Chapter 1. Introduction: Matter and Measurement
Media Resources
Figures and Tables in Transparency Pack: Section:
Figure 1.5 Molecular Comparison of Elements, 1.2 Classifications of Matter
Compounds, and Mixtures
Animations: Section:
Electrolysis of Water 1.2 Classifications of Matter
Changes of State 1.3 Properties of Matter
Movies: Section:
Sodium and Potassium in Water 1.3 Properties of Matter
Mixtures and Compounds 1.3 Properties of Matter
Paper Chromatography of Ink 1.3 Properties of Matter
3-D Models: Section:
Oxygen 1.1 The Study of Chemistry
Water 1.1 The Study of Chemistry
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Other Resources
Further Readings: Section:
Chemistry in the Real World 1.1 The Study of Chemistry
Chemicals in Everyday Life 1.1 The Study of Chemistry
Why Does Popcorn Pop? 1.3 Properties of Matter
Using History to Teach The Scientific Method: The 1.3 Properties of Matter
Role of Errors
Having Fun with the Metric System 1.4 Units of Measurement
Method for Separating or Identifying Plastics 1.4 Units of Measurement
Error, Precision, and Uncertainty 1.5 Uncertainty in Measurement
Precision and Accuracy in Measurements: A Tale 1.5 Uncertainty in Measurement
of Four Graduated Cylinders
A Simple but Effective Demonstration for 1.5 Uncertainty in Measurement
Live Demonstrations: Section:
The First Demonstration: Proof that Air Is a 1.1 The Study of Chemistry
Substance
Science Demonstrations, Experiments, and 1.1 The Study of Chemistry
Resources. A Reference List for Elementary
through College Teachers Emphasizing
Chemistry with Some Physics and Life
Science
Introduction: Matter and Measurement
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Chapter 1. Introduction: Matter and Measurement
Common Student Misconceptions
• Students often confuse mass and weight.
• Students have difficulty with algebraic manipulation. Conversion of temperatures between Celsius
and Fahrenheit scales is particularly problematic.
• Students tend to equate density with mass.
Teaching Tips
• Many students have problems using dimensional analysis (“from physics”) in chemistry. This text
bases the whole of stoichiometry on dimensional analysis: students should be encouraged to embrace
Lecture Outline
1.1 The Study of Chemistry
• Chemistry:
• is the study of properties of materials and changes that they undergo.
• can be applied to all aspects of life (e.g., development of pharmaceuticals, leaf color change
in fall, etc.).
The Atomic and Molecular Perspective of Chemistry
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Chemistry involves the study of the properties and the behavior of matter.
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• Matter:
• is the physical material of the universe.
• About 100 elements constitute all matter.
• Elements:
• are made up of unique atoms, the building blocks of matter.
• Molecules:
• are combinations of atoms held together in specific shapes.
Why Study Chemistry?
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We study chemistry because:
• it has a considerable impact on society (health care, food, clothing, conservation of natural resources,
1.2 Classifications of Matter
• Matter is classified by state (solid, liquid, or gas) or by composition (element, compound or mixture).
States of Matter
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• Solids, liquids and gases are the three forms of matter called the states of matter.
• Properties described on the macroscopic level:
• gas (vapor): no fixed volume or shape, conforms to shape of container, compressible.
Pure Substances
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• Pure substances:
• are matter with distinct properties and fixed composition..
• are elements (cannot be decomposed into simpler substances; i.e. only one kind of atom) or
compounds (consist of two or more elements).
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• Mixtures:
• are a combination of two or more pure substances.
• Each substance retains its own identity.
Elements
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• There are 117 known elements.
Compounds
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• Compounds are combinations of elements.
Example: The compound H2O is a combination of elements H and O.
Mixtures
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• A mixture is a combination of two or more pure substances.
• Each substance retains its own identity, each substance is a component of the mixture.
FORWARD REFERENCES
• States of matter will be essential in properly writing chemical reaction equations, including net
ionic equations, in Chapter 4, as well as equilibrium constant expressions in Chapters 15, 19, and
20.
• Solutions will be further discussed in Chapters 4 and 13.
• Mixtures of gases will be discussed in Chapter 10.
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“What’s the Use?” from Further Readings
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1.3 Properties of Matter
• Each substance has a unique set of physical and chemical properties.
• Physical properties are measured without changing the substance (e.g., color, density, odor,
melting point, etc.).
Physical and Chemical Changes
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• Physical change: substance changes physical appearance without altering its identity (e.g., changes
of state).
• Chemical change (or chemical reaction): substance transforms into a chemically different
substances (i.e. identity changes, e.g., reaction of hydrogen and oxygen gases to produce water).
Separation of Mixtures
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• Key: separation techniques exploit differences in properties of the components.
The Scientific Method
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• The scientific method provides guidelines for the practice of science.
• Collect data (observe, experiment, etc.).
1.4 Units of Measurement
• Many properties of matter are quantitative, i.e., associated with numbers.
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• A measured quantity must have BOTH a number and a unit.
• The units most often used for scientific measurement are those of the metric system.
SI Units
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• 1960: All scientific units use Système International d’Unités (SI Units).
base units.
Length and Mass
• SI base unit of length = meter (1 m = 1.0936 yards).
Temperature
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• Temperature is the measure of the hotness or coldness of an object.
Derived SI Units
• These are formed from the seven base units.
Volume
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• Units of volume = (units of length)3 = m3.
• Important: the liter is not an SI unit.
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Density
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• Density is defined as mass divided by volume.
FORWARD REFERENCES
• Prefixes (Table 1.5) will be heavily used in future chapters: kJ (Chapters: 5-8, 14, 19-21); nm, pm and
1.5 Uncertainty in Measurement
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• There are two types of numbers:
• exact numbers (known as counting or defined).
• inexact numbers (derived from measurement).
Precision and Accuracy
• Precision: how well measured quantities agree with each other.
Significant Figures
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• All measurements have some degree of uncertainty or error associated with them.
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“Sugar in a Can of Soft Drink: A Density Exercise” from Live Demonstrations
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“The Mysterious Sunken Ice Cube” from Live Demonstrations
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“Method for Separating or Identifying Plastics” from Further Readings
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“Densities and Miscibilities of Liquids and Liquid Mixtures” from Live Demonstrations
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• Final calculations are only as significant as the least significant measurement.
• Rules:
1. Nonzero numbers and zeros between nonzero numbers (ie. imbedded zeros) are always significant.
• Method:
1. Write the number in scientific notation.
Significant Figures in Calculations
• In calculations, the least certain measurement limits the certainty of the calculated result
• The answer is reported with only 1 uncertain digit.
• Guidelines for keeping track of significant figures:
• Addition and Subtraction:
FORWARD REFERENCES
• Working in scientific notation will be required in Chapters: 6 (wavelength, frequency, energy of
levels), 14 (rate constant), 15 (equilibrium constant calculations), 16 and 17 (acid or base ionization
constants, concentrations), 19 (G° vs K calculations), 20 (G° vs K vs E° calculations).
• Rules for significant figures will appear whenever calculations are to be performed; rules for
significant figures in calculations with logarithms will be needed in Chapters 14, 16, 17, 19, and 20.
1.6 Dimensional Analysis
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• Dimensional analysis is a method of calculation utilizing a knowledge of units.
• Given units can be multiplied and divided to give the desired units.
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Using Two or More Conversion Factors
• We often need to use more than one conversion factor in order to complete a problem.
• When identical units are found in the numerator and denominator of a conversion, they will cancel.
Conversions Involving Volume
• We often will encounter conversions from one measure to a different measure.
• For example:
Summary of Dimensional Analysis
• In dimensional analysis always ask three questions:
1. What data are we given?
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Further Readings:
1. Martin B. Jones and Christina R. Miller, “Chemistry in the Real World,” J. Chem. Educ., Vol. 78,
2001, 484–487.
7. Vivi Ringes, “Origin of the Names of Chemical Elements,” J. Chem. Educ., Vol. 66, 1989, 731–738.
8. Nicholas C. Thomas, “Connecting Element names with the Names of U.S. Towns,” J. Chem. Educ.,
Vol. 86, 2009, 181-184. A sampling of towns in the US with names based on chemical elements.
12. Jeanne M. Buccigros, “T–Shirt Chromatography: A Chromatogram You Can Wear,” J. Chem. Educ.,
Vol. 69, 1992, 977–978.
13. Frederick C. Sauls, “Why Does Popcorn Pop? An Introduction to the Scientific Method,” J. Chem.
Educ., Vol. 86, 1991, 415–416.
14. Carmen J. Guinta, “Using History to Teach Scientific Method: The Case of Argon,” J. Chem. Educ.,
Vol. 75, 1998, 1322–1325. The scientific method is introduced using a case-study approach involving the
discovery of argon.
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23. Richard S. Treptow, “Precision and Accuracy in Measurements: a Tale of Four Graduated Cylinders,”
J. Chem. Educ., Vol. 75, 1998, 992–995.
24. Charles J. Guare, “Error, Precision, and Uncertainty,” J. Chem. Educ., Vol. 68, 1991, 649–652.
25. Kenton B. Abel and William M. Hemmerlin, “Significant Figures,” J. Chem. Educ., Vol. 67, 1990,
213.
26. H. Graden Kirksey and Paul Krause, “Significant Figures: A Classroom Demonstration,” J. Chem.
Educ., Vol. 69, 1992, 497–498.
27. Ben Ruekberg, “A Joke Based on Significant Figures,” J. Chem. Educ., Vol. 71, 1994, 306.
Live Demonstrations:
1. David A. Katz, “Science Demonstrations, Experiments, and Resources. A Reference List for
Elementary through College Teachers Emphasizing Chemistry with Some Physics and Life Science,” J.
Chem. Educ., Vol. 68, 1991, 235–244. An extensive listing of various demonstrations, experiments and
resources for the science educator.
Introduction: Matter and Measurement
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6. Clarke W. Earley, “A Simple Demonstration for Introducing the Metric System to Introductory
Chemistry Classes,” J. Chem. Educ., Vol. 76, 1999, 1215–1216.
7. Lee. R. Summerlin, Christie L. Borgford, and Julie B. Ealy, “Sugar in a Can of Soft Drink: A Density
Exercise,” Chemical Demonstrations, A Sourcebook for Teachers, Volume 2 (Washington: American
Chemical Society, 1988), p. 126–127.
8. Lee. R. Summerlin, Christie L. Borgford, and Julie B. Ealy, “The Mysterious Sunken Ice Cube,”
Chemical Demonstrations, A Sourcebook for Teachers, Volume 2 (Washington: American Chemical
Society, 1988), p. 15–16.