chapters 5–8 resources - a beauty...

Chapters 5–8 Resources

Copyright © by The McGraw-Hill Companies, Inc. All rights reserved. Permission is granted to reproduce the material contained herein on the condition that such materials be reproduced only for classroom use; be provided to students, teachers, and families without charge; and be used solely in conjunction with the Glencoe Chemistry: Matter and Change program. Any other reproduction, for sale or other use, is expressly prohibited.

Send all inquiries to:Glencoe/McGraw-Hill8787 Orion PlaceColumbus, OH 43240-4027

ISBN: 978-0-07-878761-4MHID: 0-07-878761-0

Printed in the United States of America.

1 2 3 4 5 6 7 8 9 10 045 11 10 09 08 07

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To the Teacher . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . iv

Chapters 5-8 Resources

Reproducible Student Pages

Student Lab Safety Form . . . . . . . . . . . . . . . . . . . . . . . . . . vi

Chapter 5

Electrons in Atoms . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1

Chapter 6

The Periodic Table and Periodic Law . . . . . . . . . . . . . . . . . . . 27

Chapter 7

Ionic Compounds and Metals . . . . . . . . . . . . . . . . . . . . . . 55

Chapter 8

Covalent Bonding . . . . . . . . . . . . . . . . . . . . . . . . . . . . 85

Teacher Guide and Answers

Chapter 5 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 113

Chapter 6 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 117

Chapter 7 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 121

Chapter 8 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 127

Table ofContents

iii

iv

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To the Teacher

This booklet contains resource materials to help you teach more effectively. You will find the following in the chapters:

Reproducible Pages

Hands-on ActivitiesMiniLab and ChemLab Worksheets: Each activity in this book is an expanded version of each lab that appears in the Student Edition of Glencoe Chemistry: Matter and Change. All materials lists, procedures, and questions are repeated so that students can read and complete a lab in most cases without having a textbook on the lab table. All lab ques-tions are reprinted with lines on which students can write their answers. In addition, for student safety, all appropriate safety symbols and caution statements have been reproduced on these expanded pages. Answer pages for each MiniLab and ChemLab are included in the Teacher Guide and Answers section at the back of this book.

Transparency ActivitiesTeaching Transparency Masters and Worksheets: These transparencies relate to major concepts that will benefit from an extra visual learning aid. Most of the transparencies contain art or photos that extend the concepts put forth in the textbook. Others contain art or photos directly from the Student Edition. There are 73 Teaching Transparencies, provided here as black-and-white masters accompanied by worksheets that review the concepts presented in the transparencies. Answers to worksheet questions are provided in the Teacher Guide and Answers section at the back of this book.

Math Skills Transparency Masters and Worksheets: These transparencies relate to math-ematical concepts that will benefit from an extra visual learning aid. Most of the trans-parencies contain art or photos directly from the Student Edition, or extend concepts put forth in the textbook. There are 42 Math Skills Transparencies, provided here as black-and-white masters accompanied by worksheets that review the concepts presented in the transparencies. Answers to worksheet questions are provided in the Teacher Guide and Answers section at the back of this book.

Intervention and AssessmentStudy Guide: These pages help students understand, organize, and compare the main chemistry concepts in the textbook. The questions and activities also help build strong study and reading skills. There are six study guide pages for each chapter. Students will find these pages easy to follow because the section titles match those in the textbook. Italicized sentences in the study guide direct students to the related topics in the text.

The Study Guide exercises employ a variety of formats including multiple-choice, matching, true/false, labeling, completion, and short answer questions. The clear, easy-to-follow exercises and the self-pacing format are geared to build your students’ confi-dence in understanding chemistry. Answers or possible responses to all questions are provided in the Teacher Guide and Answers section at the back of this book.

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Chapter Assessment: Each chapter assessment includes several sections that assess stu-dents’ understandings at different levels.

• The Reviewing Vocabulary section tests students’ knowledge of the chapter’s vocabu-lary. A variety of formats are used, including matching, true/false, completion, and comparison of terms.

• The Understanding Main Ideas section consists of two parts: Part A tests recall and basic understanding of facts presented in the chapter, while Part B is designed to be more challenging and requires deeper comprehension of concepts than does Part A. Students may be asked to explain chemical processes and relationships or to make comparisons and generalizations.

• The Thinking Critically section requires students to use several different higher-order learning skills, such as interpreting data and discovering relationships in graphs and tables, as well as applying their understanding of concepts to solve problems, com-pare and contrast situations, and to make inferences or predictions.

• The Applying Scientific Methods section puts students into the role of researcher. They may be asked to read about an experiment, simulation, or model and then apply their understanding of chapter concepts and scientific methods to analyze and explain the procedure and results. Many of the questions in this section are open-ended, giving students the opportunity to demonstrate both reasoning and creative problem-solv-ing skills.

Answers or possible responses to all questions are provided in the Teacher Guide and Answers section at the back of this book.

STP Recording Sheet: Recording Sheets allow students to use the Standardized Test Practice questions in the Student Edition as a practice for standardized tests. STP Recording Sheets give them the opportunity to use bubble answer grids and numbers grids for recording answers. Answers for the STP Recording Sheets can be found in the Teacher Wraparound Edition on Standardized Test Practice pages.

Teacher Guide and Answers: Answers or possible answers for questions in this booklet can be found in the Teacher Guide and Answers section. Materials, teaching strate-gies, and content background, along with chapter references, are also provided where appropriate.

To the Teacher continued

vi

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Lab Safety Form

Name:

Date:

Lab type (circle one) : Launch Lab MiniLab ChemLab

Lab Title:

Read carefully the entire lab and then answer the following questions. Your teacher must initial this form before you begin the lab.

1. What is the purpose of the investigation?

2. Will you be working with a partner or on a team?

3. Is this a design-your-own procedure? Circle: Yes No

4. Describe the safety procedures and additional warnings that you must follow as you perform this investigation.

5. Are there any steps in the procedure or lab safety symbols that you do not understand? Explain.

Teacher Approval Initials

Date of Approval

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Chapter 5 Electrons in AtomsMiniLab . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2

ChemLab . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3

Teaching Transparency Masters and Worksheets . . . . . . . . . . . . . . . . . . . . . . . . . . 6

Math Skills Transparency Masters and Worksheets . . . . . . . . . . . . . . . . . . . . . . . . . . 12

Study Guide . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14

Chapter Assessment. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 20

STP Recording Sheet . . . . . . . . . . . . . . . . . . . . . . . . . . . . 26

Table ofContents

1

Reproducible Pages

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2 Chemistry: Matter and Change • Chapter 5 ChemLab and MiniLab Worksheets

mini LAB 5Identify Compounds

How do flame colors vary for different elements?

Materials Bunsen burner; cotton swabs (6); crystals of lithium chloride, sodium chloride,potassium chloride, calcium chloride, strontium chloride, unknown solution

Procedure 1. Read and complete the lab safety form.

2. Dip one of six cotton swabs into the lithiumchloride solution. Put the swab into the flameof a Bunsen burner. Observe the color of theflame, and record it in your data table.

3. Repeat Step 2 for each of the metallic chlo-ride solutions (sodium chloride, potassiumchloride, calcium chloride, and strontiumchloride). Record the color of each flame inyour data table.

4. Compare your results to the flame testsshown in the Elements Handbook.

5. Repeat Step 2 using a sample of unknownsolution obtained from your teacher. Recordthe color of the flame produced.

6. Dispose of the used cotton swabs as directedby your teacher.

Analysis

1. Suggest a reason why each compound produced a flame of a different color, eventhough they each contain chlorine.

2. Explain how an element’s flame test might be related to its atomic emission spectrum.

3. Infer the identity of the unknown crystals. Explain your reasoning.

Compound Flame color

Lithium chloride

Sodium chloride

Potassium chloride

Calcium chloride

Strontium chloride

Unknown

Flame Test Results

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ChemLab and MiniLab Worksheets Chemistry: Matter and Change • Chapter 5 3

CHEMLAB 5

Safety Precautions• Always wear safety goggles and a lab apron.• Use care around the spectrum tube power supplies.• Spectrum tubes will get hot when used.

ProblemWhat absorption and emis-sion spectra do various sub-stances produce?

Objectives• Observe emission spectra

of several gases.• Observe the absorption

spectra of various solu-tions.

• Analyze patterns ofabsorption and emissionspectra.

Materials ring stand with

clamp40-W tubular light-

bulblight socket with

grounded powercord

275-mL polystyreneculture flask

Flinn C-Spectra® orsimilar diffractiongrating

food coloring (red,green, blue, and yellow)

set of colored pencils

spectrum tubes(hydrogen, neon,and sodium)

spectrum–tubepower supplies (3)

Analyze Line SpectraEmission spectra are produced when excited atoms return to a

more stable state by emitting radiation of specific wavelengths.When white light passes through a sample, atoms in the sampleabsorb specific wavelengths. This produces dark lines in the continu-ous spectrum of white light and is called an absorption spectrum.

Pre-Lab

1. Read the entire CHEMLAB.

2. Explain how electrons in an element’s atoms produce an emission spectrum.

3. Distinguish among a continuous spectrum, anemission spectrum, and an absorption spectrum.

4. Use the data table on the next page.

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Procedure

1. Read and complete the lab safety form.

2. Use a Flinn C-Spectra® or similar diffractiongrating to view an incandescent lightbulb. Whatdo you observe? Draw the observed spectrumusing colored pencils.

3. Use the Flinn C-Spectra® to view the emissionspectra from tubes of gaseous hydrogen, neon,and sodium. Use colored pencils to make draw-ings in the data table of the spectra observed.

4. Fill a 275-mL culture flask with about 100-mLwater. Add 2 or 3 drops of red food coloring tothe water. Shake the solution.

5. Repeat step 4 for the green, blue, and yellow foodcoloring.

6. Set up the 40-W lightbulb so that it is near eyelevel. Place the flask with red food coloring about 8 cm from the lightbulb. You should be able to see light from the bulb above the solutionand light from the bulb projecting through thesolution.

7. With the room lights darkened, view the lightusing the Flinn C-Spectra®. The top spectrum

viewed will be a continuous spectrum of thewhite lightbulb. The bottom spectrum will be theabsorption spectrum of the red solution. Use col-ored pencils to make a drawing in the data tableof the absorption spectra you observed.

8. Repeat steps 6 and 7 using the green, blue, andyellow colored solutions.

9. Cleanup and Disposal Turn off the light socketand spectrum tube power supplies.Wait severalminutes to allow the incandescent lightbulb andthe spectrum tubes to cool. Follow your teacher’sinstructions on how to dispose of the liquids andhow to store the lightbulb and spectrum tubes.

CHEMLAB 5

Hydrogen

Neon

Mercury

Drawings of Emission Spectra

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Analyze and Conclude

1. Think Critically How can the single electron in a hydrogen atom produce all of the linesfound in its emission spectrum?

2. Predict How can you predict the absorption spectrum of a solution by looking at its color?

3. Apply How can spectra be used to identify the presence of specific elements in a substance?

4. Error Analysis Name a potential source of error in this experiment. Choose one of theelements you observed, and research its absorption spectrum. Compare your findings withthe results of your experiment.

Inquiry Extension

Hypothesize What would happen if you mixed more than one color of food coloringwith water and repeated the experiment? Design an experiment to test your hypothesis.

CHEMLAB 5

Red

Green

Blue

Yellow

Drawings of Absorption Spectra

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6 Chemistry: Matter and Change • Chapter 5 Teaching Transparency Masters

TEACHING TRANSPARENCY MASTER

Freq

uen

cy (

�)

in h

ertz

Rad

ioIn

frar

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ltra

vio

let

Gam

ma

rays

Elec

tro

mag

net

ic S

pec

tru

m

104

3 �

104

106

108

1010

3 �

10�

2

1012

3 �

10�

4

1014

3 �

10�

6

1016

3 �

10�

8

1018

3 �

10�

10

1020

3 �

10�

12

1022

3 �

10�

14

Wav

elen

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) in

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ers

Vis

ible

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3

TV, F

MA

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Mic

row

aves

X r

ays

3 �

102

Ener

gy

incr

ease

s

The Electromagnetic SpectrumThe Electromagnetic Spectrum


Use with Chapter 5,Section 5.1

15

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Teaching Transparency Worksheets Chemistry: Matter and Change • Chapter 5 7

1. What kinds of waves have the longest wavelength? What kinds of waves have the short-est wavelength?

2. Which waves have the lowest frequency?

3. Which has a higher frequency: microwaves or X rays?

4. Which waves can be seen by the eye?

5. Sequence the different segments of the visible spectrum in order from shortest wave-length to longest wavelength.

6. Sequence the following types of waves from lowest frequency to highest frequency:ultraviolet rays, infrared rays, gamma rays, radio waves, and green light.

7. Compare the wavelengths and frequencies of each kind of wave. What is the relationshipbetween frequency and wavelength?

8. What is the wavelength of a radio station emitting its signal at 95.5 MHz? Estimate youranswer to the nearest power of ten.

The Electromagnetic SpectrumThe Electromagnetic Spectrum

TEACHING TRANSPARENCY WORKSHEET


15

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z

x

y

2s orbital

z

x

y

1s orbital

z

x

pzpx py

p orbitals

d orbitals

y

z

x

y

z

x

y

z

x

y

dxy dxz dyz dz2dx2�y2

x

yy

x

z

x

y

z zz

x

y

Atomic OrbitalsAtomic Orbitals



16

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1. What is the shape of an s orbital?

2. What is the relationship between the size of an s orbital and the principal energy level inwhich it is found?

3. What is the shape of a p orbital? How many p orbitals are there in a sublevel?

4. How many electrons can each orbital hold?

5. Look at the diagrams of the p orbitals. What do x, y, and z refer to?

6. How many d orbitals are there in a given sublevel? How many total electrons can thed orbitals in a sublevel hold?

7. Which d orbitals have the same shape?

8. What point in each diagram represents an atom’s nucleus?

9. How likely is it that an electron occupying a p or a d orbital would be found very near anatom’s nucleus? What part of the diagram supports your conclusion?

Atomic OrbitalsAtomic Orbitals



16

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2s2p3s3p4s3d5p4f6p5f7p 5s

4p4d6s5d7s

6p7p6d

5f 4f5d 4d 3d

4p 3p 2p

6d

7s 5s 4s 3s 2s 1s6s5p

1s

Increasing Energy

Orbital filling sequence

Orbital Filling Sequence and Energy LevelsOrbital Filling Sequence and Energy Levels



17

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1. What does each small box in the diagram represent?

2. How many electrons can each orbital hold?

3. How many electrons can the d sublevel hold?

4. Which is associated with more energy: a 2s or a 2p orbital?

5. Which is associated with more energy: a 2s or a 3s orbital?

6. According to the aufbau principle, which orbital should fill first, a 4s or a 3d orbital?

7. Which orbital has the least amount of energy?

8. What is the likelihood that an atom contains a 1s orbital?

9. Sequence the following orbitals in the order that they should fill up according to the aufbau principle: 4d, 4p, 4f, 5s, 6s, 5p, 3d, 4s.

10. Write a general rule to describe the filling of orbitals in an atom.

Orbital Filling Sequence and Energy LevelsOrbital Filling Sequence and Energy Levels



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12 Chemistry: Matter and Change • Chapter 5 Math Skills Transparency Masters

Waves A and B are both electromagnetic waves.

c � �� for all electromagnetic waves.

amplitude

amplitude

A

B

Interpreting WavesInterpreting Waves

MATH SKILLS TRANSPARENCY MASTER


5

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Math Skills Transparency Worksheets Chemistry: Matter and Change • Chapter 5 13

1. Look at the two waves shown. What is the speed of each wave?

2. Look at the two waves shown. Which wave has a higher frequency? Which wave has alonger wavelength?

3. Assume that wave A has a wavelength of 699 nm. Calculate the frequency of the wave.Show your work.

4. Assume that wave B has a wavelength of 415 nm. Calculate the frequency of the wave.Show your work.

5. Compare your calculations in question 4 with your answer to question 3. Do your calcu-lations support your answer in question 2?

6. If wave A has a frequency of 4.60 � 1014 s�1, what is its wavelength in nanometers?Show your work.

Interpreting WavesInterpreting Waves

MATH SKILLS TRANSPARENCY WORKSHEET


5

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14 Chemistry: Matter and Change • Chapter 5 Study Guide

Electrons in AtomsElectrons in Atoms

Section 5.1 Light and Quantized EnergyIn your textbook, read about the wave nature of light.

Use each of the terms below just once to complete the passage.

Electromagnetic radiation is a kind of (1) that behaves like a(n)

(2) as it travels through space. (3) is one type of

electromagnetic radiation. Other examples include X rays, radio waves, and microwaves.

All waves can be characterized by their wavelength, amplitude, frequency, and

(4) . The shortest distance between equivalent points on a continuous wave is

called a(n) (5) . The height of a wave from the origin to a crest or from the

origin to a trough is the (6) . (7) is the number of

waves that pass a given point in one second. The SI unit for frequency is the

(8) , which is equivalent to one wave per second.

Use the figure to answer the following questions.

9. Which letter(s) represent one wavelength?

10. Which letter(s) represent the amplitude?

11. If twice the length of A passes a stationary point every second, what is the frequency ofthe wave?

Origin

A

D

C

B

STUDY GUIDECHAPTER 5

amplitude energy frequency hertz

light wave wavelength speed

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Study Guide Chemistry: Matter and Change • Chapter 5 15

In your textbook, read about the particle nature of light.

Circle the letter of the choice that best completes the statement or answers the question.

12. A(n) is the minimum amount of energy that can be lost or gained by an atom.

a. valence electron b. electron c. quantum d. Planck’s constant

13. According to Planck’s theory, for a given frequency, �, matter can emit or absorb energyonly in

a. units of hertz. c. entire wavelengths.

b. whole-number multiples of h�. d. multiples of �12� h�, �

14� h�, and so on.

14. The is the phenomenon in which electrons are emitted from a metal’s surfacewhen light of a certain frequency shines on it.

a. quantum b. Planck concept c. photon effect d. photoelectric effect

15. Which equation would you use to calculate the energy of a photon?

a. Ephoton � h� � Planck’s constant c. Ephoton � �12� h�

b. Ephoton � h� d. c � ��

In your textbook, read about atomic emission spectra.

For each statement below, write true or false.

16. Like the visible spectrum, an atomic emission spectrum is a continuousrange of colors.

17. Each element has a unique atomic emission spectrum.

18. A flame test can be used to identify the presence of certain elements in a compound.

19. The fact that only certain colors appear in an element’s atomic emissionspectrum indicates that only certain frequencies of light are emitted.

20. Atomic emission spectra can be explained by the wave model of light.

21. The neon atoms in a neon sign emit their characteristic color of light asthey absorb energy.

22. When an atom emits light, photons having certain specific energies arebeing emitted.

Section 5.1 continued


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Section 5.2 Quantum Theory and the AtomIn your textbook, read about the Bohr model of the atom.

Use each of the terms below to complete the statements.

1. The lowest allowable energy state of an atom is called its .

2. Bohr’s model of the atom predicted the of the lines in

hydrogen’s atomic emission spectrum.

3. According to Bohr’s atomic model, the smaller an electron’s orbit, the

the atom’s energy level.

4. According to Bohr’s atomic model, the larger an electron’s orbit, the

the atom’s energy level.

5. Bohr proposed that when energy is added to a hydrogen atom, its

moves to a higher-energy orbit.

6. According to Bohr’s atomic model, the hydrogen atom emits a photon corresponding to

the difference between the associated with the two

orbits it transitions between.

7. Bohr’s atomic model failed to explain the of elements

other than hydrogen.

In your textbook, read about the quantum mechanical model of the atom.

Answer the following questions.

8. If you looked closely, could you see the wavelength of a fast-moving car? Explain your answer.

9. Using de Broglie’s equation, � � �mh�� which would have the larger wavelength, a

slow-moving proton or a fast-moving golf ball? Explain your answer.


atomic emission spectrum electron frequencies ground state

higher energy levels lower

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In your textbook, read about the Heisenberg uncertainty principle.

For each item in Column A, write the letter of the matching item in Column B.

Column A Column B

10. The modern model of the atom that treats electrons as waves

11. States that it is impossible to know both the velocityand the position of a particle at the same time

12. A three-dimensional region around the nucleusrepresenting the probability of finding an electron

13. Originally applied to the hydrogen atom, it led to thequantum mechanical model of the atom

Answer the following question.

14. How do the Bohr model and the quantum mechanical model of the atom differ in howthey describe electrons?

In your textbook, read about hydrogen’s atomic orbitals.

In the space at the left, write the term in parentheses that correctly completes the statement.

15. Atomic orbitals (do, do not) have an exactly defined size.

16. Each orbital may contain at most (two, four) electrons.

17. All s orbitals are (spherically shaped, dumbbell shaped).

18. A principal energy has (n, n2) energy sublevels.

19. The maximum number of (electrons, orbitals) related to eachprincipal energy level equals 2n2.

20. There are (three, five) equal energy p orbitals.

21. Hydrogen’s principal energy level 2 consists of (2s and 3s, 2s and2p) orbitals.

22. Hydrogen’s principal energy level 3 consists of (nine, three)orbitals.


a. Heisenberg uncertaintyprinciple

b. Schrödinger wave equation

c. quantum mechanical modelof the atom

d. atomic orbital

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Section 5.3 Electron ConfigurationIn your textbook, read about ground-state electron configurations.


The arrangement of electrons in an atom is called the atom’s

(1) . Electrons in an atom tend to assume the arrangement

that gives the atom the (2) possible energy. This arrangement

of electrons is the most (3) arrangement and is called the

atom’s (4) .

Three rules define how electrons can be arranged in an atom’s orbitals. The

(5) states that each electron occupies the lowest energy

orbital available. The (6) states that a maximum of two

electrons may occupy a single atomic orbital, but only if the electrons have opposite

(7) . (8) states that single

electrons with the same spin must occupy each equal-energy orbital before additional

electrons with opposite spins occupy the same orbitals.

Complete the following table.


Aufbau principle electron configuration ground-state electron configuration Hund’s rule

lowest Pauli exclusion principle spins stable

Element Atomic Number Orbitals Electron Configuration

1s 2s 2px 2py 2pz

9. Helium 1s2

10. 7

11. Neon )( )( )( )( )(

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12. What is germanium’s atomic number? How many electrons does germanium have?

13. What is noble-gas notation, and why is it used to write electron configurations?

14. Write the ground-state electron configuration of a germanium atom, using noble-gasnotation.

In your textbook, read about valence electrons.


15. The electrons in an atom’s outermost orbitals are called

a. electron dots. b. quantum electrons. c. valence electrons. d. noble-gas electrons.

16. In an electron-dot structure, the element’s symbol represents the

a. nucleus of the noble gas closest to the atom in the periodic table.

b. atom’s nucleus and inner-level electrons.

c. atom’s valence electrons.

d. electrons of the noble gas closest to the atom in the periodic table.

17. How many valence electrons does a chlorine atom have if its electron configuration is [Ne]3s23p5?

a. 3 b. 21 c. 5 d. 7

18. Given boron’s electron configuration of [He]2s22p1, which of the following represents itselectron-dot structure?

a. •Be• b. •B• c. B d. Be

19. Given beryllium’s electron configuration of 1s22s2, which of the following represents itselectron-dot structure?

a. •Be• b. •B• c. B d. Be

20. Which electrons are represented by the dots in an electron-dot structure?

a. valence electrons c. only s electrons

b. inner-level electrons d. both a and c


• ••••••

• ••••••


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20 Chemistry: Matter and Change • Chapter 5 Chapter Assessment

Electrons in AtomsElectrons in Atoms

Reviewing VocabularyMatch the definition in Column A with the term in Column B.

Column A Column B

1. The set of frequencies of the electromagnetic wavesemitted by the atoms of an element

2. The minimum amount of energy that can be lost or gainedby an atom

3. A form of energy that exhibits wavelike behavior as ittravels through space

4. A three-dimensional region around the nucleus of an atomthat describes an electron’s probable location

5. The shortest distance between equivalent points on acontinuous wave

6. The lowest allowable energy state of an atom

7. A particle of electromagnetic radiation with no mass thatcarries a quantum of energy

8. The emission of electrons from a metal’s surface whenlight of a certain frequency shines on it

9. A figure indicating the relative sizes and energies of atomic orbitals

Describe how each pair is related.

10. frequency, amplitude

11. valence electron, electron-dot structure

12. principal energy levels, energy sublevels

CHAPTER ASSESSMENTCHAPTER 5

a. wavelength

b. photoelectric effect

c. photon

d. quantum

e. atomic orbital

f. atomic emissionspectrum

g. principal quantumnumber

h. ground state

i. electromagneticradiation

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Chapter Assessment Chemistry: Matter and Change • Chapter 5 21

Understanding Main Ideas (Part A)

Match the equation in Column A with its description in Column B.

Column A Column B

1. E � h�

2. c ��

3. � � h /m�

4. �E � Ehigher-energy orbit � Elower-energy orbit

Complete the table.

Write the orbital diagram and complete electron configuration for each atom.

9. nitrogen

10. fluorine

11. sodium


a. Relates the wavelength, frequency, andspeed of an electromagnetic wave

b. Describes the energy change of anelectron undergoing an orbit transition

c. Energy relationship developed by Planck

d. de Broglie’s equation

Principal Quantum Number, n Types of Orbitals Number of Orbitals Related to Principal Energy Level

5. 1 s

6.

7. 3 9

8. 4

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Understanding Main Ideas (Part B)

Circle the letter of the choice that best completes the statement or answers the question.Use the following figure to answer questions 1 and 2.

1. According to Bohr’s atomic model, which letter(s) in the figure represents a place wherean electron cannot be?

a. A b. B, C and E c. A and D d. D

2. According to the quantum mechanical model of the atom, point E in the figure represents a

a. point where an electron cannot be. c. position where an electron must be.

b. position where an electron probably is. d. point beyond which no electron can go.

3. What can you conclude from the figure on the right?

a. Hund’s rule has been violated.

b. The Pauli exclusion principle has been violated.

c. The Aufbau principle has been violated.

d. This is a valid orbital diagram.

4. What can you conclude from the figure on the right?

a. Hund’s rule has been violated.

b. The Pauli exclusion principle has been violated.

c. The Aufbau principle has been violated.

d. This is a valid orbital diagram.

5. Which of the following can you conclude based on the de Broglie equation?

a. Waves behave like particles. c. All matter has an associated wavelength.

b. Most particles are electrons. d. All matter behaves like particles.

6. Which of the following best describes the Heisenberg uncertainty principle?

a. Light behaves like a particle and like a wave.

b. The shorter the wavelength, the higher the frequency.

c. It is impossible to know both the velocity and the position of a particle at the same time.

d. You can measure an object without disturbing it.

BC

DE

A


)(

)(

2s

2p

1s

) )

)(

))

)(2s

2p

1s

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Thinking CriticallyAnswer the following questions.

1. A radio station has a frequency of 103.7 MHz. (1 MHz � 106 s�1) What is the wavelength of the radiation emitted by the station? Indicate where this wavelength falls on the electromagnetic spectrum shown below.

2. Look at the electromagnetic spectrum again. Are the microwaves used to cook foodhigher or lower in frequency than radio waves? Are microwaves longer or shorter inwavelength than radio waves?

3. Write the orbital diagram of aluminum.

4. Write the complete electron configuration and the noble-gas notation for aluminum.

5. Write the noble-gas notation for iodine.

6. Identify each atom.

a. 1s22s22p1 b. [Ar]4s1

7. Write electron-dot structures for the following atoms.

a. neon c. carbon

b. hydrogen d. sulfur

1019101810171016101510141013101210111010109

10–1110–1010–910–810–710–610–510–410–1 10–2 10–3102 101 1103 104

108107106105

UH

F-TV

Mic

row

ave

Infr

ared

Vis

ible

Ult

ravi

ole

t

X r

ays

� r

ays

� (Hz)

Electromagnetic SpectrumV

HF-

TVFM

rad

io

AM

rad

io

� (m)

Radio


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Applying Scientific MethodsA chemist isolated four samples, A, B, C, and D. She obtained the following atomic emissionspectra of the samples.

1. Examine each sample’s atomic emission spectra. Assume that each sample represents asingle element. What can you conclude by looking at the spectra? Do the samples repre-sent the same element or different elements?

2. Which part of the electromagnetic spectrum do the atomic emission spectra show?

3. Would the atomic emission spectrum for each sample change if you repeated the proce-dure? Explain your answer

4. What does each line in an atomic emission spectrum represent?

A

BC

D

400 500

Nanometers

600 700


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5. You find the following atomic emission spectrum for hydrogen in your textbook.Compare this spectrum to the spectra of the samples that the chemist obtained. What can you conclude? Explain your answer.

6. Which, if any, of the atomic emission spectra can the Bohr model explain? Explain youranswer.

7. According to Bohr’s model, how many times were photons emitted from the excitedatoms in each sample to produce its atomic emission spectrum?

A

B

C

D

8. The difference between successive energy levels becomes smaller as n becomes larger.Explain how hydrogen’s emission spectrum demonstrates this statement.

9. Assume that instead of measuring the photons emitted by each sample, the chemist meas-ured the photons absorbed by each sample. What would the absorption spectra look like?Explain your answer.

400 500

Nanometers

600 700

Applying Scientific Methods, continued


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26 Chemistry: Matter and Change • Chapter 5

Student Recording Sheet

Name Date Class

Standardized Test PracticeMultiple Choice

Select the best answer from the choices given, and fill in the corresponding circle.

1. 4. 7. 10.

2. 5. 8.

3. 6. 9.

Short Answer

Answer each question with complete sentences.

11.

12.

13.

14.

SAT Subject Test: Chemistry

15. 17. 19.

16. 18.

CHAPTER 5

Assessment

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Chapter 6 The Periodic Table and Periodic LawMiniLab . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 28

ChemLab . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 29


Math Skills Transparency Masters and Worksheets . . . . . . . . . . . . . . . . . . . . . . . . . 40

Study Guide . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 42

Chapter Assessment . . . . . . . . . . . . . . . . . . . . . . . . . . . . 46


Table ofContents

27

Reproducible Pages

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Can you find the pattern?

Materials Index cards, pencil


2. Make a set of element cards based on theinformation in the chart at right.

3. Organize the cards by increasing mass, andstart placing them into a 4 � 3 grid.

4. Place each card based on its properties, andleave gaps when necessary.

mini LAB 6Organize Elements

Ad 52.9 solid/liquid orange

Ax 108.7 ductile solid light blue

Bp 69.3 gas red

Cx 112.0 brittle solid light green

Lq 98.7 ductile solid blue

Pd 83.4 brittle solid green

Qa 68.2 ductile solid dark blue

Rx 106.9 liquid yellow

Tu 64.1 brittle solid hunter

Xn 45.0 gas crimson

Symbol Mass (g) State Color

Analysis

1. Make a table listing the placement of each element.

2. Describe the period (across) and group (down) trends for the color in your new table.

3. Describe the period and group trends for the mass in your new table. Explain your placement ofany elements that do not fit the trends.

4. Predict the placement of a newly found element, Ph, that is a fuchsia gas. What would be anexpected range for the mass of Ph?

5. Predict the properties for the element that would fill the last remaining gap in the table.

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CHEMLAB 6

Safety Precautions• Wear safety goggles and a lab apron at all times.• Do not handle elements with bare hands. Brittle samples might shatter

into sharp pieces.• 1.0M HCl is harmful to eyes and clothing.• Never test chemicals by tasting.• Follow any additional safety precautions provided by your teacher.

ProblemWhat is the pattern ofproperties of the representative elements?

Objectives• Observe properties of

various elements.• Classify elements as

metals, nonmetals, andmetalloids.

• Examine general trendswithin the periodic table.

Materialsstoppered test

tubes containingsmall samples ofelements

plastic dishes con-taining samplesof elements

conductivity apparatus

1.0M HCltest tubes (6)test-tube rack10-mL graduated

cylinderspatulasmall hammerglass marking

pencil

Investigate Descriptive ChemistryYou can observe several of the representative elements, classify them, and

compare their properties. The observation of the properties of elements iscalled descriptive chemistry.

Pre-Lab


2. Use the data table on the next page to record theobservations you make during the lab.

3. Examine the periodic table. What is the physicalstate of most metals? Nonmetals? Metalloids?

4. Look up the definitions of the terms luster, malleability, and electrical conductivity. To whatelements do they apply?

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Procedure


2. Observe and record the appearance (physicalstate, color, luster, texture, and so on) of the ele-ment sample in each test tube without removingthe stoppers.

3. Remove a small sample of each of the elementscontained in a plastic dish and place it on a hardsurface. Gently tap each element sample with asmall hammer. If the element is malleable, it willflatten. If it is brittle, it will shatter. Record yourobservations.

4. Use the conductivity tester to determine whichelements conduct electricity. Clean the electrodeswith water, and dry them before testing each element.

5. Label each test tube with the symbol for one ofthe elements in the plastic dishes. Using a gradu-ated cylinder, add 5 mL of water to each test tube.

6. Use a spatula to put a small amount of each ele-ment into the corresponding test tubes. Using agraduated cylinder, add 5 mL of 1.0M HCl to

each test tube. Observe each tube for at least 1minute. The formation of bubbles is evidence of areaction between the acid and the element.Record your observations.

7. Cleanup and Disposal Dispose of all materialsas instructed by your teacher.

CHEMLAB 6


1. Interpret Data Using the table above and your observations, list the element samples that display thegeneral characteristics of metals.

2. Interpret Data Using the table above and your observations, list the element samples that display thegeneral characteristics of nonmetals.

3. Interpret Data Using the table above and your observations, list the element samples that display thegeneral characteristics of metalloids.

Classification Properties

Metals • malleable• good conductor of electricity• lustrous• silver or white in color• many react with acids

Nonmetals • solids, liquids, or gases• do not conduct electricity• do not react with acids• likely brittle if solid

Metalloids • combine properties of metalsand nonmetals

Observation of Elements

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4. Model Construct a periodic table, and label the representative elements by group (1 through 17). Using your results and the periodic table presented in this chapter, record the identities of elementsobserved during the lab in the periodic table you have constructed.

5. Infer Describe any trends among the elements you observed in the lab.

Inquiry Extension

Investigate Were there any element samples that did not fit into one of the three categories? What additional investigations could you conduct to learn even more about these elements’ characteristics?

CHEMLAB 6

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Name Date Class


Hyd

roge

n

1 H1.

008

Lith

ium

3 Li6.

941

Sodi

um

11 Na

22.9

90

Pota

ssiu

m

19 K39

.098

Rubi

dium

37 Rb

85.4

68

Ces

ium

55 Cs

132.

905

Fran

cium

87 Fr(2

23)

Radi

um

88 Ra

(226

)

Cer

ium

58 Ce

140.

115

Thor

ium

90 Th

232.

038

Ura

nium

92 U23

8.02

9

Nep

tuni

um

93 Np

(237

)

Plut

oniu

m

94 Pu

(244

)

Am

eric

ium

95 Am

(243

)

Neo

dym

ium

60 Nd

144.

242

Prom

ethi

um

61 Pm

(145

)

Sam

ariu

m

62 Sm15

0.36

Euro

piu

m

63 Eu15

1.96

5

Pra

seod

ymiu

m

59 Pr

140.

908

Pro

tact

iniu

m

91 Pa

231.

036

Act

iniu

m

89 Ac

(227

)

Rut

herf

ordi

um

104

Rf

(261

)

Bariu

m

56 Ba

137.

327

Lant

hanu

m

57 La13

8.90

5

Haf

nium

72 Hf

178.

49

Tant

alum

73 Ta

180.

948

Dub

nium

105

Db

(262

)

Seab

orgi

um

106

Sg (266

)

Has

sium

108

Hs

(277

)

Mei

tner

ium

109

Mt

(268

)

Bohr

ium

107

Bh

(264

)

Tung

sten

74 W18

3.84

Rhen

ium

75 Re

186.

207

Osm

ium

76 Os

190.

23

Irid

ium

77 Ir19

2.21

7

Stro

ntiu

m

38 Sr87

.62

Yttr

ium

39 Y88

.906

Zirc

oniu

m

40 Zr

91.2

24

Nio

bium

41 Nb

92.9

06

Mol

ybde

num

42 Mo

95.9

4

Cal

cium

20 Ca

40.0

78

Scan

dium

21 Sc44

.956

Tita

nium

22 Ti

47.8

67

Vana

dium

23 V50

.942

Chr

omiu

m

24 Cr

51.9

96

Tech

netiu

m

43 Tc

(98)

Ruth

eniu

m

44 Ru

101.

07

Man

gane

se

25 Mn

54.9

38

Iron 26 Fe

55.8

47

Cob

alt

27 Co

58.9

33

Rhod

ium

45 Rh

102.

906

Mag

nesi

um

12 Mg

24.3

05

Bery

llium

4 Be

9.01

2

Lan

than

ide

seri

es

Act

inid

e se

ries

1

12

2 3 4 5 6 7

9

18

34

56

78

Hel

ium

2 He

4.00

3

Cur

ium

96 Cm

(247

)

Berk

eliu

m

97 Bk

(247

)

Cal

iforn

ium

98 Cf

(251

)

Eins

tein

ium

99 Es(2

52)

Ferm

ium

100

Fm (257

)

Nob

eliu

m

102

No

(259

)

Law

renc

ium

103

Lr(2

62)

Men

dele

vium

101

Md

(258

)

Gad

olin

ium

64 Gd

157.

25

Terb

ium

65 Tb

158.

925

Dys

pro

sium

66 Dy

162.

50

Hol

miu

m

67 Ho

164.

930

Erbi

um

68 Er16

7.25

9

Thul

ium

69 Tm

168.

934

Ytte

rbiu

m

70 Yb

173.

04

Lute

tium

71 Lu17

4.96

7

Plat

inum

78 Pt

195.

08

Gol

d

79 Au

196.

967

Mer

cury

80 Hg

200.

59

Thal

lium

81 Tl

204.

383

Lead 82 Pb

207.

2

Bism

uth

83 Bi

208.

980

Ast

atin

e

85 At

209.

987

Rado

n

86 Rn

222.

018

Nic

kel

28 Ni

58.6

93

Cop

per

29 Cu

63.5

46

Zin

c

30 Zn

65.3

9

Gal

lium

31 Ga

69.7

23

Ger

man

ium

32 Ge

72.6

1

Ars

enic

33 As

74.9

22

Sele

nium

34 Se78

.96

Brom

ine

35 Br

79.9

04

Kryp

ton

36 Kr

83.8

0

Palla

dium

46 Pd

106.

42

Silv

er

47 Ag

107.

868

Cad

miu

m

48 Cd

112.

411

Indi

um

49 In11

4.82

Tin

50 Sn11

8.71

0

Ant

imon

y

51 Sb12

1.75

7

Tellu

rium

52 Te

127.

60

Iodi

ne

53 I12

6.90

4

Xen

on

54 Xe

131.

290

Alu

min

um

13 Al

26.9

82

Silic

on

14 Si28

.086

Phos

pho

rus

15 P30

.974

Sulfu

r

16 S32

.066

Chl

orin

e

17 Cl

35.4

53

Arg

on

18 Ar

39.9

48

Boro

n

5 B10

.811

Car

bon

6 C12

.011

Nitr

ogen

7 N14

.007

Oxy

gen

8 O15

.999

Fluo

rine

9 F18

.998

Neo

n

10 Ne

20.1

80

10

11

12

13

14

15

16

17

Polo

nium

84 Po

208.

982

Dar

mst

adtiu

m

110

Ds

(281

)

Roen

tgen

ium

111

Rg

(272

)

Unu

nqua

dium

114

Uu

q(2

89)

Unu

nhex

ium

116

Uu

h(2

91)

Unu

ntriu

m

113

Uu

t(2

84)

Unu

npen

tium

115

Uu

p(2

88)

Unu

noct

ium

118

Uu

o(2

94)

2Hydr

ogen

1 H 1.00

8

Ele

me

nt

Ato

mic

nu

mb

er

Sta

te o

fm

atte

r

Me

tal

Me

tallo

id

No

nm

eta

l

Gas

Liq

uid

So

lid

Syn

the

tic

Sym

bo

l

Ato

mic

mas

sR

ece

ntl

yo

bse

rve

d

PER

IOD

IC T

AB

LE O

F TH

E E

LEM

EN

TS

Unu

nbiu

m

112

Uu

b(2

85)

The

num

ber i

n pa

rent

hese

s is t

he m

ass n

umbe

r of t

he lo

nges

t liv

ed is

otop

e fo

r tha

t ele

men

t.*

**

**

**

The Periodic TableThe Periodic Table



18

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1. How many elements are listed in the periodic table?

2. What is the atomic number of selenium?

3. What is the symbol for palladium?

4. What is the atomic mass of strontium?

5. How are elements that are gases at room temperature designated in the periodic table?

6. How many columns of elements does the periodic table contain?

7. What is another name for a column of elements?

8. How many rows of elements does the periodic table contain?

9. What is another name for a row of elements?

10. Which period contains the least number of elements?

11. What element is found in period 4, group 7?

12. How are metals designated in the periodic table?

13. How are metalloids designated in the periodic table?

14. How are nonmetals designated in the periodic table?

15. What is the name of the group 1 elements (excluding hydrogen)?

16. What is the name of the group 2 elements?



19. What can be said about the electron configurations of all the elements in a group?

The Periodic TableThe Periodic Table



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s1 1 H

2 He

3 Li 11 Na

19 K 37 Rb 55 Cs

4 Be 12 Mg

20 Ca 38 Sr 56 Ba

87 Fr

58 Ce 90 Th

59 Pr 91 Pa

60 Nd

92 U

61 Pm 93 Np

62 Sm 94 Pu

63 Eu 95 Am

64 Gd

96 Cm

65 Tb 97 Bk

66 Dy

98 Cf

67 Ho 99 Es

68 Er 100

Fm

69 Tm 101

Md

70 Yb

102

No

71 Lu 103

Lr

88 Ra

5 B 13 Al

31 Ga

49 In 81 Tl

6 C 14 Si 32 Ge

50 Sn 82 Pb

7 N 15 P 33 As

51 Sb 83 Bi

8 O 16 S 34 Se 52 Te 84 Po

9 F 17 Cl 35 Br 53 I 85 At

10 Ne

18 Ar

36 Kr

54 Xe

86 Rn

21 Sc 39 Y 57 La 89 Ac

22 Ti 40 Zr 72 Hf

104

Rf

23 V 41 Nb

73 Ta 105

Db

24 Cr 42 Mo

74 W 106

Sg

25 Mn

43 Tc 75 Re

107

Bh

26 Fe 44 Ru 76 Os

108

Hs

27 Co 45 Rh 77 Ir 109

Mt

110

Uun28 Ni

46 Pd 78 Pt

111

Uuu29 Cu 47 Ag

79 Au

112

Uub30 Zn 48 Cd 80 Hg

s2p

1p

2p

3p

4p

5p

6s2

s b

lock

d b

lock

p b

lock

f b

lock

The s-, p-, d-, and f-Block ElementsThe s-, p-, d-, and f-Block Elements



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1. What are the four sections, or blocks, of the periodic table?

2. What does each block represent?

3. What do elements in the s-block have in common?

4. What is the valence electron configuration of each element in group 1?

5. What is the valence electron configuration of each element in group 2?

6. Why does the s-block span two groups of elements?

7. Why does the p-block span six groups of elements?

8. Why are there no p-block elements in period 1?

9. What is the ending of the electron configuration of each element in group 4?

10. What is the electron configuration of neon?

11. In what period does the first d-energy sublevel appear?

12. Why does the d-block span ten groups of elements?

13. What is the ending of the electron configuration of each element in group 3?

14. What is the electron configuration of titanium?

15. In what period does the first f-energy sublevel appear?

16. Determine the group, period, and block for the element having the electron configuration[Xe]4f145d106s26p3.

a. group b. period c. block

The s-, p-, d-, and f-Block ElementsThe s-, p-, d-, and f-Block Elements



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Li 76

Na 102

K 138

Rb 152

Cs 167

Be 31

Mg 72

Ca 100

Sr 118

Ba 135

B 20

Al 54

Ga 62

n 81

Tl 95

C 15

Si 41

Ge 53

Sn 71

Pb 84

N 146

P 212

As 222

Sb 62

Bi 74

O 140

S 184

Se 198

Te 221

F 133

Cl 181

Br 195

220

K 138

1�

1

Chemical symbol

Charge

Ionic radius

Transitionmetals

Relative size

2

3

4

5

6

1�

1�

1�

1�

1�

2

Peri

od

13 14 15 16 17

2�

2�

2�

2�

2 �

3�

3�

3�

3�

3�

4�

4�

4�

4�

4�

3�

3�

3�

5�

5�

2�

2�

2�

2�

1�

1�

1�

1�

Radii are given in picometers (1 � 10�12 m)

Atomic and Ionic RadiiAtomic and Ionic Radii



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1. Which groups and periods of elements are shown in the table of atomic radii?

2. In what unit is atomic radius measured? Express this unit in scientific notation.

3. What are the values of the smallest and largest atomic radii shown? What elements havethese atomic radii?

4. What happens to atomic radii within a period as the atomic number increases?

5. Cite any exceptions to the generalization you stated in your answer to question 4.

6. What accounts for the trend in atomic radii within a period?

7. What happens to atomic radii within a group as the atomic number increases?

8. Cite any exceptions to the generalization you stated in your answer to question 7.

9. What accounts for the trend in atomic radii within a group?

10. In the table of ionic radii, how is the charge of the ions of elements in groups 1 and 2 related to the group number of the elements?

Atomic and Ionic RadiiAtomic and Ionic Radii



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K Rb

0 10

Period 2 Period 3 Period 4 Period 5

20 30 50 6040

Firs

t io

niz

atio

n e

ner

gy

(kJ/

mo

l)

2500

2000

1500

1000

500

0

Atomic number

First Ionization Energy of Elements in Periods 1–5

H

He

Li

Ne

Ar

Na

XeKr

Li

Be

B

C

N

O

F

Ne

1

2

3

4

5

6

7

8

1st

520

900

800

1090

1400

1310

1680

2080

2nd

7300

1760

2430

2350

2860

3390

3370

3950

3rd

14,850

3660

4620

4580

5300

6050

6120

4th

25,020

6220

7480

7470

8410

9370

5th

37,830

9440

10,980

11,020

12,180

6th

53,270

13,330

15,160

15,240

7th

71,330

17,870

20,000

8th

92,040

23,070

9th

115,380

Element

Successive Ionization Energies for the Period 2 Elements

Ionization energy (kJ/mol)*Valenceelectrons

* mol is an abbreviation for mole, a quantity of matter.

First Ionization and SuccessiveIonization EnergiesFirst Ionization and SuccessiveIonization Energies



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1. What is meant by first ionization energy?

2. Which element has the smallest first ionization energy? The largest? What are their values?

3. What generally happens to the first ionization energy of the elements within a period asthe atomic number of the elements increases?

4. What accounts for the general trend in the first ionization energy of the elements within a period?

5. What happens to the values of the successive ionization energies of an element?

6. Based on the graph, rank the group 2 elements in periods 1–5 in decreasing order of firstionization energy.

7. How is a jump in ionization energy related to the valence electrons of the element?

8. What generally happens to the first ionization energy of the elements within a group asthe atomic number of the elements increases?

9. What accounts for the general trend in the first ionization energy of the elements within a group?

First Ionization and SuccessiveIonization EnergiesFirst Ionization and SuccessiveIonization Energies



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Using the Periodic TableUsing the Periodic Table



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1. Identify the number of valence electrons in each of the following elements.

a. Ne e. O

b. K f. Cl

c. B g. P

d. Mg h. Si

2. Identify the energy level of the valence electrons in each of the following elements.

a. Br

b. N

c. Ra

d. H

e. Ar

f. I

3. Use the periodic table to write the electron configurations (using noble gas notation) foreach of the following elements.

a. Li

b. F

c. As

d. Sr

e. Bi

4. Determine the group, period, and block of the elements having the following electron configurations.

a. 1s2

b. [Ne]3s23p1

c. [Ar]4s1

d. [Kr]5s24d1

e. [Xe]6s24f145d106p4

Using the Periodic TableUsing the Periodic Table



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The Periodic Table and Periodic LawThe Periodic Table and Periodic Law

Section 6.1 Development of the Modern Periodic TableIn your textbook, reads about the history of the periodic table’s development.


The table below was developed by John Newlands and is based on a relationship called

the law of (1) . According to this law, the properties of the elements

repeated every (2) elements. Thus, for example, element two and

element (3) have similar properties. The law of octaves did not work

for all the known elements and was not generally (4) .

The first periodic table is mostly credited to (5) . In his table, the

elements were arranged according to increasing (6) . One important

result of this table was that the existence and properties of undiscovered

(7) could be predicted.

The element in the modern periodic table are arranged according to increasing

(8) , as a result of the work of (9) . This

arrangement is based on number of (10) in the nucleus of an atom of

the element. The modern form of the periodic table results in the

(11) , which states that when elements are arranged according to

increasing atomic number, there is a periodic repetition of their chemical and physical

(12) .


octaves atomic mass atomic number nine

elements properties Henry Moseley eight

protons periodic law Dmitri Mendeleev accepted

1 2 3 4 5 6 7

H Li G Bo C N O

8 9 10 11 12 13 14

F Na Mg Al Si P S

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In your textbook, read about the modern periodic table.

Use the information in the box on the left taken from the periodic table to complete thetable on the right.


Column A Column B

18. A column on the periodic table

19. A row on the periodic table

20. Elements in groups 1, 2, and 13 to 18

21. Elements that are shiny and conduct electricity

22. Elements in groups 3 to 12

In the space at the left, write true if the statement is true; if the statement is false,change the italicized word or phrase to make it true.

23. There are two main classifications of elements.

24. More than three-fourths of the elements in the periodic table arenonmetals.

25. Group 1 elements (except for hydrogen) are known as the alkalimetals.

26. Group 13 elements are the alkaline earth metals.

27. Group 17 elements are highly reactive nonmetals known ashalogens.

28. Group 18 elements are very unreactive elements known astransition metals.

29. Metalloids have properties of both metals and inner transitionmetals.



a. metals

b. group

c. period

d. representative elements

e. transition elements

Atomic Mass 13.

Atomic Number 14.

Electron Configuration 15.

Chemical Name 16.

Chemical Symbol 17.

7

N

Nitrogen

14.007

[He]2s22p3

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Section 6.2 Classification of the ElementsIn your textbook, read about organizing the elements by electron configuration.

Use the periodic table on pages 178–179 in your textbook to match each element inColumn A with the element in Column B that has the most similar chemical properties.

Column A Column B

1. arsenic (As)

2. bromine (Br)

3. cadmium (Cd)

4. gallium (Ga)

5. germanium (Ge)

6. iridium (Ir)

7. magnesium (Mg)

8. neon (Ne)

9. nickel (Ni)

10. osmium (Os)

11. sodium (Na)

12. tellurium (Te)

13. tungsten (W)

14. yttrium (Y)

15. zirconium (Zr)


16. Why do sodium and potassium, which belong to the same group in the periodic table,have similar chemical properties?

17. How is the energy level of an element’s valence electrons related to its period on the periodic table? Give an example.


a. boron (B)

b. cesium (Cs)

c. chromium (Cr)

d. cobalt (Co)

e. hafnium (Hf)

f. iodine (I)

g. iron (Fe)

h. nitrogen (N)

i. platinum (Pt)

j. scandium (Sc)

k. silicon (Si)

l. strontium (Sr)

m. sulfur (S)

n. zinc (Z)

o. xenon (Xe)

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In your textbook, read about s-, p-, d-, and f-block elements.

Use the periodic table on pages 178–179 in your textbook and the periodic table below toanswer the following questions.

18. Into how many blocks is the periodic table divided?

19. What groups of elements does the s-block contain?

20. Why does the s-block portion of the periodic table span two groups?

21. What groups of elements does the p-block contain?

22. Why are members of group 18 virtually unreactive?

23. How many d-block elements are there?

24. What groups of elements does the d-block contain?

25. Why does the f-block portion of the periodic table span 14 groups?

26. What is the electron configuration of the element in period 3, group 16?

s1

1H

2He

3Li

11Na

19K

37Rb

55Cs

4Be

12Mg

20Ca

38Sr

56Ba

87Fr

57La

89Ac

58Ce

90Th

59Pr

91Pa

60Nd

92U

61Pm

93Np

62Sm

94Pu

63Eu

95Am

64Gd

96Cm

65Tb

97Bk

66Dy

98Cf

67Ho

99Es

68Er

100Fm

69Tm

101Md

70Yb

102No

88Ra

5B

13Al

31Ga

49In

81Tl

6C

14Si

32Ge

50Sn

82Pb

7N

15P

33As

51Sb

83Bi

8O

16S

34Se

52Te

84Po

9F

17Cl

35Br

53I

85At

10Ne

18Ar

36Kr

54Xe

86Rn

21Sc

39Y

71Lu

103Lr

22Ti

40Zr

72Hf

104Rf

23V

41Nb

73Ta

105Db

24Cr

42Mo

74W

106Sg

25Mn

43Tc

75Re

107Bh

26Fe

44Ru

76Os

108Hs

27Co

45Rh

77Ir

109Mt

110Uun

28Ni

46Pd

78Pt

111Uuv

29Cu

47Ag

79Au

112Uub

30Zn

48Cd

80Hg

s2 p1 p2 p3 p4 p5 p6

s2

s block

d block

p block

f block



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Section 6.3 Periodic TrendsIn your textbook, read about atomic radius and ionic radius.


1. Atomic radii cannot be measured directly because the electron cloud surrounding thenucleus does not have a clearly defined

a. charge. b. mass. c. outer edge. d. probability.

2. Which diagram best represents the group and period trends in atomic radii in the periodictable?

a. c.

b. d.

3. The general trend in the radius of an atom moving down a group is partially accounted

for by the

a. decrease in the mass of the nucleus. c. increase in the charge of the nucleus.

b. fewer number of filled orbitals. d. shielding of the outer electrons by inner electrons.

4. A(n) is an atom, or bonded group of atoms, that has a positive or negativecharge.

a. halogen b. ion c. isotope d. molecule

5. An atom becomes negatively charged by

a. gaining an electron. b. gaining a proton. c. losing an electron. d. losing a neutron.

6. Which diagram best represents the relationship between the diameter of a sodium atomand the diameter of a positive sodium ion?

a. b. c.

Na Na� Na Na� Na Na�

Generally decrease

Gen

eral

lyd

ecre

ase

Generally increase

Gen

eral

lyd

ecre

ase

Generally decrease

Gen

eral

lyin

crea

se

Generally increase

Gen

eral

lyin

crea

se


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In your textbook, read about ionization energy and electronegativity.


7. What is ionization energy?

8. Explain why an atom with a high ionization-energy value is not likely to form a positiveion.

9. What is the period trend in the first ionization energies? Why?

10. What is the group trend in the first ionization energies? Why?

11. State the octet rule.

12. What does the electronegativity of an element indicate?

13. What are the period and group trends in electronegativities?



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The Periodic Table and Periodic LawThe Periodic Table and Periodic Law


Column A Column B

1. Statement that when the elements are arranged byincreasing atomic number, there is a periodic repetition oftheir chemical and physical properties

2. Groups 1 and 2, 13 through 18

3. Groups 3 through 12

4. Group 1 elements (except for hydrogen)

5. Group 2 elements

6. A column in the periodic table

7. A row in the periodic table

8. Group 17 elements

9. Group 18 elements

10. Atom or bonded group of atoms that has a positive ornegative charge

11. Energy required to remove an electron from a gaseousatom

12. Statement that atoms tend to gain, lose, or share electronsto acquire a full set of eight valence electrons

13. Indication of an atom’s ability to attract electrons in achemical bond

Write a sentence that uses each group of terms.

14. transition metals, inner transition metals

15. metal, nonmetal, metalloid


a. alkali metals

b. alkaline earthmetals

c. electronegativity

d. halogens

e. period

f. ion

g. ionization energy

h. noble gases

i. octet rule

j. periodic law

k. representativeelements

l. transition elements

m. group

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Briefly describe the contribution of each of the following to the development of the periodic table.

1. John Newlands:

2. Henry Moseley:

3. Dmitri Mendeleev:

Match each of the following terms with a number or chemical symbol from the periodictable below.

4. alkali metals 10. a metalloid element

5. alkaline earth metals 11. noble gases

6. a d-block element 12. a p-block element that is not a metalloid

7. an f-block element 13. an s-block element

8. halogens 14. transition metals

9. inner transition metals

1 2

3

Th6

Ni

Si

O

4 5

Be


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Circle the letter of the choice that best completes the statement or answers the question.Use the periodic table in your textbook.

1. Elements in the same group have the same

a. atomic radius. c. nuclear charge.

b. energy level of outer electrons. d. number of valence electrons.

2. Most of the elements in groups 16 through 18 are classified as

a. alkali metals. c. nonmetals.

b. inner transition metals. d. alkaline earth metals.

3. Which energy level of the period 4 transition elements is being filled with electrons?

a. third b. fourth c. fifth d. sixth

4. Identify the period and group of the element that has the electron configuration[Ne]3s23p3.

a. period 2, group 2 b. period 3, group 1 c. period 3, group 13 d. period 3, group 15

5. Which of the following classifications describes the element with the electron configuration [Ar]4s23d104p5?

a. stable metal b. stable nonmetal c. unstable nonmetal d. unstable metal

6. What is the electron configuration of the element in group 14 and period 4 of the periodic table?

a. [Ne]3s23p4 b. [Ar]4s2 c. [Ar]4s23d104p2 d. [Kr]5s24d2

7. What is the trend in atomic radii as you move from left-to-right across a period?

a. generally decreases b. generally increases c. remains the same d. varies randomly

8. The trend in the atomic radii as you move down the group 1 elements is partially due to

a. decreased distance of outer electrons.

b. increased nuclear charge.

c. increased number of electrons in outer energy level.

d. shielding by inner electrons.

9. In which of the following pair is the second particle listed larger than the first?

a. K, Ga b. Pb, C c. Br, Br� d. Li, Li�

10. How many electrons does an atom generally need in its outer level to be the most stable?

a. 4 b. 8 c. 10 d. 12

11. Which of the following electron configurations represents the most chemically stableatom?

a. [He]2s22p3 b. [Ne]3s23p5 c. [Ne]3s23p64s23d5 d. [Ne]3s23p6


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Thinking CriticallyThe graph below shows the atomic radii of the elements in the first four periods of theperiodic table, as well as the major ionic radii of the representative elements. The chargeof each ion is indicated above the plotted point representing its radius. Use the graphand the periodic table in your textbook to answer the following questions.

1. Describe the relationship between atomic radii and atomic number for the representative elements in the first four periods.

2. How can you account for the trend you described in your answer to question 1?

3. Describe the relationship between the atomic radii and the atomic numbers of the transi-tion elements in period 4.

4. Explain why the two elements in the first period do not have ionic radii listed.

5. Predict whether the arsenic ion shown in the graph has a positive or negative charge.(Arsenic has an atomic number of 33.) Explain your prediction.

representative element

transition element

representative ion0

50

100

150

200

250

0 5

Period 1Period 2

Atomicradius

Period 3 Period 4

10 15Atomic number

Rad

ius

(pm

)

20 25 30 35

1�

1�

1� 2�

2�

2�3�

3�

4�

4� ?

4�

1� 1�

1�

2� 2�

2�

3�

3�

3�

Ionic radius


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Applying Scientific MethodsAt the beginning of the nineteenth century, chemists were searching for numerical relation-ships among the elements. From these relationships, they hoped that some fundamental chem-ical principle might be revealed. One such chemist was the German, Johann WolfgangDöbereiner. In 1817, Döbereiner noted that if the three alkaline earth metals Ca, Sr, and Bawere arranged in increasing atomic mass, the atomic mass of the middle element was close tothe average of the other two atomic masses, as shown below.

In 1829, Döbereiner discovered that the halogens—Cl, Br, and I—also followed a similar pat-tern, as shown below. He named these three-member groups of elements with similar chemi-cal and physical properties triads.

1. Six of the eight elements in the table below make up two of Döbereiner’s triads. Plot theatomic mass of each element on the number line below the table. From the sequence ofthe atomic masses and your knowledge of elements with similar chemical and physicalproperties, identify the three elements in each of the two triads. Explain your choices.

0 10 20 30 40 50 60 70 80 90 100 110 120 130 140 150 160 170 180 190 200

Atomic mass (amu)

0 10 20 30 40 50 60 70 80 90 100 110 120 130 140 150

Atomic mass (u)

Chlorine (Cl) Bromine (Br) Iodine (I)

0 10 20 30 40 50 60 70 80 90 100 110 120 130 140 150

Atomic mass (u)

Calcium (Ca) Strontium (Sr) Barium (Ba)


Element Mass (amu)

Lithium (Li) 7

Carbon (C) 12

Sodium (Na) 23

Sulfur (S) 32

Potassium (K) 39

Selenium (Se) 79

Tellurium (Te) 128

Gold (Au) 197

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2. Recall that atomic mass is a quantitative physical property of an element. So are density,melting point, and boiling point. Use the table below to sequence the values of the densi-ties, melting points, and boiling points of the following triads—Ca, Sr, and Ba; and Cl,Br, and I. Then determine whether each property shows a trend similar to that of theatomic masses of the elements in the triads. Explain your reasoning.

3. If silicon (Si), germanium (Ge), and tin (Sn) are classified as a triad similar to those ofDöbereiner, predict values that will complete the following table. Record the values inthe table.



Element Density (g/mL) Melting Point (°C) Boiling Point (°C)

Barium (Ba) 3.62 726.9 1845

Bromine (Br) 3.11 �7.25 59.35

Calcium (Ca) 1.55 841.5 1500.5

Chlorine (Cl) 0.003 214 �101 �34

Iodine (I) 4.93 113.6 184.5

Strontium (Sr) 2.6 776.9 1412

Element Atomic Mass (amu) Density (g/mL) Melting Point (°C)

Silicon (Si) 28 1411

Germanium (Ge) 5.3 945

Tin (Sn) 119 7.3

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CHAPTER 6

Assessment



1. 4. 7. 10.

2. 5. 8.

3. 6. 9.

Short Answer


11.

12.

13.

Extended Response


14.

15.


16.

17.

18.

19.

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Chapter 7 Ionic Compounds and MetalsMiniLab . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 56

ChemLab . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 57


Math Skills Transparency Masters and Worksheets . . . . . . . . . . . . . . . . . . . . . . . . . . 68

Study Guide . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 72

Chapter Assessment. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 78


Table ofContents

55

Reproducible Pages

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mini LAB 7Observe Properties

Analysis

1. Analyze your results, and identify the two types of steel that appear to have their properties com-bined in tempered steel.

2. Hypothesize how the different observed properties relate to crystal size.

3. State a use for spring steel that takes advantage of its unique properties.

4. Infer the advantages and disadvantages of using softened steel for body panels on automobiles.

5. Apply What is the major disadvantage of hardened steel? Do you think hardened steel would bewear-resistant and retain a sharpened edge? Explain your reasoning.

How do the properties of steel change when it is subjected to different types ofheat treatment? For many centuries, people have treated metals with heat to changetheir properties. The final properties of the metal depend on the temperature to which themetal is heated and the rate at which it cools.

Materials laboratory burner, forceps (2), hairpins (3), 250-mL beaker, cold water

Procedure

1. Read and complete the lab safety form.2. Examine a property of spring steel by try-

ing to bend open one of three hairpins.Record your observations.

3. Next hold each end of the hairpin with apair of forceps. Place the curved centralloop portion of the hairpin in the top ofthe blue flame from a laboratory burner.When the metal turns red, pull the hair-pin open to form a straight piece ofmetal. Allow it to cool as you record yourobservations. Repeat Step 3 for theremaining two hairpins. WARNING: Donot touch the hot metal. Do not holdyour hand above the flame of the labo-ratory burner.

4. To make softened steel, use a pair of for-ceps to hold all three hairpins verticallyin the flame of the laboratory burneruntil the hairpins glow red all over.Slowly raise the three hairpins straightup and out of the flame so they coolslowly. Slow cooling results in the forma-tion of large crystals.

5. After cooling, bend each of the three hairpins into the shape of the let-ter J. Record how the metal feels as youbend it.

6. To harden the steel, use tongs to holdtwo of the bent hairpins in the flameuntil they are glowing red all over.Quickly plunge the hot metals into a250-mL beaker containing approximately200 mL of cold water. Quick-coolingcauses the crystal size to be small.

7. Attempt to straighten one of the bends.Record your observations.

8. To temper the steel, use tongs to holdthe remaining hardened metal bendabove the flame for a brief period oftime. Slowly move the metal back andforth just above the flame until the graymetal turns to an iridescent blue-graycolor. Do not allow the metal to becomehot enough to glow red. Slowly cool themetal and then try to unbend it usingthe end of your finger. Record yourobservations.

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CHEMLAB 7

Safety Precautions• Always wear safety glasses and a lab apron.• Do not look directly at the burning magnesium. The intensity of the

light can damage your eyes.• Avoid handling heated materials until they have cooled.

ProblemCan the physical propertiesof a compound indicatethat they have ionic bonds?

Objectives• Observe evidence of a

chemical reaction.• Acquire and analyze

information that willenable you to decide if acompound has an ionicbond.

• Classify the products asionic or not ionic.

Materialsmagnesium ribbon

(25 cm)cruciblering stand and ringclay triangleBunsen burnerstirring rod

crucible tongscentigram balance100-mL beakerdistilled waterconductivity tester

Synthesize an Ionic CompoundElements combine to form compounds. If energy is released as the

compound is formed, the resulting product is more stable than thereacting elements. In this investigation, you will react elements to formtwo compounds. You will test the compounds to determine several of their properties. Ionic compounds have properties that are different from those of other compounds. You will decide if theproducts you formed are ionic compounds.

Pre-Lab

1. Read the entire CHEMLAB. Identify the variable. List any conditions that must be keptconstant.

2. Write the electron configuration of the magne-sium atom.

a. Based on this configuration, will magnesiumlose or gain electrons to become a magnesiumion?

b. Write the electron configuration of the magnesium ion.

c. The magnesium ion has an electron configuration like that of which noble gas?

3. Repeat question 2 for oxygen and nitrogen.

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4. Use the data table in the next column.

5. In your data table, which mass values will bemeasured directly? Which mass values will be calculated?

6. Explain what must be done to calculate eachmass value that is not measured directly.

Procedure


2. Record all measurements in your data table.

3. Position the ring on the ring stand about 7 cmabove the top of the Bunsen burner. Place theclay triangle on the ring.

4.Measure the mass of the clean, dry crucible.

5.Roll 25 cm of magnesium ribbon into a looseball. Place it in the crucible. Measure the mass ofthe magnesium and crucible together.

6.Place the crucible on the triangle, and heat it with a hot flame (flame tip should be near thecrucible).

7.Turn off the burner as soon as the magnesiumignites and begins to burn with a bright whitelight. Allow it to cool, and measure the mass ofthe magnesium product and the crucible.

8.Place the dry, solid product in the beaker.

9.Add 10 mL of distilled water to the beaker, andstir. Check the mixture with a conductivity tester.

10.Cleanup and Disposal Dispose of the productas directed by your teacher. Wash out the cruciblewith water. Return all lab equipment to its properplace.

CHEMLAB 7

Material(s) Mass (g)

Empty crucible

Crucible and Mg ribbon before heating

Magnesium ribbon

Crucible and magnesium products after heating

Magnesium products

Mass Data


1. Analyze Data Calculate the mass of the ribbon and the product. Record these masses in your table.

2. Classify the forms of energy released. What can you conclude about the stability of products?

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3. Infer Does the magnesium react with the air?

4. Predict the ionic formulas for the two binary products formed, and write their names.

5. Analyze and Conclude The product of the magnesium-oxygen reaction is white, whereas the product ofthe magnesium-nitrogen reaction is yellow. Which compound makes up most of the product?

6. Analyze and Conclude Did the magnesium compounds conduct a current when in solution? Do theseresults verify that the compounds are ionic?

7. Error Analysis If the results show that the magnesium lost mass instead of gaining mass, cite possiblesources of the error.

Inquiry Extension

Design an Experiment If the magnesium compounds conduct a current in solution, can you affect how well they conduct electricity? If they did not conduct a current, could they? Design an experiment tofind out.

CHEMLAB 7

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20 e

lect

ron

s (2

0�)

Ener

gy

Ca

ato

m1s

2 2s2 2

p6 3

s2 3p

6 4s2

Ca

ion

(C

a2�)

O a

tom

1s2 2

s2 2p

4

20 p

roto

ns

(20�

)

20 p

roto

ns

(20�

)

8 el

ectr

on

s (8

�)

8 p

roto

ns

(8�

)

O io

n (

O2�

)

10 e

lect

ron

s (1

0�)

8 p

roto

ns

(8�

)

18 e

lect

ron

s (1

8�)

Two

elec

tro

ns

+2e

�

+

Two

elec

tro

ns

2e�

+

Ener

gy

+

Formation of IonsFormation of Ions



22

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1. What are the names of the two elements shown?

2. Are the elements shown on the left sides of the two equations neutral? How can you tell?

3. What is the name for the energy needed to remove electrons from an atom, such as theCa atom shown?

4. What kind of charge does the Ca atom take on as a result of the reaction? What is thename for an ion with that kind of charge?

5. What kind of charge does the O atom take on as a result of the reaction? What is thename for an ion with that kind of charge?

6. Is the outer electron configuration of the Ca atom before the reaction a very stable one?How can you tell?

7. Is the outer electron configuration of the O atom before the reaction a very stable one?How can you tell?

8. Is the outer electron configuration of the Ca ion after the reaction a very stable one? Howcan you tell?

9. Is the outer electron configuration of the O ion after the reaction a very stable one? Howcan you tell?

10. What is the electron configuration of the Ca ion? What neutral atom has the same config-uration, and in what chemical family is it located in the periodic table?

11. What is the electron configuration of the O ion? What neutral atom has the same config-uration, and in what chemical family is it located in the periodic table?

Formation of IonsFormation of Ions



22

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e�

e�

Mg

+

C

l

+

Cl

Mg

+

C

l

+

Cl

Mg

+

S ?

Mg

+

P ?

2��

Ionic BondsIonic Bonds



23

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1. How many valence electrons does a neutral magnesium (Mg) atom have?

2. What is the charge on a magnesium ion? What does magnesium have to do to form suchan ion, and why does it tend to do so?

3. How many valence electrons does a single neutral chlorine (Cl) atom have?

4. What is the charge on a chloride ion? What does chlorine have to do to form such an ion,and why does it tend to do so?

5. How many magnesium atoms and how many chlorine atoms react to form one formulaunit of magnesium chloride? Why? What is the formula of magnesium chloride?

6. What kind of compound is magnesium chloride? What happens to electrons during theformation of the compound? What holds the atoms together in the compound?

7. What is the formula of the ionic compound formed by magnesium and sulfur (S) atoms?Explain why, in terms of electron transfer, stability, and overall charge.

8. What is the formula of the ionic compound formed by magnesium and phosphorus (P)atoms? Explain why, in terms of electron transfer, stability, and overall charge.

Ionic BondsIonic Bonds



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Peri

od

icTa

ble

Ro

w 3

Na

Mg

3p3s

3p3s

3p3s

3p3s

3p3s

3p3s

3p3s

3p3s

Cl

Ar

SP

Al

Si

Formulas for Ionic CompoundsFormulas for Ionic Compounds



24

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1. How many valence electrons are there in an atom of sodium (Na)? What would a sodiumatom tend to do in bonding with another atom to form an ionic compound? Why?

2. How many valence electrons are there in an atom of phosphorus (P)? What would a phos-phorus atom tend to do in bonding with another atom to form an ionic compound? Why?

3. What would be the formula and name of the ionic compound formed when sodium reactswith phosphorus? What are the oxidation numbers of each of the ions present?

4. How many valence electrons are in an atom of sulfur (S)? What would a sulfur atom tendto do in bonding with another atom to form an ionic compound? Why?

5. What would be the formula and name of the ionic compound formed when sodium reactswith sulfur? What are the oxidation numbers of each of the ions present?

6. How many valence electrons are in an atom of aluminum (Al)? What would an alu-minum atom tend to do in bonding with another atom to form an ionic compound? Why?

7. How many valence electrons are in an atom of chlorine (Cl)? What would a chlorineatom tend to do in bonding with another atom to form an ionic compound? Why?

8. What would be the formula and name of the ionic compound formed when aluminumreacts with chlorine? What are the oxidation numbers of each of the ions present?

9. What would be the formula and name of the ionic compound formed when aluminumreacts with sulfur? What are the oxidation numbers of each of the ions present?

Formulas for Ionic CompoundsFormulas for Ionic Compounds



24

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��

��

��

��

��

�

2�2�

2�2�

2�2�

2�

2�2�

2�

2�

�

��

��

�

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�

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�

�

�

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�

��

�

Gro

up

2 M

etal

Ato

ms

Gro

up

1 M

etal

Ato

ms

Metallic BondingMetallic Bonding



25

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1. What is a regular, repeating three-dimensional arrangement of atoms called?

2. Do the separate electrons that are shown belong exclusively to a single atom? What wordis used to describe such electrons?

3. Are the electrons shown the only ones actually present? Explain.

4. Why are the central atoms shown as positively charged?

5. How does the number of separate electrons shown for the group 1 metal atoms compareto the number of atoms? Explain why in terms of valence electrons.

6. How does the number of separate electrons shown for the group 2 metal atoms compareto the number of atoms?

7. What holds the metal atoms together in such an arrangement?

8. What term is used to describe this model of metallic bonding?

9. How well do metals tend to conduct electricity? How does the model of metallic bondingaccount for that property?

10. Do metals tend to be brittle, or are they malleable and ductile? How does the model ofmetallic bonding account for that property?

Metallic BondingMetallic Bonding



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Gro

up

s

Periods

12

2 3

3 Li4 Be

11 Na

12 Mg

Gro

up

s13

1415

1617

18

5 B6 C

7 N8 O

9 F10 N

e

17 Cl

18 Ar

16 S15 P

13 Al

14 Si

Calculating Numbers of Electronsand Predicting Ionic ChangeCalculating Numbers of Electronsand Predicting Ionic Change



7

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1. What happens to a neutral atom if it loses one electron? Why?

2. What happens to a neutral atom if it gains two electrons? Why?

3. Write a simple word equation that shows how you can figure out the charge of an ion,given its numbers of electrons and protons.

4. What is the electron configuration of each of the following, given its position in the periodic table? In each case, also tell what charge the atom is likely to take on if it bonds,and explain why. Write the electron configuration of the ion that is formed.

a. beryllium (Be)

b. fluorine (F)

c. argon (Ar)

d. sulfur (S)

e. sodium (Na)

f. nitrogen (N)

Calculating Numbers of Electronsand Predicting Ionic ChangeCalculating Numbers of Electronsand Predicting Ionic Change



7

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Determining Numbers of IonsDetermining Numbers of Ions



8

Common Polyatomic Ions

Ion Name Ion Name

NH4� ammonium IO4

� periodate

NO2� nitrite C2H3O2

� acetate

NO3� nitrate H2PO4

� dihydrogen phosphate

HSO4� hydrogen sulfate CO3

2� carbonate

OH� hydroxide SO32� sulfite

CN� cyanide SO42� sulfate

MnO4� permanganate S2O3

2� thiosulfate

HCO3� hydrogen carbonate O2

2� peroxide

ClO� hypochlorite CrO42� chromate

ClO2� chlorite Cr2O7

2� dichromate

ClO3� chlorate HPO4

2� hydrogen phosphate

ClO4� perchlorate PO4

3� phosphate

BrO3� bromate AsO4

3� arsenate

IO3� iodate

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1. Write a simple word equation that illustrates what must be true of total positive chargeand total negative charge in an ionic compound.

2. How many potassium ions (group 1) would be related to balance the charge of each ofthe following in a compound?

a. one cyanide ion

b. one sulfite ion

c. one arsenate ion

3. How many iodate ions would be needed to balance the charge of each of the following ina compound?

a. one Fe3� ion

b. one lithium ion (group 1)

c. one barium ion (group 2)

4. What is the formula of the ionic compound formed by each of the following in combination? Demonstrate that each result is correct by figuring out total positive charge and total negative charge.

a. ammonium ions and sulfate ions

b. sodium ions (group 1) and phosphate ions

c. magnesium ions (group 2) and hydrogen sulfate ions

d. aluminum ions (group 13) and carbonate ions

e. ammonium ions and arsenate ions

f. calcium ions (group 2) and acetate ions

g. ammonium ions and nitrite ions

Determining Numbers of IonsDetermining Numbers of Ions



8

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Ionic Compounds and MetalsIonic Compounds and Metals

Section 7.1 Ion FormationIn your textbook, read about chemical bonds and formation of ions.


The force that holds two atoms together is called a(n) (1) .

Such an attachment may form by the attraction of the positively charged

(2) of one atom for the negatively charged

(3) of another atom, or by the attraction of charged atoms,

which are called (4) . The attractions may also involve

(5) electrons, which are the electrons in the outermost

(6) . The (7) are a family of elements that

have very little tendency to react. Most of these elements have a set of eight outermost

electrons, which is called a stable (8) . The relatively stable electron

structures developed by loss of electrons in certain elements of groups 3, 4, 13, and 14 are

called (9) .


10. A positively charged ion is called an anion.

11. Elements in group 1 lose their one valence electron, forming an ion with a1� charge.

12. Elements tend to react so that they acquire the electron structure of ahalogen.

13. A sodium atom tends to lose one electron when it reacts.

14. The electron structure of a zinc ion (Zn2�) is an example of a pseudo-noble gas formation.

15. A Cl� ion is an example of a cation.

16. The ending -ide is used to designate an anion.

17. Nonmetals form a stable outer electron configuration by losing electronsand becoming anions.


chemical bond electrons energy level ions noble gases

nucleus octet pseudo-noble gas formations valence

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Section 7.2 Ionic Bonds and Ionic CompoundsIn your textbook, read about forming ionic bonds and the characteristics of ionic compounds.


1. An ionic bond is

a. attraction of an atom for its electrons.

b. attraction of atoms for electrons they share.

c. a force that holds together atoms that are oppositely charged.

d. the movement of electrons from one atom to another.

2. The formula unit of an ionic compound shows the

a. total number of each kind of ion in a sample.

b. simplest ratio of the ions.

c. numbers of atoms within each molecule.

d. number of nearest neighboring ions surrounding each kind of ion.

3. The overall charge of a formula unit for an ionic compound

a. is always zero. c. is always positive.

b. is always negative. d. may have any value.

4. How many chloride (Cl�) ions are present in a formula unit of magnesium chloride,given that the charge on a Mg ion is 2�?

a. one-half b. one c. two d. four

5. Ionic bonds generally occur between

a. metals. c. a metal and a nonmetal.

b. nonmetals. d. noble gases.

6. Salts are examples of

a. nonionic compounds. b. metals. c. nonmetals. d. ionic compounds.

7. A three-dimensional arrangement of particles in an ionic solid is called a(n)

a. crystal lattice. b. sea of electrons. c. formula unit. d. electrolyte.

8. In a crystal lattice of an ionic compound,

a. ions of a given charge are clustered together, far from ions of the opposite charge.

b. ions are surrounded by ions of the opposite charge.

c. a sea of electrons surrounds the ions.

d. neutral molecules are present.


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9. What is the relationship between lattice energy and the strength of the attractive forceholding ions in place?

a. The more positive the lattice energy is, the greater the force.

b. The more negative the lattice energy is, the greater the force.

c. The closer the lattice energy is to zero, the greater the force.

d. There is no relationship between the two quantities.

10. The formation of a stable ionic compound from ions

a. is always exothermic. c. is always endothermic.

b. may be either exothermic or endothermic. d. neither absorbs nor releases energy.

11. In electron transfer involving a metallic atom and a nonmetallic atom during ion forma-tion, which of the following is correct?

a. The metallic atom gains electrons from the nonmetallic atom.

b. The nonmetallic atom gains electrons from the metallic atom.

c. Both atoms gain electrons.

d. Neither atom gains electrons.

Underline the word that correctly describes each property in ionic compounds.

12. Melting point Low High

13. Boiling point Low High

14. Hardness Hard Soft

15. Brittleness Flexible Brittle

16. Electrical conductivity in the solid state Good Poor

17. Electrical conductivity in the liquid state Good Poor

18. Electrical conductivity when dissolved in water Good Poor


19. The crystal lattice of ionic compounds affects their melting and boilingpoints.

20. The lattice energy is the energy required to separate the ions of an ioniccompound.

21. The energy of an ionic compound is higher than that of the separateelements that formed it.

22. Large ions tend to produce a more negative value for lattice energy thansmaller ions do.

23. Ions that have larger charges tend to produce a more negative latticeenergy than ions with smaller charges do.



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Section 7.3 Names and Formulas for Ionic CompoundsIn your textbook, read about communicating what is in a compound and naming ionsand ionic compounds.


A one-atom ion is called a(n) (1) ion. The charge of such an ion is

equal to the atom’s (2) , which is the number of (3)

transferred to or from the atom to form the ion. In ionic compounds, the sum of the charges of

all the ions equals (4) . Ions made up of more than one atom are called

(5) ions. If such an ion is negatively charged and includes one or more

oxygen atoms, it is called a(n) (6) . If two such ions can be formed that

contain different numbers of oxygen atoms, the name for the ion with more oxygen atoms ends

with the suffix (7) . The name for the ion with fewer oxygen atoms

ends with (8) .

In the chemical formula for any ionic compound, the chemical symbol for the

(9) is written first, followed by the chemical symbol for the

(10) . A(n) (11) is a small number used to

represent the number of ions of a given element in a chemical formula. Such numbers are

written to the (12) of the symbol for the element. If no number

appears, the assumption is that the number equals (13) .

For each formula in Column A, write the letter of the matching name in Column B.

Column A Column B

14. ClO2�

15. ClO4�

16. ClO�

17. Cl�

18. ClO3�


anion -ate cation electrons zero

lower right monatomic one oxidation number -ite

oxyanion polyatomic subscript

a. chlorate

b. hypochlorite

c. chloride

d. perchlorate

e. chlorite

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For each of the following chemical formulas, write the correct name of the ionic compound represented. You may refer to the periodic table on pages 156–157 and Table 8.7 for help.

19. NaI

20. CaCl2

21. K2S

22. MgO

23. LiHSO4

24. NH4Br

25. Ca3N2

26. Cs3P

27. KBrO3

28. Mg(ClO)2

29. Li2O2

30. Be3(PO4)2

31. (NH4)2CO3

32. NaBrO3

33. Fe2O3

34. Fe(IO3)2

For each of the following ionic compounds, write the correct formula for the compound.You may refer to the periodic table on pages 156–157 and Table 8.7 for help.

35. beryllium nitride

36. nickel(II) chloride

37. potassium chlorite

38. copper(I) oxide

39. magnesium sulfite

40. ammonium sulfide

41. calcium iodate

42. iron(III) perchlorate

43. sodium nitride



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Section 7.4 Metallic Bonds and the Properties of MetalsIn your textbook, read about metallic bonds.

Use the diagram of metallic bonding to answer the following questions.

1. What is the name of the model of metallic bonding that is illustrated?

2. Why are the electrons in a metallic solid described as delocalized?

3. Which electrons from the metal make up the delocalized electrons?

4. Are the metal atoms that are shown cations or anions? How can you tell?

5. How do the metallic ions differ from the ions that exist in ionic solids?

6. Explain what holds the metal atoms together in the solid.

In your textbook, read about the properties of metals.

For each property, write yes if the property is characteristic of most metals, or no if it isnot. If the property is a characteristic of metals, explain how metallic bonding accountsfor the property.

7. Malleable

8. Brittle

9. Lustrous

10. High melting point

11. Low boiling point

12. Ductile

13. Poor conduction of heat

14. Good conduction of electricity


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Ionic Compounds and MetalsIonic Compounds and Metals


Column A Column B

1. Electrons that are free to move in metals

2. For a monatomic ion, is equal to the charge

3. The force that holds two atoms together

4. A charged particle containing more than one atom

5. A positively charged ion

6. A negatively charged ion

7. An ionic compound whose aqueous solution conductselectricity

8. The name for most ionic compounds other than oxides

9. Represents the way electrons exist in metals

10. A charged particle containing only one atom

11. The energy needed to separate the ions of an ioniccompound

12. The electrostatic force that holds oppositely chargedparticles together

13. A mixture of elements that has metallic properties

14. A mixture formed when small atoms fill holes in ametallic crystal

15. A polyatomic ion composed of an element bonded to atleast one oxygen atom

16. Shows the simplest ratio of ions in an ionic compound

17. The attraction of a metallic cation for delocalizedelectrons


a. alloy

b. anion

c. cation

d. chemical bond

e. delocalized electrons

f. electrolyte

g. electron sea model

h. formula unit

i. interstitial alloy

j. ionic bond

k. lattice energy

l. metallic bond

m. monatomic ion

n. oxidation number

o. oxyanion

p. polyatomic ion

q. salts

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In the space at the left, write true if the statement is true; if the statement is false,change the italicized term to make it true.

1. The overall charge of a formula unit for a compound is never zero.

2. In a crystal lattice, each positive ion is surrounded by negative ions.

3. Delocalized valence electrons are typical of ionic compounds.

4. A sulfate ion contains fewer oxygen atoms than a sulfite ion does.

5. Metals tend to be malleable and ductile and to have relatively high melting points.

6. The more negative the lattice energy is, the stronger is the force of attraction between the ions of an ionic compound.

7. In naming ionic compounds, the cation is named first.

8. When a metal reacts with a nonmetal, the metal tends to gain electrons.

9. In naming a monatomic anion, the suffix -ide is used.

10. The prefix per- is used in naming the anion with the most oxygen atoms.

Circle the letter of the word or phrase that best completes the statement or answers thequestion.

11. What is the electron configuration for the noble gases other than helium?

a. ns2np6 b. ns2 c. ns2np3 d. ns2np2

12. How many outermost d electrons are there in an ion that has achieved a pseudo-noble gasconfiguration?

a. none b. five c. eight d. ten

13. The anion that has the formula ClO� is called the

a. chloride ion. b. chlorate ion. c. hypochlorite ion. d. perchlorate ion.

14. Where does a subscript that indicates the number of atoms appear, relative to a chemicalsymbol in a formula?

a. to the upper left b. to the lower left c. to the upper right d. to the lower right

15. What is the formula of calcium phosphate, which is made up of the ions Ca2� and PO43�?

a. Ca3PO4 b. Ca6PO4 c. Ca3(PO4)2 d. Ca2(PO4)3

16. Which of the following is an example of an interstitial alloy?

a. brass b. 14-carat gold c. carbon steel d. sterling silver


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The diagram on the right represents a lithium atom (group 1 of the periodic table) and a fluorine atom (group 17). Use the diagram to answer the questions that follow.

1. What is the electron configuration of the neutral lithium atom?

2. What is the electron configuration of the neutral fluorine atom?

3. What happens to the lithium atom when it reacts with the fluorine atom? What is theelectron configuration of the lithium after the change?

4. In terms of the electron arrangement, why is this change favorable for the lithium atom?

5. What happens to the fluorine atom when it reacts with the lithium atom? What is theelectron configuration of the fluorine after the change?

6. In terms of the electron arrangement, why is this change favorable for the fluorine atom?

7. What kind of compound is formed in the reaction?

8. What always happens to one or more electrons during a reaction that forms such a compound?

9. What are the formula and name of the product in this reaction?

10. What holds the atoms together in the compound?

11. What is the name for the overall three-dimensional solid structure that samples of suchcompounds form? How are the particles generally arranged in such a structure?

12. How do the physical properties of such compounds differ from those typical of metals?


3 electrons 9 electrons

3 protons 9 protonsLi F

�

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Thinking CriticallyAnswer the following questions, which refer to thegraph on the right. The graph shows the lattice ener-gies of the halide compounds of the group 1 metalssodium (atomic number 11) and potassium (atomicnumber 19).

1. How does lattice energy relate to the attractionbetween ions and the stability of an ionic compound?

2. What is the relationship between the lattice energy of the halides of sodium and the atomicnumber and size of the halides? What is the relationship for the halides of potassium?

3. How does the plot of lattice energy for the potassium halides compare with that for thesodium halides? What does this suggest about lattice energy and ionic size?

4. Given what you know about lattice energy and the stability of a crystal, how would youexpect the melting point of NaBr to compare with that of NaI? How would you expectthe melting point of NaBr to compare with that of KBr? Explain your answer.

5. What effect on lattice energy would you expect the amount of charge on an ion to have?Explain.


�1000

�900

�800

�700

�600F Cl Br

Halides

Latt

ice

ener

gy

(kJ/

mo

l)

IIncreasing atomic number and size

SodiumPotassium

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Applying Scientific MethodsFour students (A, B, C, and D) are given separate solid samples of the same unknown puresubstance. They are asked to determine whether the substance is a metallic element, a non-metallic element, or an ionic compound. In order to find out, they independently carry outexperiments to determine some of the properties of the substance. Student A observes the sub-stance to determine its luster. Student B tests the solid’s ability to conduct electricity. StudentC determines whether the solid is malleable and ductile. Student D determines its meltingpoint and tests the melted liquid’s ability to conduct electricity. The students do not communi-cate their separate findings to one another. The results of their experiments are shown in thetable below. Use the information to answer the questions that follow.

1. On the basis of his results, Student A concludes that the solid is a nonmetallic elementrather than a metallic element or an ionic compound. Comment on the soundness of hisconclusion, given only what he has determined about luster.

2. On the basis of her results, Student B also concludes that the unknown is a nonmetallicelement. Evaluate her conclusion.

3. On the basis of his results, Student C concludes that the unknown is a metallic element.Evaluate his conclusion.


Student Property Studied Result

A Luster Nonlustrous

B Ability of solid to conduct electricity Nonconducting

C Malleability and ductility Nonmalleable and nonductile (brittle)

D Melting point Approximately 800°C

D Ability of melted liquid to conduct electricity Good

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4. On the basis of her results, Student D states that she cannot conclude anything aboutwhether the unknown is a metallic element, nonmetallic element, or ionic solid. Evaluateher comment.

5. Suppose that Student B and Student C shared their results with each another. Should theythen be able to come to a definite conclusion as to the nature of the unknown? Explain.

6. Suppose that Student B and Student D shared their results with each other. Should theythen be able to come to a definite conclusion as to the nature of the unknown? Explain.

7. Would a test of a water solution of the unknown might also have been useful in deter-mining the nature of the unknown? Explain.

8. What do the individual students’ problems in coming to definite conclusions illustrateabout scientific methods?



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1. 4. 7.

2. 5. 8.

3. 6. 9.

Short Answer


10.

11.

12.

Extended Response


13.

14.


15. 17. 19.

16. 18. 20.

CHAPTER 7

Assessment

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Chapter 8 Covalent BondingMiniLab . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 86

ChemLab . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 87

Teaching Transparency Masters and Worksheets . . . . . . . . . . . . . . . . . . . . . . . . . 90

Math Skills Transparency Masters and Worksheets . . . . . . . . . . . . . . . . . . . . . . . . . 96

Study Guide . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 100

Chapter Assessment. . . . . . . . . . . . . . . . . . . . . . . . . . . . 106

STP Recording Sheet . . . . . . . . . . . . . . . . . . . . . . . . . . . 112

Table ofContents

85

Reproducible Pages

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mini LAB 8Compare Melting Points

How can you determine the relationship between bond type and meltingpoint? The properties of a compound depend on whether the bonds in the compoundare ionic or covalent.

Materials permanent marker, disposable 9-inch aluminum pie pan, hot plate, sugar crystals, salt crystals, paraffin


2. Create a data table for the experiment.

3. Using a permanent marker, draw three lines on the inside bottom of a disposable, 9-inch aluminum pie pan to create three, equal wedges. Label the wedges, A, B, and C.

4. Set the pie pan on a hot plate.

WARNING: Hot plate and metal pie pan will burn skin—handle with care.

5. Obtain samples of the following from your teacher and deposit them onto thelabeled wedges as follows: sugar crystals (C12H22O11), A; salt crystals (NaCl) B; paraffin (C23H48), C.

6. Predict the order in which the compounds will melt.

7. Turn the temperature knob on the hot plate to the highest setting. You will heat thecompounds for 5 min. Assign someone to time the heating of the compounds.

8. Observe the compounds during the 5-min period. Record which compounds melt andthe order in which they melt.

9. After 5 min, turn off the hot plate and remove the pie pan using a hot mitt or tongs.

10. Allow the pie pan to cool, and then place it in the proper waste container.

Analysis

1. State Which solid melted first? Which solid did not melt?

2. Apply Based on your observations and data, describe the melting point of each solidas low, medium, high, or very high.

3. Infer Which compounds are bonded with ionic bonds? Which are bonded with cova-lent bonds?

4. Summarize how the type of bonding affects the melting points of compounds.

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CHEMLAB 8

ProblemHow do the Lewis structureand the positions of valenceelectrons affect the shape ofthe covalent compound?

Objectives• Predict the shape of

molecular compounds• Construct molecular

models• Evaluate the strength of

covalent bonds• Identify resonance

structures• Distinguish polar and

nonpolar molecules

Materialsmolecular-model kit

Model Molecular ShapesCovalent bonding occurs when atoms share valence elec-

trons. In the Valence Shell Electron Pair Repulsion (VSEPR)theory, the way in which valence electrons of bonding atomsare positioned, is the basis for predicting a molecule’s shape.This method of visualizing shape is also based on the mole-cule’s Lewis structure.

Pre-Lab


2. Review the VSEPR model. What do the initialsVSEPR stand for? What is the basis for the pre-dictions of molecular shape made by the VSEPRmodel?

3. How do you determine whether or not a bond is apolar covalent bond? What additional information

do you need to determine whether or not a mole-cule is polar?

4. What is a resonance structure? If a molecule hasresonance structures, what can you infer about thetypes of bonds in that molecule?

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Procedure


2. Create a table to record your data.

3. Note the color used to represent each of theseatoms in the molecular modele kit: H, O, P, C, F,S, and N.

4. Draw the Lewis structures of the H2, O2, and N2molecules.

5. Obtain two hydrogen atoms and one connectorand assemble a hydrogen (H2) molecule.

6. Obtain two oxygen atoms and two connectors andassemble an oxygen (O2) molecule. Observe thatyour model represents a double-bonded diatomicoxygen molecule.

7. Obtain two nitrogen atoms and three connectorsand assemble a nitrogen (N2) molecule. Observethat your model represents a triple-bondeddiatomic nitrogen molecule.

8. Recognize that these diatomic molecules are lin-ear in shape because they have only two atoms.

9. Draw the Lewis structure of water (H2O) andconstruct its molecule.

10.Classify the shape of the H2O molecule usinginformation in Table 8.6.

11.Repeat steps 9 and 10 for PH3, CF4, CO2, SO3,HCN, and CO molecules.

CHEMLAB 8


1. Think Critically Based on the molecular models you built and observed in this lab, ranksingle, double, and triple bonds in order of increasing flexibility and increasing strength.

2. Observe and Infer Explain why H2O and CO2 molecules have different shapes.

3. Analyze and Conclude One of the molecules from this lab undergoes resonance. Identify the moleculethat has three resonance structures, draw the structures, and explain why resonance occurs.

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4. Recognize Cause and Effect Use the electronegativity difference to determine thepolarity of the molecules in Steps 9-11. Based on their calculated bond polarities and themodels constructed in this lab, determine the molecular polarity of each structure.

Inquiry Extension

Model Use a molecular model kit to build the two resonance structures of ozone (O3).Then, use Lewis structures to explain how you can convert between the two resonancestructures by interchanging a lone pair for a covalent bond.

CHEMLAB 8

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Form

ula

N2H

4

SiO

2O

OSi

Step 1

Step 2

Step 3

Step 4

Step 5

Step 6

Un

nec

essa

ry(o

ctet

sco

mp

lete

)H

N

N

H

H

H

2

H

N

N

H

H

H

147

H N

N H

H H

OO

Si

Lewis StructuresLewis Structures



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1. Step 1 in drawing the Lewis structure for a molecule is to decide which atoms of themolecule are most likely the terminal ones. In the transparency, why are the hydrogen(H) atoms in hydrazine (N2H4) shown as the terminal atoms?

2. Step 2 in drawing a Lewis structure involves determining the total number of valenceelectrons in the atoms in the molecule. Explain why the total number of valence electronsin N2H4 is 14.

3. Step 3 in drawing a Lewis structure requires finding the number of bonding pairs. Whatmust be done to the result of step 2 to find the number of bonding pairs? Verify that thisis so in the case of N2H4 in the transparency.

4. In step 4 in the transparency, one bonding pair has been placed between each pair ofbonded atoms in N2H4. How many such bonding pairs are shown in step 4, and whatsymbol is used to represent them?

5. Step 5 requires subtraction of the number of bonding pairs used in step 4 from the number of bonding pairs determined in step 3. Verify that the result is 2 for N2H4. Lonepairs are then placed around each terminal atom to achieve a full outer level, and anyremaining pairs are assigned to the central atom(s). Explain the drawing that has resultedfor N2H4.

6. In step 6, if any central atom drawn in step 5 does not have an octet, lone pairs from theterminal atoms must be converted to double or triple bonds involving the central atom.Why was this extra step unnecessary in the case of N2H4?

7. What number should be placed in the blank for step 2 for the silicon dioxide (SiO2) molecule?

8. What number should be placed in the blank for step 3 for SiO2?

Lewis StructuresLewis Structures



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H

HH

B H

HH

B

H

HH

HC H H

HH

C

HH

HN H

HH

N

HHO H

HO

HF

HF

VSEPR Model and MolecularShapeVSEPR Model and MolecularShape



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1. The shapes of the molecules shown have been determined by means of the VSEPRmodel. What is the basic assumption of this model?

2. How many lone pairs and how many shared pairs of electrons surround the boron (B)atom in the borane (BH3) molecule shown?

3. What is the shape of the BH3 molecule? Explain why.

4. How many lone pairs and how many shared pairs of electrons surround the carbon (C)atom in the methane (CH4) molecule shown? What is the shape of the molecule?

5. How many lone pairs and how many shared pairs of electrons surround the nitrogen (N)atom in the ammonia (NH3) molecule shown? What is the shape of the molecule?

6. How many lone pairs and how many shared pairs of electrons surround the oxygen (O)atom in the water (H2O) molecule shown? What is the shape of the molecule?

7. How many lone pairs and how many shared pairs of electrons surround the fluorine (F)atom in the hydrogen fluoride (HF) molecule shown? What is the shape of the molecule?

VSEPR Model and MolecularShapeVSEPR Model and MolecularShape



27

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Peri

od

2

Gro

up

1G

rou

p2

Gro

up

3G

rou

p4

Gro

up

5G

rou

p6

Gro

up

7

Li 0.98

Be

1.57

B2.

04C

2.55

N3.

04O 3.44

F3.

98 Cl

3.16 Br

2.96 I

2.66 At

2.2

Elec

tron

egat

ivit

ies

Electronegativity and PolarityElectronegativity and Polarity



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1. What is electronegativity?

2. The electronegativities of the elements of period 2 and of group 17 of the periodic table areshown. What trends in electronegativity do you see across the period? Down the group?

3. When there is an electronegativity difference between two covalently bonded atoms,what is true of the bond between them? Toward which of the atoms are the shared electrons more attracted?

4. What kind of bond exists between a carbon (C) atom and a chlorine (Cl) atom? (Assumethat a bond is nonpolar covalent if the electronegativity difference is 0, polar covalent ifthe difference is greater than 0 but not more than 1.70, and ionic if the difference is morethan 1.70.)

5. Given your answer to question 4 and your knowledge of molecular shapes, is a carbontetrachloride (CCl4) molecule polar or nonpolar? Explain.

6. What kind of bond exists between a nitrogen (N) atom and a fluorine (F) atom? Is anitrogen trifluoride (NF3) molecule polar or nonpolar? Explain.

7. What kind of bond exists between a beryllium (Be) atom and a bromine (Br) atom? Is aberyllium bromide (BeBr2) molecule polar or nonpolar? Explain.

8. What kind of bond exists between a beryllium (Be) atom and a fluorine (F) atom?

9. What kind of bond exists between a boron (B) atom and an iodine (I) atom? Is a borontriiodide (BI3) molecule polar or nonpolar? Explain.

Electronegativity and PolarityElectronegativity and Polarity



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O O

O O

I

Num

ber

of A

tom

s

1 2 3 4 5

(a)

(c)

(e)

(b)

(d)

Pref

ix

mo

no

-d

i-tr

i-te

tra-

pen

ta-

(nev

er f

or

firs

tel

emen

t in

th

efo

rmu

la)

Num

ber

of A

tom

s

6 7 8 9 10

Pref

ix

hex

a-h

epta

-o

cta-

no

na-

dec

a-

I

ON

NN

CS

S

IF

F

FSiF

FF

FF

Determining the Names ofBinary Compounds and TheirNumbers of Atoms




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1. The table shows the prefixes used in naming binary compounds. Notice that each prefixcorresponds to a certain number of atoms. These prefixes are also used in naming manygeometric figures with which you probably are familiar. For example, a triangle has threesides. The prefix tri- signifies three. How many sides does a pentagon have? A hexagon?An octagon?

2. Look at the drawings of molecules (labeled a–e) below the table. For each molecule, tellhow many atoms of each kind are present. Then write the name of the compound, usingthe prefixes in the table. In each case, the central atom is named first.

a.

b.

c.

d.

e.

3. For each compound listed below, tell how many of each kind of atom are present in amolecule of the compound and the total number of atoms per molecule. Then name thecompound.

a. P4O10

b. S2O7

c. Si3H8

4. What is the formula for each compound listed below?

a. tetrasulfur dinitride

b. dichlorine monoxide





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100

Peri

od

3 E

lem

ent

Elec

tro

neg

ativ

ity

Na

0.93

Mg

1.31

Al

1.61

Si 1.90

P

2.19

S

2.58

Cl 3.16

75 50 25 00

1.0

2.0

3.0

Elec

tron

egat

ivit

y di

ffer

ence

Percent ionic character

Ioni

c

Cova

lent

Pola

r co

vale

nt

Determining ElectronegativityDifference and Percent IonicCharacter




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1. The electronegativities of the period-3 elements are listed on the transparency. Calculatethe electronegativity differences for the following pairs of bonded period-3 atoms.

a. Na and Cl d. Si and Cl

b. Mg and S e. Si and S

c. Al and P

2. An electronegativity difference greater than 1.70 generally corresponds to a bond that isconsidered ionic. For which of the bonded-atoms combinations from question 1 wouldthe bonds be considered ionic according to this rule?

3. Bonds are rarely completely covalent or completely ionic. Rather, they have a percentionic character. The graph illustrates that fact. What two variables are plotted on thegraph, and what is the overall relationship between the variables?

4. Use the graph to find the approximate percent ionic character for the five pairs of bondedatoms listed in question 1.



c. Al and P

5. Given what you know about percentages in general, what must be the relationshipbetween the percent ionic character and the percent covalent character for a given bond?

6. Calculate the percent covalent character for the five pairs of bonded atoms listed in question 1.



c. Al and P





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Covalent BondingCovalent Bonding

Section 8.1 The Covalent BondIn your textbook, read about the nature of covalent bonds.


When sharing of electrons occurs, the attachment between atoms that results is called

a(n) (1) . When such an attachment is formed, bond dissociation

energy is released, and the process is (2) . When two or more

atoms bond by means of electron sharing, the resulting particle is called a(n)

(3) . If the electrons shared are centered between the two atoms, the

attachment is called a(n) (4) . If the sharing involves the overlap of

parallel orbitals, the attachment is called a(n) (5) .

In your textbook, read about single and multiple bonds and bond strength.


6. In what form do elements such as hydrogen, nitrogen, and oxygen normally occur?

a. as single atoms c. as molecules containing three atoms

b. as molecules containing two atoms d. as molecules containing four atoms

7. How many electrons are shared in a double covalent bond?

a. none b. one c. two d. four

8. Bond length is the distance between

a. two molecules of the same substance. c. the nuclei of two attached atoms.

b. the electrons in two attached atoms. d. the orbitals of two attached atoms.

9. Which of the following relationships relating to bond length is generally correct?

a. the shorter the bond, the stronger the bond

b. the shorter the bond, the weaker the bond

c. the shorter the bond, the fewer the electrons in it

d. the shorter the bond, the lower the bond dissociation energy


covalent bond molecule sigma bond exothermic pi bond

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Section 8.2 Naming MoleculesIn your textbook, read about how binary compounds and acids are named from theirformulas.


1. Binary molecular compounds are generally composed of a metal and anonmetal.

2. The second element in the formula of a binary compound is named usingthe suffix -ite.

3. The prefix tetra- indicates three atoms.

4. The prefix hexa- indicates six atoms.

5. In naming the first element in a formula, the prefix mono- is not used.

6. For binary acids, the hydrogen part of the compound is named using theprefix hydro-.

7. An oxyacid contains only two elements.

8. If the name of the anion of an oxyacid ends in -ate, the acid namecontains the suffix -ous.

In your textbook, read about naming molecular compounds and oxyacids.


Column A Column B

9. CO

10. CO2

11. H2CO3

12. NH3

13. N2O4

14. HNO2

15. HNO3

16. HBr

17. HBrO3


a. hydrobromic acid

b. dinitrogen tetroxide

c. carbon monoxide

d. nitrous acid

e. ammonia

f. nitric acid

g. carbonic acid

h. bromic acid

i. carbon dioxide

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Section 8.3 Molecular StructuresIn your textbook, read about Lewis structures.


1. A structural formula shows the arrangement of the atoms in a molecule.

2. The central atom in a molecule is the one with the highest electronaffinity.

3. In molecules, hydrogen is always a terminal atom.

4. The number of bonding pairs in a molecule is equal to the number ofelectrons.

5. To find the total number of electrons available for bonding in a positiveion, you should add the ion charge to the total number of valenceelectrons of the atoms present.

6. The electrons in a coordinate covalent bond are donated by both thebonded atoms.

7. Resonance occurs when more than one valid Lewis structure can bewritten for a molecule.

8. Nitrate is an example of an ion that forms resonance structures.

9. The carbon dioxide molecule contains two double bonds.

10. All electrons in an atom are available for bonding.

11. In the sulfate ion (SO42�), 32 electrons are available for bonding.

12. When carbon and oxygen bond, the molecule contains ten pairs ofbonding electrons.

In your textbook, read about resonance structures and exceptions to the octet rule.


Column A Column B

13. Odd number of valence electrons

14. Fewer than 8 electrons around an atom

15. More than 8 electrons around central atom

16. More than one valid Lewis structure


a. O3

b. BF3

c. NO

d. SF6

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Section 8.4 Molecular ShapeIn your textbook, read about the VSEPR model.

Circle the letter of the choice that best completes the statement.

1. The VSEPR model is used mainly to

a. determine molecular shape. c. determine ionic charge.

b. write resonance structures. d. measure intermolecular distances.

2. The bond angle is the angle between

a. the sigma and pi bonds in a double bond. c. two terminal atoms and the central atom.

b. the nucleus and the bonding electrons. d. the orbitals of a bonding atom.

3. The VSEPR model is based on the idea that

a. there is always an octet of electrons around an atom in a molecule.

b. electrons are attracted to the nucleus.

c. molecules repel one another.

d. shared and unshared electron pairs repel each other as much as possible.

4. The shape of a molecule whose central atom has four pairs of bonding electrons is

a. tetrahedral. b. trigonal planar. c. trigonal pyramidal. d. linear.

5. The shape of a molecule that has two covalent single bonds and no lone pairs on the central atom is


6. The shape of a molecule that has three single covalent bonds and one lone pair on thecentral atom is


In your textbook, read about hybridization.


The formation of new orbitals from a combination or rearrangement of valence electrons

is called (7) . The orbitals that are produced in this way are

(8) to one another. An example of an element that commonly

undergoes such formation is (9) . When this atom combines its three

p orbitals and its one s orbital, the orbitals that result are called (10)

orbitals. An example of a molecule that has this type of orbital is (11) .


carbon hybridization sp3 identical methane

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Section 8.5 Electronegativity and PolarityIn your textbook, read about electronegativity.

Use the table of electronegativities below to answer the following questions.

1. What is the meaning of the term electronegativity?

2. Which element has the highest electronegativity? What is the numerical value? What are thename and group number of the chemical family that has the highest overall electronegativities?

3. Which element has the lowest electronegativity? What is the numerical value? What are thename and group number of the chemical family that has the lowest overall electronegativities?

4. What general trend in electronegativity do you note going down a group? Across a period?

5. How are the electronegativity values used to determine the type of bond that existsbetween two atoms?

In your textbook, read about the properties of covalent compounds.


6. Ionic compounds are usually soluble in polar substances.

7. In a covalent molecular compound, the attraction between molecules tendsto be strong.

Electronegativities of Some Elements

MetalMetalloidNonmetal

78Pt2.2

79Au2.4

80Hg1.9

81Tl

1.8

82Pb1.8

83Bi1.9

85At2.2

28Ni

1.91

29Cu

1.90

30Zn

1.65

31Ga

1.81

32Ge

2.01

33As

2.18

34Se

2.55

35Br

2.96

46Pd

2.20

47Ag

1.93

48Cd

1.69

49In

1.78

50Sn

1.96

51Sb

2.05

52Te2.1

53I

2.66

13Al

1.61

14Si

1.90

15P

2.19

16S

2.58

17Cl

3.16

5B

2.04

6C

2.55

7N

3.04

8O

3.44

9F

3.98

84Po2.0

1H

2.203Li

0.98

11Na

0.9319K

0.82

37Rb

0.82

55Cs

0.7987Fr0.7

88Ra0.9

89Ac1.1

56Ba

0.89

57La

1.10

72Hf1.3

73Ta1.5

74W1.7

75Re1.9

76Os2.2

77Ir

2.2

38Sr

0.95

39Y

1.22

40Zr

1.33

41Nb1.6

42Mo2.16

20Ca

1.00

21Sc

1.36

22Ti

1.54

23V

1.63

24Cr

1.66

43Tc

2.10

44Ru2.2

25Mn1.55

26Fe

1.83

27Co

1.88

45Rh

2.28

12Mg1.31

4Be

1.57


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In your textbook, read about bond polarity.

Using the table of electronegativities on the preceding page, circle the letter of the choicethat best completes the statement or answers the question.

8. Unequal sharing of electrons between two bonded atoms always indicates

a. a nonpolar covalent bond. c. a polar covalent bond.

b. an ionic bond. d. a polar molecule.

9. When electronegativities of two bonded atoms differ greatly, the bond is

a. polar covalent. b. coordinate covalent. c. polar covalent. d. ionic.

10. What is the electronegativity difference that usually is the dividing line between covalentand ionic bonds?

a. 1.0 b. 1.7 c. 2.7 d. 4.0

11. The symbol �� is placed next to which of the following?

a. the less electronegative atom in a polar covalent bond c. a positive ion

b. the more electronegative atom in a polar covalent bond d. the nucleus

12. A nonpolar covalent bond is one in which

a. electrons are transferred. c. electrons are shared equally.

b. electrons are shared unequally. d. both electrons are provided by the same atom.

13. Molecules containing only polar covalent bonds

a. are always polar. c. are always ionic.

b. may or may not be polar. d. are always nonpolar.

14. What factor other than electronegativity determines whether a molecule as a whole ispolar or not?

a. temperature b. its geometry c. its physical state d. its mass

15. Which of the following correctly describes the compound water, H2O?

a. ionic c. polar overall, with nonpolar covalent bonds

b. nonpolar overall, with polar covalent bonds d. polar overall, with polar covalent bonds

16. Which of the following correctly describes the compound carbon tetrachloride, CCl4?

a. ionic c. polar overall, with nonpolar covalent bonds

b. nonpolar overall, with polar covalent bonds d. polar overall, with polar covalent bonds

17. A molecule of ammonia, NH3, is

a. nonpolar because it is linear.

b. polar because it is linear.

c. nonpolar because there is no electronegativity difference.

d. polar because there is an electronegativity difference and the molecule is trigonal pyramidal.



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Covalent BondingCovalent Bonding


Column A Column B

1. The tendency of an atom in a compound to attractelectrons

2. A kind of bond in which there is unequal sharing ofelectrons

3. Any bond in which there is electron sharing

4. The particle formed when two or more atoms bondcovalently

5. Reactions that occur when more energy is releasedforming new bonds than is required to break bonds in theinitial reactants

6. A kind of bond in which electrons are shared in an areacentered between the two atoms

7. A kind of bond formed by overlap of parallel orbitals

8. Any acidic compound that contains oxygen

9. A model that shows how the atoms are arranged in amolecule

10. Reactions that occur when more energy is required tobreak existing bonds in reactants than is released whennew bonds form in the product molecules

11. A condition that occurs when more than one valid Lewisstructure can be drawn for a molecule

12. A kind of bond in which one of the atoms provides bothelectrons for sharing

13. A model used to determine molecular shape

14. The combining of orbitals in an atom to form new,identical orbitals


a. structural formula

b. molecule

c. VSEPR model

d. coordinate covalent bond

e. hybridization

f. oxyacid

g. electronegativity

h. sigma bond

i. polar covalent

j. pi bond

k. covalent bond

l. resonance

m. endothermic

n. exothermic

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1. In the formation of a covalent bond, electrons are

a. shared. b. lost. c. gained. d. transferred.

2. Which of the following elements normally exists in the form of diatomic molecules?

a. helium b. argon c. iron d. nitrogen

3. Four electrons are shared in a

a. single covalent bond. c. triple covalent bond.

b. double covalent bond. d. quadruple covalent bond.

4. Which of the following molecules contains only sigma bonds?

a. methane b. oxygen c. carbon dioxide d. nitrogen

5. Which of the following molecules contains a triple bond?

a. methane b. oxygen c. carbon dioxide d. nitrogen

6. How many pi bonds are there in a triple bond?

a. none b. one c. two d. three

7. Which of the following molecules would be expected to have the greatest bond dissociation energy?

a. F2 b. O2 c. N2 d. Cl2

In the space at the left, write true if the statement is true; if the statement is false,change the italicized term to make it true.

8. In a chemical name, the prefix used to indicate the presence oftwo atoms of a given kind is bi-.

9. The prefix hydro- is used in naming binary acids.

10. The oxyacid suffix for an acid that contains an anion ending in -ate is -ic.

11. In Lewis structures, hydrogen is always a terminal atom.

12. In the carbon dioxide molecule, the central atom is a carbon atom.

13. In the compound BH3, the boron atom has more than an octet ofelectrons.

14. The VSEPR model is based on the idea that in a molecule, nucleirepel each other as much as possible.


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The diagram below represents different hybridizations of the orbitals of three carbonatoms, labeled X, Y, and Z. Use the diagram to answer the questions that follow.

1. Write the electron configuration of an unbonded carbon atom (atomic number 6) beforehybridization occurs.

2. Look at carbon atom X in the diagram. What is the symbol for the hybrid orbitals formedby carbon atom X? How many of them are there in that atom?

3. Does carbon atom X have any unhybridized orbitals? If so, tell how many and write thesymbol for them.

4. To how many other atoms would carbon atom X be attached? What types of bonds—single, double, or triple—would they be? State whether each attachment would involvesigma bonds, pi bonds, or both.

5. What would be the shape of the molecule formed by carbon atom X when it bonds inthat way?

6. Suppose that the attachments in that molecule are to atoms of the same kind with an electronegativity greater than that of carbon. Would each bond be polar or nonpolar?Would the molecule as a whole be polar or nonpolar? Explain.

7. Look at carbon atom Y in the diagram. What is the symbol for the hybrid orbitals formedby carbon atom Y? How many of them are there in that atom?

sp3

hybridsp3

hybrid

sp3

hybridsp2

hybrid

sp2

hybrid

sp2

hybridsphybrid

sphybrid

p orbital(above andbelow)

p orbital(above andbelow)

p orbital(above and

below)

sp3 hybrid

CC C

X Y Z


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Thinking CriticallyThe graph shows the electronegativities of the elements of periods 2 and 3 of the periodictable, except for the noble gases. Use the graph to answer the following questions.

1. If two atoms of differing electronegativity are bonded covalently to each other, what istrue of the electrons they share? What type of bond results in such a case?

2. How does the electronegativity of boron (B) compare with that of nitrogen (N)? In termsof the periodic table, how are these two elements related? Compare aluminum (Al) withphosphorus (P) in the same way.

3. Describe the trend in electronegativity illustrated for period-2 and period-3 elements.

4. How does the electronegativity of boron (B) compare with that of aluminum (Al)? Interms of the periodic table, how are these two elements related? Compare nitrogen (N)with phosphorus (P) in the same way.

5. Describe the trend in electronegativity within a group of the periodic table, as suggestedby the graph.

0.0

0.5

1.0

1.5

2.0

2.5

3.0

3.5

4.0

1 2 13 14 15 16 17Group number

Elec

tro

neg

ativ

ity

Li

Be

B

C

N

O

F

Na

MgAl

SiP

S

ClPeriod 2

Period 3


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Applying Scientific MethodsA college chemistry student is studying the properties of four unknown compounds, W, X, Y,and Z. She has been informed that one of them is ionic and that the other three are covalent.Of the latter, she has been told that the attractions between the formula units are dispersionforces in one case, hydrogen bonds in another, and covalent (network) bonds in another. Shehas been asked to determine the type of attractions for each substance. In an attempt to do so,she carries out experiments that reveal information about the substances’ properties. Use thisinformation to answer the questions that follow.

1. On the basis of the results of her hardness test, the student concludes that unknowns Wand X are covalent substances, and that either Y or Z is the ionic substance. Comment onthe soundness of her conclusion, given only the hardness information.

2. The student examines the melting-point data, but claims that it is not useful in determiningthe substances’ types of attractions. How would you reply to that claim?

3. On the basis of the results of the boiling-point and melting-point tests, the student concludes that W is the hydrogen-bonded substance and that X must therefore be thesubstance that has the dispersion forces. Is that conclusion valid? Explain.


Property Unknown W Unknown X Unknown Y Unknown Z

Hardness of solid soft soft brittle brittle

Melting point (°C) �10 �200 1500 3000

Boiling point (°C) 120 �150 2500 4500

Solubility in polar solvent very soluble insoluble very soluble insoluble

Solubility in nonpolar solvent insoluble very soluble insoluble insoluble

Conductivity of solid nonconducting nonconducting nonconducting nonconducting

Conductivity of liquid nonconducting nonconducting conducting nonconducting

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4. On the basis of the results of her solubility test, the student concludes that X must becovalent. Comment, and state any further conclusions that can be derived about the othersubstances.

5. On the basis of her electrical conductivity test, the student states that there must be someerror because unknown Y acts like a covalent substance when solid but like an ionic sub-stance when melted. Comment on her statement.

6. The student examines the entire set of data but is unable to come to final, definite con-clusions about the substances’ types of attractions. Is it possible to do so, given the data?Explain, and state your own conclusions if any.

7. How do the overall procedure and your reasoning illustrate scientific methods?



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CHAPTER 8

Assessment



1. 4. 7. 10.

2. 5. 8.

3. 6. 9.

Short Answer


11.

12.

13.

Extended Response


14.


15. 17. 19.

16. 18.


Chemistry: Matter and Change Teacher Guide and Answers 113Fast Files, Chapters 5-8 Resources

TEACHER GUIDE AND ANSWERS

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CHAPTER 5

MiniLab 5 – Identify Compounds

Analysis

1. The colors are due primarily to electrontransitions of the metal atoms. The colors arecharacteristic of lithium, sodium, potassium,calcium, and strontium.

2. The colors are a composite of each element’svisible spectrum.

3. Answers will vary depending on the identity ofthe unknown sample.

Expected Results:

ChemLab 5 – Analyze Line Spectra

Pre-Lab

2. When electrons drop from higher-energyorbitals to lower-energy orbitals, the atom emitsenergy in the form of light. Each orbitaltransition is associated with a characteristicspectral line.

3. A continuous spectrum contains a continuumof visible colors from red to violet. Anabsorption spectrum is a continuous spectrumcontaining black lines at wavelengths associatedwith the atoms’ energy absorptions. Anemission spectrum consists of colored linesassociated with the atoms’ energy-leveltransitions.


1. At any given time, the electron occupies asingle orbital. However, it can move into

other, vacant orbitals as the atom absorbs oremits energy.

2. The color of a solution is due to the color oflight it transmits. The colors not transmittedare absorbed, and these colors comprise theabsorption spectrum.

3. The spectrum of each element is unique. Thus,the presence of a unique atomic spectrumindicates the presence of that element.

4. Answers will vary.

Inquiry Extension

Answers will vary.

Expected Results:

For each colored solution listed below, all colorsare visible except as noted.Red solution: blue and greenGreen solution: red and orangeBlue solution: yellow, orange, and some redYellow solution: blue

Teaching Transparency 15 – The Electromagnetic Spectrum

1. Radio waves are the longest waves. Gamma raysare the shortest waves.

2. Radio waves have the lowest frequency.

3. X rays (1018 s–1) have a higher frequency thanmicrowaves (1011 s–1).

4. The waves in the visible portion of thespectrum can be seen by the eye.

5. violet, blue, green, yellow, orange, red light

6. radio waves, infrared waves, green light,ultraviolet waves, gamma rays

7. Frequency and wavelength are inverselyproportional. This means that as wavelengthincreases, frequency decreases and as frequencyincreases, wavelength decreases.

8. The wavelength of a radio signal at 95.5 MHz isabout 1 m long, or 10° m.

Teaching Transparency 16 – Atomic Orbitals

1. An s orbital is spherical.

2. The size of an s orbital increases with increasingprincipal energy level number.

Compound Flame color

lithium chloride red

sodium chloride yellow

potassium chloride violet

calcium chloride red-orange

strontium chloride bright red

unknown depends on compound

Flame Test Results

113_131_CMC_FF_878761.qxd 04/14/2007 12:06 PM Page 113 Printer PDF

114 Teacher Guide and Answers Chemistry: Matter and ChangeFast Files, Chapters 5-8 Resources

3. A p orbital is dumbbell shaped. There are threep orbitals in a given sublevel.

4. Each orbital can hold two electrons.

5. These letters refer to the three perpendicularaxes; p orbitals are situated along these threeaxes in space.

6. There are five d orbitals in a given sublevel.Therefore, the d orbitals in one sublevel canhold 10 electrons.

7. dxy, dxz, dyz, and dx2–y2

8. The point where the x, y, and z axes intersectrepresents the location of an atom’s nucleus.

9. Very unlikely; the shapes of the orbitals come toa point at the intersection of the three axes,making the possibility of an electron beinglocated there very unlikely.

Teaching Transparency 17 – OrbitalFilling Sequence and Energy Levels

1. Each box represents an orbital.

2. Each orbital can hold two electrons.

3. A d sublevel can hold 10 electrons.

4. A 2p orbital has more energy than a 2s orbital.

5. A 3s orbital has more energy than a 2s orbital.

6. A 3d orbital has more energy than a 4s orbital,thus, the 4s orbital fills first.

7. The 1s orbital has the least amount of energy.

8. All atoms have 1s orbitals.

9. 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f

10. Answers will vary. An orbital with lower energyis generally occupied by an electron before anorbital of higher energy is.

Math Skills Transparency 5 – Interpreting Waves

1. All electromagnetic waves travel at the speed oflight, c. (c = 3.00 108 m/s)

2. B has a higher frequency. A has a longerwavelength.

3. 4.29 � 1014 s–1

Solution: First, convert 699 nm to meters:(699 nm) � (1 meter/109 nm) = 6.99 � 10–7 mc = ��, where c = 3.00 � 108 m/s� = c/� = (3.00 � 108 m/s)/(6.99 � 10–7 m) = 4.29 � 1014 s–1

4. 7.23 � 1014 s–1

Solution:First, convert 415 nm to meters:(415 nm) � (1 meter/109 nm) = 4.15 � 10–7 mc = ��, where c = 3.00 � 108 m/s� = c/� = (3.00 � 108 m/s)/(4.15 � 10–7 m) = 7.23 � 1014 s–1

5. Answers to question 3 should be supported bystudents’ calculations. Wave B has a higherfrequency than Wave A does.

6. 652 nmSolution:c = ��, where c = 3.00 � 108 m/s� = c/� = (3.00 � 108 m/s)/(4.60 � 1014 s–1) = 6.52 � 10–7 mConvert meters to nanometers:(6.52 � 10–7 m)(109 nm/1 m) = 652 nm

Study Guide - Chapter 5 – Quantum Theory and the Atom

Section 5.1 Light and Quantized Energy

1. energy

2. wave

3. Light

4. speed

5. wavelength

6. amplitude

7. Frequency

8. hertz

9. both A and C

10. B

11. The frequency is 2 waves/s or 2 Hz

12. c

13. b

14. c

15. b

16. false

17. true

18. true

19. true

20. false

21. false

22. true


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Section 5.2 Quantum Theory and theAtom

1. ground state

2. frequencies

3. lower

4. higher

5. electron

6. energy levels

7. atomic emission spectrum

8. No; the wavelength is far too small to be seen ormeasured even with the most sensitive scientificinstrument.

9. The proton would have the larger wavelengthbecause wavelength increases with decreasingmass and velocity.

10. c

11. a

12. d

13. b

14. The quantum mechanical model treatselectrons as waves and does not describe theelectrons’ path around the nucleus. The Bohrmodel treats electrons as particles traveling inspecific circular orbits.

15. do not

16. two

17. spherically shaped

18. n

19. electrons

20. three

21. 2s and 2p

22. nine

Section 5.3 Electron Configurations

1. electron configuration

2. lowest

3. stable

4. ground-state electron configuration

5. Aufbau principle

6. Pauli exclusion principle

7. spins

8. Hund’s rule

9. 2

10. Nitrogen; 1s22s22p3

11. 10; 1s22s22p6

12. 32; 32

13. Noble-gas notation uses the bracketed symbolof the nearest preceding noble-gas atom in theperiodic table in the electron configurations of an atom. Using noble-gas notation allowsyou to represent the complete electronconfiguration of an atom with many electronsin a shorthand form.

14. [Ar]4s23d104p2

15. c

16. b

17. d

18. b

19. a

20. a

Chapter Assessment - Chapter 5

Reviewing Vocabulary

1. f

2. d

3. i

4. e

5. a

6. h

7. c

8. b

9. g

10. Frequency is the number of waves that pass agiven point per second. Amplitude is a wave’sheight from the origin to a crest or trough.

11. Valence electrons are electrons in an atom’soutermost orbitals. An electron-dot structurerepresents an atom’s valence electrons using dots.


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)( )( ) ) )

)( )( )( )( )

)( )( )( )( )( )


12. Principal energy levels are an atom’s majorenergy levels, ranging in value from 1 to 7.Energy sublevels are contained within principalenergy levels, and their number increases as thevalue of the principal energy level increases.


1. c

2. a

3. d

4. b

5. 1

6. 2; s, p; 4

7. s, p, d

8. s, p, d, f; 16

9. 1s 2s 2px 2py 2pz

1s2 2s2 2p3

10. 1s 2s 2px 2py 2pz

1s2 2s2 2p5

11. 1s 2s 2px 2py 2pz 3s

1s2 2s2 2p6 3s1


1. c

2. b

3. c

4. b

5. c

6. c

Thinking Critically

1. � = c/�� = (3.00 � 108 m/s)/(103.7 � 106 s–1) = 2.892 mStudents should label the electromagneticspectrum between 101 m and 1 m.

2. Microwaves are higher in frequency and shorterin wavelength than radio waves.

3. 1s 2s 2px 2py 2pz

3s 3px 3py 3pz

4. complete electron configuration:1s22s22p63s23p1; noble-gas notation: [Ne]3s23p1

5. [Kr]5s24d105p5

6. a. boron

b. potassium

7. a. Ne

b. H

c. C

d. S

Applying Scientific Methods

1. Students should recognize that the atomicemission spectra of samples A and C areidentical; hence, those samples are the sameelement. Students should also recognize that the spectra of samples B and D are different;therefore, samples B and D are differentelements. Students should conclude that thefour samples represent three different elements.

2. The atomic emission spectra showdiscontinuous parts of the visible portion of theelectromagnetic spectrum. The other portionsof the electromagnetic spectrum are not visible,although they are involved.

3. The atomic emission spectrum would notchange. Like a fingerprint, the atomic emissionspectrum of each element has a characteristic,specific pattern of lines.

4. Each line represents a change in energy of oneof the atom’s electrons. Students may also saythat each line represents a photon of a specificenergy being emitted (or absorbed).

5. Students should conclude that sample B ishydrogen because the atomic emissionspectrum of sample B and that of hydrogen are identical.

6. The Bohr model explains only hydrogen’satomic emission spectrum (or more correctlythe spectra of atoms with only one electron);


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)( )( )( )( )(

)( )

��

��

��

��

�



therefore, it can be used to explain the atomicemission spectrum of sample B only.

7. A: 6; each line represents a photon being emitted B: 4; each line represents a photon being emittedC: 6; each line represents a photon being emittedD: 2; each line represents a photon being emitted

8. By looking at the spectrum, you can see that theemission lines get closer together as thewavelength becomes shorter. The lines gettingcloser together demonstrate the differencebetween successive energy levels.

9. Answers may vary. Dark absorption lines wouldbe seen in the absorption spectra at the samewavelengths where bright emission lines are inthe emission spectra. Students should realizethat the spectral pattern for each element wouldnot change because the energies would

CHAPTER 6

MiniLab 6 – Organize Elements

Analysis

1. Students will have organized the known cardsaccording to the following grid.

2. The wavelength decreases across the period andcolor becomes lighter down the group.

3. The mass increases across the period and downthe group. Cx does not fit the period trend formass, but it fits in the third column with othergreen, brittle solids.

4. Ph would fit in the third period, first columnbased on color and stated trends. The masswould be between 99 g and 106 g.

5. The remaining gap would be a yellow-coloredliquid with a mass most likely between 70 g and82 g.

ChemLab 6 – Investigate DescriptiveChemistry

Pre-Lab

3. All naturally occurring metals are solids, exceptfor mercury, which is a liquid at roomtemperature. All metalloids are solids.Nonmetals are primarily gases and solids, withbromine being the only liquid.

4. Luster: shininess; malleability: capable ofbeing flattened into sheets or formed intoshapes; electrical conductivity: capable oftransmitting an electric current; they areproperties commonly associated with metals.


1.–4. Answers will vary depending on the samplesprovided to students.

5. Students might note that the metalliccharacteristic increases from right-to-left,and from top-to-bottom.

Inquiry Extension

Answers will vary.

Teaching Transparency 18 – The Periodic Table

1. 117

2. 34

3. Pd

4. 87.62 amu

5. Their boxes contain a red balloon.

6. 18

7. group or family

8. 7

9. period

10. period 1

11. manganese

12. Their boxes are tinted blue.

13. Their boxes are tinted green.

Xn Ad Tu Qa

Bp Pd Lq

Rx Cx Ax


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14. Their boxes are tinted yellow.

15. alkali metals

16. alkaline earth metals

17. halogens

18. noble gases

19. Their valence electron configurations areidentical.

Teaching Transparency 19 – The s-, p-, d-, and f-Block Elements

1. s-, p-, d-, and f-blocks

2. Each block represents the energy sublevel being filled by valence electrons.

3. They have valence electrons only in the s orbitals.

4. s1

5. s2

6. The single s orbital can hold a maximum of twovalence electrons.

7. The three p orbitals can each hold a maximumof two electrons, thus, the p orbitals can containa maximum of six valence electrons, whichcorresponds to the six columns spanned by thep-block.

8. There are no p-block elements in period 1because the p sublevel does not exist for thefirst principal energy level.

9. p2

10. [He]2s22p6

11. period 4

12. The five d orbitals can each hold a maximum oftwo electrons, resulting in a total of ten possiblevalence electrons.

13. d1

14. [Ar]4s23d2

15. period 6

16. a. 15b. 6c. p

Teaching Transparency 20 – Atomic and Ionic Radii

1. groups 1 and 2 and 13 through 17; periods 1through 6

2. picometer (pm); 10–12 m

3. 31 pm and 265 pm; helium and cesium,respectively

4. The atomic radius of the elements within aperiod generally decreases as the atomicnumber of the elements increases.

5. Exceptions are antimony (Sb) and tellurium (Te) in period 5, and bismuth (Bi)and polonium (Po) in period 6.

6. With increasing atomic number, the increasedpositive charge of the nucleus pulls morestrongly on the outermost electrons, pullingthem closer to the nucleus. Consequently, theatomic radius decreases.

7. The atomic radius of the elements within agroup generally increases as the atomic numberof the elements increases.

8. There are no exceptions.

9. With increasing atomic number, the increasedpull by the larger positive charge of the nucleusis offset by the outer electrons’ larger orbitalsand by shielding by inner electrons.Consequently, the atomic radius increases.

10. The charge of the ion of each element is thesame as the element’s group number.

Teaching Transparency 21 – First Ionization and Successive Ionization Energies

1. First ionization energy is the energy required toremove the first electron from a gaseous atom.

2. rubidium; helium; about 400 and 2375 kJ/mol,respectively

3. The first ionization energy of the elementswithin a period generally increases as theatomic number of the elements increases.

4. With increasing atomic number, the increasedpositive charge of the nucleus produces anincreased hold on the valence electrons.Consequently, the first ionization energyincreases.


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5. The values of the successive ionization energies increase.

6. beryllium, magnesium, calcium, strontium

7. The jump occurs after the valence electronshave been removed.

8. The first ionization energy of the elementswithin a group generally decreases as the atomicnumber of the elements increases.

9. With increasing atomic number, the size of theatom increases and the valence electrons arefather from the nucleus. Consequently, lessenergy is needed to remove them, and the firstionization energy decreases.

Math Skills Transparency 6 – Using the Periodic Table

1. a. 8 b. 1 c. 3 d. 2 e. 6 f. 7 g. 5 h. 4

2. a. fourth energy levelb. second energy levelc. seventh energy leveld. first energy levele. third energy levelf. fifth energy level

3. a. [He]2s1

b. [He]2s22p5

c. [Ar]4s23d104p3

d. [Kr]5s2

e. [Xe]6s24f145d106p3

4. a. group 18, period 1, s-blockb. group 13, period 3, p-blockc. group 1, period 4, s-blockd. group 3, period 5, d-blocke. group 16, period 6, p-block

Study Guide - Chapter 6 – The Periodic Table and Periodic Law

Section 6.1 Development of the Modern Periodic Table

1. octaves

2. eight

3. nine

4. accepted

5. Dmitri Mendeleev

6. atomic mass

7. elements

8. atomic number

9. Henry Moseley

10. protons

11. periodic law

12. properties

13. 14.007 u

14. 7

15. [He]2s22p3

16. Nitrogen

17. N

18. b

19. c

20. d

21. a

22. e

23. three

24. metals

25. true

26. Group 2

27. true

28. noble gases

29. nonmetals

Section 6.2 Classification of the Elements

1. h

2. f

3. n

4. a

5. k

6. d

7. l

8. o

9. i

10. g

11. b

12. m

13. c

14. j

15. e


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16. Sodium and potassium have the same numberof valence electrons.

17. The energy level indicates the period.For example, lithium’s valence electron is in the second energy level and lithium is found in period 2.

18. four

19. groups 1 and 2

20. The s orbital holds a maximum of twoelectrons.

21. groups 13-18

22. group 18 elements have both their s orbitalsand p orbitals completely filled with electrons.This configuration is very stable, thus, thegroup 18 elements are very unreactive.

23. 40

24. groups 3-8

25. The seven f orbitals hold a maximum of14 electrons.

26. 1s22s22p63s23p4

Section 6.3 Periodic Trends

1. c

2. c

3. d

4. b

5. a

6. a

7. Ionization energy is the energy required toremove an electron from a gaseous atom.

8. A high ionization-energy value indicates thatthe atom has a strong hold on its electrons andis not likely to lose an outer electron and form apositive ion.

9. The first ionization energies generally increaseas you move left-to-right across a period.The increased nuclear charge of each successiveelement produces an increased hold on thevalence electrons.

10. The first ionization energies generally decreaseas you move down a group. Because atomic sizeincreases down a group, the valence electronsare farther from the nucleus and, therefore, lessstrongly attracted to the nucleus. As a result,

less energy is required to remove the valenceelectrons.

11. Atoms tend to gain, lose, or share electrons toacquire a full set of eight valence electrons.

12. The electronegativity of an element indicates its atom’s ability to attract electrons in achemical bond.

13. Electronegativities generally increase as youmove left-to-right across a period and decreaseas you move down a group.



1. j

2. k

3. l

4. a

5. b

6. m

7. e

8. d

9. h

10. f

11. g

12. i

13. c

14. Answers will vary.

15. Answers will vary. Every element can beclassified as a metal, nonmetal, or metalloid.


1. Newlands developed the law of octaves,which correctly demonstrated the concept ofperiodic behavior.

2. Moseley arranged the elements by increasingatomic number, resulting in the modernperiodic table and periodic law.

3. Mendeleev developed the first widely acceptedperiodic table of elements by arranging theelements by increasing atomic mass intocolumns with similar properties.

4. 1

5. 2


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6. Ni

7. Th

8. 4

9. 6

10. Si

11. 5

12. O

13. Be

14. 3


1. d

2. c

3. a

4. d

5. c

6. c

7. a

8. d

9. c

10. b

11. d

Thinking Critically

1. The atomic radii of the elements in groups 1 and 2 and groups 13-17 decrease as theatomic numbers of the elements increase withina period.

2. Increasing the nuclear charge of the nucleusincreases the attractive force on the outerelectrons, pulling them closer to the nucleus.

3. The atomic radii of the transition elements inperiod 4 decrease, then increase, as the atomicnumbers of the elements increase.

4. The two elements in the first period arehydrogen and helium. Hydrogen atoms form apositive ion, H+, a proton, which has noelectron cloud associated with it and, therefore,no radius associated with it. Helium is a noblegas and does not form ions.

5. The arsenic ion has a 3+ charge. The charge ispositive since the radius of the arsenic ion issmaller than the radius of the arsenic atom.


1. One triad, in which the atomic mass of thesecond element is about the average of theatomic masses of the first and third elements,includes Li, Na, and K—the alkali metals ofgroup 1, which have common properties.Similarly, the elements S, Se, and Te make up asecond triad and have similar properties, as theyall are group 16 elements.

2. Densities of Ca, Sr, and Ba: 1.55 g/mL,2.6 g/mL, 3.62 g/mLMelting points of Ca, Sr, and Ba: 841.5°C,776.9°C, 726.9°CBoiling points of Ca, Sr, and Ba: 1500.5°C,1412°C, 1845°CDensities of Cl, Br, and I: 0.003 214 g/mL,3.11 g/mL, 4.93 g/mLMelting points of Cl, Br, and I: -101°C, –7.25°C,113.6°CBoiling points of Cl, Br, and I: -34°C, 59.35°C,184.5°C

In the Cl, Br, and I triad, the density, meltingpoint, and boiling point sequences each show atrend similar to that of the sequence of atomicmasses. That is, the value of the middle memberof the triad is close to the average of the valuesof the other two members. In the Ca, Sr, and Batriad, the density sequence follows a trendsimilar to that of the atomic mass sequence,however, the melting point sequence is reversed.There is no sequence in the boiling points ofthe elements in this triad.

3. Silicon: 3.0Germanium: 74Tin: 480

CHAPTER 7

MiniLab 7 – Observe Properties

Analysis

1. spring and hardened steel

2. Possible hypothesis: Soft steel has large crystals;tempered steel has intermediate-sized crystals.


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0 10 20 30 40 50 60 70 80 90 100 110 120 130 140 150 160 170 180 190 200

Atomic mass (amu)

Li C Na S K Se Te Au



3. Answers may include pop-up tent, spring toys,and clips.

4. Smooth curves are possible, but they dent easily.

5. it is brittle and breaks easily; yes

Expected Results:

The metal can be straightened when heatedbecause the layers of atoms separate and slide overone another easily. The hard steel breaks whenstudents attempt to bend it. The tempered steel ishard and has a springlike feel.

Chemlab 7 – Synthesize an Ionic Compound

Pre-Lab

1. Variable: mass of Mg; Constant: there must bean excess of oxygen

2. Mg 1s22s22p63s2

a. lose electronsb. Mg2+1s22s22p6

c. neon

3. O 1s22s22p4, N 1s22s22p3

a. Both will gain electrons.b. O2– 1s22s22p6, N3– 1s22s22p6

c. Both have the configuration of neon.

5. The mass of the magnesium and the mass ofthe magnesium products are calculated. Othermass values are measured directly.

6. The mass of magnesium ribbon is calculated by subtracting the mass of the crucible from the mass of the crucible and magnesium.The mass of the magnesium products iscalculated by subtracting the mass of thecrucible from the mass of the crucible and itscontents after heating.


1. Refer to Expected Results.

2. heat and light; It is more stable than thereacting elements.

3. There is an increase in mass from 0.29 g to 0.37 g.

4. MgO, magnesium oxide; Mg3N2, magnesiumnitride

5. MgO; The product appears white.

6. yes; Yes, because ionic compounds conduct anelectric current in solution.

7. Possible answers include that some of theproduct blew away or that the reaction wasincomplete.

Inquiry Extension

Student experimental designs will vary. However,the basic point students should investigate is thatmore concentrated ionic solutions are moreconductive than less concentrated ones.

Expected Results

Sample Data

Mass of empty crucible: 7.56 g

Mass of crucible + Mg ribbon before heating:7.85 g

Mass of Mg ribbon: 0.29 g

Mass of crucible + Mg ribbon after heating: 7.93 g

Mass of Mg products: 0.37 g

Teaching Transparency 22 – Formation of Ions

1. calcium and oxygen

2. Yes; each contains equal numbers of protonsand electrons.

3. ionization energy

4. positive; cation

5. negative; anion

6. No; it is not a stable octet of electrons.

7. No; it is not a stable octet of electrons.

8. Yes; it is a stable octet of electrons.

9. Yes; it is a stable octet of electrons.

10. 1s2 2s2 2p6 3s2 3p6; argon, a noble gas, has thesame configuration.

11. 1s2 2s2 2p6; neon, a noble gas, has the sameconfiguration.

Teaching Transparency 23 – Ionic Bonds

1. two

2. 2+; it must lose its two valence electrons.It tends to do so to achieve the stable octetconfiguration of a noble gas.

3. seven


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4. 1-; it must gain one electron. It tends to do soto achieve the stable octet configuration of anoble gas.

5. One Mg atom and two Cl atoms react becauseone 2+ ion and two 1- ions produce therequired overall charge of zero. The formula isMgCl2.

6. Ionic; the two valence electrons of themagnesium atom are transferred, one to each ofthe chlorine atoms. The attraction of oppositelycharged particles (the Mg2+ and the Cl–) holdsthe atoms together, forming an ionic bond.

7. MgS; the Mg atom transfers its two valenceelectrons to the S atom. In that way, eachachieves a stable noble-gas configuration.One Mg ion with a charge of 2– is balanced byone S ion with a charge of 2+, producing a netoverall charge of zero.

8. Mg3P2; each Mg atom transfers its two valenceelectrons to the P atoms, each of which mustgain three electrons. In that way, each achieves astable noble-gas configuration. Three Mg ions,each with a charge of 2+, are balanced by two P ions, each with a charge of 3–, producing anet overall charge of zero.

Teaching Transparency 24 – Formulas for Ionic Compounds

1. One; it would tend to lose an electron becausethe loss would leave it with a stable octet ofelectrons, like that of a noble gas.

2. Five; it would tend to gain three electronsbecause the gain would give it a stable octet ofelectrons, like that of a noble gas.

3. Na3P, sodium phosphide; Na+, +1; P3–, –3

4. Six; it would tend to gain two electrons becausethe gain would give it a stable octet of electrons,like that of a noble gas.

5. Na2S, sodium sulfide; Na+, +1; S2–, –2

6. Three; it would tend to lose three electronsbecause the loss would leave it with a stableoctet of electrons, like that of a noble gas.

7. Seven; it would tend to gain one electronbecause the gain would give it a stable octet ofelectrons, like that of a noble gas.

8. AlCl3, aluminum chloride; Al3+, +3; Cl–, –1

9. Al2S3, aluminum sulfide; Al3+, +3; S2–, –2

Teaching Transparency 25 – Metallic Bonding

1. a crystal lattice

2. no; delocalized

3. No; they are the valence electrons from themetal atoms.

4. The delocalized negative electrons came fromneutral atoms, thus leaving the atoms with apositive charge.

5. They are equal. Group 1 atoms have only onevalence electron and thus only one electron thatcan become delocalized.

6. There are twice as many electrons as group 2atoms.

7. The delocalized electrons are simultaneouslyattracted to more than one metal cation.

8. electron sea model

9. Metals tend to conduct electricity well.The model’s delocalized electrons are not heldstrongly by individual atoms and are thus ableto move easily throughout the metal.

10. Metals are malleable and ductile. The model’sdelocalized electrons are able to move aroundthe positive metal core atoms and keep thecrystal from breaking during hammering ordrawing into wire.

Math Skills Transparency 7 – Calculating Numbers of Electronsand Predicting Ionic Change

1. The atom becomes a cation with a charge of1+ because it then has one fewer negativelycharged particles than it has protons, which arepositively charged.

2. The atom becomes an anion with a charge of 2–because it then has two more negatively chargedparticles than it has protons.

3. Charge of ion = Number of protons – Numberof electrons.


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4. a. 1s22s2; it is likely to form an ion with a 2+charge because by losing two valenceelectrons, it achieves the stable noble-gasconfiguration 1s2.

b. 1s22s22p5; it is likely to form an ion with a 1– charge because by gaining one electron,it achieves the stable noble-gasconfiguration 1s22s22p6.

c. 1s22s22p63s23p6; it is not likely to bond orform ions because it already has a stablenoble-gas configuration.

d. 1s22s22p63s23p4; it is likely to form an ionwith a 2– charge because by gaining twoelectrons, it achieves the stable noble-gasconfiguration 1s22s22p63s23p6.

e. 1s22s22p63s1; it is likely to form an ion with a 1+ charge because by losing its onevalence electron, it achieves the stablenoble-gas configuration 1s22s22p6.

f. 1s22s22p3; it is likely to form an ion with a 3– charge because by gaining threeelectrons, it achieves the stable noble-gasconfiguration 1s22s22p6.

Math Skills Transparency 8 – Determining Numbers of Ions

1. Total positive charge + Total negative charge =Zero

2. a. oneb. twoc. three

3. a. threeb. onec. two

4. a. (NH4)2SO4; (2 � 1+) + (1 � 2–) = 0

b. Na3PO4; (3 � 1+) + (1 � 3–) = 0

c. Mg(HSO4)2; (1 � 2+) + (2 � 1–) = 0

d. Al2(CO3)3; (2 � 3+) + (3 � 2–) = 0

e. (NH4)3AsO4; (3 � 11–) + (1 � 3–) = 0

f. Ca(C2H3O2)2; (1 � 2+) + (2 � 1–) = 0

g. NH4NO2; (1 � 1+) + (1 + 1–) = 0

Study Guide - Chapter 7 – Ionic Compounds and Metals

Section 7.1 Ion Formation

1. chemical bond

2. nucleus

3. electrons

4. ions

5. valence

6. energy level

7. noble gases

8. octet

9. pseudo-noble gas formations

10. false

11. true

12. false

13. true

14. true

15. false

16. true

17. false

Section 7.2 Ionic Bonds and IonicCompounds

1. c

2. b

3. a

4. c

5. c

6. d

7. a

8. b

9. b

10. b

11. b

12. high

13. high

14. hard

15. brittle

16. poor


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17. good

18. good

19. true

20. true

21. false

22. false

23. true

Section 7.3 Names and Formulas for Ionic Compounds

1. monatomic

2. oxidation number

3. electrons

4. zero

5. polyatomic

6. oxyanion

7. –ate

8. –ite

9. cation

10. anion

11. subscript

12. lower right

13. one

14. e

15. d

16. b

17. c

18. a

19. sodium iodide

20. calcium chloride

21. potassium sulfide

22. magnesium oxide

23. lithium hydrogen sulfate

24. ammonium bromide

25. calcium nitride

26. cesium phosphide

27. potassium bromate

28. magnesium hypochlorite

29. lithium peroxide

30. beryllium phosphate

31. ammonium carbonate

32. sodium bromate

33. iron(III) oxide

34. iron(II) iodate

35. Be3N2

36. NiCl237. KClO2

38. Cu2O

39. MgSO3

40. (NH4)2S

41. Ca(IO3)2

42. Fe(ClO4)3

43. Na3N

Section 7.4 Metallic Bonds and theProperties of Metals

1. electron sea model

2. They are free to move from one atom toanother.

3. the valence electrons

4. Cations; they are positively charged.

5. The electrons are not completely lost by themetal atoms, as they are in an ionic solid.

6. They are bonded by the oppositely chargedelectron sea that surrounds them.

7. yes; when the metal is hammered, thedelocalized electrons move, keeping the metallic bonds intact.

8. no

9. Yes; the delocalized electrons move, absorb andrelease protons.

10. Yes; the metallic bonds are strong.

11. no

12. Yes; when the metal is pulled, the delocalizedelectrons move, keeping the metallic bondsintact.

13. no

14. Yes; the delocalized electrons are mobile.


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1. e

2. n

3. d

4. p

5. c

6. b

7. f

8. q

9. g

10. m

11. k

12. j

13. a

14. i

15. o

16. h

17. l


1. always

2. true

3. metals

4. more

5. true

6. true

7. true

8. lose

9. true

10. true

11. a

12. d

13. c

14. d

15. c

16. c


1. 1s22s1

2. 1s22s22p5

3. The lithium atom loses one electron, to form an Li+ ion The configuration is 1s2.

4. The lithium has achieved the stableconfiguration of a noble gas.

5. The fluorine atom gains one electron, to forman F– ion. The configuration is 1s22s22p6.

6. The fluorine has achieved a stable octetconfiguration, like that of a noble gas.

7. ionic

8. There is a transfer of electrons.

9. LiF; lithium fluoride

10. electrostatic forces of attraction between theoppositely charged ions

11. A crystal lattice; the arrangement is regular andrepeating, with positive ions surrounded bynegative ions, and vice versa.

12. Ionic solids tend to be brittle and are not goodconductors of electricity, whereas metals aremalleable and ductile and are good conductors.

Thinking Critically

1. The more negative the lattice energy is, thestronger is the attraction between ions and themore stable is the ionic compound.

2. The lattice energy decreases (becomes lessnegative) as atomic number and size increasefor the halides of sodium. The samerelationship exists for the potassium halides.

3. The lattice energy values are lower for thehalides of potassium than they are for thecorresponding halides of sodium. As the ionicsize increases, the lattice energy becomes less negative.

4. The melting point of NaBr should be higherthan that of NaI because NaBr has higher latticeenergy and therefore more energy would berequired to separate the ions.

The melting point of NaBr should be higherthan that of KBr for the same reason.

5. The lattice energy should be greater (morenegative) for ions of greater charge because theelectrostatic force of attraction would be higher.


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1. His conclusion is questionable. Although it iscorrect that nonmetallic elements tend to benonlustrous, so do ionic compounds becauseboth types of substances do not havedelocalized electrons.

2. Her conclusion is questionable. Although it istrue that most nonmetallic elements are poorconductors of electricity, so are ioniccompounds when in the solid state.

3. His conclusion is almost definitely faulty. Metals tend to be malleable and ductile. On the basisof his results, C should have concluded that the unknown is probably either a nonmetallicelement or an ionic compound.

4. A limited conclusion is possible. Nonmetallicelements tend to be poor conductors ofelectricity, even when melted, so they can beruled out. Metals and melted ionic solids bothtend to conduct well. The melting point iswithin the range of that of many metallicelements and ionic solids, so it is not very usefulin further narrowing the conclusion.

5. No; neither separate conclusion can rule outeither nonmetals or ionic solids.

6. Yes; only ionic solids are nonconductors assolids but good conductors as liquids.

7. A test of conductivity of the solution wouldhave been useful because dissolved ionic solidstend to conduct electricity.

8. Communicating findings and pooling theresults of different tests are helpful in allowingproper conclusions to be drawn.

CHAPTER 8

MiniLab 8 – Compare Melting Points

Analysis

1. The paraffin melted first. The salt crystals did not melt.

2. paraffin, low; sugar, medium; salt crystals, veryhigh

3. ionic bonds: salt covalent bonds: paraffin and sugar

4. Ionic compounds have lower melting pointsthan covalently bonded compounds.

Expected Results:

The lone pairs of electrons take up more spacethan paired electrons. Refer to Figure 9-3 to see Lewis structures for CH4, NH3, and H2O. SeeSolutions Manual for sketches.

ChemLab 8 – Model Molecular Shapes

Pre-Lab

2. Valence Shell Electron Pair Repulsion; TheVSEPR model bases its predictions of molecularshape on the arrangement that minimizes therepulsion of shared and unshared electron pairsaround a central atom.

3. To determine whether or not a bond is polar,compare the electronegativities of the atomsinvolved in the bond. To determine whether ornot a molecule is polar, you also need to knowthe shape of the molecule.

4. A resonance structure occurs when more thanone valid Lewis structure can be written for amolecule. If a molecule has resonancestructures, you can infer that the molecule hasat least one single bond one double bond.


1. increasing flexibility: triple, double, single;increasing strength: single, double, triple

2. The H2O molecule has two bonds and two lone pairs around the central atom.The lone pairs take up space around the centralatom and repel the bonding electrons, causingthe bent shape. The CO2 molecule has twodouble bonds with no lone pairs. The bondingelectrons repel to form the linear shape,which maximizes the distance between electron densities.

3. The SO3 molecule has a central S and threeterminal O atoms. One of the terminal O atomsforms a double bond. Three resonancestructures exist, one for each possible locationof the double bond.

4. The following molecules are polar: H2O, PH3,HCN, and CO. All others are nonpolar.


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Inquiry Extension

Students should assemble two models, each with acentral oxygen atom joined to two terminaloxygen atoms. One terminal atom is joined by asingle bond, the other by a double bond. The loca-tion of these two bonds determines the two reso-nance structures. The Lewis structures shouldshow that you can convert between the two reso-nance structures by swapping the position of alone pair and a covalent bond.

Teaching Transparency 26 – Lewis Structures

1. A hydrogen atom can form only one bond, so itcannot be a central atom.

2. Each nitrogen (N) atom has five valenceelectrons and each hydrogen (H) atom has onevalence electron, resulting in a total of (2 � 5) +(4 – 1) = 14.

3. The total number of valence electrons from step 2 must be divided by 2; in the case ofN2H4, 14/2 = 7.

4. five; a line

5. 7 – 5 = 2. Because the H atoms already had acomplete outer level of electrons, a lone pair ofelectrons was positioned next to each N atom.

6. The central N atoms already had completeoctets, and the Lewis structure was alreadycorrect.

7. 4 + 6 + 6 = 16, the total number of valenceelectrons

8. 16/2 = 8, the number of bonding pairs

Teaching Transparency 27 – VSEPR Model and Molecular Shape

1. Pairs of electrons, either shared or unshared,repel each other as much as possible around acentral atom.

2. no lone pairs; three shared pairs

3. The BH3 molecule is trigonal planar. There arethree electron pairs that repel so that they are asfar as possible from one another. The shape thatmaximizes the distance is trigonal planar.

4. no lone pairs; four shared pairs; tetrahedral

5. one lone pair; three shared pairs; trigonalpyramidal

6. two lone pairs; two shared pairs; bent

7. three lone pairs; one shared pair; linear

Teaching Transparency 28 –Electronegativity and Polarity

1. Electronegativity is the tendency of an atom to attract electrons.

2. Electronegativity increases from left to right across the period and decreases down the group.

3. The covalent bond is polar; toward the moreelectronegative atom.

4. 3.16 – 2.55 = 0.61, polar covalent bond

5. CCl4 is a nonpolar molecule because itssymmetrical tetrahedral shape results in abalancing of the partial charges resulting fromeach polar covalent bond.

6. 3.98 – 3.04 = 0.94, polar covalent bond; NF3 is apolar molecule because its asymmetrical,trigonal pyramidal shape does not balance itspartial charges.

7. 2.96 – 1.57 = 1.39, polar covalent bond; BeBr2 isa nonpolar molecule because its symmetrical,linear shape balances its partial charges.

8. 3.98 – 1.57 = 2.41, ionic bond

9. 2.66 – 2.04 = 0.62, polar covalent bond;BI3 is a nonpolar molecule because itssymmetrical, trigonal planar shape balances its partial charges.

Math Skills Transparency 9 – Determining the Names of BinaryCompounds and Their Numbers ofAtoms

1. five; six; eight

2. a. one nitrogen, three iodine;nitrogen triiodide

b. one carbon, two sulfur; carbon disulfide

c. two nitrogen, four oxygen;dinitrogen tetroxide

d. one silicon, six fluorine; silicon hexafluoride

e. one oxygen, two fluorine; oxygen difluoride


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3. a. P4O10four phosphorus, ten oxygen, a total of4 + 10 = 14 atoms;tetraphosphorus decoxide

b. S2O7two sulfur, seven oxygen, a total of 2 + 7 = 9 atoms; disulfur heptoxide

c. Si3H8three silicon, eight hydrogen, a total of3 + 8 = 11 atoms;trisilicon octahydride

4. a. tetrasulfur dinitride

b. dichlorine monoxide

Math Skills Transparency 10 –Determining ElectronegativityDifference and Percent IonicCharacter

1. a. 3.16 – 0.93 = 2.23

b. 2.58 – 1.31 = 1.27

c. 2.19 – 1.61 = 0.58

d. 3.16 – 1.90 = 1.26

e. 2.58 – 1.90 = 0.68

2. Na and Cl only

3. Percent ionic character and electronegativitydifference; as electronegativity differenceincreases, percent ionic character increases.

4. a. 70% b. 32% c. 10% d. 32% e. 12%

5. % ionic character + % covalent character =100%

6. a. 100% – 70% = 30%

b. 100% – 32% = 68%

c. 100% – 10% = 90%

d. 100% – 32% = 68%

e. 100% – 12% = 88%

Study Guide - Chapter 8 – Covalent Bonding

Section 8.1 The Covalent Bond

1. covalent bond

2. exothermic

3. molecule

4. sigma bond

5. pi bond

6. b

7. d

8. c

9. a

Section 8.2 Naming Molecules

1. false

2. false

3. false

4. true

5. true

6. true

7. false

8. false

9. c

10. i

11. g

12. e

13. b

14. d

15. f

16. a

17. h

Section 8.3 Molecular Structures

1. true

2. false

3. true

4. false

5. false

6. false

7. true

8. true

9. true

10. false

11. true

12. false

13. c


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14. b

15. d

16. a

Section 8.4 Molecular Shape

1. a

2. c

3. d

4. a

5. d

6. c

7. hybridization

8. identical

9. carbon

10. sp3

11. methane

Section 8.5 Electronegativity and Polarity

1. the tendency of an atom to attract electrons

2. fluorine; 3.98; halogens; group 17

3. francium; 0.7; alkali metals; group 1

4. Electronegativity tends to decrease.Electronegativity tends to increase.

5. The values are subtracted.

6. true

7. false

8. c

9. d

10. b

11. a

12. c

13. b

14. b

15. d

16. b

17. d



1. g

2. i

3. k

4. b

5. n

6. h

7. j

8. f

9. a

10. m

11. l

12. d

13. c

14. e


1. a

2. d

3. b

4. a

5. d

6. c

7. c

8. di-

9. true

10. true

11. true

12. true

13. less

14. electrons


1. 1s22s22p2

2. sp3; four

3. It has no unhybridized orbitals.

4. four other atoms; four single bonds,each sigma only

5. tetrahedral


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6. The bonds would be polar, but the molecule asa whole would be nonpolar because of its shapeand the fact that the polarities would “cancel”one another.

7. sp2; three

Thinking Critically

1. The electrons are not shared equally, but aremore attracted to the more electronegativeatom. The bond is polar covalent.

2. N is more electronegative than B. Both are inperiod 2, with N to the right of B. P is moreelectronegative than Al. Both are in period 3,with P to the right of Al.

3. Electronegativity increases from left to rightacross both periods.

4. B is more electronegative than Al. Both are ingroup 13, with B above Al. N is moreelectronegative than P. Both are in group 15,with N above P.

5. Electronegativity decreases going down a groupin the periodic table.


1. The conclusion is sound. Only covalentsubstances are soft, although some covalentsubstances (network solids) are hard and brittle.Ionic solids also are brittle, so there is no way totell whether Y or Z is the ionic substance.

2. The claim is not valid. The low melting pointsof W and X strongly suggest that they arecovalent molecular substances. The highmelting points of Y and Z suggest that one mustbe ionic, the other covalent network, with Zmore likely to be the network solid.

3. The conclusion is valid. Hydrogen-bondedsubstances have relatively high boiling pointsfor molecular substances, and that is true of X.Because Y has very low melting and boilingpoints, and the remaining unknowns do not, itmust be the substance that has the very weakdispersion forces.

4. The student’s conclusion is valid because onlynonpolar covalent substances are highly solublein nonpolar solvents. The high solubility ofW and Y in the polar solvent suggests that oneis the ionic substance and the other thehydrogen-bonded (highly polar) one, but moreinformation would be needed to derive furtherconclusions.

5. The statement is not valid. The two pieces ofinformation together suggest that Y is ionicbecause ionic substances conduct when in theliquid state, but not when in the solid state.

6. The entire set of data does allow identificationof each substance’s type of attractions. W mustbe hydrogen-bonded covalent because of itsrelatively high boiling point for a covalentmolecular compound and its solubility in thepolar solvent. The fact that it does not conductas liquid and that its melting point is not veryhigh allows one to rule out that it is ionic. Xmust be covalent with inter-moleculardispersion forces because of its very low meltingand boiling points, lack of conductivity, andsolubility in the nonpolar solvent. Y must beionic because of its high melting and boilingpoints coupled with its conductivity as liquidand its solubility in the polar solvent. Z must bethe covalent network solid. Its very highmelting and boiling points, brittleness, and lackof conductivity support that conclusion.

7. The steps involve those often used ininvestigating a problem scientifically. A question was initially posed, experimentsplanned and carried out, observations made,data analyzed, and conclusions reached on thebasis of the data.


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