common Themes

Students will understand the relationships and common themes that connect mathematics, science, and technology

and apply the themes to these and other areas of learning.

Key Idea 1:

STANDARD 6

Through systems thinking, people can recognize the commonalities that exist among all

Interconnectedness:

systems and how parts of a system interrelate and combine to perform specific

Common Themes

functions.

Examples include:

SYSTEMS

• use the concept of systems and surroundings to describe heat flow in a chemical

THINKING:

or physical change, e.g., dissolving process

Chemistry 8

Key Idea 2:

STANDARD 6

Models are simplified representations of objects, structures, or systems used in analysis,

Interconnectedness:

explanation, interpretation, or design.

Common Themes

2.1 Revise a model to create a more complete or improved representation of the system.

• show how models are revised in response to experimental evidence, e.g., atomic

MODELS:

theory, Periodic Table

2.2 Collect information about the behavior of a system and use modeling tools to

represent the operation of the system.

• show how information about a system is used to create a model, e.g., kinetic

molecular theory (KMT)

2.3 Find and use mathematical models that behave in the same manner as the

processes under investigation.

• show how mathematical models (equations) describe a process, e.g., combined

gas law

2.4 Compare predictions to actual observations, using test models.

• compare experimental results to a predicted value, e.g., percent error

STANDARD 6

Key Idea 3:

The grouping of magnitudes of size, time, frequency, and pressures or other units of

Interconnectedness:

Common Themes

measurement into a series of relative order provides a useful way to deal with the

immense range and the changes in scale that affect the behavior and design of systems.

3.1 Describe the effects of changes in scale on the functioning of physical, biological, or

MAGNITUDE AND

SCALE:

designed information systems.

• show how microscale processes can resemble or differ from real-world

processes, e.g., microscale chemistry

3.2 Extend the use of powers of ten notation to understanding the exponential

function and performing operations with exponential factors.

• use powers often to represent a large range of values for a physical quantity,

e.g., pH scale

STANDARD 6

Key Idea 4:

Equilibrium is a state of stability due either to a lack of change (static equilibrium) or a

Interconnectedness:

balance between opposing forces (dynamic equilibrium).

Common Themes

4.1 Describe specific instances of how disturbances might affect a system’s equilib-

rium, from small disturbances that do not upset the equilibrium to larger distur-

EQUILIBRIUM AND

STABILITY:

bances (threshold level) that cause the system to become unstable.

• explain how a small change might not affect a system, e.g., activation energy

4.2 Cite specific examples of how dynamic equilibrium is achieved by equality of

change in opposing directions.

• explain how a system returns to equilibrium in response to a stress, e.g.,

LeChatelier’s principle

Chemistry 9

Key Idea 5:

STANDARD 6

Identifying patterns of change is necessary for making predictions about future

Interconnectedness:

behavior and conditions.

Common Themes

Examples include:

• use graphs to make predictions, e.g., half-life, solubility

PATTERNS OF

• use graphs to identify patterns and interpret experimental data, e.g., heating

CHANGE:

and cooling curves

STANDARD 7—Interdisciplinary Problem Solving

Students will apply the knowledge and thinking skills of mathematics, science, and technology to address real-life

problems and make informed decisions.

STANDARD 7

Key Idea 1:

The knowledge and skills of mathematics, science, and technology are used together to

Interdisciplinary

make informed decisions and solve problems, especially those relating to issues of sci-

Problem Solving

ence/technology/society, consumer decision making, design, and inquiry into

phenomena.

CONNECTIONS:

1.1 Analyze science/technology/society problems and issues on a community,

national, or global scale and plan and carry out a remedial course of action.

• carry out a remedial course of action by communicating the plan to others, e.g.,

writing and sending “a letter to the editor”

1.2 Analyze and quantify consumer product data, understand environmental and eco-

nomic impacts, develop a method for judging the value and efficacy of competing

products, and discuss cost-benefit and risk-benefit trade-offs made in arriving at

the optimal choice.

• compare and analyze specific consumer products, e.g., antacids, vitamin C

1.3 Design solutions to real-world problems on a community, national, or global scale,

using a technological design process that integrates scientific investigation and rig-

orous mathematical analysis of the problem and of the solution.

• design a potential solution to a regional problem, e.g., suggest a plan to adjust

the acidity of a lake in the Adirondacks

1.4 Explain and evaluate phenomena mathematically and scientifically by formulating

a testable hypothesis, demonstrating the logical connections between the scientific

concepts guiding the hypothesis and the design of an experiment, applying and

inquiring into the mathematical ideas relating to investigation of phenomena, and

using (and if needed, designing) technological tools and procedures to assist in the

investigation and in the communication of results.

• design an experiment that requires the use of a mathematical concept to solve a

scientific problem, e.g., an experiment to compare the density of different types

of soda pop

Chemistry 10

STANDARD 7

Interdisciplinary

Problem Solving

STRATEGIES:

Key Idea 2:

Solving interdisciplinary problems involves a variety of skills and strategies, including

effective work habits; gathering and processing information; generating and analyzing

ideas; realizing ideas; making connections among the common themes of mathematics,

science, and technology; and presenting results.

If students are asked to do a project, then the project would require students to:

• work effectively

• gather and process information

• generate and analyze ideas

• observe common themes

• realize ideas

• present results

Chemistry 11

PROCESS SKILLS

BASED ON STANDARD 4

STANDARD 4—The Physical Setting

Students will understand and apply scientific concepts, principles, and theories pertaining to the physical setting

and living environment and recognize the historical development of ideas in science.

Note: The use of e.g. denotes examples which may be used for in-depth study. The terms for example and such as denote

material which is testable. Items in parentheses denote further definition of the word(s) preceding the item and are testable.

STANDARD 4

Key Idea 3:

The Physical Setting

Matter is made up of particles whose properties determine the observ-

able characteristics of matter and its reactivity.

3.1 Explain the properties of materials in terms of the arrange-

ment and properties of the atoms that compose them.

i use models to describe the structure of an atom

3.1b, 3.1c

ii relate experimental evidence (given in the introduction

3.1a

of Key Idea 3) to models of the atom

iii determine the number of protons or electrons in an atom

3.1e

or ion when given one of these values

iv calculate the mass of an atom, the number of neutrons or

3.1f

the number of protons, given the other two values

v distinguish between ground state and excited state

3.1j

electron configurations, e.g., 2-8-2 vs. 2-7-3

vi identify an element by comparing its bright-line

3.1k

spectrum to given spectra

vii distinguish between valence and non-valence electrons,

3.1l

given an electron configuration, e.g., 2-8-2

viii draw a Lewis electron-dot structure of an atom

3.1l

ix determine decay mode and write nuclear equations

3.1p, 4.4b

showing alpha and beta decay

x interpret and write isotopic notation

3.1g

xi given an atomic mass, determine the most abundant

3.1n

isotope

xii calculate the atomic mass of an element, given the

3.1n

masses and ratios of naturally occurring isotopes

xiii classify elements as metals, nonmetals, metalloids, or

3.1v, 3.1w, 3.1x, 3.1y

noble gases by their properties

xiv compare and contrast properties of elements within a group

3.1aa, 3.1bb

or a period for Groups 1, 2, 13-18 on the Periodic Table

xv determine the group of an element, given the chemical

3.1z

formula of a compound, e.g., XCl or XCl

xvi explain the placement of an unknown element on the

3.1v, 3.1w, 3.1x, 3.1y

Periodic Table based on its properties

xvii classify an organic compound based on its structural or

3.1ff, 3.1gg, 3.1hh

condensed structural formula

(i.e., CH

COOH or -C-C-OH)

xviii describe the states of the elements at STP

3.1jj

xix distinguish among ionic, molecular, and metallic sub-

3.1dd, 3.1w, 5.2g, 5.2h

stances, given their properties

xx draw a structural formula with the functional group(s)

3.1ff, 3.1hh

on a straight chain hydrocarbon backbone, when given

the IUPAC name for the compound

Chemistry 12

STANDARD 4

The Physical Setting

continued

xxi draw structural formulas for alkanes, alkenes, and

alkynes containing a maximum of ten carbon atoms

xxii use a simple particle model to differentiate among prop-

erties of solids, liquids, and gases

xxiii compare the entropy of phases of matter

xxiv describe the processes and uses of filtration, distillation,

and chromatography in the separation of a mixture

xxv interpret and construct solubility curves

xxvi apply the adage “like dissolves like” to real-world

situations

xxviiinterpret solution concentration data

xxviii use solubility curves to distinguish among saturated,

supersaturated, and unsaturated solutions

xxix calculate solution concentration in molarity (M), percent

mass, and parts per million (ppm)

xxx describe the preparation of a solution, given the molarity

xxxi given properties, identify substances as Arrhenius acids

or Arrhenius bases

xxxii identify solutions as acid, base, or neutral based upon

the pH

xxxiii interpret changes in acid-base indicator color

xxxiv write simple neutralization reactions when given the

reactants

xxxv calculate the concentration or volume of a solution,

using titration data

xxxvi use particle models/diagrams to differentiate among

elements, compounds, and mixtures

3.2 Use atomic and molecular models to explain common chemi-

cal reactions.

i distinguish between chemical and physical changes

ii identify types of chemical reactions

iii determine a missing reactant or product in a balanced

equation

iv identify organic reactions

v balance equations, given the formulas of reactants and

products

vi write and balance half-reactions for oxidation and

reduction of free elements and their monatomic ions

vii identify and label the parts of a voltaic cell (cathode,

anode, salt bridge) and direction of electron flow, given

the reaction equation

viii identify and label the parts of an electrolytic cell (cath-

ode, anode) and direction of electron flow, given the

reaction equation

ix compare and contrast voltaic and electrolytic cells

x use an activity series to determine whether a redox

reaction is spontaneous

3.3 Apply the principle of conservation of mass to chemical

reactions.

i balance equations, given the formulas for reactants and

products

ii interpret balanced chemical equations in terms of

conservation of matter and energy

3.1ff, 3.1gg

3.1jj, 3.1kk

3.1mm

3.1nn

3.1oo

3.1pp

3.1oo

3.1pp

3.1uu

3.1ss

3.1xx

3.1zz

3.1r

3.2a

3.2b, 3.2c

3.2c, 3.2d

3.2c

3.2a, 3.3a, 3.3c

3.2f, 3.2h

3.2k

3.2l

3.2j

3.2k

3.3c

3.3a, 3.3c

Chemistry 13

STANDARD 4

iii create and use models of particles to demonstrate bal-

The Physical Setting

anced equations

iv calculate simple mole-mole stoichiometry problems,

continued

given a balanced equation

v determine the empirical formula from a molecular

formula

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

substance

vii determine the molecular formula, given the empirical

formula and the molecular mass

viii calculate the formula mass and gram-formula mass

ix determine the number of moles of a substance, given

its mass

3.4 Use kinetic molecular theory (KMT) to explain rates of reac-

tions and the relationships among temperature, pressure,

and volume of a substance.

i explain the gas laws in terms of KMT

ii solve problems, using the combined gas laws

iii convert temperatures in Celsius degrees (

C) to

kelvins (K), and kelvins to Celsius degrees

iv describe the concentration of particles and rates of

opposing reactions in an equilibrium system

v qualitatively describe the effect of stress on equilib-

rium, using LeChatelier’s principle

vi use collision theory to explain how various factors,

such as temperature, surface area, and concentration,

influence the rate of reaction

vii identify examples of physical equilibria as solution

equilibrium and phase equilibrium, including the con-

cept that a saturated solution is at equilibrium

Key Idea 4:

Energy exists in many forms, and when these forms change, energy is

conserved.

4.1 Observe and describe transmission of various forms of

energy.

i distinguish between endothermic and exothermic

reactions, using energy terms in a reaction equation,

∆H, potential energy diagrams, or experimental data

ii read and interpret potential energy diagrams: PE reac-

tants, PE products, activation energy (with or without

a catalyst), heat of reaction

4.2 Explain heat in terms of kinetic molecular theory.

i distinguish between heat energy and temperature in terms

of molecular motion and amount of matter

ii explain phase change in terms of the changes in energy and

intermolecular distances

iii qualitatively interpret heating and cooling curves in terms

of changes in kinetic and potential energy, heat of

vaporization, heat of fusion, and phase changes

iv calculate the heat involved in a phase or temperature

change for a given sample of matter

3.3a, 3.3c

3.3c

3.3d

3.3f

3.3d

3.3f

3.4c

3.4e

3.4i

3.4j

3.4d

3.4h

4.1b

4.1c, 4.1d

4.2a, 4.2b

4.2b

4.2a, 4.2c

4.2c

Chemistry 14

STANDARD 4

4.4 Explain the benefits and risks of radioactivity.

The Physical Setting

i calculate the initial amount, the fraction remaining, or the half-

4.4a

life of a radioactive isotope, given two of the three variables

continued

ii compare and contrast fission and fusion reactions

4.4b, 4.4f, 5.3b

iii complete nuclear equations; predict missing particles from

4.4c

nuclear equations

iv identify specific uses of some common radioisotopes, such as

4.4d

I-131 in diagnosing and treating thyroid disorders, C-14 to C-12

ratio in dating once-living organisms, U-238 to Pb-206 ratio in

dating geological formations, and Co-60 in treating cancer

Key Idea 5:

Energy and matter interact through forces that result in changes in

motion.

5.2 Students will explain chemical bonding in terms of the behavior

of electrons.

i demonstrate bonding concepts, using Lewis dot structures rep-

5.2a, 5.2d

resenting valence electrons:

§ transferred (ionic bonding)

§ shared (covalent bonding)

§ in a stable octet

Example:

atom ion

:Cl:

.. -

:Cl:

[ ]

ii compare the physical properties of substances based on chemi-

cal bonds and intermolecular forces, e.g., conductivity, mal-

5.2n

leability, solubility, hardness, melting point, and boiling point

iii explain vapor pressure, evaporation rate, and phase changes in

terms of intermolecular forces

5.2m

iv determine the noble gas configuration an atom will achieve by

bonding

5.2b

v distinguish between nonpolar covalent bonds (two of the same

nonmetals) and polar covalent bonds

5.2k

Chemistry 15

STANDARD 4: The Physical Setting

Students will understand and apply scientific concepts, principles, and theories pertaining to the physical

setting and living environment and recognize the historical development of ideas in science.

Key Idea 3:

Matter is made up of particles whose properties determine the observable characteristics of matter and its

reactivity.

Chemistry is the study of matter—its properties and its changes. The idea that matter is made up of particles is over

2000 years old, but the idea of using properties of these particles to explain observable characteristics of matter has

more recent origins. In ancient Greece, it was proposed that matter is composed of particles of four elements (earth,

air, water, and fire) and that these particles are in continual motion. The idea that particles could explain properties

of matter was not used for about 2000 years. In the late 1600s the properties of air were attributed to its particulate

nature; however, these particles were not thought to be fundamental. Instead, it was thought that they could

change into other particles with different properties.

In the late 1700s solid evidence about the nature of matter, gained through quantitative scientific experiments, accu-

mulated. Such evidence included the finding that during a chemical reaction matter was conserved. In the early

1800s a theory was proposed to explain these experimental facts. In this theory, atoms were hard, indivisible

spheres of different sizes and they combined in simple whole-number ratios to form compounds. The further treat-

ment of particles of matter as hard spheres in continual motion resulted in the 1800s in the kinetic molecular theory

of matter, which was used to explain the properties of gases.

In the late 1800s evidence was discovered that particles of matter could not be considered hard spheres; instead, par-

ticles were found to have an internal structure. The development of cathode ray tubes, and subsequent experiments

with them in the 1860s, led to the proposal that small, negatively charged particles—electrons—are part of the inter-

nal structure of atoms. In the early 1900s, to explain the results of the "gold foil experiment," a small, dense nucleus

was proposed to be at the center of the atom with electrons moving about in the empty space surrounding the

nucleus. Around this time, energy was proposed to exist in small, indivisible packets called quanta. This theory was

used to develop a model of the atom which had a central nucleus surrounded by shells of electrons. The model was

successful in explaining the spectra of the hydrogen atom and was used to explain aspects of chemical bonding.

Additional experiments with radioactivity provided evidence that atomic nuclei contained protons and neutrons.

Further investigation into the nature of the electron determined that it has wave-like properties. This feature was

incorporated into the wave-mechanical model of the atom, our most sophisticated model, and is necessary to

explain the spectra of multi-electron atoms.

Note: The use of e.g. denotes examples which may be used for in-depth study. The terms for example and such as denote

material which is testable. Items in parentheses denote further definition of the word(s) preceding the item and are testable.

PERFORMANCE

INDICATOR 3.1

Explain the properties of materials in terms of the arrangement and properties of the atoms that

compose them.

Major Understandings:

3.1a The modern model of the atom has evolved over a long period of time through the

work of many scientists.

3.1b Each atom has a nucleus, with an overall positive charge, surrounded by

negatively charged electrons.

3.1c Subatomic particles contained in the nucleus include protons and neutrons.

Chemistry 16

PERFORMANCE

INDICATOR 3.1

continued

3.1d The proton is positively charged, and the neutron has no charge. The electron is

negatively charged.

3.1e Protons and electrons have equal but opposite charges. The number of protons

equals the number of electrons in an atom.

3.1f The mass of each proton and each neutron is approximately equal to one atomic

mass unit. An electron is much less massive than a proton or a neutron.

3.1g The number of protons in an atom (atomic number) identifies the element. The sum

of the protons and neutrons in an atom (mass number) identifies an isotope. Common

notations that represent isotopes include:

C, carbon-14, C-14.

3.1h In the wave-mechanical model (electron cloud model) the electrons are in orbitals,

which are defined as the regions of the most probable electron location (ground state).

3.1i Each electron in an atom has its own distinct amount of energy.

3.1j When an electron in an atom gains a specific amount of energy, the electron is at a

higher energy state (excited state).

3.1k When an electron returns from a higher energy state to a lower energy state, a

specific amount of energy is emitted. This emitted energy can be used to identify an

element.

3.1l The outermost electrons in an atom are called the valence electrons. In general, the

number of valence electrons affects the chemical properties of an element.

3.1m Atoms of an element that contain the same number of protons but a different num-

ber of neutrons are called isotopes of that element.

3.1n The average atomic mass of an element is the weighted average of the masses of

its naturally occurring isotopes.

3.1o Stability of an isotope is based on the ratio of neutrons and protons in its nucleus.

Although most nuclei are stable, some are unstable and spontaneously decay, emitting

radiation.

3.1p Spontaneous decay can involve the release of alpha particles, beta particles,

positrons, and/or gamma radiation from the nucleus of an unstable isotope. These

emissions differ in mass, charge, ionizing power, and penetrating power.

3.1q Matter is classified as a pure substance or as a mixture of substances.

3.1r A pure substance (element or compound) has a constant composition and constant

properties throughout a given sample, and from sample to sample.

3.1s Mixtures are composed of two or more different substances that can be separated

by physical means. When different substances are mixed together, a homogeneous or

heterogeneous mixture is formed.

3.1t The proportions of components in a mixture can be varied. Each component in a

mixture retains its original properties.

Chemistry 17

PERFORMANCE

INDICATOR 3.1

continued

3.1u Elements are substances that are composed of atoms that have the same atomic

number. Elements cannot be broken down by chemical change.

3.1v Elements can be classified by their properties and located on the Periodic Table as

metals, nonmetals, metalloids (B, Si, Ge, As, Sb, Te), and noble gases.

3.1w Elements can be differentiated by physical properties. Physical properties of sub-

stances, such as density, conductivity, malleability, solubility, and hardness, differ

among elements.

3.1x Elements can also be differentiated by chemical properties. Chemical properties

describe how an element behaves during a chemical reaction.

3.1y The placement or location of an element on the Periodic Table gives an indication

of the physical and chemical properties of that element. The elements on the Periodic

Table are arranged in order of increasing atomic number.

3.1z For Groups 1, 2, and 13-18 on the Periodic Table, elements within the same group

have the same number of valence electrons (helium is an exception) and therefore simi-

lar chemical properties.

3.1aaThe succession of elements within the same group demonstrates characteristic

trends: differences in atomic radius, ionic radius, electronegativity, first ionization

energy, metallic/nonmetallic properties.

3.1bb The succession of elements across the same period demonstrates characteristic

trends: differences in atomic radius, ionic radius, electronegativity, first ionization

energy, metallic/nonmetallic properties.

3.1cc A compound is a substance composed of two or more different elements that are

chemically combined in a fixed proportion. A chemical compound can be broken down

by chemical means. A chemical compound can be represented by a specific chemical

formula and assigned a name based on the IUPAC system.

3.1dd Compounds can be differentiated by their physical and chemical properties.

3.1eeTypes of chemical formulas include empirical, molecular, and structural.

3.1ff Organic compounds contain carbon atoms, which bond to one another in chains,

rings, and networks to form a variety of structures. Organic compounds can be named

using the IUPAC system.

3.1gg Hydrocarbons are compounds that contain only carbon and hydrogen. Saturated

hydrocarbons contain only single carbon-carbon bonds. Unsaturated hydrocarbons

contain at least one multiple carbon-carbon bond.

3.1hh Organic acids, alcohols, esters, aldehydes, ketones, ethers, halides, amines,

amides, and amino acids are categories of organic compounds that differ in their struc-

tures. Functional groups impart distinctive physical and chemical properties to organic

compounds.

3.1ii Isomers of organic compounds have the same molecular formula, but different

structures and properties.

Chemistry 18

PERFORMANCE

INDICATOR 3.1

continued

3.1jj The structure and arrangement of particles and their interactions determine the

physical state of a substance at a given temperature and pressure.

3.1kkThe three phases of matter (solids, liquids, and gases) have different properties.

3.1ll Entropy is a measure of the randomness or disorder of a system. A system with

greater disorder has greater entropy.

3.1mm Systems in nature tend to undergo changes toward lower energy and higher

entropy.

3.1nnDifferences in properties such as density, particle size, molecular polarity, boiling

and freezing points, and solubility permit physical separation of the components of the

mixture.

3.1ooA solution is a homogeneous mixture of a solute dissolved in a solvent. The solu-

bility of a solute in a given amount of solvent is dependent on the temperature, the

pressure, and the chemical natures of the solute and solvent.

3.1ppThe concentration of a solution may be expressed in molarity (M), percent by vol-

ume, percent by mass, or parts per million (ppm).

3.1qqThe addition of a nonvolatile solute to a solvent causes the boiling point of the sol-

vent to increase and the freezing point of the solvent to decrease. The greater the con-

centration of solute particles, the greater the effect.

3.1rr An electrolyte is a substance which, when dissolved in water, forms a solution

capable of conducting an electric current. The ability of a solution to conduct an electric

current depends on the concentration of ions.

3.1ss The acidity or alkalinity of an aqueous solution can be measured by its pH value.

The relative level of acidity or alkalinity of these solutions can be shown by using

indicators.

3.1tt On the pH scale, each decrease of one unit of pH represents a tenfold increase in

hydronium ion concentration.

3.1uuBehavior of many acids and bases can be explained by the Arrhenius theory.

Arrhenius acids and bases are electrolytes.

3.1vvArrhenius acids yield H

(aq), hydrogen ion as the only positive ion in an aqueous

solution. The hydrogen ion may also be written as H

(aq), hydronium ion.

3.1ww Arrhenius bases yield OH

(aq), hydroxide ion as the only negative ion in an

aqueous solution.

3.1xx In the process of neutralization, an Arrhenius acid and an Arrhenius base react to

form a salt and water.

3.1yy There are alternate acid-base theories. One theory states that an acid is an H

donor and a base is an H

acceptor.

3.1zz Titration is a laboratory process in which a volume of a solution of known

concentration is used to determine the concentration of another solution.

Chemistry 19

PERFORMANCE

INDICATOR 3.2

PERFORMANCE

INDICATOR 3.3

Use atomic and molecular models to explain common chemical reactions.

Major Understandings:

3.2a A physical change results in the rearrangement of existing particles in a substance. A

chemical change results in the formation of different substances with changed properties.

3.2b Types of chemical reactions include synthesis, decomposition, single replacement,

and double replacement.

3.2c Types of organic reactions include addition, substitution, polymerization, esterifi-

cation, fermentation, saponification, and combustion.

3.2d An oxidation-reduction (redox) reaction involves the transfer of electrons (e

3.2e Reduction is the gain of electrons.

3.2f A half-reaction can be written to represent reduction.

3.2g Oxidation is the loss of electrons.

3.2h A half-reaction can be written to represent oxidation.

3.2i Oxidation numbers (states) can be assigned to atoms and ions. Changes in

oxidation numbers indicate that oxidation and reduction have occurred.

3.2j An electrochemical cell can be either voltaic or electrolytic. In an electrochemical

cell, oxidation occurs at the anode and reduction at the cathode.

3.2k A voltaic cell spontaneously converts chemical energy to electrical energy.

3.2l An electrolytic cell requires electrical energy to produce a chemical change. This

process is known as electrolysis.

Apply the principle of conservation of mass to chemical reactions.

Major Understandings:

3.3a In all chemical reactions there is a conservation of mass, energy, and charge.

3.3b In a redox reaction the number of electrons lost is equal to the number of electrons

gained.

3.3c A balanced chemical equation represents conservation of atoms. The coefficients

in a balanced chemical equation can be used to determine mole ratios in the reaction.

3.3d The empirical formula of a compound is the simplest whole-number ratio of

atoms of the elements in a compound. It may be different from the molecular formula,

which is the actual ratio of atoms in a molecule of that compound.

3.3e The formula mass of a substance is the sum of the atomic masses of its atoms. The

molar mass (gram-formula mass) of a substance equals one mole of that substance.

3.3f The percent composition by mass of each element in a compound can be

calculated mathematically.

Chemistry 20

PERFORMANCE

INDICATOR 3.4

Use kinetic molecular theory (KMT) to explain rates of reactions and the relationships among

temperature, pressure, and volume of a substance.

Major Understandings:

3.4a The concept of an ideal gas is a model to explain the behavior of gases. A real gas

is most like an ideal gas when the real gas is at low pressure and high temperature.

3.4b Kinetic molecular theory (KMT) for an ideal gas states that all gas particles:

• are in random, constant, straight-line motion.

• are separated by great distances relative to their size; the volume of the gas

particles is considered negligible.

• have no attractive forces between them.

• have collisions that may result in a transfer of energy between gas particles, but

the total energy of the system remains constant.

3.4c Kinetic molecular theory describes the relationships of pressure, volume, tempera-

ture, velocity, and frequency and force of collisions among gas molecules.

3.4d Collision theory states that a reaction is most likely to occur if reactant particles

collide with the proper energy and orientation.

3.4e Equal volumes of gases at the same temperature and pressure contain an equal

number of particles.

3.4f The rate of a chemical reaction depends on several factors: temperature, concentra-

tion, nature of the reactants, surface area, and the presence of a catalyst.

3.4g A catalyst provides an alternate reaction pathway, which has a lower activation

energy than an uncatalyzed reaction.

3.4h Some chemical and physical changes can reach equilibrium.

3.4i At equilibrium the rate of the forward reaction equals the rate of the reverse

reaction. The measurable quantities of reactants and products remain constant at

equilibrium.

3.4j LeChatelier's principle can be used to predict the effect of stress (change in

pressure, volume, concentration, and temperature) on a system at equilibrium.

Chemistry 21

Key Idea 4:

Energy exists in many forms, and when these forms change energy is conserved.

Throughout history, humankind has tried to effectively use and convert various forms of energy. Energy is used to

do work that makes life more productive and enjoyable. The Law of Conservation of Matter and Energy applies to

phase changes, chemical changes, and nuclear changes that help run our modern world. With a complete under-

standing of these processes and their application to the modern world comes a responsibility to take care of waste,

limit pollution, and decrease potential risks.

PERFORMANCE

INDICATOR 4.1

Observe and describe transmission of various forms of energy.

Major Understandings:

4.1a Energy can exist in different forms, such as chemical, electrical, electromagnetic,

thermal, mechanical, nuclear.

4.1b Chemical and physical changes can be exothermic or endothermic.

4.1c Energy released or absorbed during a chemical reaction can be represented by a

potential energy diagram.

4.1d Energy released or absorbed during a chemical reaction (heat of reaction) is equal

to the difference between the potential energy of the products and potential energy of

the reactants.

PERFORMANCE

INDICATOR 4.2

Explain heat in terms of kinetic molecular theory.

Major Understandings:

4.2a Heat is a transfer of energy (usually thermal energy) from a body of higher tem-

perature to a body of lower temperature. Thermal energy is the energy associated with

the random motion of atoms and molecules.

4.2b Temperature is a measurement of the average kinetic energy of the particles in a

sample of material. Temperature is not a form of energy.

4.2c The concepts of kinetic and potential energy can be used to explain physical

processes that include: fusion (melting), solidification (freezing), vaporization (boiling,

evaporation), condensation, sublimation, and deposition.

Chemistry 22

PERFORMANCE

INDICATOR 4.4

Explain the benefits and risks of radioactivity.

Major Understandings:

4.4a Each radioactive isotope has a specific mode and rate of decay (half-life).

4.4b Nuclear reactions include natural and artificial transmutation, fission, and fusion.

4.4c Nuclear reactions can be represented by equations that include symbols which

represent atomic nuclei (with mass number and atomic number), subatomic particles

(with mass number and charge), and/or emissions such as gamma radiation.

4.4d Radioactive isotopes have many beneficial uses. Radioactive isotopes are used in

medicine and industrial chemistry for radioactive dating, tracing chemical and biologi-

cal processes, industrial measurement, nuclear power, and detection and treatment of

diseases.

4.4e There are inherent risks associated with radioactivity and the use of radioactive

isotopes. Risks can include biological exposure, long-term storage and disposal, and

nuclear accidents.

4.4f There are benefits and risks associated with fission and fusion reactions.

Key Idea 5:

Energy and matter interact through forces that result in changes in motion.

Atoms and molecules are in constant motion. Chemical bonding between atoms involves energy and the interac-

tion of electrons with atomic nuclei. Intermolecular attractions, which may be temporary, occur when there is an

asymmetric distribution of charge.

Within all chemical interactions, matter and energy are conserved according to the Law of Conservation of Matter

and Energy. During a chemical change energy is absorbed or released as bonds are broken or formed. In maintain-

ing conservation of matter and energy, nuclear changes convert matter into energy. The energy released during a

nuclear change is much greater than the energy released during a chemical change.

The discovery of the energy stored in the nucleus of an atom, its uses, and its inherent benefits and risks is a contin-

uing process that began with the serendipitous detection of the first radioactive isotope. Early researchers added to

this knowledge and expanded our ability to utilize this newly discovered phenomenon. Using radioactivity, the

inner structure of the atom was defined by other researchers. Scientists involved in the development of nuclear fis-

sion and the atomic bomb explored both peaceful and destructive uses of nuclear energy. Modern researchers con-

tinue to search for ways in which the power of the nucleus can be used for the betterment of the world.

Chemistry 23

PERFORMANCE

INDICATOR 5.2

Explain chemical bonding in terms of the behavior of electrons.

Major Understandings:

5.2a Chemical bonds are formed when valence electrons are:

• transferred from one atom to another (ionic)

• shared between atoms (covalent)

• mobile within a metal (metallic)

5.2b Atoms attain a stable valence electron configuration by bonding with other atoms.

Noble gases have stable valence configurations and tend not to bond.

5.2c When an atom gains one or more electrons, it becomes a negative ion and its

radius increases. When an atom loses one or more electrons, it becomes a positive ion

and its radius decreases.

5.2d Electron-dot diagrams (Lewis structures) can represent the valence electron

arrangement in elements, compounds, and ions.

5.2e In a multiple covalent bond, more than one pair of electrons are shared between

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

5.2f Some elements exist in two or more forms in the same phase. These forms differ in

their molecular or crystal structure, and hence in their properties.

5.2g Two major categories of compounds are ionic and molecular (covalent)

compounds.

5.2h Metals tend to react with nonmetals to form ionic compounds. Nonmetals tend to

react with other nonmetals to form molecular (covalent) compounds. Ionic compounds

containing polyatomic ions have both ionic and covalent bonding.

5.2i When a bond is broken, energy is absorbed. When a bond is formed, energy is

released.

5.2j Electronegativity indicates how strongly an atom of an element attracts electrons

in a chemical bond. Electronegativity values are assigned according to arbitrary scales.

5.2k The electronegativity difference between two bonded atoms is used to assess the

degree of polarity in the bond.

5.2l Molecular polarity can be determined by the shape of the molecule and distribu-

tion of charge. Symmetrical (nonpolar) molecules include CO

, CH

, and diatomic ele-

ments. Asymmetrical (polar) molecules include HCl, NH

, and H

5.2m Intermolecular forces created by the unequal distribution of charge result in vary-

ing degrees of attraction between molecules. Hydrogen bonding is an example of a

strong intermolecular force.

5.2n Physical properties of substances can be explained in terms of chemical bonds and

intermolecular forces. These properties include conductivity, malleability, solubility,

hardness, melting point, and boiling point.

Chemistry 24

PERFORMANCE

INDICATOR 5.3

Compare energy relationships within an atom's nucleus to those outside the nucleus.

Major Understandings:

5.3a A change in the nucleus of an atom that converts it from one element to another is

called transmutation. This can occur naturally or can be induced by the bombardment of

the nucleus with high-energy particles.

5.3b Energy released in a nuclear reaction (fission or fusion) comes from the fractional

amount of mass that is converted into energy. Nuclear changes convert matter into

energy.

5.3c Energy released during nuclear reactions is much greater than the energy released

during chemical reactions.

Chemistry 25

APPENDIX A

CHEMISTRY CORE TOPICS

This section contains ten topic areas in which the major understandings found in the core are sorted by content

topic. These ten topic areas may be used for ease in curriculum writing; however, they do not connote a suggested

scope and sequence.

I. Atomic Concepts

I.1 The modern model of the atom has evolved over a long period of time through the work of many scien-

tists. (3.1a)

I.2 Each atom has a nucleus, with an overall positive charge, surrounded by one or more negatively charged

electrons. (3.1b)

I.3 Subatomic particles contained in the nucleus include protons and neutrons. (3.1c)

I.4 The proton is positively charged, and the neutron has no charge. The electron is negatively charged.

(3.1d)

I.5 Protons and electrons have equal but opposite charges. The number of protons equals the number of

electrons in an atom. (3.1e)

I.6 The mass of each proton and each neutron is approximately equal to one atomic mass unit. An electron is

much less massive than a proton or a neutron. (3.1f)

I.7 In the wave-mechanical model (electron cloud model), the electrons are in orbitals, which are defined as

the regions of the most probable electron location (ground state). (3.1h)

I.8 Each electron in an atom has its own distinct amount of energy. (3.1i)

I.9 When an electron in an atom gains a specific amount of energy, the electron is at a higher energy state

(excited state). (3.1j)

I.10 When an electron returns from a higher energy state to a lower energy state, a specific amount of energy

is emitted. This emitted energy can be used to identify an element. (3.1k)

I.11 The outermost electrons in an atom are called the valence electrons. In general, the number of valence

electrons affects the chemical properties of an element. (3.1l)

I.12 Atoms of an element that contain the same number of protons but a different number of neutrons are

called isotopes of that element. (3.1m)

I.13 The average atomic mass of an element is the weighted average of the masses of its naturally occurring

isotopes. (3.1n)

Chemistry 26

II. Periodic Table

II.1 The placement or location of elements on the Periodic Table gives an indication of physical and chemical

properties of that element. The elements on the Periodic Table are arranged in order of increasing atomic

number. (3.1y)

II.2 The number of protons in an atom (atomic number) identifies the element. The sum of the protons and

neutrons in an atom (mass number) identifies an isotope. Common notations that represent

isotopes include:

C, carbon-14, C-14. (3.1g)

II.3 Elements can be classified by their properties and located on the Periodic Table as metals, nonmetals,

metalloids (B, Si, Ge, As, Sb, Te), and noble gases. (3.1v)

II.4 Elements can be differentiated by their physical properties. Physical properties of substances, such as

density, conductivity, malleability, solubility, and hardness, differ among elements. (3.1w)

II.5 Elements can be differentiated by chemical properties. Chemical properties describe how an element

behaves during a chemical reaction. (3.1x)

II.6 Some elements exist in two or more forms in the same phase. These forms differ in their molecular or

crystal structure, and hence in their properties. (5.2f)

II.7 For Groups 1, 2, and 13-18 on the Periodic Table, elements within the same group have the same number

of valence electrons (helium is an exception) and therefore similar chemical properties. (3.1z)

II.8 The succession of elements within the same group demonstrates characteristic trends: differences in atomic

radius, ionic radius, electronegativity, first ionization energy, metallic/nonmetallic properties. (3.1aa)

II.9 The succession of elements across the same period demonstrates characteristic trends: differences in atomic

radius, ionic radius, electronegativity, first ionization energy, metallic/nonmetallic properties. (3.1bb)

III. Moles/Stoichiometry

III.1 A compound is a substance composed of two or more different elements that are chemically combined

in a fixed proportion. A chemical compound can be broken down by chemical means. A chemical com-

pound can be represented by a specific chemical formula and assigned a name based on the IUPAC

system. (3.1cc)

III.2 Types of chemical formulas include empirical, molecular, and structural. (3.1ee)

III.3 The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a

compound. It may be different from the molecular formula, which is the actual ratio of atoms in a mole-

cule of that compound. (3.3d)

III.4 In all chemical reactions there is a conservation of mass, energy, and charge. (3.3a)

III.5 A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical

equation can be used to determine mole ratios in the reaction. (3.3c)

III.6 The formula mass of a substance is the sum of the atomic masses of its atoms. The molar mass (gram-

formula mass) of a substance equals one mole of that substance. (3.3e)

Chemistry 27

III.7 The percent composition by mass of each element in a compound can be calculated mathematically. (3.3f)

III.8 Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement. (3.2b)

IV. Chemical Bonding

IV.1 Compounds can be differentiated by their chemical and physical properties. (3.1dd)

IV.2 Two major categories of compounds are ionic and molecular (covalent) compounds. (5.2g)

IV.3 Chemical bonds are formed when valence electrons are (5.2a):

• transferred from one atom to another (ionic)

• shared between atoms (covalent)

• mobile within a metal (metallic)

IV.4 In a multiple covalent bond, more than one pair of electrons are shared between two atoms. (5.2e)

IV.5 Molecular polarity can be determined by the shape of the molecule and the distribution of charge.

Symmetrical (nonpolar) molecules include CO

, CH

and diatomic elements. Asymmetrical (polar)

molecules include HCl, NH

, and H

O. (5.2l)

IV.6 When an atom gains one or more electrons, it becomes a negative ion and its radius increases. When an atom

loses one or more electrons, it becomes a positive ion and its radius decreases. (5.2c)

IV.7 When a bond is broken, energy is absorbed. When a bond is formed, energy is released. (5.2i)

IV.8 Atoms attain a stable valence electron configuration by bonding with other atoms. Noble gases have stable

valence configurations and tend not to bond. (5.2b)

IV.9 Physical properties of substances can be explained in terms of chemical bonds and intermolecular forces. These

properties include conductivity, malleability, solubility, hardness, melting point, and boiling point. (5.2n)

IV.10 Electron-dot diagrams (Lewis structures) can represent the valence electron arrangement in elements, com-

pounds, and ions. (5.2d)

IV.11 Electronegativity indicates how strongly an atom of an element attracts electrons in a chemical bond.

Electronegativity values are assigned according to arbitrary scales. (5.2j)

IV.12 The electronegativity difference between two bonded atoms is used to assess the degree of polarity in the

bond. (5.2k)

IV.13 Metals tend to react with nonmetals to form ionic compounds. Nonmetals tend to react with other nonmetals

to form molecular (covalent) compounds. Ionic compounds containing polyatomic ions have both ionic and

covalent bonding. (5.2h)

V. Physical Behavior of Matter

Chemistry 28

V. 1 Matter is classified as a pure substance or as a mixture of substances. (3.1q)

V. 2 The three phases of matter (solids, liquids, and gases) have different properties. (3.1kk)

V. 3 A pure substance (element or compound) has a constant composition and constant properties through-

out a given sample, and from sample to sample. (3.1r)

V. 4 Elements are substances that are composed of atoms that have the same atomic number. Elements can-

not be broken down by chemical change. (3.1u)

V. 5 Mixtures are composed of two or more different substances that can be separated by physical means.

When different substances are mixed together, a homogeneous or heterogeneous mixture is formed.

(3.1s)

V. 6 The proportions of components in a mixture can be varied. Each component in a mixture retains its

original properties. (3.1t)

V. 7 Differences in properties such as density, particle size, molecular polarity, boiling point and freezing

point, and solubility permit physical separation of the components of the mixture. (3.1nn)

V. 8 A solution is a homogeneous mixture of a solute dissolved in a solvent. The solubility of a solute in a

given amount of solvent is dependent on the temperature, the pressure, and the chemical natures of

the solute and solvent. (3.1oo)

V. 9 The concentration of a solution may be expressed as molarity (M), percent by volume, percent by

mass, or parts per million (ppm). (3.1pp)

V.10 The addition of a nonvolatile solute to a solvent causes the boiling point of the solvent to increase and

the freezing point of the solvent to decrease. The greater the concentration of particles, the greater the

effect. (3.1qq)

V. 1 1 Energy can exist in different forms, such as chemical, electrical, electromagnetic, thermal, mechanical,

and nuclear. (4.1a)

V.12 Heat is a transfer of energy (usually thermal energy) from a body of higher temperature to a body of

lower temperature. Thermal energy is the energy associated with the random motion of atoms and

molecules. (4.2a)

V.13 Temperature is a measurement of the average kinetic energy of the particles in a sample of material.

Temperature is not a form of energy. (4.2b)

V.14 The concept of an ideal gas is a model to explain the behavior of gases. A real gas is most like an ideal

gas when the real gas is at low pressure and high temperature. (3.4a)

V.15 Kinetic molecular theory (KMT) for an ideal gas states that all gas particles (3.4b):

1. are in random, constant, straight-line motion.

2. are separated by great distances relative to their size; the volume of the gas particles is

considered negligible.

3. have no attractive forces between them.

4. have collisions that may result in the transfer of energy between gas particles, but the total

energy of the system remains constant.

Chemistry 29

V.16 Collision theory states that a reaction is most likely to occur if reactant particles collide with the proper

energy and orientation. (3.4d)

V.17 Kinetic molecular theory describes the relationships of pressure, volume, temperature, velocity, and fre-

quency and force of collisions among gas molecules. (3.4c)

V.18 Equal volumes of different gases at the same temperature and pressure contain an equal number of particles.

(3.4e)

V.19 The concepts of kinetic and potential energy can be used to explain physical processes that include: fusion

(melting), solidification (freezing), vaporization (boiling, evaporation), condensation, sublimation, and depo-

sition. (4.2c)

V.20 A physical change results in the rearrangement of existing particles in a substance. A chemical change results

in the formation of different substances with changed properties. (3.2a)

V.21 Chemical and physical changes can be exothermic or endothermic. (4.1b)

V.22 The structure and arrangement of particles and their interactions determine the physical state of a substance

at a given temperature and pressure. (3.1jj)

V.23 Intermolecular forces created by the unequal distribution of charge result in varying degrees of attraction

between molecules. Hydrogen bonding is an example of a strong intermolecular force. (5.2m)

V.24 Physical properties of substances can be explained in terms of chemical bonds and intermolecular forces.

These properties include conductivity, malleability, solubility, hardness, melting point, and boiling point.

(5.2n)

VI. Kinetics/Equilibrium

VI.1 Collision theory states that a reaction is most likely to occur if reactant particles collide with the proper

energy and orientation. (3.4d)

VI.2 The rate of a chemical reaction depends on several factors: temperature, concentration, nature of reactants,

surface area, and the presence of a catalyst. (3.4f)

VI.3 Some chemical and physical changes can reach equilibrium. (3.4h)

VI.4 At equilibrium the rate of the forward reaction equals the rate of the reverse reaction.The measurable

quantities of reactants and products remain constant at equilibrium. (3.4i)

VI.5 LeChatelier’s principle can be used to predict the effect of stress (change in pressure, volume, concentration,

and temperature) on a system at equilibrium. (3.4j)

VI.6 Energy released or absorbed by a chemical reaction can be represented by a potential energy diagram. (4.1c)

VI.7 Energy released or absorbed during a chemical reaction (heat of reaction) is equal to the difference between

the potential energy of the products and the potential energy of the reactants. (4.1d)

VI.8 A catalyst provides an alternate reaction pathway, which has a lower activation energy than an uncatalyzed

reaction. (3.4g)

Chemistry 30

VI.9 Entropy is a measure of the randomness or disorder of a system. A system with greater disorder has

greater entropy. (3.1ll)

VI.10 Systems in nature tend to undergo changes toward lower energy and higher entropy. (3.1mm)

VII. Organic Chemistry

VII.1 Organic compounds contain carbon atoms which bond to one another in chains, rings, and networks to

form a variety of structures. Organic compounds can be named using the IUPAC system. (3.1ff)

VII.2 Hydrocarbons are compounds that contain only carbon and hydrogen. Saturated hydrocarbons contain

only single carbon-carbon bonds. Unsaturated hydrocarbons contain at least one multiple carbon-carbon

bond. (3.1gg)

VII.3 Organic acids, alcohols, esters, aldehydes, ketones, ethers, halides, amines, amides, and amino acids are

categories of organic molecules that differ in their structures. Functional groups impart distinctive physi-

cal and chemical properties to organic compounds. (3.1hh)

VII.4 Isomers of organic compounds have the same molecular formula but different structures and proper-

ties. (3.1ii)

VII.5 In a multiple covalent bond, more than one pair of electrons are shared between two atoms.

Unsaturated organic compounds contain at least one double or triple bond. (5.2e)

VII.6 Types of organic reactions include: addition, substitution, polymerization, esterification, fermentation,

saponification, and combustion. (3.2c)

VIII. Oxidation-Reduction

VIII.1 An oxidation-reduction (redox) reaction involves the transfer of electrons (e

). (3.2d)

VIII.2 Reduction is the gain of electrons. (3.2e)

VIII.3 A half-reaction can be written to represent reduction. (3.2f)

VIII.4 Oxidation is the loss of electrons. (3.2g)

VIII.5 A half-reaction can be written to represent oxidation. (3.2h)

VIII.6 In a redox reaction the number of electrons lost is equal to the number of electrons gained. (3.3b)

VIII.7 Oxidation numbers (states) can be assigned to atoms and ions. Changes in oxidation numbers indicate

that oxidation and reduction have occurred. (3.2i)

VIII.8 An electrochemical cell can be either voltaic or electrolytic. In an electrochemical cell, oxidation occurs

at the anode and reduction at the cathode. (3.2j)

VIII.9 A voltaic cell spontaneously converts chemical energy to electrical energy. (3.2k)

VIII.10 An electrolytic cell requires electrical energy to produce chemical change. This process is known as

electrolysis. (3.2l)

Chemistry 31

IX. Acids, Bases, and Salts

IX.1 Behavior of many acids and bases can be explained by the Arrhenius theory. Arrhenius acids and bases are

electrolytes. (3.1uu)

IX.2 An electrolyte is a substance which, when dissolved in water, forms a solution capable of conducting an elec-

tric current. The ability of a solution to conduct an electric current depends on the concentration of ions.

(3.1rr)

IX.3 Arrhenius acids yield H

(aq), hydrogen ion as the only positive ion in an aqueous solution. The hydrogen

ion may also be written as H

(aq), hydronium ion. (3.1vv)

IX.4 Arrhenius bases yield OH

(aq), hydroxide ion as the only negative ion in an aqueous solution. (3.1ww)

IX.5 In the process of neutralization, an Arrhenius acid and an Arrhenius base react to form a salt and water. (3.1xx)

IX.6 Titration is a laboratory process in which a volume of solution of known concentration is used to determine

the concentration of another solution. (3.1zz)

IX.7 There are alternate acid-base theories. One theory states that an acid is an H

donor and a base is an H

acceptor. (3.1yy)

IX.8 The acidity or alkalinity of a solution can be measured by its pH value. The relative level of acidity or

alkalinity of a solution can be shown by using indicators. (3.1ss)

IX.9 On the pH scale, each decrease of one unit of pH represents a tenfold increase in hydronium ion

concentration. (3.1tt)

X. Nuclear Chemistry

X.1 Stability of isotopes is based on the ratio of neutrons and protons in its nucleus. Although most nuclei are sta-

ble, some are unstable and spontaneously decay, emitting radiation. (3.1o)

X.2 Each radioactive isotope has a specific mode and rate of decay (half-life). (4.4a)

X.3 A change in the nucleus of an atom that converts it from one element to another is called transmutation. This

can occur naturally or can be induced by the bombardment of the nucleus by high-energy particles. (5.3a)

X.4 Spontaneous decay can involve the release of alpha particles, beta particles, positrons and/or gamma radia-

tion from the nucleus of an unstable isotope. These emissions differ in mass, charge, and ionizing power, and

penetrating power. (3.1p)

X.5 Nuclear reactions include natural and artificial transmutation, fission, and fusion. (4.4b)

X.6 There are benefits and risks associated with fission and fusion reactions. (4.4f)

X.7 Nuclear reactions can be represented by equations that include symbols which represent atomic nuclei (with

the mass number and atomic number), subatomic particles (with mass number and charge), and/or

emissions such as gamma radiation. (4.4c).

Chemistry 32

X.8 Energy released in a nuclear reaction (fission or fusion) comes from the fractional amount of mass con-

verted into energy. Nuclear changes convert matter into energy. (5.3b)

X.9 Energy released during nuclear reactions is much greater than the energy released during chemical

reactions. (5.3c)

X.10 There are inherent risks associated with radioactivity and the use of radioactive isotopes. Risks can

include biological exposure, long-term storage and disposal, and nuclear accidents. (4.4e)

X.11 Radioactive isotopes have many beneficial uses. Radioactive isotopes are used in medicine and indus-

trial chemistry, e.g., radioactive dating, tracing chemical and biological processes, industrial measure-

ment, nuclear power, and detection and treatment of disease. (4.4d)

Chemistry 33

APPENDIX B

PHYSICAL SETTING/CHEMISTRY CONTENT CONNECTIONS TABLE

STANDARD 4: The Physical Setting

The Content Connections Table has been designed to assist teachers in curriculum writing and lesson planning.

Some of the listed major understandings have a related skill and/or real-world connection to a specific content

focus area. The scope of the content connections and skills is not meant to be limited; i.e., a skill may be connected

to more than one major understanding.

Additionally, real-world connections have been identified only to assist teachers in planning and are not meant to

link these connections to any assessment.

Students will understand and apply scientific concepts, principles, and theories pertaining to the physical

setting and living environment and recognize the historical development of ideas in science.

SKILLS

The student should be able to:

relate experimental evidence

(given in the introduction of

Key Idea 3) to models of the

atom (3.1ii)

use models to describe the

structure of an atom (3.1i)

determine the number of pro-

tons or electrons in an atom or

ion when given one of these

values (3.1iii)

calculate the mass of an atom,

the number of neutrons or the

number of protons, given the

other two values (3.1iv)

REAL-WORLD

CONNECTIONS

 lasers

I. Atomic Concepts

KEY

3.1a

3.1b

3.1c

3.1d

3.1e

3.1f

MAJOR

UNDERSTANDINGS

3.1a The modern model of

the atom has evolved over a

long period of time through

the work of many scientists.

3.1b Each atom has a

nucleus, with an overall posi-

tive charge, surrounded by

negatively charged electrons.

3.1c Subatomic particles

contained in the nucleus

include protons and neutrons.

3.1d The proton is positively

charged, and the neutron has

no charge. The electron is

negatively charged.

3.1e Protons and electrons

have equal but opposite

charges. The number of pro-

tons is equal to the number of

electrons in an atom.

3.1f The mass of each pro-

ton and each neutron is

approximately equal to one

atomic mass unit. An electron

is much less massive than a

proton or neutron.

LINK TO

APPENDIX

I.1

I.2

I.3

I.4

I.5

I.6

Chemistry 34

I. Atomic Concepts

SKILLS

LINK TO

MAJOR

KEY

REAL-WORLD

The student should be able to:

APPENDIX

UNDERSTANDINGS

CONNECTIONS

I.7

3.1h

In the wave-mechanical

model (electron cloud), the

electrons are in orbitals,

which are defined as regions

of most probable electron

location (ground state).

I.8

Each electron in an atom has

its own distinct amount of

energy.

3.1i

I.9

3.1j

I.10

3.1k

I.11

3.1l

I.12

3.1m

I.13

3.1n

When an electron in an atom

gains a specific amount of

energy, the electron is at a

higher energy state (excited

state).

When an electron returns

from a higher energy state to a

lower energy state, a specific

amount of energy is emitted.

This emitted energy can be

used to identify an element.

The outermost electrons in an

atom are called the valence

electrons. In general, the

number of valence electrons

affects the chemical proper-

ties of an element.

Atoms of an element that

contain the same number of

protons but a different num-

ber of neutrons are called iso-

topes of that element.

The average atomic mass of

an element is the weighted

average of the masses of its

naturally occurring isotopes.

distinguish between ground

state and excited state electron

configurations, e.g., 2-8-2 vs. 2-

7-3 (3.1v)

identify an element by compar-

ing its bright-line spectrum to

given spectra (3.1vi)

draw a Lewis electron-dot

structure of an atom (3.1viii)

distinguish between valence

and non-valence electrons,

given an electron configuration,

e.g., 2-8-2 (3.1vii)

given an atomic mass, deter-

mine the most abundant isotope

(3.1xi)

calculate the atomic mass of an

element, given the masses and

ratios of naturally occurring iso-

topes (3.1xii)

 flame tests

 neon lights

 fireworks

 forensic analysis

 spectral analysis of stars

Chemistry 35

II. Periodic Table

KEY

LINK TO

MAJOR

SKILLS

REAL-WORLD

APPENDIX

UNDERSTANDINGS

The student should be able to:

CONNECTIONS

II.1

3.1y

II.2

3.1g

II.3

3.1v

II.4

3.1w

II.5

3.1x

II.6

5.2f

The placement or location of

an element on the Periodic

Table gives an indication of

physical and chemical prop-

erties of that element. The

elements on the Periodic

Table are arranged in order of

increasing atomic number.

The number of protons in an

interpret and write isotopic

atom (atomic number) identi-

notation (3.1x)

fies the element. The sum of

the protons and neutrons in

an atom (mass number) iden-

tifies an isotope. Common

notations that represent iso-

topes include:

C, carbon-14, C-14.

Elements can be classified by

classify elements as metals,

 similar properties

their properties, and located

nonmetals, metalloids, or noble

and uses for elements in the

on the Periodic Table, as met-

gases by their properties

same family

als, nonmetals, metalloids (B,

(3.1xiii)

Si, Ge, As, Sb, Te) , and noble

gases.

Elements can be differenti-

ated by their physical proper-

ties. Physical properties of

substances, such as density,

conductivity, malleability,

solubility, and hardness,

differ among elements.

Elements can be differenti-

ated by chemical properties.

Chemical properties describe

how an element behaves dur-

ing a chemical reaction.

Some elements exist as two

or more forms in the same

phase. These forms differ in

their molecular or crystal

structure, and hence in their

properties.

explain the placement of an

unknown element in the

Periodic Table based on its

properties (3.1xvi)

describe the states of the ele-

ments at STP (3.1xviii)

 similar properties

and uses for elements in the

same family

 characteristics of a class

of elements are similar

 uses of different

elements, e.g., use of semicon-

ductors in solid state electron-

ics and computer technology

 alloys as superconduc-

tors

 metallurgy

 recovery of metals

 different properties for

each allotrope:

∞ oxygen gas vs. ozone

∞ coal vs. graphite vs.

diamond vs. buck-

minsterfullerene

Chemistry 36

II. Periodic Table

KEY

LINK TO

APPENDIX

3.1z

II.7

II.8

3.1aa

II.9

3.1bb

III.1

3.1cc

MAJOR

UNDERSTANDINGS

For Groups 1, 2, and 13-18 on

the Periodic Table, elements

within the same group have

the same number of valence

electrons (helium is an excep-

tion) and therefore similar

chemical properties.

The succession of elements

within the same group

demonstrates characteristic

trends: differences in atomic

radius, ionic radius, elec-

tronegativity, first ionization

energy, metallic/nonmetallic

properties.

The succession of elements

across the same period

demonstrates characteristic

trends: differences in atomic

radius, ionic radius, elec-

tronegativity, first ionization

energy, metallic/nonmetallic

properties.

SKILLS

The student should be able to:

determine the group of an ele-

ment, given the chemical for-

mula of a compound, e.g., XCl

or XCl

(3.1xv)

compare and contrast proper-

ties of elements within a group

or a period for Groups 1, 2, 13-

18 on the Periodic Table (3.1xiv)

III. Moles/Stoichiometry

A compound is a substance

composed of two or more dif-

ferent elements that are

chemically combined in a

fixed proportion. A chemical

compound can be broken

down by chemical means. A

chemical compound can be

represented by a specific

chemical formula and

assigned a name based on the

IUPAC system.

REAL-WORLD

CONNECTIONS

 reading food and bever-

age labels (consumer Chemistry)

III.2

Types of chemical formulas

include: empirical, molecular,

and structural.

3.1ee

Chemistry 37

III. Moles/Stoichiometry

KEY

LINK TO

APPENDIX

3.3d

III.3

MAJOR

UNDERSTANDINGS

The empirical formula of a

compound is the simplest

whole-number ratio of atoms

of the elements in a com-

pound. It may be different

from the molecular formula,

which is the actual ratio of

atoms in a molecule of that

compound.

III.4

In all chemical reactions there

is a conservation of mass,

energy, and charge.

3.3a

SKILLS

The student should be able to:

determine the molecular for-

mula, given the empirical for-

mula and molecular mass

(3.3vii)

determine the empirical for-

mula from a molecular formula

(3.3v)

interpret balanced chemical

equations in terms of conserva-

tion of matter and energy (3.3ii)

balance equations, given the

formulas for reactants and

products (3.3i)

interpret balanced chemical

equations in terms of conserva-

tion of matter and energy (3.3ii)

create and use models of parti-

cles to demonstrate balanced

equations (3.3iii)

calculate simple mole-mole stoi-

chiometry problems, given a

balanced equation (3.3iv)

calculate the formula mass and

the gram-formula mass (3.3viii)

determine the number of moles

of a substance, given its mass

(3.3ix)

determine the mass of a given

number of moles of a substance

(3.3vi)

REAL-WORLD

CONNECTIONS

III.5

3.3c

III.6

3.3e

A balanced chemical equa-

tion represents conservation

of atoms. The coefficients in a

balanced chemical equation

can be used to determine

mole ratios in the reaction.

The formula mass of a sub-

stance is the sum of the

atomic masses of its atoms.

The molar mass (gram-

formula mass) of a substance

equals one mole of that

substance.

III.7

The percent composition by

mass of each element in a

compound can be calculated

mathematically.

3.3f

Chemistry 38

III.8

III. Moles/Stoichiometry

SKILLS

LINK TO

MAJOR

KEY

REAL-WORLD

The student should be able to:

APPENDIX

UNDERSTANDINGS

CONNECTIONS

identify types of chemical reac-

Types of chemical reactions

3.2b

 recovery of metals from

tions (3.2ii)

include synthesis, decompo-

ores

sition, single replacement,

 electroplating

and double replacement.

 corrosion

 precipitation reactions

 dangers of

mixing household chemicals

together, e.g., bleach and

ammonia

 electrolysis of active

metal compounds

 explosives (inflation of

air bags)

IV. Chemical Bonding

distinguish among ionic, molec-

tiated by their chemical and

IV.1

Compounds can be differen-

3.1dd

ular, and metallic substances,

physical properties.

given their properties (3.1xix)

IV.2

Two major categories of com-

pounds are ionic and molec-

ular (covalent) compounds.

5.2g

demonstrate bonding concepts

IV.3

Chemical bonds are formed

5.2a

 photosynthesis

using Lewis dot structures rep-

when valence electrons are:

 DNA bonding

resenting valence electrons:

another (ionic); shared

transferred from one atom to

transferred (ionic bonding);

between atoms (covalent);

shared (covalent bonding); in a

mobile within a metal

stable octet (5.2i)

(metallic).

IV.4

In a multiple covalent bond,

more than one pair of elec-

trons are shared between two

atoms. Unsaturated organic

compounds contain at least

one double or triple bond.

5.2e

Molecular polarity can be

determined by the shape and

distribution of the charge.

Symmetrical (nonpolar) mole-

cules include CO

, CH

, and

diatomic elements.

Asymmetrical (polar) mole-

cules include HCl, NH

, H

5.2l

Chemistry 39

IV.5

IV. Chemical Bonding

KEY

LINK TO

MAJOR

SKILLS

REAL-WORLD

APPENDIX

UNDERSTANDINGS

The student should be able to:

CONNECTIONS

IV.6

5.2c

When an atom gains one or

more electrons, it becomes a

negative ion and its radius

increases. When an atom

loses one or more electrons, it

becomes a positive ion and

its radius decreases.

IV.7

When a bond is broken,

energy is absorbed. When a

bond is formed, energy is

released.

5.2i

IV.8

5.2b

IV.9

5.2n

IV.10

5.2d

IV.11

5.2j

Atoms attain a stable valence

electron configuration by

bonding with other atoms.

Noble gases have stable

valence electron configura-

tions and tend not to bond.

Physical properties of sub-

stances can be explained in

terms of chemical bonds and

intermolecular forces. These

properties include conductiv-

ity, malleability, solubility,

hardness, melting point, and

boiling point.

Electron-dot diagrams (Lewis

structures) can represent the

valence electron arrangement

in elements, compounds, and

ions.

Electronegativity indicates

how strongly an atom of an

element attracts electrons in a

chemical bond. Electronega-

tivity values are assigned

according to arbitrary scales.

determine the noble gas config-

uration an atom will achieve

when bonding (5.2iv)

demonstrate bonding concepts,

using Lewis dot structures rep-

resenting valence electrons:

transferred (ionic bonding);

shared (covalent bonding); in a

stable octet (5.2i)

 saturated vs. unsatu-

rated compounds—health

connections

 free radicals

Chemistry 40

IV.12

IV. Chemical Bonding

SKILLS

LINK TO

MAJOR

KEY

REAL-WORLD

The student should be able to:

APPENDIX

UNDERSTANDINGS

CONNECTIONS

distinguish between nonpolar

ence between two bonded

The electronegativity differ-

5.2k

covalent bonds (two of the same

atoms is used to assess the

nonmetals) and polar covalent

degree of polarity in the

bonds (5.2v)

bond.

IV.13

Metals tend to react with

nonmetals to form ionic com-

pounds. Nonmetals tend to

react with other nonmetals to

form molecular (covalent)

compounds. Ionic com-

pounds containing poly-

atomic ions have both ionic

and covalent bonding.

5.2h

V. Physical Behavior of Matter

V. 1

Matter is classified as a pure

substance or as a mixture of

substances.

3.1q

use a simple particle model to

V. 2

The three phases of matter

3.1kk

 common

(solids, liquids, and gases)

differentiate among properties

everyday examples of solids,

of a solid, a liquid, and a gas

have different properties.

liquids, and gases

(3.1xxii)

 nature of H

O in

our environment

 solids

∞ metallic

∞ crystalline

∞ amorphous (quartz

glass, opals)

∞ solid state

 liquids

∞ surface tension

∞ capillary

∞ viscosity

 gases

∞ real and ideal gases

use particle models/diagrams

compound) has a constant

A pure substance (element or

3.1r

to differentiate among elements,

composition and constant

compounds, and mixtures

properties throughout a

(3.1xxxvi)

given sample, and from sam-

ple to sample.

Chemistry 41

V. 3

V. Physical Behavior of Matter

SKILLS

LINK TO

MAJOR

KEY

REAL-WORLD

The student should be able to:

APPENDIX

UNDERSTANDINGS

CONNECTIONS

V. 4

3.1u

V. 5

3.1s

V. 6

3.1t

V. 7

3.1nn

V. 8

3.1oo

Elements are substances that

are composed of atoms that

have the same atomic num-

ber. Elements cannot be bro-

ken down by chemical

change.

Mixtures are composed of

two or more different sub-

stances that can be separated

by physical means. When dif-

ferent substances are mixed

together, a homogeneous or

heterogeneous mixture is

formed.

The proportions of compo-

nents in a mixture can be var-

ied. Each component in a

mixture retains its original

properties.

Differences in properties such

as density, particle size, mole-

cular polarity, boiling point

and freezing point, and solu-

bility permit physical separa-

tion of the components of the

mixture.

A solution is a homogeneous

mixture of a solute dissolved

in a solvent. The solubility of

a solute in a given amount of

solvent is dependent on the

temperature, the pressure,

and the chemical natures of

the solute and solvent.

describe the process and use of

filtration, distillation, and chro-

matography in the separation of

a mixture (3.1xxiv)

interpret and construct solubil-

ity curves (3.1xxv)

use solubility curves to distin-

guish among saturated, super-

saturated and unsaturated solu-

tions (3.1xxviii)

apply the adage "like dissolves

like" to real-world situations

(3.1xxvi)

 alloys

 separation by filtration,

distillation, desalination, crys-

tallization, extraction, chro-

matography

 water quality testing

 colloids

 emulsifiers (making ice

cream)

 sewage treatment

 degrees of saturation of

solutions

 dry cleaning

Chemistry 42

V. Physical Behavior of Matter

KEY

LINK TO

MAJOR

SKILLS

REAL-WORLD

APPENDIX

UNDERSTANDINGS

The student should be able to:

CONNECTIONS

V. 9

3.1pp

V.10

3.1qq

V. 11

Energy can exist in different

forms, such as chemical, elec-

trical, electromagnetic, ther-

mal, mechanical, and nuclear.

4.1a

V.12

4.2a

V.13

4.2b

V.14

3.4a

The concentration of a solu-

tion may be expressed as:

molarity (M), percent by vol-

ume, percent by mass, or

parts per million (ppm).

The addition of a nonvolatile

solute to a solvent causes the

boiling point of the solvent to

increase and the freezing

point of the solvent to

decrease. The greater the con-

centration of solute particles

the greater the effect.

Heat is a transfer of energy

(usually thermal energy)

from a body of higher tem-

perature to a body of lower

temperature. Thermal energy

is associated with the ran-

dom motion of atoms and

molecules.

Temperature is a measure of

the average kinetic energy of

the particles in a sample of

matter. Temperature is not a

form of energy.

The concept of an ideal gas is

a model to explain behavior

of gases. A real gas is most

like an ideal gas when the

real gas is at low pressure

and high temperature.

describe the preparation of a solu-

tion, given the molarity (3.1xxx)

interpret solution concentration

data (3.1xxx)

calculate solution concentrations

in molarity (M), percent mass,

and parts per million (ppm)

(3.1xxix)

distinguish between heat energy

and temperature in terms of

molecular motion and amount

of matter (4.2i)

qualitatively interpret heating and

cooling curves in terms of changes

in kinetic and potential energy,

heat of vaporization, heat of

fusion, and phase changes (4.2iii)

distinguish between heat energy

and temperature in terms of

molecular motion and amount

of matter (4.2i)

explain phase changes in terms

of the changes in energy and

intermolecular distance (4.2ii)

 salting an icy sidewalk

 ice cream making

 antifreeze/engine

coolant

 airplane deicing

 cooking pasta

 Earth's primitive

atmosphere

 use of models to explain

something that cannot be seen

Chemistry 43

V. Physical Behavior of Matter

SKILLS

LINK TO

MAJOR

KEY

REAL-WORLD

The student should be able to:

APPENDIX

UNDERSTANDINGS

CONNECTIONS

V.15

3.4b

V.16

3.4d

V.17

3.4c

Kinetic molecular theory

(KMT) for an ideal gas states

all gas particles:

 are in random, con-

stant, straight-line motion

 are separated by great

distances relative to their

size; the volume of gas parti-

cles is considered negligible

 have no attractive

forces between them

 have collisions that

may result in a transfer of

energy between particles, but

the total energy of the system

remains constant.

Collision theory states that a

reaction is most likely to

occur if reactant particles col-

lide with the proper energy

and orientation.

Kinetic molecular theory

describes the relationships of

pressure, volume, tempera-

ture, velocity, and frequency

and force of collisions among

gas molecules.

explain the gas laws in terms of

KMT (3.4i)

solve problems, using the com-

bined gas law (3.4ii)

convert temperatures in Celsius

V.18

3.4e

Equal volumes of gases at the

same temperature and pres-

degrees (

C) to kelvins (K), and

sure contain an equal number

kelvins to Celsius degrees

of particles.

(3.4iii)

V.19

4.2c

The concepts of kinetic and

potential energy can be used

to explain physical processes

that include: fusion (melting);

solidification (freezing);

vaporization (boiling, evapo-

ration), condensation, subli-

mation, and deposition.

qualitatively interpret heating

and cooling curves in terms of

changes in kinetic and potential

energy, heat of vaporization,

heat of fusion, and phase

changes (4.2iii)

calculate the heat involved in a

phase or temperature change

for a given sample of matter

(4.2iv)

explain phase change in terms

of the changes in energy and

intermolecular distances (4.2ii)

 structure and composi-

tion of Earth's atmosphere

(variations in pressure and

temperature)

 weather processes

 greenhouse gases

Chemistry 44

V. Physical Behavior of Matter

REAL-WORLD

SKILLS

LINK TO

MAJOR

KEY

CONNECTIONS

The student should be able to:

APPENDIX

UNDERSTANDINGS

V.20

3.2a

A physical change results in

the rearrangement of existing

particles in a substance. A

chemical change results in

the formation of different

substances with changed

properties.

V.21

Chemical and physical

changes can be exothermic or

endothermic.

4.1b

V.22

3.1jj

V.23

5.2m

V.24

5.2n

The structure and arrange-

ment of particles and their

interactions determine the

physical state of a substance

at a given temperature and

pressure.

Intermolecular forces created

by the unequal distribution

of charge result in varying

degrees of attraction between

molecules. Hydrogen bond-

ing is an example of a strong

intermolecular force.

Physical properties of sub-

stances can be explained in

terms of chemical bonds and

intermolecular forces. These

properties include conductiv-

ity, malleability, solubility,

hardness, melting point, and

boiling point.

distinguish between endother-

mic and exothermic reactions,

using energy terms in a reaction

equation, sH, potential energy

diagrams or experimental data

(4.1i)

use a simple particle model to

differentiate among properties

of solids, liquids, and gases

(3.1xxii)

explain vapor pressure, evapo-

ration rate, and phase changes

in terms of intermolecular

forces (5.2iii)

compare the physical properties

of substances based upon chem-

ical bonds and intermolecular

forces (5.2ii)

 calorimetry

 refrigeration

 meniscus (concave/-

convex)

 capillary action

 surface tension

Chemistry 45

VI Kinetics/Equilibrium

KEY

LINK TO

MAJOR

SKILLS

REAL-WORLD

APPENDIX

UNDERSTANDINGS

The student should be able to:

CONNECTIONS

VI.1

3.4d

VI.2

3.4f

VI.3

Some chemical and physical

changes can reach equilib-

rium.

3.4h

VI.4

3.4i

VI.5

3.4j

VI.6

4.1c

VI.7

4.1d

Collision theory states that a

reaction is most likely to

occur if reactant particles col-

lide with the proper energy

and orientation.

The rate of a chemical reac-

tion depends on several fac-

tors: temperature, concentra-

tion, nature of reactants,

surface area, and the

presence of a catalyst.

At equilibrium the rate of the

forward reaction equals the

rate of the reverse reaction.

The measurable quantities of

reactants and products

remain constant at equilib-

rium.

LeChatelier's principle can be

used to predict the effect of

stress (change in pressure,

volume, concentration, and

temperature) on a system at

equilibrium.

Energy released or absorbed

by a chemical reaction can be

represented by a potential

energy diagram.

Energy released or absorbed

by a chemical reaction (heat

of reaction) is equal to the

difference between the poten-

tial energy of the products

and the potential energy of

the reactants.

use collision theory to explain

how various factors, such as

temperature, surface area, and

concentration, influence the rate

of reaction (3.4vi)

identify examples of physical

equilibria as solution equilib-

rium and phase equilibrium,

including the concept that a sat-

urated solution is at equilibrium

(3.4 vii)

describe the concentration of

particles and rates of opposing

reactions in an equilibrium sys-

tem (3.4iv)

qualitatively describe the effect

 Haber process

of stress on equilibrium, using

LeChatelier's principle (3.4v)

read and interpret potential

energy diagrams: PE of reac-

tants and products, activation

energy (with or without a cata-

lyst), heat of reaction (4.1ii)

 synthesis of compounds

 catalysts and inhibitors

 balloons

 burning fossil fuels

 photosynthesis

 production of photo-

chemical smog

Chemistry 46

VI Kinetics/Equilibrium

KEY

LINK TO

APPENDIX

MAJOR

UNDERSTANDINGS

SKILLS

The student should be able to:

REAL-WORLD

CONNECTIONS

3.4g

VI.8

A catalyst provides an alter-

nate reaction pathway which

has a lower activation energy

than an uncatalyzed reaction.

 enzymes in the human

body

3.1ll

VI.9

Entropy is a measure of the

randomness or disorder of a

system. A system with greater

disorder has greater entropy.

compare the entropy of phases

of matter (3.1xxiii)

 relationship to phase

change

VI.10

3.1mm

VII.1

3.1ff

VII.2

3.1gg

VII.3

3.1hh

Systems in nature tend to

undergo changes toward lower

energy and higher entropy.

VII. Organic Chemistry

classify an organic compound

carbon atoms which bond to

Organic compounds contain

based on its structural or con-

one another in chains, rings,

densed structural formula

and networks to form a vari-

(3.1xvii)

ety of structures. Organic

compounds can be named

using the IUPAC system.

Hydrocarbons are com-

draw structural formulas for

pounds that contain only car-

alkanes, alkenes, and alkynes

bon and hydrogen. Saturated

containing a maximum of ten

hydrocarbons contain only

carbon atoms (3.1xxi)

single carbon-carbon bonds.

Unsaturated hydrocarbons

contain at least one multiple

carbon-carbon bond.

 chaos theory—random-

ness vs. order

 biochemical molecules-

formation of carbohydrates,

proteins, starches, fats, nucleic

acids

 synthetic polymers-

polyethylene (plastic bags,

toys), polystyrene (cups, insu-

lation), polypropylene ( car-

pets, bottles) polytetrafluoro-

ethylene (nonstick surfaces—

Teflon™), polyacrilonitrile

(yarns, fabrics, wigs)

 disposal problems of

synthetic polymers

Organic acids, alcohols,

esters, aldehydes, ketones,

ethers, halides, amines,

amides, and amino acids are

types of organic compounds

that differ in their structures.

Functional groups impart

distinctive physical and

chemical properties to

organic compounds.

classify an organic compound

based on its structural or con-

densed structural formula (3.1xvii)

draw a structural formula with

the functional group(s) on a

straight chain hydrocarbon

backbone, when given the cor-

rect IUPAC name for the com-

pound (3.1xx)

 making perfume

 wine production

 nuclear magnetic reso-

nance spectroscopy (NMR),

(MRI)

 dyes

 cosmetics

 odors (esters)

Chemistry 47

VII. Organic Chemistry

KEY

SKILLS

REAL-WORLD

APPENDIX

LINK TO

MAJOR

The student should be able to:

CONNECTIONS

UNDERSTANDINGS

3.1ii

 types, varieties, uses of

pounds have the same molec-

VII.4

Isomers of organic com-

organic compounds

ular formula, but different

 organic isomers

structures and properties.

5.2e

VII.5

 saturated vs. unsatu-

more than one pair of elec-

In a multiple covalent bond,

rated compounds—health

trons are shared between two

connections

atoms. Unsaturated organic

compounds contain at least

one double or triple bond.

3.2c

identify types of organic reac-

VII.6

 saponification—making

include: addition, substitu-

Types of organic reactions

tions (3.2iv)

soap

tion, polymerization, esterifi-

 polymerization– forma-

cation, fermentation, saponi-

determine a missing reactant or

tion of starches

fication, and combustion.

product in a balanced equation

 fermentation—alcohol

(3.2iii)

production

 combustion of fossil

fuels

 cellular respiration

VIII. Oxidation-Reduction

3.2d

determine a missing reactant or

VIII.1

An oxidation-reduction

 electrochemical

(redox) reaction involves

product in a balanced equation

cells

3.2iii)

 corrosion

transfer of electrons (e

 electrolysis

 photography

 rusting

Reduction is the gain of elec-

3.2e

VIII.2

 smelting

trons.

 leaching (refining of

gold)

 thermite reactions

(reduction of metal oxides,

e.g., aluminum)

A half-reaction can be written

3.2f

write and balance half-reactions

VIII.3

to represent reduction.

for oxidation and reduction of

free elements and their

monatomic ions (3.2vi)

Oxidation is the loss of elec-

3.2g

VIII.4

 recovery of active non-

trons.

metals (I

)

48 Chemistry

VIII. Oxidation-Reduction

REAL-WORLD

APPENDIX

SKILLS

LINK TO

MAJOR

KEY

CONNECTIONS

The student should be able to:

UNDERSTANDINGS

VIII.5

A half-reaction can be written

to represent oxidation.

3.2h

VIII.6

In a redox reaction the num-

ber of electrons lost is equal

to the number of electrons

gained.

3.3b

VIII.7

3.2i

VIII.8

3.2j

Oxidation numbers (states)

can be assigned to atoms and

ions. Changes in oxidation

numbers indicate that oxida-

tion and reduction have

occurred.

An electrochemical cell can

be either voltaic or elec-

trolytic. In an electrochemical

cell, oxidation occurs at the

anode and reduction at the

cathode.

VIII.9

A voltaic cell spontaneously

converts chemical energy to

electrical energy.

3.2k

 patina (copper—Statue

compare and contrast voltaic

of Liberty)

and electrolytic cells (3.2ix)

identify and label the parts of a

voltaic cell (cathode, anode, salt

bridge) and direction of electron

flow, given the reaction equa-

tion (3.2vii)

use an activity series to deter-

mine whether a redox reaction

is spontaneous (3.2x)

 metallurgy of

identify and label the parts of

VIII.10

An electrolytic cell requires

3.2l

an electrolytic cell (anode, cath-

iron and steel

electrical energy to produce

 electroplating

ode) and direction of electron

chemical change. This

flow, given the reaction

process is known as

equation (3.2viii)

electrolysis.

Chemistry 49

IX. Acids, Bases, and Salts

KEY

LINK TO

MAJOR

SKILLS

REAL-WORLD

APPENDIX

UNDERSTANDINGS

The student should be able to:

CONNECTIONS

IX.1

Behavior of many acids and

bases can be explained by the

Arrhenius theory. Arrhenius

acids and bases are elec-

trolytes.

3.1uu

given properties, identify sub-

stances as Arrhenius acids or

Arrhenius bases (3.1xxxi)

IX.2

An electrolyte is a substance

which, when dissolved in

water, forms a solution capa-

ble of conducting an electric

current. The ability of a solu-

tion to conduct an electric

current depends on the con-

centration of ions.

3.1rr

Arrhenius acids yield H

(aq),

hydrogen ion as the only pos-

itive ion in aqueous solution.

The hydrogen ion may also

be written as H

O+

(aq)

hydronium ion.

IX.3

3.1vv

Arrhenius bases yield

IX.4

3.1ww

 cleaning agents

(aq), hydroxide ion as

the only negative ion in an

aqueous solution.

write simple neutralization

tion, an Arrhenius acid and

IX.5

In the process of neutraliza-

3.1xx

reactions when given the

an Arrhenius base react to

reactants (3.1xxxiv)

form salt and water.

3.1zz

IX.6

Titration is a laboratory

process in which a volume of

solution of known concentra-

tion is used to determine the

concentration of another

solution.

3.1yy

IX.7

There are alternate acid-base

theories. One such theory

states that an acid is an H

donor and a base is an H

acceptor.

calculate the concentration or

volume of a solution, using

titration data (3.1xxxv)

Chemistry 50

IX.8

IX. Acids, Bases, and Salts

SKILLS

LINK TO

MAJOR

KEY

REAL-WORLD

The student should be able to:

UNDERSTANDINGS

APPENDIX

CONNECTIONS

interpret changes in acid-base

The acidity and alkalinity of

3.1ss

 acid rain

an aqueous solution can be

indicator color (3.1xxxiii)

 household chemicals

measured by its pH value.

 buffers

identify solutions as acid, base,

The relative level of acidity or

 swimming pool

or neutral based upon the pH

alkalinity of a solution can be

chemistry

(3.1xxxii)

shown by using indicators.

 blood acidosis/alkalosis

IX.9

On the pH scale, each

decrease of one unit of pH

represents a tenfold increase

in hydronium ion

concentration.

3.1tt

X. Nuclear Chemistry

X.1

Stability of isotopes is based

on the ratio of the neutrons

and protons in its nucleus.

Although most nuclei are sta-

ble, some are unstable and

spontaneously decay emit-

ting radiation.

3.1o

calculate the initial amount, the

X.2

Each radioactive isotope has

4.4a

 radioactive dating

fraction remaining, or the half-

decay (half-life).

a specific mode and rate of

life of a radioactive isotope,

given two of the three variables

(4.4i)

X.3

A change in the nucleus of an

5.3a

 nuclear fission and

atom that converts it from

fusion reactions that release

one element to another is

energy

called transmutation. This

 radioisotopes,

can occur naturally or can be

tracers, transmutation

induced by the bombardment

 man-made elements

of the nucleus by high-energy

particles.

determine decay mode and

involve the release of alpha

X.4

Spontaneous decay can

3.1p

write nuclear equations show-

particles, beta particles,

ing alpha and beta decay (3.1ix)

positrons, and/or gamma

radiation from the nucleus of

an unstable isotope. These

emissions differ in mass,

charge, ionizing power, and

penetrating power.

Chemistry 51

X. Nuclear Chemistry

KEY

LINK TO

APPENDIX

MAJOR

UNDERSTANDINGS

SKILLS

The student should be able to:

REAL-WORLD

CONNECTIONS

4.4b

X.5

Nuclear reactions include

natural and artificial trans-

mutation, fission, and fusion.

compare and contrast fission

and fusion reactions (4.4ii)

4.4f

X.6

There are benefits and risks

associated with fission and

fusion reactions.

X.7

4.4c

X.8

5.3b

X.9

5.3c

X.10

4.4e

X.11

4.4d

Nuclear reactions can be rep-

resented by equations that

include symbols which repre-

sent atomic nuclei (with the

mass number and atomic

number), subatomic particles

(with mass number and

charge), and/or emissions

such as gamma radiation.

Energy released in a nuclear

reaction (fission or fusion)

comes from the fractional

amount of mass converted

into energy. Nuclear changes

convert matter into energy.

Energy released during

nuclear reactions is much

greater than the energy

released during chemical

reactions.

There are inherent risks asso-

ciated with radioactivity and

the use of radioactive iso-

topes. Risks can include bio-

logical exposure, long-term

storage and disposal, and

nuclear accidents.

Radioactive isotopes have

many beneficial uses.

Radioactive isotopes are used

in medicine and industrial

chemistry, e.g., radioactive

dating, tracing chemical and

biological processes, indus-

trial measurement, nuclear

power, and detection and

treatment of diseases.

complete nuclear equations;

predict missing particles from

nuclear equations (4.4iii)

identify specific uses of some

common radioisotopes, such as:

I-131 in diagnosing and treating

thyroid disorders; C-14 to C-12

ratio in dating living organisms;

U-238 to Pb-206 ratio in dating

geological formations; Co-60 in

treating cancer (4.4iv)

 production of

nuclear power

∞ fission

∞ fusion (breeder reac-

tors)

 cost-benefit analysis

among various types of

power production

 nuclear waste

 radioactive pollution

 use of radioactive trac-

ers

 radiation therapy

 irradiated food

Chemistry 52

Chemistry 53

number of

particles (3.4e)

gas behavior

Naming (3.1cc)

(3.4c)

related to

explains

Physical and

Formula Mass (3.3e),

Metals,nonmetals,

Chemical Properties

Percent Composition (3.3f)

metalloids (3.1v)

Heat (4.2a)

(3.1w,x)

there is a

can be

explains

has a method

procedure for

classified as

for calculating

have

ideal gases (3.4a)

Formulas-chemi-

Temperature

cal, empirical, mol-

Elements (3.1u)

(4.2b)

ecular (3.3ee, 3.3d)

can be

works exactly

described with

explains

includes

for

Compound

includes

Kinetic

(3.1cc,3.1dd)

Pure

Physical

molecular

Substances

Processes (4.2c)

theory

Separation (3.1nn)

(3.1q,r)

(3.4b,c)

freezing

explains

point/boiling

have

defines

explains

Phases of matter

point (3.1qq)

techniques for

have properties

have changes

used in

Simple

explains

(3.1kk)

Dalton's

Mixtures

(3.1s, 3.1t)

atomic

Particle

theory

Model

can be

used in

solution (3.1oo)

classified as

can be used to

(3.1jj)

Particles as

have

understand

hard spheres

used in

concentrations

Reactions

(3.1pp)

and explains

using

Chemical vs. physical

conservation

change (3.2a)

of mass (3.3c)

Students will explain

properties of materials

in terms of the arrange-

ment and properties of

using

the atoms that com-

Structure of atom

(3.1a-g)

pose them (3.1)

describes the

can be

represented

Wave

Bohr-

Particles as

mechanical

Electron

Rutherford

have a

model

Energies (4.3a,

Model

having a

include

Organic (3.1ff-ii)

modern

3.1i - 3.1k)

(3.1h)

describes

structure

model of

the

include

can be

have

Types

represented

help

Properties

Lewis

Acid/bases

helps

explain

of compounds

are explained

(3.1rr-zz)

Stability (5.2b)

Electron-dot

explain

Model

include

helps

(5.2d)

Principles

explain

Properties of

Polarity (5.2l)

uses

helps

include

elements

explain some

include

Ion formation

explained

classifying as

(5.2c)

having trends

valence electrons

Molecular and

include

(3.1aa,bb)

(3.1l,z)

Ionic (5.2g)

include

Bond Formation

Atomic mass

examples

include

are

has

(5.2a,h)

(3.1n), Isotopes

Structure and

properties

forces (5.2m)

(3.1m,o)

related to

(3.1u,v,w,x,y)

Energy change

can

include change in

(5.2i)

electronegativity

radius (5.2c)

lead to

having different

(5.2j, 5.2k)

forms (5.2f))

Multiple bonds

5.2e

Chemistry 54

Note: This is an example of how

the chemistry core might be pre-

sented during the year. It is not a

suggested format from the New

York State Education Department.

Benefits

and risks

(4.4e)

Uses (4.4e)

Conservation

of energy &

forms of energy

(4.1a)

Conservation

of mass

Isotopes (3.1m-p)

Energy

released 5.3b,c

Half-life (4.4a)

Equations

(4.4b, c,3.1p)

Transmutation

(5.3a)

acid-base 3.1xx

Redox 3.2d-3.2o,

3.3b

Organic 3.2c

Equilibrium

3.4i-3.4j

T. Shiland

Saratoga Springs

Senior High School

effect of a

catalyst (3.4g)

entropy

(3.1ll)

PE diagrams

(4.1c)

Energy

changes during

bonding (5.2i)

Moles of reactants

or products can be

calculated (3.3c)

Charge and energy

conserved (3.3 a)

Exothermic or

endothermic (4.1b)

factors that

influence reaction

rate (3.4f)

tendency toward

lower energy

(3.1mm)

Students will explain

properties of materials in

terms of the arrangement

and properties of the

atoms that compose

them (3.1)

Properties of

matter and

energy

Nuclear

reactions

(related to 5.3)

Chemical

reactions (3.2)

Principles

Types of

chemical

reactions

(3.2b)

Principles

of chemical

reactions

Conservation

of Mass

(3.3a,3.3c)

Energy

Changes during a

reaction (4.1d)

Driving Forces

(3.1mm)

Collision theory

(3.4e)

including

includes

include

includes idea

includes

that

include

can be

also

involves

classified as

represented by

explains

using

understanding

includes

explaining

using

and should

understand

need to

understand basic

Chemistry 55