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Leaming-Focused® Notebook ©2005 Learning Concepts
Reproduction of this page for classroom and planning use by owner of this book is permitted.
Student Learning Map Unit Topic: How do scientists study Chemistry?
Key Learning(s): A. Scientific Community B. Quantifying and measuring matter.
Unit Essential Questions: A. What is the difference between observation and interpretation? B. How are scientific measurements calculated accurately? C. What is matter?
Concept: Concept: Concept: Concept:
Scientific Method
SI System Significant Figures Matter
Lesson EQ:
How is the scientific method used? What are the parts of the scientific method? How does the scientific community communicate? How do variables affect an observed result?
Lesson EQ:
What is the SI system? What are the units of the SI system? How are units converted in the SI system?
Lesson EQ:
List four steps for solving problems. What is uncertainty? What is a significant figure? How are significant figures used in calculations?
Lesson EQ:
What are physical properties and changes? What are chemical properties and changes? What is density? How is matter classified? Explain the difference between a substance and a mixture?
Vocabulary: Scientific method (parts)
Theory
Law
Observation
Interpretation
Variable and Control
Vocabulary: Prefix
Base unit
Derived unit
Conversion factors
Dimensional Analysis
Vocabulary: Accuracy
Precision
Uncertainty
Significant digit
Percent Error
Scientific Notation
Vocabulary: Mass
Volume
Density
Matter
Mixture
Filtration/Distillation
Instructional Tools
Chemical Changes Lab
Island of Rapone Dimensional Analysis Game
Density Demonstration
Chapter 2 Matter and Change 11
SECTION 2.1 PROPERTIES OF MATTER (pages 39–42)
This section helps you distinguish extensive from intensive properties andidentify substances by their properties. It teaches you how to differentiate thethree states of matter. It also defines a physical property and lists examples ofphysical properties and physical changes.
Describing Matter (page 39)
1. The ______________________ of an object is a measure of the amount of matter
the object contains.
2. How does an extensive property differ from an intensive property?
Identifying Substances (page 40)
3. Matter that has a uniform and definite composition is called a
______________________ .
4. Is the following sentence true or false? All samples of a substance have
different physical properties. ______________________
5. A physical property is a quality or condition of a substance that can be
______________________ or ______________________ without
changing the substance’s composition.
6. Circle the letter of the term that is NOT a physical property.
a. hardness c. boiling point
b. color d. melting
7. Look at Table 2.1 on page 40. What is the melting point of bromine? ______
8. Look at Table 2.1 on page 40. Circle the letter of the substance that is a yellow solid and melts at 115°C.
a. sulfur
b. chlorine
c. gold
d. copper
Name ___________________________ Date ___________________ Class __________________
MATTER AND CHANGE2
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substance
false
observed measured
�7°C
mass
An extensive property depends on the amount of matter, while an intensive property
depends on the type of matter.
05_Chem_GRSW_Ch02.SE/TE 6/11/04 3:27 PM Page 11
Name ___________________________ Date ___________________ Class __________________©
Pearson Education, Inc., publishing as Pearson Prentice Hall. All rights reserved.
12 Guided Reading and Study Workbook
CHAPTER 2, Matter and Change (continued)
9. Is the following sentence true or false? Physical properties can help a chemist
identify a substance. ______________________
States of Matter (pages 41–42)
10. Circle the letter of the term that is NOT a physical state of matter.
a. water
b. gas
c. liquid
d. solid
11. Complete the table about properties of three states of matter. Use these terms:definite, indefinite, easily, and not easily.
12. Match each arrangement of the particles in matter with a physical state.
Physical State Arrangement
______ gas a. packed tightly together
______ liquid b. close, but free to flow
______ solid c. spaced relatively far apart
13. Is the following sentence true or false? The words gas and vapor can be used
interchangeably. ______________________
14. The term gas is limited to those substances that exist in the gaseous state at
______________________ .
Properties of the States of Matter
Property Solid Liquid Gas or Vapor
Shape indefinite
Volume definite indefinite
Can be compressed easily
definite
noteasily
noteasily
definite
indefinite
false
room temperature
c
b
a
true
05_Chem_GRSW_Ch02.SE/TE 6/11/04 3:27 PM Page 12
Chapter 2 Matter and Change 13
15. What does vapor describe?
Physical Changes (page 42)
16. A physical change alters a given material without changing its chemical
______________________ .
17. What are some words that describe physical changes?
18. What is true about all physical changes that involve a change of state?
____________________________________________________________________________
SECTION 2.2 MIXTURES (pages 44–47)
This section explains how to classify a mixture as heterogeneous or homogeneous. Italso describes ways to separate mixtures.
Classifying Mixtures (pages 44–45)
1. Is the following sentence true or false? Most samples of matter are mixtures.
______________________
2. What is a mixture?
3. Is the following sentence true or false? A heterogeneous mixture is one that
has a completely uniform composition. ______________________
4. What is another name for a homogeneous mixture?
______________________
5. Circle the letter of the term that describes a part of a sample with uniform composition and properties.
a. solution
b. mixture
c. state
d. phase
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boil, freeze, dissolve, melt, condense, break, split, crack, grind, cut, and crush
composition
They are all reversible.
true
false
a solution
A mixture is a physical blend of two or more components.
Vapor describes the gaseous state of a substance that is generally a liquid or solid
at room temperature.
05_Chem_GRSW_Ch02.SE/TE 6/11/04 3:27 PM Page 13
14 Guided Reading and Study Workbook
6. How many phases exist in these types of mixtures?
a. Homogeneous ______________________
b. Heterogeneous ______________________
Separating Mixtures (pages 46–47)
7. In general, what is used to separate mixtures?
8. The process that separates a solid from a liquid in a heterogeneous mixture is
called ______________________ .
9. What happens during a distillation?
Match each term with its location in the diagram.
______ 15. condenser
______ 16. heat source
______ 17. thermometer
______ 18. tap water
______ 19. distilled water
CHAPTER 2, Matter and Change (continued)
Name ___________________________ Date ___________________ Class __________________©
Pearson Education, Inc., publishing as Pearson Prentice Hall. All rights reserved.
A liquid is boiled to produce a vapor that is then condensed into a liquid.
differences in physical properties
two or more
filtration
one
C
A
Coldwater
in
Cold water
out
D
E
Receiverflask
Distillationflask
B
D
A
B
E
C
05_Chem_GRSW_Ch02.SE/TE 6/11/04 3:27 PM Page 14
Chapter 2 Matter and Change 15
SECTION 2.3 ELEMENTS AND COMPOUNDS (pages 48–52)
This section explains a key difference between an element and a compound, anddescribes how chemical symbols and formulas are used to represent elements and compounds. It also summarizes the process for classifying substances and mixtures.
Distinguishing Elements and Compounds (pages 48–49)
1. All living and nonliving things are made up of building blocks called
______________________ .
2. What are the two groups into which substances can be classified?
3. Is the following sentence true or false? Elements can be easily separated into
simpler substances. ______________________
4. Compounds are substances that can be separated into simpler substances
only by ______________________ means.
5. Is the following sentence true or false? The properties of compounds are
different from those of their component elements. ______________________
6. Complete this sentence.
Sodium chloride (table salt) is a ______________________ of sodium, which
is a soft ______________________ , and chlorine, which is a pale yellow
______________________ .
Distinguishing Substances and Mixtures (page 50)
7. Describe one way to decide whether a sample of matter is a substance or amixture.
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false
By looking carefully at photographs and drawings in textbooks, you can better understand what you have read. Look carefully at Figure 2.8 on page 44. What important idea does this drawingcommunicate?
Reading Skill Practice
This drawing shows that the components of a mixture can be separated based on differences in their physical properties.
Substances can be classified are elements and compounds.
true
chemical
compound
metal
If the composition of a sample is fixed, the sample is a substance. If the composition
of a sample may vary, the sample is a mixture.
gas
elements
05_Chem_GRSW_Ch02.SE/TE 6/11/04 3:27 PM Page 15
© Pearson Education, Inc., publishing as Pearson Prentice Hall. All rights reserved.
Definite composition(homogeneous)
Mixture ofsubstances
Variable composition
(Examples: water, sodium chloride)
Can be separatedchemically
Can be separatedphysically
Uniform;also called a solution
a. _________
b. _________
(Examples: carbon, hydrogen, oxygen)
d. __________c. __________Nonuniform;
distinct phases
e. __________ f. __________
16 Guided Reading and Study Workbook
Name ____________________________________ Date __________ Class ___________________
8. Complete the labels in the diagram below.
Symbols and Formulas (pages 51–52)
9. What is used to represent an element?
10. What are chemical symbols used for?
11. Subscripts in chemical formulas are used to indicate the relative proportions
of the elements in the ______________________ .
12. Is the following sentence true or false? The elements that make
up a compound are always present in the same
proportions. ______________________
CHAPTER 2, Matter and Change (continued)
Chemical symbols are used to write the chemical formulas of compounds.
compound
true
Matter
Substance
Heterogeneous HomogeneousElementCompound
An element is represented by a one- or two-letter chemical symbol.
05_Chem_GRSW_Ch02.SE/TE 6/11/04 3:27 PM Page 16
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Chapter 2 Matter and Change 17
Name ____________________________________ Date __________ Class ___________________
13. Use Table 2.2 on page 52 to answer the following questions.
a. Pb is the symbol for what element? ______________________
b. What is the symbol for gold? ______________________
c. Stibium is the Latin name for which element? ______________________
SECTION 2.4 CHEMICAL REACTIONS (pages 53–55)
This section provides clues to help you recognize a chemical change. It alsoteaches the law of conservation of mass.
Chemical Changes (page 53)
1. What is a chemical property?
2. Is the following sentence true or false? Chemical properties are observed only
when a substance undergoes a chemical change. ______________________
3. What happens during a chemical reaction?
4. In chemical reactions, the substances present at the start of the reaction are
called ______________________ and the substances produced are
called ______________________ .
5. Circle the letter of the term that best completes the sentence. A change in the
composition of matter _______ occurs during a chemical reaction.
a. sometimes
b. rarely
c. always
d. never
6. Which representation of a chemical reaction is correct?
a. products reactants
b. reactants products➞
➞
lead
Au
antimony
One or more substances change into one or more new substances.
A chemical property is the ability of a substance to undergo a specific
chemical change.
reactants
products
true
05_Chem_GRSW_Ch02.SE/TE 6/11/04 3:27 PM Page 17
18 Guided Reading and Study Workbook
Recognizing Chemical Changes (page 54)
7. List the four possible clues to a chemical change?
8. Is the following statement true or false? If you observe a clue for chemical change,
you can be certain that a chemical change has taken place. ______________
9. Define a precipitate.
Conservation of Mass (page 55)
10. During a chemical reaction, the mass of the products is always equal to
the mass of the ______________________ .
11. The law of conservation of mass states that in any physical change or
chemical reaction, mass is neither ______________________ nor
______________________ .
12. Look at Figure 2.15 on page 55. How do you know that mass was conserved?
____________________________________________________________________________________
CHAPTER 2, Matter and Change (continued)
Name ___________________________ Date ___________________ Class __________________©
Pearson Education, Inc., publishing as Pearson Prentice Hall. All rights reserved.
created
destroyed
reactants
A precipitate is a solid that forms and settles out of a liquid mixture.
false
The mass of the product is the same as the mass of the reactants.
transfer of energy, a change in color, the production of a gas, or the formation
of a precipitate
05_Chem_GRSW_Ch02.SE/TE 6/11/04 3:27 PM Page 18
Chapter 3 Scientific Measurement 19
SECTION 3.1 MEASUREMENTS AND THEIR UNCERTAINTY (pages 63–72)
This section describes the concepts of accuracy, precision, and error in measure-ments. It also explains the proper use of significant figures in measurements and calculations.
Using and Expressing Measurements (page 63)
1. Why are numbers used in chemistry often expressed in scientific notation?
2. Circle the letter of each sentence that is true about numbers expressed inscientific notation.
a. A number expressed in scientific notation is written as the product of acoefficient and a power of 10.
b. The power of 10 is called the exponent.
c. The coefficient is always a number greater than or equal to one and less than ten.
d. For numbers less than one, the exponent is positive.
3. Circle the letter of the answer in which 503,000,000 is written correctly inscientific notation.
a. 5.03 � 10–7
b. 503 � 106
c. 5.03 � 108
d. 503 million
Accuracy, Precision, and Error (pages 64–65)
4. Is the following sentence true or false? To decide whether a measurement has
good precision or poor precision, the measurement must be made more than
once. ______________________
Name ___________________________ Date ___________________ Class __________________
SCIENTIFIC MEASUREMENT3
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Numbers used in chemical calculations are often very large or very small. Writing
out all the zeros in such numbers can be very cumbersome. Scientific notation
makes it easier to work with these numbers.
true
05_Chem_GRSW_Ch03.SE/TE 6/11/04 3:29 PM Page 19
Name ___________________________ Date ___________________ Class __________________©
Pearson Education, Inc., publishing as Pearson Prentice Hall. All rights reserved.
20 Guided Reading and Study Workbook
CHAPTER 3, Scientific Measurement (continued)
Label each of the three following sentences that describes accuracy with an A. Labeleach sentence that describes precision with a P.
__________ 5. Four of five repetitions of a measurement were numerically identical, and the fifth varied from the others in value by less than 1%.
__________ 6. Eight measurements were spread over a wide range.
__________ 7. A single measurement is within 1% of the correct value.
8. On a dartboard, darts that are closest to the bull’s-eye have been thrown with the greatest accuracy. On the second target, draw three darts to represent threetosses of lower precision, but higher accuracy than the darts on the first target.
9. What is the meaning of “accepted value” with respect to an experimentalmeasurement?
10. Complete the following sentence. For an experimental measurement, the
experimental value minus the accepted value is called the _________________ .
11. Is the following sentence true or false? The value of an error must be
positive. _________________
12. Relative error is also called _____________________________ .
13. The accepted value of a length measurement is 200 cm, and the experimentalvalue is 198 cm. Circle the letter of the value that shows the percent error of thismeasurement.
a. 2%
b. �2%
c. 1%
d. �1%
First target Second target
P
P
A
The accepted value is the correct value based on reliable references.
error
false
percent error
05_Chem_GRSW_Ch03.SE/TE 6/11/04 3:29 PM Page 20
Chapter 3 Scientific Measurement 21
Significant Figures in Measurements (pages 66–67)
14. If a thermometer is calibrated to the nearest degree, to what part of a degree
can you estimate the temperature it measures? ____________________________
15. Circle the letter of the correct digit. In the measurement 43.52 cm, which digitis the most uncertain?
a. 4 c. 5
b. 3 d. 2
16. Circle the letter of the correct number of significant figures in themeasurement 6.80 m.
a. 2 c. 4
b. 3 d. 5
17. List two situations in which measurements have an unlimited number ofsignificant figures.
18. Circle the letter of each sentence that is true about significant figures.
a. Every nonzero digit in a reported measurement is assumed to besignificant.
b. Zeros appearing between nonzero digits are never significant.
c. Leftmost zeros acting as placeholders in front of nonzero digits in numbersless than one are not significant.
d. All rightmost zeros to the right of the decimal point are always significant.
e. Zeros to the left of the decimal point that act as placeholders for the firstnonzero digit to the left of the decimal point are not significant.
Significant Figures in Calculations (pages 68–71)
19. Is the following sentence true or false? An answer is as precise as the most
precise measurement from which it was calculated. ______________________
Round the following measurements as indicated.
20. Round 65.145 meters to 4 significant figures. ______________________
21. Round 100.1°C to 1 significant figure. ______________________
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one tenth of a degree
When the measurement involves counting.
When the measurement involves exactly defined quantities.
a.
b.
false
65.15 meters
100°C
05_Chem_GRSW_Ch03.SE/TE 6/11/04 3:29 PM Page 21
22 Guided Reading and Study Workbook
22. Round 155 cm to two significant figures. ______________________
23. Round 0.000 718 kilograms to two significant figures. ______________________
24. Round 65.145 meters to three significant figures. ______________________
SECTION 3.2 THE INTERNATIONAL SYSTEM OF UNITS (pages 73–79)
This section defines units of measurement for length, volume, mass, temperature, andenergy in the International System of Units (SI).
Units and Quantities (pages 74–79)
1. Complete the table showing selected SI base units of measurement.
2. All metric units of length are based on multiples of _______ .
3. The International System of Units (SI) is a revised version of the
______________________ .
4. Explain what is meant by a “derived unit.”
5. Give at least one example of a derived unit.
CHAPTER 3, Scientific Measurement (continued)
Name ___________________________ Date ___________________ Class __________________©
Pearson Education, Inc., publishing as Pearson Prentice Hall. All rights reserved.
0.00072 kilograms
160 cm
65.1 meters
Units of Measurement
Quantity SI base unit Symbol
Length
Mass
Temperature
Time
meter
K
s
kilogram
kelvin
second
m
kg
10
Derived units are combinations of base units. All SI units are base units, or are
derived from base units.
Students’ responses will vary. Possible responses are units of volume, density,
or speed.
metric system
05_Chem_GRSW_Ch03.SE/TE 6/11/04 3:29 PM Page 22
Chapter 3 Scientific Measurement 23
6. Complete the following table showing some metric units of length. Rememberthat the meter is the SI base unit for length.
Match each metric unit with the best estimate of its length or distance.
_______ 7. Height of a stove top above the floor a. 1 km
_______ 8. Thickness of about 10 sheets of paper b. 1 m
_______ 9. Distance along a road spanning about 10 telephone poles c. 1 cm
_______ 10. Width of a key on a computer keyboard d. 1 mm
11. The space occupied by any sample of matter is called its ___________________ .
12. Circle the letter of each sentence that is true about units of volume.
a. The SI unit for volume is derived from the meter, the SI unit for length.
b. The liter (L) is a unit of volume.
c. The liter is an SI unit.
d. There are 1000 cm3 in 1 L, and there are also 1000 mL in 1 L, so 1 cm3 isequal to 1 mL.
Match each of the three descriptions of a volume to the appropriate metric unit ofvolume.
Example Unit of Volume
_______ 13. Interior of an oven a. 1 L
_______ 14. A box of cookies b. 1 m3
_______ 15. One-quarter teaspoon c. 1 mL
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Metric Units of Length
Unit Symbol Factor Multiplying Base Unit
Meter m 1
Kilometer km 1000
Centimeter cm 10�2
Millimeter mm 10�3
Nanometer nm 10�9
b
d
a
c
b
a
c
volume
05_Chem_GRSW_Ch03.SE/TE 6/11/04 3:29 PM Page 23
24 Guided Reading and Study Workbook
16. A volume of 1 L is also equal to
a. 1000 mL
b. 1 dm3
c. 1000 cm
17. The volume of any solid, liquid, or gas will change with
______________________ .
18. A kilogram was originally defined as the mass of __________________________ .
19. Circle the letter of the unit of mass commonly used in chemistry that equals1/1000 kilogram.
a. gram b. milligram c. milliliter
Match each unit of mass with the object whose mass would be closest to that unit.
Mass Unit of Mass
_______ 20. A few grains of sand a. 1 kg
_______ 21. A liter bottle of soda b. 1 g
_______ 22. Five aspirin tablets c. 1 mg
23. Circle the letter of the instrument shown that is used to measure mass.
a. scale
b. balance beam
c. platform balance
d. analytical balance
24. Is the following sentence true or false? The mass of an object changes with
location. ______________________
25. When brought to the surface of the moon, will a mass have more or less weight
than it did on the surface of Earth, or will it be the same weight? Explain.
26. Draw an arrow below the diagram, showing the direction of heat transferbetween two objects.
CHAPTER 3, Scientific Measurement (continued)
Name ___________________________ Date ___________________ Class __________________©
Pearson Education, Inc., publishing as Pearson Prentice Hall. All rights reserved.
temperature
1 L of liquid water at 4°C
false
Its weight will be less, because weight is a measure of gravitational force, and the
force of gravity on the moon is one-sixth what it is on Earth.
c
a
b
lowertemperature
highertemperature
➞
05_Chem_GRSW_Ch03.SE/TE 6/11/04 3:29 PM Page 24
________
Chapter 3 Scientific Measurement 25
27. What properties explain the behavior of liquid-filled thermometers?
28. What are the two reference temperatures on the Celsius scale?
29. What is the zero point, 0 K, on the Kelvin scale called?
______________________
30. A change of temperature equal to one Kelvin is equal to a change of
temperature of how many degrees Celsius? ______________________
31. Complete the diagram to show the reference temperatures in the Celsius and Kelvin scales.
32. One calorie is the quantity of heat that raises the temperature of
______________________ of pure water by ______________________ .
SECTION 3.3 CONVERSION PROBLEMS (pages 80–87)
This section explains how to construct conversion factors from equivalentmeasurements. It also describes how to apply the techniques of dimensionalanalysis to a variety of conversion problems.
Conversion Factors (pages 80–81)
1. How are the two parts of a conversion factor related?
________________________________________________
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Almost all liquids expand in volume with an increase in temperature. This expansion
forces the liquid higher in the thermometer tube as the temperature rises. As the
temperature falls, the liquid contracts and its level in the tube drops.
The freezing point of water is 0°C and the boiling point of water is 100°C.
absolute zero
1°C
Kelvin
Celsius
100divisions
Freezing pointof water
Boiling pointof water
100divisions
________
________
________
1°C1 g
They are equivalent.
0°C 100°C
273.15 K 373.15 K
05_Chem_GRSW_Ch03.SE/TE 6/11/04 3:29 PM Page 25
26 Guided Reading and Study Workbook
2. Look at Figure 3.11. In a conversion factor, the smaller number is part of the
quantity that has the ______________________ unit. The larger number is part
of the quantity that has the ______________________ unit.
3. Is the following sentence true or false? The actual size of a measurement
multiplied by a conversion factor remains the same, because the measurement
being converted is multiplied by unity. ______________________
4. Write two conversion factors based on the relationship between hours andminutes.
5. The average lead for a mechanical pencil is 6.0 cm long when it is new. Circlethe letter of the conversion factor you would use to find its length in inches.
a.
b.
c.
d.
6. A student is asked to calculate the volume, in milliliters, of 2 cups of oil. Thereare 225 mL per cup. The student calculates the volume as follows:
Volume � 2 cups � � 0.08 cup
List three errors the student made.
Dimensional Analysis (pages 81–83)
7. What is dimensional analysis?
8. Reread Sample Problem 3.5. The correct conversion factor has the ______________ unit
in the denominator and the ______________ unit in the numerator.
1 cup�25 mL
6.0 cm�
1 in.
1 in.�6.0 cm
1 in.�2.54 cm
2.54 cm�
1 in.
CHAPTER 3, Scientific Measurement (continued)
Name ___________________________ Date ___________________ Class __________________©
Pearson Education, Inc., publishing as Pearson Prentice Hall. All rights reserved.
larger
smaller
and 1 hour
��60 minutes
60 minutes��
1 hour
true
The conversion factor was inverted.
25 mL was used instead of 225 mL.
The unit in the answer is incorrect.
Dimensional analysis is a way to analyze and solve problems, using the units,
or dimensions, of the measurements.
known
unknown
05_Chem_GRSW_Ch03.SE/TE 6/11/04 3:29 PM Page 26
Chapter 3 Scientific Measurement 27
9. A container can hold 65 g of water. Circle the conversion factor needed to findthe mass of water that 5 identical containers can hold.
a.
b.
c.
d.
Converting Between Units (pages 84–85)
10. Converting between units is easily done using ____________________________ .
11. Circle the letter of the conversion factor that you would use to converttablespoons to milliliters.
a.
b.
c.
d.
12. Show the calculation you would use to convert the following:
a. 0.25 m to centimeters_______________________________________
b. 9.8 g to kilograms_______________________________________
c. 35 ms to seconds_______________________________________
d. 4.2 dL to liters_______________________________________
13. Complex conversions between units may require using __________________________conversion factor.
14. How many conversion factors would you need to use to find the number ofliters in a cubic decimeter? What are they?
15 mL��1 tablespoon
1 tablespoon��
15 mL
1 tablespoon��4 fluid ounces
4 fluid ounces��1 tablespoon
65 g water��5 containers
65 g water��1 container
1 container��65 g water
5 containers��
65 g water
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dimensional analysis
Three conversion factors are needed: 1000 cm3/1 dm3, 1 mL/1 cm3, 1 L/1000 mL
0.25 m �100 cm�
1 m
4.2 dL �1 L�10 dL
9.8 g �1 kg�1000 g
35 ms �1 s
��1000 ms
more than one
05_Chem_GRSW_Ch03.SE/TE 6/11/04 3:29 PM Page 27
28 Guided Reading and Study Workbook
15. How would you calculate the number of nanometers in 8.1 cm?
16. What is the equivalent of 0.35 lb in grams?
17. A scientist has 0.46 mL of a solution. How would she convert this volume tomicroliters?
18. Describe the steps you would use to solve this problem. In a scale drawing of adining room floor plan, 10 mm equals 2 meters. If the homeowners wanted topurchase flooring that costs $10.89 per square yard, how much would theyspend on flooring for the dining room? The dimensions of the dining room onthe floor plan are 40 mm � 32 mm.
19. Name three common measurements that are expressed as a ratio of two units.
20. What technique can be used to convert complex units?
21. A normal concentration of glucose, or sugar, in the blood is 95 mg/dL. Howmany grams of sugar would be present per liter of blood? Show the conversionfactors you use.
22. Replace each question mark in the table with the conversion factors needed toobtain the given units of density.
23. A man can run a mile in 4 minutes. Calculate his average speed in kilometersper hour. Show your work. (1 mile � 1.61 km)
CHAPTER 3, Scientific Measurement (continued)
Name ___________________________ Date ___________________ Class __________________©
Pearson Education, Inc., publishing as Pearson Prentice Hall. All rights reserved.
1. Convert each dimension to meters, then to yards. 2. Multiply the lengths in yards to
find the number of square yards. 3. Multiply by the cost per yard to find the total cost.
8.1 cm � 1 m/102 cm � 109 nm/1 m � 8.1 � 107 nm
0.35 lb � 1 kg/2.2 lb � 1000 g/1 kg � 770 g
0.46 mL � 1 L/103 mL � 106 �L/1 L � 460 �L
speed, density, gas mileage
1 mi/4 min � 1.61 km/1 mi � 60 min/1 h � 24.2 km/h
Mass Volume Density
g mm3 � ? g/m3
kg � ? cm3 � ? g/m31000 g��
1 kg
1 m3
��109 mm3
Complex units can be converted using dimensional analysis.
95 mg/dL � 10 dL/1 L � 1 g/1000 mg � 0.95 g/L
1 m3
�106 cm3
05_Chem_GRSW_Ch03.SE/TE 6/11/04 3:29 PM Page 28
Chapter 3 Scientific Measurement 29
24. A baseball player’s batting average is .254 (254 hits per 1000 at bats). If she is atbat an average of 3 times per game, how many hits will she make in 52 games?Show your work.
SECTION 3.4 DENSITY (pages 89–93)
This section defines density. It explains that density is a characteristic propertythat depends on the composition of a substance, not on the size of the sample.
Determining Density (page 89–90)
1. Is the mass of one pound of lead greater than, less than, or equal to the mass
of one pound of feathers? ______________________
2. Which material has a greater density, lead or feathers? ______________________
3. How is density defined?
4. The mass of a sample is measured in grams, and its volume is measured incubic centimeters. In what units would its density be reported?
5. Look at Table 3.6 on page 90. Circle the letter of the material that will sink inliquid water at 4°C.
a. aluminum
b. corn oil
c. ice
d. gasoline
Density and Temperature (pages 91–92)
6. The density of a substance generally decreases as its temperature increases.Are there any exceptions to this statement? Explain.
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52 games � 3 at bats/1 game � 254 hits/1000 at bats � 39 or 40 hits
equal to
lead
Density is the ratio of the mass of an object to its volume.
Yes. Over a small range of temperatures near the freezing point, the density of water
decreases as the temperature decreases. As a result, ice floats on liquid water.
grams per cubic centimeter (g/cm3)
05_Chem_GRSW_Ch03.SE/TE 6/11/04 3:29 PM Page 29
30 Guided Reading and Study Workbook
GUIDED PRACTICE PROBLEMS
GUIDED PRACTICE PROBLEM 3 (page 69)
3. Round 87.073 meters to three significant figures. Write your answer inscientific notation.
Step 1. To round to three significant figures,round to the nearest tenth.
Step 2. Write the number in scientific notation.
GUIDED PRACTICE PROBLEM 34 (page 85)
34. The radius of a potassium atom is 0.227 nm. Express this radius in centimeters.Complete the following steps to solve the problem.
Step 1. Use the conversion factors 0.227 nm � �for nanometers and centimeters. 1 � 109 nm
Step 2. Simplify. �
Step 3. Divide. � cm
EXTRA PRACTICE (similar to Practice Problem 36, page 86)
36. Gold has a density of about 20 g/cm3. Estimate this density in kg/m3.
0.227 � 102
��109
CHAPTER 3, Scientific Measurement (continued)
Name ___________________________ Date ___________________ Class __________________©
Pearson Education, Inc., publishing as Pearson Prentice Hall. All rights reserved.
87.073 rounds to _______
______________________ meters
87.1
8.71 �101
� � � �
� 2.0 � 104 kg/m3
2.0�107 kg��
103 m320 �106 kg��
103 m3
(102 cm)3
��1 m3
1 kg�103 g
20 g�1 cm3
1 m 102 cm�
1 m
2.27 � 10–8
centimeter
05_Chem_GRSW_Ch03.SE/TE 6/11/04 3:29 PM Page 30
Chapter 3 Scientific Measurement 31
GUIDED PRACTICE PROBLEM 46 (page 91)
46. A student finds a shiny piece of metal that she thinks is aluminum. In the lab,she determines that the metal has a volume of 245 cm3 and a mass of 612 g.Calculate the density. Is the metal aluminum?
Analyze
Step 1. List the known values. Volume � 245 cm3
Mass � _______ g
Step 2. List the unknown.
Calculate
Step 3. Use the following relationship to find the density. Remember to round your answer to three significant figures.
�
Step 4. To determine whether the piece of metal is aluminum, compare the density of the metal to the density of aluminum given in Table 3.7 on
page 90. Is the metal aluminum? ______________________
Evaluate
Step 5. Underline the correct word(s) that complete(s) this statement. Becausethe mass of the metal is about two and one-half times the volume, adensity of about 2.5 g/cm3 is reasonable. Because a density of 2.50 g/cm3
is nearly 10% less than 2.7 g/cm3, the density of aluminum, the metal (is,is not) aluminum.
EXTRA PRACTICE (similar to Practice Problem 48a, page 92)
48a. Use dimensional analysis to convert 4.68 g of boron to cubic centimeters ofboron. The density of boron is 2.34 g/cm3.
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______________________
612
245
2.50
Density �mass�volume
� _____ g/cm3
no
density
612 g��_____ cm3
4.68 g of boron � �2.
c
3
m
43
g� � 4.68 g �
� 2 cm3 of boron
1 cm3
�2.34 g
05_Chem_GRSW_Ch03.SE/TE 6/11/04 3:29 PM Page 31
Leaming-Focused® Notebook ©2005 Learning Concepts
Reproduction of this page for classroom and planning use by owner of this book is permitted.
Student Learning Map Unit Topic: What is the Atom?
Key Learning: A. Atomic Theory B. Electrons C. Applying the Scientific Method.
Unit Essential Questions: A. How are the steps in the scientific method applied? B. Decribe how atomic theory changed through time. C. How is electronic structure (e-) described in an atom?
Concept: Concept: Concept: Concept:
Scientific Method
Atomic Theory Subatomic Particles Nuclear
Lesson EQ: How is the scientific method used?
What are the parts of the scientific method?
How is a scientific model used?
What is a theory?
What are the main parts of lab report?
Lesson EQ: What is Dalton’s atomic theory? How did the cathode ray and gold foil experiment effect the description of atomic structure? How did our understanding of the atom evolve?
How is the Bohr atom represented?
Lesson EQ: What is the difference between mass number and atomic mass?
What is an isotope?
What are the two naming conventions for isotopes?
What is an ion?
Lesson EQ: How is average atomic mass calculated? What is radioactivity? How are nuclear equations written? What is half life? What are the major components of a nuclear reactor?
Vocabulary: Scientific Method (parts)
Theory
Model
Independent Variable
Dependent Variable
Control
Vocabulary: Cathode Ray Tube
Gold Foil Experiment
Planetary Model
Vocabulary: Atomic number
Atomic mass
Vocabulary: Radioactivity
Radioactive Decay
Fission/Fusion
Instructional Tools
Chemical Concoctions Inc Lab
Atomic Theory
Cooperative Group Exercise
Chapter 4 Atomic Structure 33
SECTION 4.1 DEFINING THE ATOM (pages 101–103)
This section describes early atomic theories of matter and provides ways tounderstand the tiny size of individual atoms.
Early Models of the Atom (pages 101–102)
1. Democritus, who lived in Greece during the fourth century B.C.,
suggested that matter is made up of tiny particles that cannot be divided. He
called these particles ______________________ .
2. List two reasons why the ideas of Democritus were not useful in a scientific
sense.
3. The modern process of discovery about atoms began with the theories of an
English schoolteacher named ______________________ .
4. Circle the letter of each sentence that is true about Dalton’s atomic theory.
a. All elements are composed of tiny, indivisible particles called atoms.
b. An element is composed of several types of atoms.
c. Atoms of different elements can physically mix together, or can chemicallycombine in simple, whole-number ratios to form compounds.
d. Chemical reactions occur when atoms are separated, joined, or rearranged;however, atoms of one element are never changed into atoms of anotherelement by a chemical reaction.
5. In the diagram, use the labels mixture and compound to identify the mixtureof elements A and B and the compound that forms when the atoms ofelements A and B combine chemically.
Name ___________________________ Date ___________________ Class __________________
ATOMIC STRUCTURE4
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atoms
They did not explain chemical behavior, and they lacked experimental support
because Democritus’s approach was not based on the scientific method.
John Dalton
(a) Atoms of element A (b) Atoms of element B mixturecompound
05_Chem_GRSW_Ch04.SE/TE 6/11/04 3:29 PM Page 33
Name ___________________________ Date ___________________ Class __________________©
Pearson Education, Inc., publishing as Pearson Prentice Hall. All rights reserved.
34 Guided Reading and Study Workbook
CHAPTER 4, Atomic Structure (continued)
Sizing up the Atom (page 103)
6. Suppose you could grind a sample of the element copper into smaller and
smaller particles. The smallest particle that could no longer be divided, yet still
has the chemical properties of copper, is _________________________________ .
7. About how many atoms of copper when placed side by side would form a line
1 cm long? ______________________
SECTION 4.2 STRUCTURE OF THE NUCLEAR ATOM (pages 104–108)
This section describes the experiments that led to the discovery of subatomic particles and their properties.
Subatomic Particles (pages 104–106)
1. How is the atomic theory that is accepted today different from Dalton’s atomic
theory?
2. Which subatomic particles carry a negative charge? ______________________
Match each term from the experiments of J. J. Thomson with the correct description.
_______ 3. anode a. an electrode with a negative charge
_______ 4. cathode b. a glowing beam traveling between charged electrodes
_______ 5. cathode ray c. an electrode with a positive charge
_______ 6. electron d. a negatively charged particle
7. The diagram shows electrons moving from left to right in a cathode-ray tube.Draw an arrow showing how the path of the electrons will be affected by theplacement of the negatively and positively charged plates.
an atom of copper
100,000,000
Atoms are now known to be divisible. They can be broken down into even
smaller, more fundamental particles.
electrons
c
a
b
d
�
Cathode
Slit
Vacuum pump
Negativeplate
�
Positiveplate
High voltage
Anode
05_Chem_GRSW_Ch04.SE/TE 6/11/04 3:29 PM Page 34
Chapter 4 Atomic Structure 35
8. Thomson observed that the production of cathode rays did not depend on thekind of gas in the tube or the type of metal used for the electrodes. Whatconclusion did he draw from these observations?
9. What two properties of an electron did Robert Millikan determine from hisexperiments?
10. Circle the letter of each sentence that is true about atoms, matter, and electriccharge.
a. All atoms have an electric charge.
b. Electric charges are carried by particles of matter.
c. Electric charges always exist in whole-number multiples of a single basic unit.
d. When a given number of positively charged particles combines with anequal number of negatively charged particles, an electrically neutralparticle is formed.
11. Circle the letter next to the number of units of positive charge that remain if ahydrogen atom loses an electron.
a. 0 b. 1 c. 2 d. 3
12. The positively charged subatomic particle that remains when a hydrogen atom
loses an electron is called ______________________ .
13. What charge does a neutron carry? ______________________ .
14. Complete the table about the properties of subatomic particles.
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He concluded that electrons must be parts of the atoms of all elements.
He determined the quantity of charge carried by an electron and its mass.
a proton
none
Properties of Subatomic Particles
Particle Symbol Relative Electrical Charge Relative Mass Actual Mass(mass of proton � 1) (g)
Electron e– 9.11 � 10–28
Proton p+ 1.67 � 10–24
Neutron n0 1.67 � 10–24
1�18401�
1� 1
10
05_Chem_GRSW_Ch04.SE/TE 6/11/04 3:29 PM Page 35
36 Guided Reading and Study Workbook
The Atomic Nucleus (pages 106–108)
15. Is the following sentence true or false? An alpha particle has a double positive
charge because it is a helium atom that has lost two electrons. ______________________
16. Explain why in 1911 Rutherford and his coworkers were surprised when theyshot a narrow beam of alpha particles through a thin sheet of gold foil.
17. Circle the letter of each sentence that is true about the nuclear theory of atomssuggested by Rutherford’s experimental results.
a. An atom is mostly empty space.
b. All the positive charge of an atom is concentrated in a small central regioncalled the nucleus.
c. The nucleus is composed of protons.
d. The nucleus is large compared with the atom as a whole.
e. Nearly all the mass of an atom is in its nucleus.
SECTION 4.3 DISTINGUISHING AMONG ATOMS (pages 110–119)
This section explains how atomic number identifies an element; how to useatomic number and mass number to find the number of protons, neutrons,and electrons in an atom; how isotopes differ; and how to calculate averageatomic mass.
Atomic Number (page 110)
1. Circle the letter of the term that correctly completes the sentence. Elementsare different because their atoms contain different numbers of _______ .
a. electrons
b. protons
c. neutrons
d. nuclei
CHAPTER 4, Atomic Structure (continued)
Name ___________________________ Date ___________________ Class __________________©
Pearson Education, Inc., publishing as Pearson Prentice Hall. All rights reserved.
true
They expected most of the alpha particles to pass easily through the foil with only a
slight deflection. They found that most alpha particles were not deflected at all, and
a few were deflected at very large angles.
05_Chem_GRSW_Ch04.SE/TE 6/11/04 3:29 PM Page 36
Chapter 4 Atomic Structure 37
2. Complete the table showing the number of protons and electrons in atoms ofsix elements.
Mass Number (pages 111–112)
3. The total number of protons and neutrons in an atom is its
______________________ .
4. What is the mass number of a helium atom that has two protons and two
neutrons? ______________________
5. How many neutrons does a beryllium atom with four protons and a mass
number of nine have? ______________________
6. Place the labels chemical symbol, atomic number, and mass number in theshorthand notation below.
7. Designate the atom shown in Question 6 in the form “name of element”-“mass
number.” ______________________
8. How many protons, neutrons, and electrons are in the atom discussed in
Questions 6 and 7? Protons: Neutrons: Electrons:
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Atoms of Six Elements
Name Symbol Atomic Number Number of Protons Number of Electrons
Hydrogen H 1
Helium He 2
Lithium Li 3
Boron B 5
Carbon C 6
Oxygen O 8
2
1 1
2
5
3
6
3
5
6
8 8
mass number
4
5
neon-22
10 12 10
22
Ne10
mass number
atomic number
chemical symbol
05_Chem_GRSW_Ch04.SE/TE 6/11/04 3:29 PM Page 37
38 Guided Reading and Study Workbook
Isotopes (pages 112–113)
9. How do atoms of neon-20 and neon-22 differ?
10. Neon-20 and neon-22 are called ______________________ .
11. Is the following sentence true or false? Isotopes are chemically alike because
they have identical numbers of protons and electrons. ______________________
Match the designation of each hydrogen isotope with its commonly used name.
_______ 12. hydrogen-1 a. tritium
_______ 13. hydrogen-2 b. hydrogen
_______ 14. hydrogen-3 c. deuterium
Atomic Mass (pages 114–117)
15. Why is the atomic mass unit (amu), rather than the gram, usually used toexpress atomic mass?
16. What isotope of carbon has been chosen as the reference isotope for atomicmass units? What is the defined atomic mass in amu of this isotope?
17. Is the following sentence true or false? The atomic mass of an element is
always a whole number of atomic mass units. ______________________
18. Circle the letter of each statement that is true about the average atomic massof an element and the relative abundance of its isotopes.
a. In nature, most elements occur as a mixture of two or more isotopes.
b. Isotopes of an element do not have a specific natural percent abundance.
c. The average atomic mass of an element is usually closest to that of theisotope with the highest natural abundance.
d. Because hydrogen has three isotopes with atomic masses of about 1 amu, 2 amu, and 3 amu, respectively, the average atomic mass of naturalhydrogen is 2 amu.
19. Circle the letter of the correct answer. When chlorine occurs in nature, thereare three atoms of chlorine-35 for every one atom of chlorine-37. Whichatomic mass number is closer to the average atomic mass of chlorine?
a. 35 amu b. 37 amu
CHAPTER 4, Atomic Structure (continued)
Name ___________________________ Date ___________________ Class __________________©
Pearson Education, Inc., publishing as Pearson Prentice Hall. All rights reserved.
An atom of neon-22 has two more neutrons in its nucleus than an atom of neon-20.
isotopes
true
b
c
a
carbon-12; 12 amu
false
The values of atomic masses measured in grams are inconveniently small and
impractical to work with.
05_Chem_GRSW_Ch04.SE/TE 6/11/04 3:29 PM Page 38
Chapter 4 Atomic Structure 39
20. In the periodic table, the elements are organized into groups based on
______________________________________________________________ .
The Periodic Table—A Preview (page 118)
21. What are the horizontal rows in the periodic table called?
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Outlining can help you understand and remember what you have read. Prepare an outline of Section 4.3, Distinguishing Among Atoms. Begin with the headings in the textbook. Under eachheading, write the main idea. Then list the details that support the main idea. Do your work on aseparate sheet of paper.
Reading Skill Practice
Main heads should be Atomic Number, Mass Number, Isotopes, and Atomic Mass. Subheads shouldindicate an understanding of each concept.
periods
a set of repeating properties
05_Chem_GRSW_Ch04.SE/TE 6/11/04 3:29 PM Page 39
40 Guided Reading and Study Workbook
GUIDED PRACTICE PROBLEMSFill in the write-on lines and boxes provided as you work through the guided practiceproblems.
GUIDED PRACTICE PROBLEM 18 (page 112)
18. Use Table 4.2 to express the compositions of carbon-12, fluorine-19, andberyllium-9 in shorthand notation.
AnalyzeCarbon-12
Step 1. The number of protons in an atom is called its ______________________ number.
The number of protons in an atom of carbon-12 is ___________ .
Calculate
Step 2. The number of protons plus the number of neutrons in an atom is called
its ______________________ number. For carbon-12, this number is ___________ .
Step 3. The shorthand notation for carbon-12 is:
Evaluate
Step 4. Except for hydrogen-1, the mass number of an isotope is always greater
than its atomic number. Is the mass number reasonable? ____________
Fluorine-19
Step 1. The atomic number of fluorine-19 is ____________ .
Step 2. Its mass number is ____________ .
Step 3. The shorthand notation for fluorine-19 is:
Step 4. Is your answer reasonable? Why?
mass number
Catomic number
CHAPTER 4, Atomic Structure (continued)
Name ___________________________ Date ___________________ Class __________________©
Pearson Education, Inc., publishing as Pearson Prentice Hall. All rights reserved.
9
19
yes
Yes it is, because the mass number is greater than the atomic number.
F19
9
atomic
6
mass 12
12
6
05_Chem_GRSW_Ch04.SE/TE 6/11/04 3:29 PM Page 40
Chapter 4 Atomic Structure 41
Beryllium-9
Step 1. The atomic number of beryllium-9 is ____________ .
Step 2. Its mass number is ____________ .
Step 3. The shorthand notation for beryllium-9 is:
Step 4. Is your answer reasonable? Why?
EXTRA PRACTICE (similar to Practice Problem 19, page 113)
19. Three isotopes of sulfur are sulfur-32, sulfur-33, and sulfur-34. Write thecomplete symbol for each isotope, including the atomic number and the mass number.
GUIDED PRACTICE PROBLEM 23 (page 117)
23. The element copper has naturally occurring isotopes with mass numbers of 63 and 65. The relative abundance and atomic masses are 69.2% for mass �
62.93 amu and 30.8% for mass � 64.93 amu. Calculate the average atomicmass of copper.
Analyze
Step 1. Will the average atomic mass be closer to 63 or to 65? Explain.
Solve
Step 2. For Cu-63: 69.2% � 62.93 amu � 0.692 � 62.93 amu �
Step 3. For Cu-65: 30.8% � 64.93 amu � � �
Step 4. Average mass: 43.6 amu � �
Evaluate
Step 5. Explain why your answer is reasonable.
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4
9
Yes it is, because the mass number is greater than the atomic number.
Be9
4
Ssulfur-32
Ssulfur-33
Ssulfur-34
32 33 34
16 16 16
Because about 70% of the element consists of the lighter isotope, the average atomic
mass should be closer to the atomic mass number of 63 than 65.
It is closer to 63 than 65.
43.6 amu
0.308 20.0 amu
20.0 amu 63.6 amu
64.93 amu
05_Chem_GRSW_Ch04.SE/TE 6/11/04 3:29 PM Page 41
42 Guided Reading and Study Workbook
EXTRA PRACTICE (similar to Practice Problem 24, page 117)
24. Calculate the atomic mass of rubidium. The two isotopes of rubidium have atomic masses and relative abundancies of 84.91 amu (72.16%) and 86.91 amu (27.84%). _________________________
CHAPTER 4, Atomic Structure (continued)
Name ___________________________ Date ___________________ Class __________________©
Pearson Education, Inc., publishing as Pearson Prentice Hall. All rights reserved.
85.47 amu
05_Chem_GRSW_Ch04.SE/TE 6/11/04 3:29 PM Page 42
Leaming-Focused® Notebook ©2005 Learning Concepts
Reproduction of this page for classroom and planning use by owner of this book is permitted.
Student Learning Map Unit Topic: How do electrons behave in an atom?
Key Learning: A. Electron Configurations B. Electromagnetic Spectrum
Unit Essential Questions: A. How is electronic structure (e-) described in an atom? B. Describe the dual wave-particle nature of electrons.
Concept: Concept: Concept: Concept:
Wave-Particle Duality
Electron Configurations
Atomic Charge
Lesson EQ: What is EMS?
What is the relationship between wavelength and frequency?
What is the quantum mechanical model of atomic structure?
What are electron orbitals?
Lesson EQ: How is the Bohr atom represented?
What rules do electrons follow when filling their orbitals?
How are electron configurations used to describe location of electrons?
Lesson EQ: What is a cation?
What is an anion?
How are ions formed?
Lesson EQ:
Vocabulary: Plank’s constant
Wavelength
Frequency
Photoelectric effect
Quantum numbers
Vocabulary: Bohr atom
Aufbau Principle
Pauli Exclusion Principle
Hund’s Rule
Vocabulary: Cation
Anion
Valence Electrons
Vocabulary:
Instructional Tools
Atomic Emission Spectra Flame Tests
Lab
Hog Hotel Activity
ELECTRONS IN ATOMS5
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Name ____________________________________ Date __________ Class ___________________
Chapter 5 Electrons in Atoms 43
SECTION 5.1 MODELS OF THE ATOM (pages 127–132)
This section summarizes the development of atomic theory. It also explains the significance of quantized energies of electrons as they relate to the quan–tum mechanical model of the atom.
The Development of Atomic Models (pages 127–128)
1. Complete the table about atomic models and the scientists who developedthem.
2. Is the following sentence true or false? The electrons in an atom can exist
between energy levels. ______________________
The Bohr Model (pages 128–129)
3. What is a small, discrete unit of energy called?
4. Circle the letter of the term that completes the sentence correctly. A quantumof energy is the amount of energy required to
a. move an electron from its present energy level to the next lower one
b. maintain an electron in its present energy level
c. move an electron from its present energy level to the next higher one
5. In general, the higher the electron is on the energy ladder, the
______________________ it is from the nucleus.
Scientist Model of Atom
Dalton
Thomson
Rutherford
Bohr
The atom is a solid indivisible mass.
The atom is a ball of positive charge with electrons stuck into the ball.
Electrons are arranged in concentric circular paths around the nucleus.
Most of an atom’s mass is concentrated in the small, positively chargednucleus. The electrons surround the nucleus and the rest of the atom isempty space.
false
farther
It is called a quantum.
05_Chem_GRSW_Ch05.SE/TE 6/11/04 3:30 PM Page 43
Name ___________________________ Date ___________________ Class __________________©
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44 Guided Reading and Study Workbook
CHAPTER 5, Electrons in Atoms (continued)
The Quantum Mechanical Model (page 130)
6. What is the difference between the previous models of the atom and the modern
quantum mechanical model? __________________________________________________
7. Is the following sentence true or false? The quantum mechanical model of theatom estimates the probability of finding an electron in a certain position.
______________________
Atomic Orbitals (pages 131–132)
8. Circle the letter of the term that correctly answers this question. Which namedescribes the major energy levels of electrons?
a. atomic orbitals b. quantum mechanical numbers
c. quantas d. principal quantum numbers (n)
9. Principal energy levels are assigned values in order of ______________________energy: n � 1, 2, 3, 4, and so forth.
10. In the quantum mechanical model the regions where electrons are likely
to be found are called ______________________ and are denoted by
______________________ .
11. Match each diagram below with the name of its p orbital, the px, py, or pz.
12. Use the diagram above. Describe how the px, py, and pz orbitals are similar.
13. Describe how the px, py, and pz orbitals are different. _____________________________
14. Circle the letter of the formula for the maximum number of electrons that canoccupy a principal energy level. Use n for the principal quantum number.
a. 2n2 b. n2 c. 2n d. n
y
x
z
y
x
z
p orbitals
y
x
z
Previous models described the motion of electrons the
same way as the motion of large objects. The quantum mechanical model is not based
on the exact path an electron follows around the nucleus.
true
letters
atomic orbitals
generally increasing
The p orbitals are similar because they are all dumbbell shaped.
The p orbitals have different
orientations in space. They are all perpendicular to each other.
px py pz
05_Chem_GRSW_Ch05.SE/TE 6/11/04 3:30 PM Page 44
Chapter 5 Electrons in Atoms 45
SECTION 5.2 ELECTRON ARRANGEMENT IN ATOMS (pages 133–136)
This section shows you how to apply the aufbau principle, the Pauli exclusionprinciple, and Hund’s rule to help you write the electron configurations of elements. It also explains why the electron configurations for some elementsdiffer from those assigned using the aufbau principle.
Electron Configurations (pages 133–135)
1. The ways in which electrons are arranged around the nuclei of atoms are
called ___________________________________________ .
Match the name of the rule used to find the electron configurations of atoms withthe rule itself.
_______ 2. aufbau principle
_______ 3. Pauli exclusion principle
_______ 4. Hund’s rule
5. Look at the aufbau diagram, Figure 5.7 on page 133. Which atomic orbital
is of higher energy, a 4f or a 5p orbital? _______
6. Fill in the electron configurations for the elements given in the table. Use the orbital filling diagrams to complete the table.
1s1
1s22s1
1s22s22p3
1s22s22p5
1s22s22p63s1
He
C
O
Ne
1sElement
Electron Configurations for Some Selected Elements
Orbital filling
2s 2px 2py 2pz 3sElectron configuration
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electron configurations
b
c
a
4f
a. When electrons occupy orbitals ofequal energy, one electron enters eachorbital until all the orbitals containone electron with parallel spins.
b. Electrons enter orbitals of lowestenergy first.
c. An atomic orbital may describe atmost two electrons.
H
1s 2
1s 22s 22p 2
1s 22s 22p4
1s 22s 22p6
Li
N
F
Na
05_Chem_GRSW_Ch05.SE/TE 6/11/04 3:30 PM Page 45
46 Guided Reading and Study Workbook
7. In the shorthand method for writing an electron configuration, what does asuperscript stand for?
8. In the shorthand method for writing an electron configuration, what does thesum of the superscripts equal?
Exceptional Electron Configurations (page 136)
9. Is the following sentence true or false? The aufbau principle works for
every element in the periodic table. ______________________
10. Filled energy sublevels are more ______________________ than partially filledsublevels.
11. Half-filled levels are not as stable as ______________________ levels, but aremore stable than other configurations.
SECTION 5.3 PHYSICS AND THE QUANTUM MECHANICALMODEL (pages 138–146)
This section explains how to calculate the wavelength, frequency, or energy of light, given two of these values. It also explains the origin of the atomicemission spectrum of an element.
Light (pages 138–140)
1. Match each term describing waves to its definition.
_______ amplitude a. the distance between two crests
_______ wavelength b. the wave’s height from the origin to the crest
_______ frequency c. the number of wave cycles to pass a given point perunit of time
CHAPTER 5, Electrons in Atoms (continued)
Name ___________________________ Date ___________________ Class __________________©
Pearson Education, Inc., publishing as Pearson Prentice Hall. All rights reserved.
The superscript stands for the number of electrons occupying a given sublevel.
The sum equals the number of electrons in the atom.
false
stable
filled
Outlining can help you understand and remember what you have read. Prepare an outline of Section5.2, Electron Arrangement in Atoms. Begin your outline by copying the headings from the textbook.Under each heading, write the main idea. Then list the details that support, or back up, the main idea.Do your work on a separate sheet of paper.
Reading Skill Practice
The subheads of students’ outlines of the section should be Electron Configurations and ExceptionalElectron Configurations. The section’s main ideas should form the next level of the outline.
b
a
c
05_Chem_GRSW_Ch05.SE/TE 6/11/04 3:30 PM Page 46
Chapter 5 Electrons in Atoms 47
2. The units of frequency are usually cycles per second. The SI unit of
cycles per second is called a(n) ______________________ .
3. Label the parts of a wave in this drawing. Label the wavelength, the amplitude,the crest, and the origin.
4. The product of wavelength and frequency always equals a(n)
______________________ , the speed of light.
5. Is the following sentence true or false? The wavelength and frequency of all
waves are inversely proportional to each other. ______________________
6. Light consists of electromagnetic waves. What kinds of visible and invisibleradiation are included in the electromagnetic spectrum?
7. When sunlight passes through a prism, the different wavelengths separate into
a(n) ______________________ of colors.
8. Put the visible colors in order of frequency.
_______ orange _______ violet
_______ green _______ yellow
_______ blue _______ red
9. Look at Figure 5.10 on page 139. The electromagnetic spectrum consists ofradiation over a broad band of wavelengths. What type of radiation has thelowest frequency? The highest frequency?
Atomic Spectra (page 141)
10. What happens when an electric current is passed through the gas or vapor ofan element?
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The following kinds of radiation are included in the electromagnetic spectrum: radio
waves, microwaves, infrared waves, visible light, ultraviolet waves, x-ray, and gamma
rays.
Wavelength λ
Amplitude Origin
Crest
true
constant
hertz (Hz)
spectrum
2
4
5
6
3
1
Radio waves have the lowest frequency. Gamma rays have the highest frequency.
The element emits light as it is excited by the passage of the electric current.
05_Chem_GRSW_Ch05.SE/TE 6/11/04 3:30 PM Page 47
48 Guided Reading and Study Workbook
11. Passing the light emitted by an element through a prism gives the
_________________________________________________ of the element.
12. Is the following sentence true or false? The emission spectrum of an element can be the same as the emission spectrum of another element.
______________________
An Explanation of Atomic Spectra (pages 142–143)
13. What is the lowest possible energy of an electron called? _____________________________
14. Only electrons moving from ______________________ to
______________________ energy levels lose energy and emit light.
Quantum Mechanics (pages 381–382)
15. What did Albert Einstein call the quanta of energy that is light?
______________________
16. What did de Broglie’s equation predict about the behavior of particles?
17. Is the following sentence true or false? The new method of describing the
motions of subatomic particles, atoms, and molecules is called quantum
mechanics. ______________________
18. Is the following sentence true or false? de Broglie’s conclusions were supported
by experimental evidence. ______________________
19. Does the Heisenberg uncertainty principle apply to cars and airplanes?
CHAPTER 5, Electrons in Atoms (continued)
Name ___________________________ Date ___________________ Class __________________©
Pearson Education, Inc., publishing as Pearson Prentice Hall. All rights reserved.
false
photons
true
ground state
higher
lower
It predicts that all moving objects have wavelike motions.
true
No, only to small particles.
atomic emission spectrum
05_Chem_GRSW_Ch05.SE/TE 6/11/04 3:30 PM Page 48
Chapter 5 Electrons in Atoms 49
GUIDED PRACTICE PROBLEM
GUIDED PRACTICE PROBLEM 14 (page 140)
14. What is the wavelength of radiation with a frequency of 1.50�1013 Hz (1.50 �1013 s–1 )? Does this radiation have a longer or shorter wavelength than red light?
Analyze
Step 1. What is the equation for the relationship between frequency and
wavelength? ______________________
Step 2. What does c represent and what is its value?
Step 3. What is the wavelength of red light in cm?
Solve
Step 4. Solve the equation for the unknown. λ � ___________
Step 5. Substitute the known quantities into the equation and solve.
3.00 �108 m/s �
Step 6. Compare the answer with the wavelength of red light. Does the givenradiation have a wavelength longer or shorter than that of red light?
Evaluate
Step 7. Explain why you think your result makes sense?
Step 8. Are the units in your answer correct? How do you know?
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c � λυ
It represents the speed of light, which is 3.00 � 108 m/s.
Red light has a wavelength of about 700 nm, which can be converted to m by
dividing by 109 nm/m, producing 7 � 10–7 m.
1.50 � 1013 s
2.00 �10–5 m
The answer, 2.00 � 10–5 m, is greater than the wavelength of red light, 7 � 10–7 m.
Therefore, this radiation has a longer wavelength than red light.
The magnitude of the frequency is about 105 times the magnitude of c. Because
frequency and wavelength are inversely proportional to each other, the answer
should be about 10–5.
Yes, because wavelength is measured in meters or fractions of a meter.
c�υ
05_Chem_GRSW_Ch05.SE/TE 6/11/04 3:30 PM Page 49
Leaming-Focused® Notebook ©2005 Learning Concepts
Reproduction of this page for classroom and planning use by owner of this book is permitted.
Student Learning Map Unit Topic: How is the Periodic Table be used?
Key Learning: A. Periodic Table B. Periodic Trends
Unit Essential Questions: A. How is the periodic table organized? B. What is a periodic trend?
Concept: Concept: Concept: Concept:
History of the Periodic Table
Periodic Table Organization
Trends
Lesson EQ: Compare Mendeleev’s Periodic Table to a shoe store.
How has the periodic table changed through time? How has it stayed the same?
Lesson EQ: What are the major groupings on the periodic table?
What is periodicity and how can it be used?
Lesson EQ: How does electron configuration explain periodic trends?
What is shielding?
What is the trend on the periodic table for atomic radius, ionization energy, electron affinity, and electronegativity?
Lesson EQ:
Vocabulary: Mendeleev
Moseley
Periodic law
Noble Gases
Lanthanides
Actinides
Vocabulary: Periodicity
Blocks
Group
Family
Row
Period
Vocabulary: Periodic Trend
Radii
Shielding Effect
Ionization Energy
Electronegativity
Vocabulary:
Instructional Tools
Alien Zoo Activity
Density Trends in a Group Lab
Chapter 6 The Periodic Table 51
SECTION 6.1 ORGANIZING THE ELEMENTS (pages 155–160)This section describes the development of the periodic table and explains the periodic law.It also describes the classification of elements into metals, nonmetals, and metalloids.
Searching For An Organizing Principle (page 155)
1. How many elements had been identified by the year 1700? ________________
2. What caused the rate of discovery to increase after 1700?
3. What did chemists use to sort elements into groups?
Mendeleev’s Periodic Table (page 156)
4. Who was Dmitri Mendeleev? __________________________________________________
5. What property did Mendeleev use to organize the elements into a periodic table?
6. Is the following sentence true or false? Mendeleev used his periodic table to
predict the properties of undiscovered elements. ______________________
The Periodic Law (page 157)
7. How are the elements arranged in the modern periodic table?
8. Is the following statement true or false? The periodic law states that when
elements are arranged in order of increasing atomic number, there is a
periodic repetition of physical and chemical properties. ______________________
Metals, Nonmetals, and Metalloids (pages 158–160)
9. Explain the color coding of the squares in the periodic table in Figure 6.5.
Name ___________________________ Date ___________________ Class __________________
THE PERIODIC TABLE6
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13
Dmitri Mendeleev was a Russian chemist and teacher
who developed a periodic table of elements.
Mendeleev arranged the elements in order of increasing atomic mass.
true
Yellow squares contain metals, blue squares contain nonmetals, green squares
contain metalloids.
Chemists began to use scientific methods to search for elements.
Chemists used the properties of elements.
The elements are arranged in order by increasing atomic number.
true
05_Chem_GRSW_Ch06.SE/TE 6/12/04 9:57 AM Page 51
Name ___________________________ Date ___________________ Class __________________©
Pearson Education, Inc., publishing as Pearson Prentice Hall. All rights reserved.
52 Guided Reading and Study Workbook
CHAPTER 6, The Periodic Table (continued)
10. Which property below is not a general property of metals.
a. ductile c. malleable
b. poor conductor of heat d. high luster
11. Is the following statement true or false? The variation in properties among metals
is greater than the variation in properties among nonmetals. __________________
12. Under some conditions, a metalloid may behave like a __________________ .
Under other conditions, a metalloid may behave like a __________________ .
SECTION 6.2 CLASSIFYING THE ELEMENTS (pages 161–167)
This section explains why you can infer the properties of an element based on the properties of other elements in the periodic table. It also describes the use of electron configurations to classify elements.
Squares In The Periodic Table (pages 161–163)
1. Label the sample square from the periodic table below. Use the labels elementname, element symbol, atomic number, and average atomic mass.
2. List three things, other than the name, symbol, atomic number, and averageatomic mass, you can discover about an element using the periodic table inFigure 6.9.
a. _________________________________________________________________
b. _________________________________________________________________
c. _________________________________________________________________
Electron Configurations In Groups (pages 164–165)
3. Is the following sentence true or false? The subatomic particles that play the key
role in determining the properties of an element are electrons. ________________
4. Why are Group A elements called representative elements?
12
MgMagnesium
24.305
They exhibit a wide range of physical and chemical properties.
state at room temperature
electrons in each energy level
whether an element is found in nature
true
atomic number
average atomic mass
element symbol
element name
false
metal
nonmetal
05_Chem_GRSW_Ch06.SE/TE 6/12/04 9:57 AM Page 52
Chapter 6 The Periodic Table 53
5. Classify each of the following elements as a (an) alkali metal, alkaline earthmetal, halogen, or noble gas.
a. sodium ______________________ e. xenon ______________________
b. chlorine ______________________ f. potassium ______________________
c. calcium ______________________ g. magnesium ______________________
d. fluorine ______________________
6. For elements in each of the following groups, how many electrons are in thehighest occupied energy level?
a. Group 3A ____________________
b. Group 1A ____________________
c. Group 8A ____________________
Transition Elements (page 166)
7. Complete the table about classifying elements according to the electronconfiguration of their highest occupied energy level.
8. Circle the letter of the elements found in the p block.
a. Groups 1A and 2A and helium
b. Groups 3A, 4A, 5A, 6A, 7A, and 8A except for helium
c. transition metals
d. inner transition metals
Match the category of elements with an element from that category.
_______ 9. Noble gases a. gallium
_______ 10. Representative elements b. nobelium
_______ 11. Transition metals c. argon
_______ 12. Inner transition metals d. vanadium
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alkali metal
halogen
alkaline earth metal
halogen
noble gas
alkali metal
alkaline earth metal
Category Description of Electron Configuration
Noble gases
Representative elements
s sublevel and nearby d sublevel contain electrons
s sublevel and nearby f sublevel contain electrons
s or p sublevels are filled
s or p sublevels are only partially filled
Transition metals
Inner transition metals
3
1
8
c
a
d
b
05_Chem_GRSW_Ch06.SE/TE 6/12/04 9:57 AM Page 53
54 Guided Reading and Study Workbook
13. Use Figure 6.12 on page 166. Write the electron configurations for thefollowing elements.
a. magnesium _____________________________
b. cobalt __________________________________
c. sulfur __________________________________
SECTION 6.3 PERIODIC TRENDS (pages 170–178)
This section explains how to interpret group trends and periodic trends inatomic size, ionization energy, ionic size, and electronegativity.
Trends in Atomic Size (pages 170–171)
1. Is the following sentence true or false? The radius of an atom can be measured
directly. ______________________
2. What are the atomic radii for the following molecules?
__________ __________ __________ __________
3. What is the general trend in atomic size within a group? Within a period?
4. What are the two variables that affect atomic size within a group?
a. _________________________________________________________________
b. _________________________________________________________________
5. For each pair of elements, pick the element with the largest atom.
a. Helium and argon __________________________
b. Potassium and argon __________________________
Ions (page 172)
6. What is an ion?
Hydrogenatomic radius �
Oxygenatomic radius �
Nitrogenatomic radius �
Chlorineatomic radius �
CHAPTER 6, The Periodic Table (continued)
Name ___________________________ Date ___________________ Class __________________©
Pearson Education, Inc., publishing as Pearson Prentice Hall. All rights reserved.
1s22s22p63s2
1s22s22p63s23p63d74s2
1s22s22p63s23p4
false
30 pm 68 pm 70 pm 102 pm
The atomic size increases within a group as atomic number increases. The atomic
size decreases from left to right across a period.
An ion is an atom or group of atoms that has a positive or negative charge.
the charge on the nucleus
the number of occupied energy levels
argon
potassium
05_Chem_GRSW_Ch06.SE/TE 6/12/04 9:57 AM Page 54
Chapter 6 The Periodic Table 55
7. How are ions formed?
8. An ion with a positive charge is called a(n) ______________________ ; an ion with
a negative charge is called a(n) ______________________ .
9. Complete the table about anions and cations.
Trends in Ionization Energy (pages 173–175)
10. ______________________ is the energy required to overcome the attraction ofprotons in the nucleus and remove an electron from a gaseous atom.
11. Why does ionization energy tend to decrease from top to bottom within a group?
12. Why does ionization energy tend to increase as you move across a period?
13. There is a large increase in ionization energy between the second and the third ionizationenergies of a metal. What kind of ion is the metal likely to form? Include the charge inyour answer.
Trends in Ionic Size (page 176)
14. Metallic elements tend to ______________________ electrons and form
______________________ ions.
Nonmetallic elements tend to ______________________ electrons and
form ______________________ ions.
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anion
cation
Anions Cations
Charge
Metal/Nonmetal
Minus sign/Plus sign
negative
nonmetal
plus sign
positive
metal
minus sign
An ion is formed when electrons are transferred between atoms.
Ionization energy
The nuclear charge increases across a period but the shielding effect remains
constant. There is greater attraction of the electrons to the nucleus and more
energy is required to remove an electron. Atomic size increases from top to bottom
within the group.
an ion with a 2� charge
Atomic size increases from top to bottom within the group. The nuclear charge has a
smaller effect on the electrons in the highest occupied energy level and less energy
is required to remove an electron.
positive
lose
gain
negative
05_Chem_GRSW_Ch06.SE/TE 6/12/04 9:57 AM Page 55
56 Guided Reading and Study Workbook
15. Circle the letter of the statement that is true about ion size.
a. Cations are always smaller than the neutral atoms from which they form.
b. Anions are always smaller than the neutral atoms from which they form.
c. Within a period, a cation with a greater charge has a larger ionic radius.
d. Within a group, a cation with a higher atomic number has a smaller ionic radius.
16. Which ion has the larger ionic radius: Ca2+or Cl– ? ________
Trends in Electronegativity (page 177)
17. What property of an element represents its tendency to attract electrons when
it chemically combines with another element? _____________________________
18. Use Table 6.2 on page 177. What trend do you see in the relativeelectronegativity values of elements within a group? Within a period?
19. Circle the letter of each statement that is true about electronegativity values.
a. The electronegativity values of the transition elements are all zero.
b. The element with the highest electronegativity value is sodium.
c. Nonmetals have higher electronegativity values than metals.
d. Electronegativity values can help predict the types of bonds atoms form.
Summary of Trends (page 178)
20. Use Figure 6.22 on page 178. Circle the letter of each property for whichaluminum has a higher value than silicon.
a. first ionization energy c. electronegativity
b. atomic radius d. ionic radius
CHAPTER 6, The Periodic Table (continued)
Name ___________________________ Date ___________________ Class __________________©
Pearson Education, Inc., publishing as Pearson Prentice Hall. All rights reserved.
Cl–
The electronegativity values decrease as you move down a group, but increase as
you move across a period.
electronegativity
Students’ graphs should show a trend of increasing electronegativity values within a period as atomicnumber increases, and a dramatic decrease in the electronegativity value between the Group 7Aelement in one period and the group 1A element in the next period.
A graph can help you understand comparisons of data at a glance. Use graph paper to make a graphof the data in Table 6.2 on page 177. Plot electronegativity values on the vertical axis. Use a rangefrom 0 to 4. Plot atomic number on the horizontal axis. Label each period and the first element ineach period.
Reading Skill Practice
05_Chem_GRSW_Ch06.SE/TE 6/12/04 9:57 AM Page 56
Chapter 6 The Periodic Table 57
GUIDED PRACTICE PROBLEM
GUIDED PRACTICE PROBLEM 8 (page 167)
8. Use Figure 6.9 and Figure 6.12 to write the electron configurations of theseelements.
a. carbon b. strontium c. vanadium
Analyze
a. What is the number of electrons for each element?
C _______ Sr _______ V _______
b. What is the highest occupied energy sublevel for each element, accordingto its position on the periodic table? Remember that the energy level for thed block is always one less than the period.
C _______ Sr _______ V _______
c. According to its position on the periodic table, how many electrons doeseach element have in the sublevel listed above?
C _______ Sr _______ V _______
Solve
d. Begin filling in electron sublevels. Start from the top left and move right across each period in Figure 6.12 until you reach the highest occupied sub-level for each element. Make sure the d-block is in the correct energy level.
C _________________________ Sr ____________________________________________
V _____________________________________________
e. How can you check whether your answers are correct?
f. Check your answers as outlined above.
C ________________________________________________________________________
Sr ________________________________________________________________________
V ________________________________________________________________________
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Add all the superscripts in the electron configurations. This sum should be equal
to the atomic number for that element.
2�2 � 2 � 6, equal to the atomic number
2 � 2 � 6 � 2 � 6 � 3 � 2 � 23, equal to the atomic number
2 � 2 � 6 � 2 � 6 � 10 � 2 � 6 � 2 � 38, equal to the atomic number
1s 22s22p2
1s22s22p63s23p63d 34s2
1s22s22p63s23p63d104s24p65s2
6
2p
2 2 3
3d5s
2338
05_Chem_GRSW_Ch06.SE/TE 6/12/04 9:57 AM Page 57
Leaming-Focused® Notebook ©2005 Learning Concepts
Reproduction of this page for classroom and planning use by owner of this book is permitted.
Student Learning Map Unit Topic: How do atoms bond together to form molecules and compounds?
Key Learning(s): A. Atom Bonding B. Molecular Shapes
Unit Essential Question: A. What are the main bonding theories? B. What do compounds look like three dimensionally?
Concept: Concept: Concept: Concept:
Electron Dot & Lewis Structure, VESPR
Ionic Bonds Metal + Nonmetal
Covalent Bonds Nonmetal + Nonmetal
Metallic Bonds Metal + Metal
Lesson EQ:
How many valence electrons are in an element? How are electrons in a covalent bond diagrammed? How are multiple bonds diagrammed? How is molecular shape determined?
Lesson EQ:
How are ionic bonds formed? How are ionic compound formulas named and written?
Lesson EQ:
How are covalent bonds formed? How are covalent molecules/compound formulas written? What is polarity?
Lesson EQ:
How are metallic bonds formed? How are metal compound formulas written? What is an alloy and how are they used?
Vocabulary: Valence Electrons
Octet Rule
Electron Dot Structure
Lewis Structure
resonance
Vocabulary: Ionic Bonds
Cation
Anion
Polyatomic ions
Empirical formula
Vocabulary: Covalent Bonds
Polarity
Hybrid Orbital
Molecular formula
Vocabulary: Alloy
Instructional Tools
Conductivity Demonstration
Atom Modeling
Chapter 7 Ionic and Metallic Bonding 59
SECTION 7.1 IONS (pages 187–193)
This section explains how to use the periodic table to infer the number ofvalence electrons in an atom and draw its electron dot structure. It alsodescribes the formation of cations from metals and anions from nonmetals.
Valence Electrons (pages 187–188)
1. What are valence electrons?
2. The valence electrons largely determine the __________________________ of an
element and are usually the only electrons used in ________________________ .
3. Is the following sentence true or false? The group number of an element in the
periodic table is related to the number of valence electrons it has. _________________
4. What is an electron dot structure?
5. Draw the electron dot structure of each of the following atoms.
a. argon
b. calcium
c. iodine
The Octet Rule (page 188)
6. What is the octet rule?
7. Metallic atoms tend to lose their valence electrons to produce a(n) ______________ ,
or a positively charged ion. Most nonmetallic atoms achieve a complete octet by
gaining or ________________ electrons.
Name ___________________________ Date ___________________ Class __________________
IONIC AND METALLIC BONDING7
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Valence electrons are the electrons in the highest
occupied energy level of an element’s atoms.
chemical properties
chemical bonds
true
It is a diagram that shows the valence
electrons as dots.
In forming compounds, atoms tend to achieve the electron
configuration of a noble gas, which (except for helium) has eight electrons in its
highest energy level.
cation
Ar
sharing
Ca
I
05_Chem_GRSW_Ch07.SE/TE 6/11/04 3:31 PM Page 59
Name ___________________________ Date ___________________ Class __________________©
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60 Guided Reading and Study Workbook
CHAPTER 7, Ionic and Metallic Bonding (continued)
Formation of Cations (pages 188–190)
8. Write the electron configurations for these metals and circle the electrons lostwhen each metal forms a cation.
a. Mg
b. Al
c. K
Match the noble gas with its electron configuration.
_______ 9. argon a. 1s 2
_______ 10. helium b. 1s 22s 22p 6
_______ 11. neon c. 1s 22s 22p 63s 23p 6
_______ 12. krypton d. 1s 22s 22p 63s 23p 63d104s 24p 6
13. What is the electron configuration called that has 18 electrons in the outerenergy level and all of the orbitals filled?
14. Write the electron configuration for zinc.
15. Fill in the electron configuration diagram for the copper(I) ion.
Formation of Anions (pages 191–192)
16. Atoms of most nonmetallic elements achieve noble-gas electron
configurations by gaining electrons to become ____________________ , or
negatively charged ions.
17. What property of nonmetallic elements makes them more likely to gainelectrons than lose electrons?
Copper atomCu
3d
4s
3p
3s
2p
2s
1s
Ener
gy le
vel
Copper(I) ionCu�
1s22s22p6 2s2
1s22s22p63s23p63d104s2
anions
1s22s22p63s23p6 4s1
c
a
b
d
pseudo noble-gas configuration
They have relatively full valence shells.
1s22s22p6 3s23p1
05_Chem_GRSW_Ch07.SE/TE 6/11/04 3:31 PM Page 60
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18. Is the following sentence true or false? Elements of the halogen family lose one
electron to become halide ions. _________________
19. How many electrons will each element gain in forming an ion?
a. nitrogen _____________
b. oxygen _____________
c. sulfur _____________
d. bromine _____________
20. Write the symbol and electron configuration for each ion from Question 19,and name the noble gas with the same configuration.
a. nitride
b. oxide
c. sulfide
d. bromide
SECTION 7.2 IONIC BONDS AND IONIC COMPOUNDS (pages 194–199)
This section lists the characteristics of an ionic bond. It also describes the use of these characteristics to explain the electrical conductivity of ionic compounds when melted and when in aqueous solutions.
Formation of Ionic Compounds (pages 194–195)
1. What is an ionic bond?
2. In an ionic compound, the charges of the ___________________ and ______________
must balance to produce an electrically ___________________ substance.
3. Complete the electron dot structures below to show how beryllium fluoride (BeF2)is formed. Use figure on page 194 as a model.
Be �F
FBe
F
F
Chapter 7 Ionic and Metallic Bonding 61
N3–; 1s22s22p6; neon
O2–; 1s22s22p6; neon
S2–; 1s22s22p63s23p6; argon
Br–; 1s22s22p63s23p63d104s24p6; krypton
It is the electrostatic force of attraction that binds oppositely
charged ions.
cations anions
neutral
2�
�
�
false
3
2
2
1
05_Chem_GRSW_Ch07.SE/TE 6/11/04 3:31 PM Page 61
62 Guided Reading and Study Workbook
4. Why do beryllium and fluorine combine in a 1 : 2 ratio?
5. A chemical formula shows the types and ______________________ of atoms inthe smallest representative unit of a substance.
6. List the numbers and types of atoms represented by these chemical formulas.
a. Fe2O3 _____________________________________________________________________
b. KMnO4 ___________________________________________________________________
c. CH3 ______________________________________________________________________
d. NH4NO3 __________________________________________________________________
7. What is a formula unit?
8. Explain why the ratio of magnesium ions to chloride ions in MgCl2 is 1 : 2.
9. Describe the structure of ionic compounds.
Properties of Ionic Compounds (pages 196–198)
10. Most ionic compounds are _____________________________ at room temperature.
11. Is the following sentence true or false? Ionic compounds generally have low
melting points. _________________
12. What does a coordination number tell you?
13. What is the coordination number of the ions in a crystal of NaCl? _________________
CHAPTER 7, Ionic and Metallic Bonding (continued)
Name ___________________________ Date ___________________ Class __________________©
Pearson Education, Inc., publishing as Pearson Prentice Hall. All rights reserved.
They combine in a 1 : 2 ratio because each beryllium atom gives up two electrons, but
each fluorine atom only accepts one electron. Therefore, two fluorine atoms are
needed to react with each beryllium atom.
crystalline solids
false
It is the number of ions of opposite charge that surround an ion in a crystal.
6
2 atoms of iron, 3 atoms of oxygen
1 atom of potassium, 1 atom of manganese, 4 atoms of oxygen
1 atom of carbon, 4 atoms of hydrogen
2 atoms of nitrogen, 4 atoms of hydrogen, 3 atoms of oxygen
numbers
Ionic compounds exist as collections of positively and negatively charged ions
arranged in repeating three-dimensional patterns.
A formula unit is the lowest whole-number ratio of ions in an ionic compound.
There must be twice as many chloride ions (Cl–) to balance the magnesium
ions (Mg2+) to form an electrically neutral compound.
05_Chem_GRSW_Ch07.SE/TE 6/11/04 3:31 PM Page 62
Chapter 7 Ionic and Metallic Bonding 63
14. Circle the letter of each statement that is true about ionic compounds.
a. When dissolved in water, ionic compounds can conduct electricity.
b. When melted, ionic compounds do not conduct electricity.
c. Ionic compounds have very unstable structures.
d. Ionic compounds are electrically neutral.
SECTION 7.3 BONDING IN METALS (pages 201–203)
This section uses the theory of metallic bonds to explain the physical proper-ties of metals. It also describes the arrangements of atoms in some commonmetallic crystal structures.
Metallic Bonds and Metallic Properties (pages 201–202)
1. Is the following sentence true or false? Metals are made up of cations, not
neutral atoms. _________________
2. What are metallic bonds?
3. Name three properties of metals that can be explained by metallic bonding.
a.
b.
c.
4. What happens to an ionic crystal when a force is applied to it?
Crystalline Structure of Metals (page 202)
5. Metal atoms in crystals are arranged into very ____________________ and orderly patterns.
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By looking carefully at photographs and drawings in textbooks, you can better understand what youhave read. Look carefully at Figure 7.9 on page 197. What important idea does this drawingcommunicate? Do your work on a separate sheet of paper.
Reading Skill Practice
The shape of a crystalline solid depends on the composition of the ionic compound that forms thecrystal. Different combinations of ions produce different repeating three-dimensional patterns.
true
Metallic bonds are the forces of attraction between
free-floating valence electrons and positively charged metal ions.
electrical conductivity
ductility
malleability
The force tends to push ions of like charge into contact. They repel each other and
the crystal shatters.
compact
05_Chem_GRSW_Ch07.SE/TE 6/11/04 3:31 PM Page 63
64 Guided Reading and Study Workbook
6. Label each of the following arrangements of atoms with the correct name.
7. Circle the letter of each metal whose atoms form a face-centered cubicpattern.
a. magnesium c. sodium
b. copper d. aluminum
Match the arrangement with the number of neighbors each atom in thearrangement has.
______ 8. body-centered cubic a. 12
______ 9. face-centered cubic b. 8
______ 10. hexagonal close-packed
Alloys (page 203)
11. A mixture of two or more elements, at least one of which is a metal, is called
a(n) _________________ .
12. Is the following sentence true or false? Pure metals are usually harder and
more durable than alloys. _________________
13. The most common use of nonferrous alloys is in _________________ .
CHAPTER 7, Ionic and Metallic Bonding (continued)
Name ___________________________ Date ___________________ Class __________________©
Pearson Education, Inc., publishing as Pearson Prentice Hall. All rights reserved.
b
a
a
alloy
false
coins
body-centered cubic
hexagonal close-packed
face-centered cubic
__________________________________
__________________________________
__________________________________
05_Chem_GRSW_Ch07.SE/TE 6/11/04 3:31 PM Page 64
Chapter 7 Ionic and Metallic Bonding 65
14. What four properties make steel an important alloy?
a.
b.
c.
d.
15. What are the component elements for the following alloys?
a. sterling silver
b. brass
c. surgical steel
d. cast iron
16. _________________ alloys have smaller atoms that fit into the spaces between
larger atoms. _________________ alloys have component atoms that are
roughly equal in size.
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silver and copper
copper and zinc
iron, chromium, nickel, and molybdenum
iron and carbon
Substitutional
Interstitial
corrosion resistance
ductility
hardness
toughness
05_Chem_GRSW_Ch07.SE/TE 6/11/04 3:31 PM Page 65
66 Guided Reading and Study Workbook
GUIDED PRACTICE PROBLEM
GUIDED PRACTICE PROBLEM 12 (page 196)
12. Use electron dot structures to determine formulas of the ionic compoundsformed whena. potassium reacts with iodine.b. aluminum reacts with oxygen.
a. potassium reacts with iodine.
Analyze
Step 1. Is one of the elements a metal? If so, which one? ____________________
Step 2. Metal atoms _________ their valence electrons when forming ionic compounds.
Nonmetal atoms _________ electrons when forming ionic compounds.
Solve
Step 3. Write the electron dot structures for the two atoms
potassium _________________ iodine _________________
Step 4. The metal atom, ______________ , must lose ______ electron(s) in order
to achieve an octet in the next-lowest energy level. The nonmetal atom, _________ ,
must gain ______ electron(s) in order to achieve a complete octet.
Step 5. Using electron dot structures, write an equation that shows theformation of the ionic compound from the two elements. Make sure that theelectrons lost equals the electrons gained.
Step 6. The chemical formula for the ionic compound formed is _____________ .
I IK + K+ –→
CHAPTER 7, Ionic and Metallic Bonding (continued)
Name ___________________________ Date ___________________ Class __________________©
Pearson Education, Inc., publishing as Pearson Prentice Hall. All rights reserved.
KI
lose
gain
yes; potassium
K I1
1
potassium
iodine
05_Chem_GRSW_Ch07.SE/TE 6/11/04 3:31 PM Page 66
Chapter 7 Ionic and Metallic Bonding 67
b. aluminum reacts with oxygen.
Analyze
Step 1. Is one of the elements a metal? If so, which one? ____________________
Step 2. Metal atoms _________ their valence electrons when forming ionic compounds.
Nonmetal atoms _________ electrons when forming ionic compounds.
Solve
Step 3. Write the electron dot structures for the two atoms
aluminum _________________ oxygen _________________
Step 4. The metal atom, ______________ , must lose ______ electron(s) in order
to achieve an octet in the next-lowest energy level. The nonmetal atom, _________ ,
must gain ______ electron(s) in order to achieve a complete octet.
Step 5. Using electron dot structures, write an equation that shows theformation of the ionic compound from the two elements. Make sure that theelectrons lost equals the electrons gained.
Step 6. The chemical formula for the ionic compound formed is _____________ .
O
O
O
O
O
O
3+
3+
2–
2–
2–
Al
Al
Al
Al+ →
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Al2O3
lose
3
2
gain
yes; aluminum
aluminum
oxygen
Al O
05_Chem_GRSW_Ch07.SE/TE 6/11/04 3:31 PM Page 67
Name ___________________________ Date ___________________ Class __________________
COVALENT BONDING8
© Pe
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Educ
ation
, Inc
., pu
blish
ing as
Pears
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entic
e Hall
. All r
ights
reserv
ed.
SECTION 8.1 MOLECULAR COMPOUNDS (pages 213–216)
This section explains how to distinguish between ionic and molecular compounds.
Molecules and Molecular Compounds (pages 213–214)
1. What is a covalent bond?
2. Most elements found in nature, with the exception of the ______________________ ,
exist as molecules.
3. What is a molecule?
4. Compounds that are formed when two or more atoms combine to form
molecules are called ______________________ .
5. Circle the letter of the substances that do NOT exist as molecules in nature.
a. oxygen
b. water
c. neon
d. ozone
e. helium
6. List two properties of molecular compounds.
Chapter 8 Covalent Bonding 69
molecular compounds
noble gases
A molecule is a neutral group of atoms joined together by covalent bonds.
low melting and boiling points
exist as gases or liquids at room temperature
a.
b.
A covalent bond is a bond formed when two atoms share electrons.
05_Chem_GRSW_Ch08.SE/TE 6/12/04 11:10 AM Page 69
Name ___________________________ Date ___________________ Class __________________©
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Molecular Formulas (pages 215–216)
7. What is a molecular formula?
Match each compound with its molecular formula.
_______ 8. carbon dioxide a. C2H6
_______ 9. ethane b. CO2
_______10. ammonia c. NH3
11. Is the following sentence true or false? A molecular formula shows the
arrangement of the atoms in a molecule. ______________________
In the diagram, match the type of model or formula with its representation.
a. ball-and-stick model d. space-filling molecular model
b. molecular formula e. structural formula
c. perspective drawing
17. What is the arrangement of atoms within a molecule called?
___________________________
70 Guided Reading and Study Workbook
CHAPTER 8, Covalent Bonding (continued)
false
molecular structure
b
a
c
A molecular formula shows the types and numbers of atoms present in a molecule
of a compound.
12. _______________
15. _______________
16. _______________13. _______________
14. _______________
NH3(g)
H N
H
H
H
N
H HH
N
H H
HN
H
H
b
c
d
ea
05_Chem_GRSW_Ch08.SE/TE 6/12/04 11:10 AM Page 70
Chapter 8 Covalent Bonding 71
SECTION 8.2 THE NATURE OF COVALENT BONDING (pages 217–220)
This section uses electron dot structures to show the formation of single, double,and triple covalent bonds. It also describes and gives examples of coordinatecovalent bonding, resonance structures, and exceptions to the octet rule.
The Octet Rule in Covalent Bonding (page 217)
1. How does electron sharing occur in forming covalent bonds?
Single Covalent Bonds (pages 217–220)
2. Is the following sentence true or false? A shared pair of electrons is represented
by a double dash. ______________________
3. Structural formulas show the arrangement of ______________________in molecules.
4. Use the electron dot diagram below. Circle each unshared pair of electrons in awater molecule.
5. Complete the electron dot structure for each molecule. Each moleculecontains only single covalent bonds.
a. NH3 b. H2O2 c. CH4
Double and Triple Covalent Bonds (pages 221–222)
6. A chemical bond formed when atoms share two pairs of electrons is called a(n)
.
7. How many covalent bonds are in the nitrogen molecule?
______________________
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false
atoms
Sharing of electrons usually occurs so that atoms acquire the electron configurations
of noble gases.
OH
H
H
HO O
H
N
H
H
double covalent bond
H
C
H
H H
three
05_Chem_GRSW_Ch08.SE/TE 6/12/04 11:10 AM Page 71
SECTION 8.4 POLAR BONDS AND MOLECULES (pages 237–244)
This section explains the use of electronegativity values to classify a bond asnonpolar covalent, polar covalent, or ionic. It also names and describes theweak attractive forces that hold groups of molecules together.
Bond Polarity (pages 237–238)
1. Is the following statement true or false? Covalent bonds differ in the way
electrons are shared by the bonded atoms, depending on the kind and number
of atoms joined together. ______________________
2. Describe how electrons are shared in each type of bond. Write equally orunequally.
a. Nonpolar bond ______________________ b. Polar bond ______________________
3. Why does the chlorine atom in hydrogen chloride acquire a slightly negative
charge? _______________________________________________________________________
4. What symbols are used to represent the charges on atoms in a polar covalent
bond? The polarity of the bond? ______________________
Match the electronegativity difference range with the most probable type of bondthat will form.
_______ 5. 0.0–0.4 a. ionic
_______ 6. 0.4–1.0 b. nonpolar covalent
_______ 7. 1.0–2.0 c. very polar covalent
_______ 8. � 2.0 d. moderately polar covalent
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You can increase your understanding of what you have read by making comparisons. A compare/contrasttable can help you do this. On a separate sheet of paper, draw a table to compare the three types ofhybrid orbitals as explained on pages 235 and 236. The three heads for the rows should be sp, sp2,and sp3. Then list the characteristics that will form the basis of your comparison above each column.The column heads should be Number of Hybrid Orbitals, Component Orbitals, Number of Bonds, andBond Angle.
Reading Skill Practice
Students’ tables should include the basic information about the three types of hybrid orbitalscontained on pages 235–236.
Chapter 8 Covalent Bonding 75
equally unequally
true
Chlorine has a higher electronegativity than hydrogen, and,
therefore, attracts the bonding electrons slightly away from hydrogen.
�� or ��;
b
d
c
a
JKL
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76 Guided Reading and Study Workbook
Polar Molecules (pages 239–240)
9. Circle the letter of each sentence that is true about polar molecules.
a. Some regions of a polar molecule are slightly negative and some are slightlypositive.
b. A molecule containing a polar bond is always polar.
c. A molecule that has two poles is called a dipolar molecule.
d. When polar molecules are placed in an electric field, they all line up withthe same orientation in relation to the charged plates.
10. Are the following molecules polar or nonpolar?
a. H2O ______________________ c. NH3 ______________________
b. CO2 ______________________ d. HCl ______________________
Attractions Between Molecules (pages 240–241)
11. What causes dispersion?
12. Is the following sentence true or false? Dispersion forces generally increase
in strength as the number of electrons in a molecule increases. ___________________
13. The strongest of the intermolecular forces are ______________________ .
Intermolecular Attractions and Molecular Properties (pages 243–244)
14. What determines the physical properties of a compound?
15. Use Table 8.4 on page 244. Complete the following table comparing ionic andcovalent compounds.
CHAPTER 8, Covalent Bonding (continued)
Name ___________________________ Date ___________________ Class __________________©
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true
polar
hydrogen bonds
polar
polarnonpolar
Dispersion forces are caused by the motion of electrons.
The physical properties depend on the type of bonding in the compound.
Characteristic Ionic Compound Covalent Compound
Representative unit
Physical state
Melting point
Solubility in water
formula unit molecule
solid solid, liquid, or gas
high low
usually high high to low
05_Chem_GRSW_Ch08.SE/TE 6/12/04 11:10 AM Page 76
Chapter 8 Covalent Bonding 77
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GUIDED PRACTICE PROBLEM 19 (page 239)
19. Identify the bonds between atoms of each pair of elements as nonpolarcovalent, moderately polar covalent, very polar covalent, or ionic.
a. H and Br b. K and Cl c. C and O d. Br and Br
Analyze
Step 1. What is the most probable type of bond for each electronegativitydifference range?
Electronegativity Difference Range Most Probable Type of Bond
0.0–0.4 ___________________________
0.4–1.0 ___________________________
1.0–2.0 ___________________________
� 2.0 ___________________________
Calculate
Step 2. From Table 6.2 on page 177, determine the electronegativity valuesand differences for each pair of elements.
a. H � 2.1, Br � ; difference �
b. K � , Cl � 3.0; difference �
c. C � , O � 3.5; difference �
d. Br � 2.8, Br � ; difference �
Step 3. Refer to Table 8.3 on page 238 to determine the most probable type of bond for each compound.
a. ______________________________
b. ______________________________
c. ______________________________
d. ______________________________
Evaluate
Step 4. How do you know that your answers are correct?
nonpolar covalent
2.8
0.8
2.5
2.8 0
1
2.2
0.7
moderately polar covalent
very polar covalent
ionic
moderately polar covalent
The answers are consistent with the magnitudes of the electronegativity differences.
ionic
very polar covalent
nonpolar covalent
GUIDED PRACTICE PROBLEM
05_Chem_GRSW_Ch08.SE/TE 6/12/04 11:10 AM Page 77
Chapter 9 Chemical Names and Formulas 79
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SECTION 9.1 NAMING IONS (pages 253–258)
This section explains the use of the periodic table to determine the charge of an ion. It also defines polyatomic ion and gives the names and formulas for the most common polyatomic ions.
Monatomic Ions (pages 253–256)
1. What are monatomic ions?
2. How is the ionic charge of a Group 1A, 2A, or 3A ion determined?
3. How is the ionic charge of a Group 5A, 6A, or 7A ion determined?
4. Circle the letter of the type of element that often has more than one commonionic charge.
a. alkali metal
b. alkaline earth metal
c. transition metal
d. nonmetal
5. The ______________________ of naming transition metal cations uses a Romannumeral in parentheses to indicate the numeric value of the ionic charge.
6. An older naming system uses the suffix -ous to name the cation with the
______________________ charge, and the suffix -ic to name the cation with the
______________________ charge.
7. What is a major advantage of the Stock system over the old naming system?
CHEMICAL NAMES AND FORMULAS9
The ionic charge is numerically equal to the group number.
The charge of an ion in Groups 5A, 6A, or 7A is determined by subtracting 8 from
the group number.
Stock system
lesser
greater
The Stock system gives the actual charge of the ion.
Monatomic ions are ions consisting of only one atom.
05_Chem_GRSW_Ch09.SE/TE 6/11/04 3:33 PM Page 79
Name ___________________________ Date ___________________ Class __________________©
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80 Guided Reading and Study Workbook
CHAPTER 9, Chemical Names and Formulas (continued)
8. Use the periodic table to write the name and formula (including charge) foreach ion in the table below.
Polyatomic Ions (pages 257–258)
9. What is a polyatomic ion?
10. Is the following sentence true or false? The names of polyatomic anions always
end in -ide. ______________________
11. What is the difference between the anions sulfite and sulfate?
12. Look at Table 9.3 on page 257. Circle the letter of a polyatomic ion that is a cation.
a. ammonium
b. acetate
c. oxalate
d. phosphate
13. How many atoms make up the oxalate ion and what is its charge?
14. What three hydrogen-containing polyatomic anions are essential componentsof living systems?
a. ____________________________________________
b. ____________________________________________
c. ____________________________________________
A polyatomic ion is a tightly bound group of atoms that behaves as a unit and carries
a charge.
Element Name Formula
Fluorine
Calcium
Oxygen
fluoride ion
calcium ion
oxide ion
F–
Ca2+
O2–
false
The sulfite ion has one less oxygen atom than the sulfate ion.
It is made up of 6 atoms (2 carbon atoms and 4 oxygen atoms) and it has a charge
of 2�.
hydrogen carbonate ion
hydrogen phosphate ion
dihydrogen phosphate ion
05_Chem_GRSW_Ch09.SE/TE 6/11/04 3:33 PM Page 80
Chapter 9 Chemical Names and Formulas 81
15. Look at Figure 9.5 on page 257. Identify each of the ions shown below.
a. __________________ b. ______________________ c. ______________________
SECTION 9.2 NAMING AND WRITING FORMULAS FOR IONIC COMPOUNDS (pages 260–266)
This section explains the rules for naming and writing formulas for binaryionic compounds and compounds containing a polyatomic ion.
Binary Ionic Compounds (pages 260–263)
1. Traditionally, common names were based on some ______________________ of
a compound or its ______________________ .
2. What is the general name for compounds composed of two elements?
__________________________________________________
3. When writing the formula for any ionic compound, the charges of the ions
must ______________________ .
4. What are two methods for writing a balanced formula?
a. ____________________________________________________________________
b. ____________________________________________________________________
5. What are the formulas for the compounds formed by the following pairs of ions?
a. Fe2�, Cl� ______________________
b. Cr3�, O2� ______________________
c. Na�, S2� ______________________
6. What are the formulas for these compounds?
a. lithium bromide ______________________
b. cupric nitride ______________________
c. magnesium chloride ______________________
7. The name of a binary ionic compound is written with the name of the
______________________ first followed by the name of the ______________________ .
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ammonium ion phosphate ion nitrate ion
N P N1�
3� 1�
property
source
They are binary compounds.
balance
finding the least common multiple of the charges
using the crisscross method
FeCl2
Cr2O3
Na2S
LiBr
Cu3N2
MgCl2
cation anion
05_Chem_GRSW_Ch09.SE/TE 6/11/04 3:33 PM Page 81
82 Guided Reading and Study Workbook
8. How can you tell that cobalt(II) iodide is a binary ionic compound formed by a transition metal with more than one ionic charge?
9. Write the names for these binary ionic compounds.
a. PbS __________________________________
b. MgCl2 __________________________________
c. Al2Se3 __________________________________
Compounds with Polyatomic Ions (pages 264–266)
10. What is a polyatomic ion?
11. How do you write the formula for a compound containing a polyatomic ion?
12. Why are parentheses used to write the formula Al(OH)3?
13. Complete the table for these ionic compounds containing polyatomic ions.
CHAPTER 9, Chemical Names and Formulas (continued)
Name ___________________________ Date ___________________ Class __________________©
Pearson Education, Inc., publishing as Pearson Prentice Hall. All rights reserved.
lead(II) sulfide
magnesium chloride
aluminum selenide
A polyatomic ion contains more than one element.
The name includes a Roman numeral representing the ionic charge of the transition
metal cation.
Cation Anion Name Formula
NH4+ S2�
Fe3+ iron(III) carbonate
NO3� AgNO3
potassium cyanide KCN
The parentheses indicate how many polyatomic ions are needed in the formula.
Write the symbol for the cation followed by the formula for the polyatomic ion
and balance the charges.
ammonium sulfide
CO32–
Ag+
K+CN–
silver nitrate
Fe2(CO3)3
(NH4)2S
05_Chem_GRSW_Ch09.SE/TE 6/11/04 3:33 PM Page 82
Chapter 9 Chemical Names and Formulas 83
SECTION 9.3 NAMING AND WRITING FORMULAS FORMOLECULAR COMPOUNDS (pages 268–270)
This section explains the rules for naming and writing formulas for binarymolecular compounds.
Naming Binary Molecular Compounds (pages 268–269)
1. Circle the letter of the type(s) of elements that form binary molecular compounds.
a. two nonmetallic elements
b. a metal and a nonmetal
c. two metals
2. Is the following sentence true or false? Two nonmetallic elements can combine
in only one way. ______________________
3. What method is used to distinguish between different molecular compounds
that contain the same elements? _________________________________________
Match the prefix with the number it indicates.
_______ 4. octa- a. 4
_______ 5. tetra- b. 7
_______ 6. hepta- c. 8
_______ 7. nona- d. 9
8. What are the names of the following compounds?
a. BF3 ____________________________________________
b. N2O4 ____________________________________________
c. P4S7 ____________________________________________
Writing Formulas for Binary Molecular Compounds (page 270)
9. What are the formulas for the following compounds?
a. carbon tetrabromide ______________________
b. nitrogen triiodide ______________________
c. iodine monochloride ______________________
d. tetraiodine nonaoxide ______________________
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boron trifluoride
dinitrogen tetroxide
tetraphosphorus heptasulfide
c
a
b
d
Prefixes are used.
CBr4
NI3
ICl
I4O9
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SECTION 9.4 NAMING AND WRITING FORMULAS FOR ACIDS AND BASES (pages 271–273)
This section explains the three rules for naming acids and shows how theserules can also be used to write the formulas for acids. Names and formulasfor bases are also explained.
Naming Common Acids (pages 271–272)
1. Acids produce ______________________ ions when dissolved in water.
2. When naming acids, you can consider them to be combinations of
______________________ connected to as many ______________________ ions
as are necessary to create an electrically neutral compound.
3. What is the formula for hydrobromic acid? ______________________
4. What are the components of phosphorous acid? What is its formula?
Writing Formulas for Acids (page 272)
5. Use Table 9.5 on page 272 to help you complete the table about acids.
CHAPTER 9, Chemical Names and Formulas (continued)
Name ___________________________ Date ___________________ Class __________________©
Pearson Education, Inc., publishing as Pearson Prentice Hall. All rights reserved.
Acid Name Formula Anion Name
acetic acid
carbonic acid
hydrochloric acid
nitric acid
phosphoric acid
sulfuric acid
HC2H3O2
H2CO3
HCl
HNO3
H3PO4
H2SO4
acetate
carbonate
chloride
nitrate
phosphate
sulfate
hydrogen
anions
HBr
hydrogen ion and phosphite ion; H3PO3
hydrogen
Writing a summary can help you remember the information you have read. When you write asummary, include only the most important points. Write a summary of the information in Section 9.3on pages 268–269. Your summary should be shorter than the text on which it is based. Do your workon a separate sheet of paper.
Students’ summaries should focus on the main points of each subsection and include the informationgiven in Table 9.4.
Reading Skill Practice
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Chapter 9 Chemical Names and Formulas 85
Names and Formulas for Bases (page 273)
6. A base is a compound that produces ______________________ when dissolvedin water.
7. How are bases named?
SECTION 9.5 THE LAWS GOVERNING FORMULAS AND NAMES (pages 274–279)
This section uses data to demonstrate that a compound obeys the law of definite proportions. It also explains how to use flow charts to write the name and formula of a compound.
The Laws of Definite and Multiple Proportions (pages 274–275)
1. What is the law of definite proportions?
2. Circle the whole-number mass ratio of Li to Cl in LiCl. The atomic mass of Li is6.9; the atomic mass of Cl is 35.5.
a. 42 : 1
b. 5 : 1
c. 1 : 5
3. Circle the whole-number mass ratio of carbon to hydrogen in C2H4. Theatomic mass of C is 12.0; the atomic mass of H is 1.0.
a. 1 : 6 c. 1 : 12
b. 6 : 1 d. 12 : 1
4. In the compound sulfur dioxide, a food preservative, the mass ratio of sulfur tooxygen is 1 : 1. An 80-g sample of a compound composed of sulfur and oxygencontains 48 g of oxygen. Is the sample sulfur dioxide? Explain.
5. What is the law of multiple proportions?
Name ___________________________ Date ___________________ Class __________________©
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hydroxide ions
No; If the sample contains 48 g of oxygen, it contains 32 g of sulfur. The ratio 32 : 48
is equivalent to 2 : 3, not 1 : 1.
When two elements form more than one compound, the different masses of one
element that combine with the same mass of the other element are in a ratio of
small whole numbers.
In different samples of the same chemical compound, the masses of the elements
are always present in the same proportions.
Name the cation first followed by the anion (hydroxide ion).
05_Chem_GRSW_Ch09.SE/TE 6/11/04 3:33 PM Page 85
86 Guided Reading and Study Workbook
6. Complete the table using the law of multiple proportions.
Practicing Skills: Naming Chemical Compounds (pages 276–277)
7. How can a flowchart help you to name chemical compounds?
8. Use the flowchart in Figure 9.20 on page 277 to write the names of thefollowing compounds:
a. CsCl ___________________________________
b. SnSe2 ___________________________________
c. NH4OH ___________________________________
d. HF ___________________________________
e. Si3N4 ___________________________________
9. Complete the following five rules for writing a chemical formula from achemical name.
a. In an ionic compound, the net ionic charge is ______________________ .
b. An -ide ending generally indicates a ______________________ compound.
c. An -ite or -ate ending means there is a ______________________ ion that
includes oxygen in the formula.
d. ______________________ in a name generally indicate that the compound is
molecular and show the number of each kind of atom in the molecule.
e. A ______________________ after the name of a cation shows the ionic charge
of the cation.
CHAPTER 9, Chemical Names and Formulas (continued)
Name ___________________________ Date ___________________ Class __________________©
Pearson Education, Inc., publishing as Pearson Prentice Hall. All rights reserved.
It gives step-by-step directions for naming a compound.
cesium chloride
tin(IV) selenide
ammonium hydroxide
trisilicon tetranitride
hydrofluoric acid
zero
binary
polyatomic
Prefixes
Roman numeral
10.55
1.1 2
Mass of Cu Mass of Cl Mass Ratio Whole-number Cl : Cu Ratio of Cl
Compound A 8.3 g 4.6 g
Compound B 3.3 g 3.6 g
05_Chem_GRSW_Ch09.SE/TE 6/11/04 3:33 PM Page 86
Chapter 9 Chemical Names and Formulas 87
Practicing Skills: Writing Chemical Formulas (page 278)
10. Fill in the missing labels from Figure 9.22 on page 278.
11. Use the flowchart in Figure 9.22 to write the formulas of the followingcompounds:
a. potassium silicate ______________________
b. phosphorus pentachloride ______________________
c. manganese(II) chromate ______________________
d. lithium hydride ______________________
e. diiodine pentoxide ______________________
Name ___________________________ Date ___________________ Class __________________©
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Use Table 9.1for charges.
Use Table 9.3for charges.
Uses prefixes to write formula.
Give charges for cations.
Use crisscross method. Add parentheses for any multiple polyatomic ions.
Identify symbols
no
yesContainsprefixes?
Group Aelements
Name of Compound
K2SiO3
PCl5
MnCrO4
LiH
I2O5
Molecular compound
Ioniccompound
Romannumerals
Polyatomicions
Balancecharges
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Name ___________________________ Date ___________________ Class __________________©
Pearson Education, Inc., publishing as Pearson Prentice Hall. All rights reserved.
GUIDED PRACTICE PROBLEM 2 (page 256)
2. How many electrons were lost or gained to form these ions?a. Fe3� b. O2� c. Cu�
Step 1. Determine the number of electrons based on the size of the charge.
Step 2. Determine whether the electrons were lost or gained based on the signof the charge.
a. ________________________________________________
b. ________________________________________________
c. ________________________________________________
GUIDED PRACTICE PROBLEMS 10B AND 10C (page 263)
10. Write formulas for compounds formed from these pairs of ions.
b. Li+, O2–
Analyze
Step 1. Do the ions combine in a 1:1 ratio?
Solve
Step 2. Use the crisscross method to balance the formula.
Write the formula. ___________
Evaluate
Step 3. How do you know your formula is reasonable?
CHAPTER 9, Chemical Names and Formulas (continued)
No, the charges on the ions are not equal.
The positive and negative charges are equal.
GUIDED PRACTICE PROBLEMS
�
Li2�
O2 1
Fe3� lost 3 electrons.
O2� gained 2 electrons.
Cu� lost 1 electron.
Li20
05_Chem_GRSW_Ch09.SE/TE 6/11/04 3:33 PM Page 88
Chapter 9 Chemical Names and Formulas 89
c. Ca2� , N3�
Analyze
Step 1. Will the calcium (Ca2+) and nitride (N3–) ions combine in a 1 : 1 ratio?How do you know?
Solve
Step 2. Use the crisscross method to balance the formula.
Write the formula. ___________
Evaluate
Step 3. How do you know this formula is reasonable?
GUIDED PRACTICE PROBLEM 13B (page 265)
13b. Write the formula for chromium(III) nitrate.
• Is the compound ionic or molecular? Explain.
• Use Table 9.3 on page 257 to write the formula for
the nitrate ion. ________
• Use the crisscross method to balance the formula.
• Write the formula. ____________________________
Name ___________________________ Date ___________________ Class __________________©
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No, because then the total charge would be negative, instead of neutral.
2�
Ca3�
N3 2
The positive and negative charges are equal.
It is ionic because it has no prefixes and it contains a metal.
3�
Cr�
(NO3)NO3
–
1 3Cr(NO3)3
Ca3N2
05_Chem_GRSW_Ch09.SE/TE 6/11/04 3:33 PM Page 89