All Chemistry Worksheets Chapter 5-8

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-Hill 8787 Orion Place Columbus, OH 43240-4027 ISBN: 978-0-07-878761-4 MHID: 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

Table of Contents 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

Copyright © Glencoe/McGraw-Hill, a division of The McGraw-Hill Companies, Inc.

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

Teacher Guide and Answers Chapter 5 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 113 Chapter 6 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 117 Chapter 7 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 121 Chapter 8 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 127

iii

Teacher Approval Initials Date of Approval

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.

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.

vi


1. What is the purpose of the investigation?

Table of Contents

Reproducible Pages

Chapter 5 Electrons in Atoms MiniLab . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 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

1

Name

Date

Class

mini LAB 5

Identify 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 Flame Test Results Compound

Flame color

Lithium chloride Sodium chloride Potassium chloride Calcium chloride Strontium chloride Unknown

Analysis 1. Suggest a reason why each compound produced a flame of a different color, even though 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.

2

Chemistry: Matter and Change • Chapter 5

ChemLab and MiniLab Worksheets


1. Read and complete the lab safety form. 2. Dip one of six cotton swabs into the lithium chloride solution. Put the swab into the flame of a Bunsen burner. Observe the color of the flame, and record it in your data table. 3. Repeat Step 2 for each of the metallic chloride solutions (sodium chloride, potassium chloride, calcium chloride, and strontium chloride). Record the color of each flame in your data table. 4. Compare your results to the flame tests shown in the Elements Handbook. 5. Repeat Step 2 using a sample of unknown solution obtained from your teacher. Record the color of the flame produced. 6. Dispose of the used cotton swabs as directed by your teacher.

Name

CHEMLAB

Date

Class

5

Analyze Line Spectra

E

mission 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 sample absorb specific wavelengths. This produces dark lines in the continuous spectrum of white light and is called an absorption spectrum.

Problem

Objectives

Materials

What absorption and emission spectra do various substances produce?

• Observe emission spectra of several gases. • Observe the absorption spectra of various solutions. • Analyze patterns of absorption and emission spectra.

ring stand with clamp 40-W tubular lightbulb light socket with grounded power cord 275-mL polystyrene culture flask Flinn C-Spectra® or similar diffraction grating

food coloring (red, green, blue, and yellow) set of colored pencils spectrum tubes (hydrogen, neon, and sodium) spectrum–tube power supplies (3)


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.

Pre-Lab Read the entire CHEMLAB. 2. Explain how electrons in an element’s atoms produce an emission spectrum. 1.


3.

Distinguish among a continuous spectrum, an emission spectrum, and an absorption spectrum.

4.

Use the data table on the next page.


3

Name

Date

5

Procedure 1. 2.

3.

4.

5. 6.

7.

Read and complete the lab safety form. Use a Flinn C-Spectra® or similar diffraction grating to view an incandescent lightbulb. What do you observe? Draw the observed spectrum using colored pencils. Use the Flinn C-Spectra® to view the emission spectra from tubes of gaseous hydrogen, neon, and sodium. Use colored pencils to make drawings in the data table of the spectra observed. Fill a 275-mL culture flask with about 100-mL water. Add 2 or 3 drops of red food coloring to the water. Shake the solution. Repeat step 4 for the green, blue, and yellow food coloring. Set up the 40-W lightbulb so that it is near eye level. 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 solution and light from the bulb projecting through the solution. With the room lights darkened, view the light using the Flinn C-Spectra®. The top spectrum

viewed will be a continuous spectrum of the white lightbulb. The bottom spectrum will be the absorption spectrum of the red solution. Use colored pencils to make a drawing in the data table of the absorption spectra you observed. 8. Repeat steps 6 and 7 using the green, blue, and yellow colored solutions. 9. Cleanup and Disposal Turn off the light socket and spectrum tube power supplies.Wait several minutes to allow the incandescent lightbulb and the spectrum tubes to cool. Follow your teacher’s instructions on how to dispose of the liquids and how to store the lightbulb and spectrum tubes.


CHEMLAB

Class

Drawings of Emission Spectra Hydrogen

Neon

Mercury

4



Name

CHEMLAB

Date

Class

5 Drawings of Absorption Spectra

Red Green Blue Yellow


Analyze and Conclude 1.

Think Critically How can the single electron in a hydrogen atom produce all of the lines found 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 the elements you observed, and research its absorption spectrum. Compare your findings with the results of your experiment.

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



5

6


4

AM

6

2

3

2

10

10

10

6

Energy increases

12

14

3 10

8

10

16

10

18

10

12

10

20

10

22

3 10 Gamma rays

3 10

X rays

3 10

Ultraviolet

3 10

14

The Electromagnetic Spectrum

Electromagnetic Spectrum

10

4

Infrared

3 10

Microwaves

3 10

Date

TEACHING TRANSPARENCY MASTER

Teaching Transparency Masters


8

TV, FM

Radio

3 10

10 10 10 Frequency () in hertz

4

3 10

Wavelengths () in meters

Visible light

Name Class

15

Use with Chapter 5, Section 5.1

Name

Date

Class

15

TEACHING TRANSPARENCY WORKSHEET

The Electromagnetic Spectrum


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 relationship

between frequency and wavelength?

8. What is the wavelength of a radio station emitting its signal at 95.5 MHz? Estimate your

answer to the nearest power of ten.

Teaching Transparency Worksheets


7

Name

Date

Class

16


Atomic Orbitals


z z x x y

y

1s orbital

z

z

x

x

y

x y

y

px

py

pz

p orbitals

z z

z

z

z y

dxy

y

y

y

x

x

x

dxz

dyz

x y

x

dx 2y 2

dz2

d orbitals

8




z

2s orbital

Name

Date

Class

16


Atomic Orbitals


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 in

which 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 the

d 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 an

atom’s nucleus? What part of the diagram supports your conclusion?



9

Name

Date

Class

17


1s

2p 2s

3s

Increasing Energy Orbital filling sequence 10




2s

1s

2p 3s

4s 3p

4s

4p

3d

5s

4p

3p

3d

4d 5p

5d 6s

5d 4f 6s 5p 4d 5s

6p

7s

6p

6d 7p 5f 7s

7p

6d


4f

5f

Orbital Filling Sequence and Energy Levels

Name

Date

Class

17


Orbital Filling Sequence and Energy Levels


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.



11

Name

Date

Class

5

MATH SKILLS TRANSPARENCY MASTER

Interpreting Waves


A

amplitude

amplitude

Waves A and B are both electromagnetic waves. c for all electromagnetic waves.

12


Math Skills Transparency Masters


B

Name

Date

Class

5

MATH SKILLS TRANSPARENCY WORKSHEET

Interpreting Waves


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 a

longer 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 s1, what is its wavelength in nanometers?

Show your work.

Math Skills Transparency Worksheets


13

Name

Date

CHAPTER

Class

5

STUDY GUIDE

Electrons in Atoms Section 5.1

Light and Quantized Energy

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

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

energy

frequency

hertz

light

wave

wavelength

speed

Electromagnetic radiation is a kind of (1)

that behaves like a(n)

as it travels through space. (3)

(2)

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

(4)

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 Copyright © Glencoe/McGraw-Hill, a division of The McGraw-Hill Companies, Inc.

waves that pass a given point in one second. The SI unit for frequency is the , which is equivalent to one wave per second.

(8)

Use the figure to answer the following questions. A D

B

Origin

C

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 of

the wave?

14


Study Guide

Name

CHAPTER

Date

Class

5

STUDY GUIDE

Section 5.1 continued 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 energy

only in a. units of hertz.

c. entire wavelengths.

1 1 2 4 14. The is the phenomenon in which electrons are emitted from a metal’s surface when light of a certain frequency shines on it. b. whole-number multiples of h.

d. multiples of h, h, and so on.

a. quantum

c. photon effect

b. Planck concept

d. photoelectric effect

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

1 2

a. Ephoton h Planck’s constant

c. Ephoton h

b. Ephoton h

d. c

In your textbook, read about atomic emission spectra.

For each statement below, write true or false. Copyright © Glencoe/McGraw-Hill, a division of The McGraw-Hill Companies, Inc.

16. Like the visible spectrum, an atomic emission spectrum is a continuous

range 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 emission

spectrum 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 as

they absorb energy. 22. When an atom emits light, photons having certain specific energies are

being emitted.

Study Guide


15

Name

Date

5

CHAPTER

Section 5.2

Class

STUDY GUIDE

Quantum Theory and the Atom

In your textbook, read about the Bohr model of the atom.

Use each of the terms below to complete the statements. atomic emission spectrum

electron

frequencies

higher

energy levels

lower

ground state

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.

the difference between the


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

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.

h

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

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

16


Study Guide

Name

Date

CHAPTER

Class

5

STUDY GUIDE

Section 5.2 continued 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

a. Heisenberg uncertainty

principle

as waves 11. States that it is impossible to know both the velocity

and the position of a particle at the same time

b. Schrödinger wave equation c. quantum mechanical model

of the atom

12. A three-dimensional region around the nucleus

representing the probability of finding an electron

d. atomic orbital

13. Originally applied to the hydrogen atom, it led to the

quantum 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 how


they 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 each

principal 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 and

2p) orbitals. 22. Hydrogen’s principal energy level 3 consists of (nine, three)

orbitals.

Study Guide


17

Name

Date

CHAPTER

Section 5.3

Class

5

STUDY GUIDE

Electron Configuration

In your textbook, read about ground-state electron configurations.

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

electron configuration

ground-state electron configuration

Hund’s rule

lowest

Pauli exclusion principle

spins

stable

The arrangement of electrons in an atom is called the atom’s . Electrons in an atom tend to assume the arrangement

(1)

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 states that each electron occupies the lowest energy

(5)

orbital available. The (6)

states that a maximum of two

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

Atomic Number

Orbitals 1s

2s

2px

Electron Configuration 2py

2pz 1s2

9. Helium 10.

7

11. Neon

18


)( )( )( )( )(

Study Guide


. (8)

(7)

Name

CHAPTER

Date

Class

5

STUDY GUIDE

Section 5.3 continued Answer the following questions. 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-gas

notation.

In your textbook, read about valence electrons.

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

electron-dot structure? a. • Be •

• b. •B•

•• c. B••

•• d. Be

19. Given beryllium’s electron configuration of 1s22s2, which of the following represents its

electron-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

Study Guide


19

Name

Date

Class

CHAPTER 5

Student Recording Sheet

Assessment

Standardized Test Practice Multiple Choice

Select the best answer from the choices given, and fill in the corresponding circle. 1.

4.

7.

2.

5.

8.

3.

6.

9.

10.

Short Answer

Answer each question with complete sentences. 11.

12. 13. 14. SAT Subject Test: Chemistry 17.

16.

18.

19.


15.

26


Table of Contents

Reproducible Pages

Chapter 6 The Periodic Table and Periodic Law MiniLab . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 28 ChemLab . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 29 Teaching Transparency Masters and Worksheets . . . . . . . . . . . . . . . . . . . . . . . . . . 32 Math Skills Transparency Masters and Worksheets . . . . . . . . . . . . . . . . . . . . . . . . .

40

Study Guide . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 42 Chapter Assessment . . . . . . . . . . . . . . . . . . . . . . . . . . . . 46



27

Name

Date

Class

mini LAB 6

Organize Elements Can you find the pattern? Materials Index cards, pencil Procedure 1. Read and complete the lab safety form. 2. Make a set of element cards based on the information in the chart at right. 3. Organize the cards by increasing mass, and start placing them into a 4 3 grid. 4. Place each card based on its properties, and leave gaps when necessary.

Symbol

Mass (g)

State

Color

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

Analysis

3. Describe the period and group trends for the mass in your new table. Explain your placement of any 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 an expected range for the mass of Ph?

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

28




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.

Name

Date

CHEMLAB

Class

6

Investigate Descriptive Chemistry

Y

ou can observe several of the representative elements, classify them, and compare their properties. The observation of the properties of elements is called descriptive chemistry.

Problem

Objectives

Materials

What is the pattern of properties of the representative elements?

• Observe properties of various elements. • Classify elements as metals, nonmetals, and metalloids. • Examine general trends within the periodic table.

stoppered test tubes containing small samples of elements plastic dishes containing samples of elements conductivity apparatus

1.0M HCl test tubes (6) test-tube rack 10-mL graduated cylinder spatula small hammer glass marking pencil

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.

Pre-Lab Read the entire CHEMLAB. 2. Use the data table on the next page to record the observations you make during the lab. 3. Examine the periodic table. What is the physical state of most metals? Nonmetals? Metalloids? 1.


4.

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


29

Name

Date

6

Procedure 1. 2.

3.

4.

5.

6.

Read and complete the lab safety form. Observe and record the appearance (physical state, color, luster, texture, and so on) of the element sample in each test tube without removing the stoppers. Remove a small sample of each of the elements contained in a plastic dish and place it on a hard surface. Gently tap each element sample with a small hammer. If the element is malleable, it will flatten. If it is brittle, it will shatter. Record your observations. Use the conductivity tester to determine which elements conduct electricity. Clean the electrodes with water, and dry them before testing each element. Label each test tube with the symbol for one of the elements in the plastic dishes. Using a graduated cylinder, add 5 mL of water to each test tube. Use a spatula to put a small amount of each element into the corresponding test tubes. Using a graduated cylinder, add 5 mL of 1.0M HCl to

each test tube. Observe each tube for at least 1 minute. The formation of bubbles is evidence of a reaction between the acid and the element. Record your observations. 7. Cleanup and Disposal Dispose of all materials as instructed by your teacher.

Observation of Elements 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 metals and nonmetals


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

2.

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

3.

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

30




CHEMLAB

Class

Name

CHEMLAB

Date

Class

6

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 elements observed 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?



31

32


7

6

5

4

3

2

1

5

7

8 10

Cobalt

11

12

38

Sr

37

Rb

40

Zirconium

26

(266)

Sg

106

Seaborgium

(264)

Bh

107

Bohrium

Zinc

94

U 238.029

231.036

232.038 (244)

(243)

Am

95

Americium

(247)

Cm

96

Curium

157.25

Gd

64

Gadolinium

(247)

Bk

97

Berkelium

158.925

Tb

65

Terbium

51

Sb 121.757

Sn 118.710

(284)

99

98 (251)

(252)

Es

Einsteinium

Californium

Cf

164.930

Ho

67 162.50

Dy

66

Dysprosium

Holmium

* Uut

(285)

Ununtrium 113

Ununbium

* Uub

204.383

200.59

81

Thallium

114.82

In

208.980

Bi

83

Bismuth

(257)

Fm

100

Fermium

167.259

Er

68

Erbium

(289)

* Uuq

(258)

Md

101

Mendelevium

168.934

Tm

69

Thulium

(288)

* Uup

Ununquadium Ununpentium 114 115

207.2

Pb

82

Lead

Antimony

50

49

Indium

Tin

33

As 74.922

32

Ge 72.61

31

Ga

Tl

112

30.974 Arsenic

69.723

Hg

80

Mercury

112.411

Cd

48

Cadmium

28.086 Germanium

(259)

No

102

Nobelium

173.04

Yb

70

Ytterbium

(291)

* Uuh

116

Ununhexium

208.982

Po

84

Polonium

127.60

Te

52

Tellurium

78.96

Se

34

Selenium

32.066

S

16

(262)

Lr

103

Lawrencium

174.967

Lu

71

Lutetium

209.987

At

85

Astatine

126.904

I

53

Iodine

79.904

Br

35

Bromine

35.453

Cl

17

Chlorine

18.998

F

9

Fluorine

17

(294)

* Uuo

Ununoctium 118

222.018

Rn

86

Radon

131.290

Xe

54

Xenon

83.80

Kr

36

Krypton

39.948

Ar

18

Argon

20.180

Ne

10

Neon

4.003

He

2

Helium




(237)

Pu

Plutonium

93

92

Np

Neptunium

Uranium

91

Pa

90

Protactinium

Eu 151.965

Sm 150.36

63

Europium

62

Samarium

*

(272)

Rg

Ds (281)

Mt (268)

Hs

30

Zn 65.39

26.982 Gallium

P

15

Sulfur

15.999

O

8

Oxygen

16

The Periodic Table

Th

Thorium

Pt

(277)

108

(145)

Pm

Nd 144.242

Pr 140.908

61

Promethium

60

Ce

59

Praseodymium Neodymium

140.115

58

Cerium

79

Au

78 196.967

Gold

195.08

107.868

106.42

Ag

47

Silver

Platinum

Pd

46

Palladium

Meitnerium Darmstadtium Roentgenium 111 109 110

192.217

Ir

77

Iridium

102.906

Rh

45

Rhodium

63.546

Cu

29

Ni

28 58.693

Hassium

190.23

Os

76

27

Co 58.933

Copper

Nickel

Si

14

Phosphorus

14.007

N

7

Nitrogen

15

18

Date

Actinide series

Lanthanide series

Re 186.207

W 183.84

75

74

Osmium

101.07

(98) Rhenium

Ru

Tc

95.94

44

Ruthenium

43

Technetium

Tungsten

Mo

42

Molybdenum

The number in parentheses is the mass number of the longest lived isotope for that element.

Db (262)

Rf (261)

Ac

Ra

(226)

Fr

(223)

(227)

105

104

Dubnium

Rutherfordium

89

87

Actinium

180.948

Ta

73

Tantalum

92.906

Nb

41

Niobium

55.847

Fe

25 54.938

Mn

24

Cr

V 51.996

23 50.942

Iron

Manganese

Chromium

Vanadium

88

Radium

Francium

Hf

178.49

La

138.905

Ba

137.327

72

Cs

57

56

55

Hafnium

132.905

88.906

Lanthanum

87.62

Barium

85.468

Cesium

91.224

39

Strontium

Rubidium

22

Ti

47.867

Zr

Yttrium

Ca

Titanium

Y

21

Sc

44.956

20

40.078

K

19

39.098

Scandium

24.305

Calcium

22.990

Potassium

Al

Mg

9

Na

6

13

4

12

11

Silicon

12.011

10.811 Aluminum

9.012

Magnesium

6.941

B

Sodium

6

C

5

Carbon

14

Boron

Recently observed

13

4

Synthetic

Nonmetal

Metalloid

Solid

Liquid

Be

1.008

State of matter

Beryllium

3

1

H

3

Atomic mass

2

Hydrogen

Li

2

Symbol

Element Atomic number

Metal

Gas

PERIODIC TABLE OF THE ELEMENTS

Lithium

1.008

H

1

Hydrogen

1

Name Class

18


Name

Date

Class

18


The Periodic Table


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? 17. What is the name of the group 17 elements? 18. What is the name of the group 18 elements? 19. What can be said about the electron configurations of all the elements in a group?



33

34


57 La

38 Sr

56 Ba

Rb

Cs

Fr

87

55

37

104 Rf

72 Hf

40 Zr

22 Ti

91

Pa

90

Th

60

U

92

Nd

59

Pr

58

Ce

75 Re

43 Tc

25 Mn

107 Bh

74 W

42 Mo

24 Cr

106 Sg

105 Db

73 Ta

41 Nb

23 V

Np

93

Pm

61

108 Hs

76 Os

44 Ru

26 Fe

d block

Pu

94

Sm

62

109 Mt

77 Ir

45 Rh

27 Co

Am

95

Eu

63

110 Uun

78 Pt

46 Pd

28 Ni

112 Uub

80 Hg

48 Cd

30 Zn

Cm

96

Gd

64

Bk

97

Tb

65

f block

111 Uuu

79 Au

47 Ag

29 Cu

Cf

98

Dy

66

81 Tl

49 In

31 Ga

13 Al

5 B

p1

Es

99

Ho

67

82 Pb

50 Sn

32 Ge

14 Si

6 C

p2

Fm

100

Er

68

83 Bi

51 Sb

33 As

15 P

7 N

Md

101

Tm

69

84 Po

52 Te

34 Se

16 S

8 O

p block p3 p4

No

102

Yb

70

85 At

53 I

35 Br

17 Cl

9 F

p5

103 Lr

71 Lu

86 Rn

54 Xe

36 Kr

18 Ar

10 Ne

p6

2 He

Date


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



89 Ac

39 Y

20 Ca

19

88 Ra

21 Sc

12 Mg

11 Na

K

4 Be

s2

3 Li

1 H

s block s1

s2

Name Class

19


Name

Date

Class

19


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


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 Chemistry: Matter and Change • Chapter 6

35

Name

Date

Class

20


Atomic and Ionic Radii


Chemical symbol

K

Charge

1

2

Li 76

Be 31

B

1

2

3

4

Na 102

Mg 72

Al 54

Si

1

2

3

4

Ca 100

Ga 62

1

2

3

Rb 152

Sr 118

13

Relative size

1

14

20

Ionic radius

138

15

16

17

N 146

O 140

F 133

3

2

1

P 212

S 184

Cl 181

3

2

1

Ge 53

As 222

Se 198

Br 195

4

3

2

1

Sn 71

Sb 62

Te 221

4

5

2

Pb 84

Bi

4

5

C

15

2

41

K

138

4

5

6

n Transition metals

1

2

Cs 167

Ba 135

Tl

1

2

3

81

3 95

220 1

74

Radii are given in picometers (1 1012 m)

36




Period

3

Name

Date

Class

20


Atomic and Ionic Radii


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 have

these 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?



37

Name

Date

Class

21


First Ionization and Successive Ionization Energies


First Ionization Energy of Elements in Periods 1–5 Period 2 Period 3

First ionization energy (kJ/mol)

2500

Period 4

Period 5

He Ne

2000

Ar

1500

Kr H

Xe

1000

500

Li

Na

Rb

K

0 10

20

30

40

50

60

Atomic number

Successive Ionization Energies for the Period 2 Elements Ionization energy (kJ/mol)*

Element

Valence electrons

Li

1

520

7300

Be

2

900

1760

14,850

B

3

800

2430

3660

25,020

C

4

1090

2350

4620

6220

37,830

N

5

1400

2860

4580

7480

9440

53,270

O

6

1310

3390

5300

7470

10,980

13,330

71,330

F

7

1680

3370

6050

8410

11,020

15,160

17,870

92,040

Ne

8

2080

3950

6120

9370

12,180

15,240

20,000

23,070 115,380

1st

2nd

3rd

4th

5th

6th

7th

8th

9th

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

38




0

Name

Date

Class

21


First Ionization and Successive Ionization Energies


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 as

the 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 first

ionization 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 as

the atomic number of the elements increases?

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



39

Name

Date


6 Use with Chapter 6, Section 6.2


Using the Periodic Table

Class

40



Name

Date

Class

6


Using the Periodic Table


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) for

each 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 Math Skills Transparency Worksheets


41

Name

Date

CHAPTER

6

Class

STUDY GUIDE

The Periodic Table and Periodic Law Section 6.1

Development of the Modern Periodic Table

In your textbook, reads about the history of the periodic table’s development.

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

atomic mass

atomic number

nine

elements

properties

Henry Moseley

eight

protons

periodic law

Dmitri Mendeleev

accepted

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) 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

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 could be predicted.

(7)

The element in the modern periodic table are arranged according to increasing (8)

, as a result of the work of (9)

arrangement is based on number of (10)

. This 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)

42

.


Study Guide

Name

Date

CHAPTER

Class

6

STUDY GUIDE

Section 6.1 continued 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 the table on the right. 7

Atomic Mass

13.

N

Atomic Number

14.

Electron Configuration

15.

Chemical Name

16.

Chemical Symbol

17.

Nitrogen 14.007 [He]2s22p3

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


Column A

Column B

18. A column on the periodic table

a. metals

19. A row on the periodic table

b. group

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

c. period

21. Elements that are shiny and conduct electricity

d. representative elements

22. Elements in groups 3 to 12

e. transition elements

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 are

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

metals. 26. Group 13 elements are the alkaline earth metals. 27. Group 17 elements are highly reactive nonmetals known as

halogens. 28. Group 18 elements are very unreactive elements known as

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

metals.

Study Guide


43

Name

Date

CHAPTER

Section 6.2

6

Class

STUDY GUIDE

Classification of the Elements

In 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 in Column A with the element in Column B that has the most similar chemical properties. Column A

Column B

1. arsenic (As)

a. boron (B)

2. bromine (Br)

b. cesium (Cs)

3. cadmium (Cd)

c. chromium (Cr)

4. gallium (Ga)

d. cobalt (Co)

5. germanium (Ge)

e. hafnium (Hf)

6. iridium (Ir)

f.

7. magnesium (Mg)

g. iron (Fe)

8. neon (Ne)

h. nitrogen (N)

9. nickel (Ni)

i.

platinum (Pt)

10. osmium (Os)

j.

scandium (Sc)

11. sodium (Na)

k. silicon (Si)

12. tellurium (Te)

l.

13. tungsten (W)

m. sulfur (S)

14. yttrium (Y)

n. zinc (Z)

15. zirconium (Zr)

o. xenon (Xe)


iodine (I)

strontium (Sr)

Answer the following questions. 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.

44


Study Guide

Name

Date

CHAPTER

Class

6

STUDY GUIDE

Section 6.2 continued 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 to answer the following questions. s2 2 He

s block s1 1 H

s2

p1

p2

p block p3 p4

p5

p6

3 Li

4 Be

5 B

6 C

7 N

8 O

9 F

10 Ne

11 Na

12 Mg

13 Al

14 Si

15 P

16 S

17 Cl

18 Ar

19

20 Ca

21 Sc

22 Ti

23 V

24 Cr

25 Mn

26 Fe

27 Co

28 Ni

29 Cu

30 Zn

31 Ga

32 Ge

33 As

34 Se

35 Br

36 Kr

38 Sr

39 Y

40 Zr

41 Nb

42 Mo

43 Tc

44 Ru

45 Rh

46 Pd

47 Ag

48 Cd

49 In

50 Sn

51 Sb

52 Te

53 I

54 Xe

56 Ba

57

La

72 Hf

73 Ta

74 W

75 Re

76 Os

77 Ir

78 Pt

79 Au

80 Hg

81 Tl

82 Pb

83 Bi

84 Po

85 At

86 Rn

88 Ra

Ac

104 Rf

105 Db

106 Sg

107 Bh

108 Hs

109 Mt

110 Uun

111 Uuv

112 Uub

71 Lu

K 37

Rb 55

Cs 87

Fr

d block

89


f block 58

59

60

61

62

63

64

65

66

67

68

69

70

Ce

Pr

Nd

Pm

Sm

Eu

Gd

Tb

Dy

Ho

Er

Tm

Yb

90

91

92

93

94

95

96

97

98

99

100

101

102

Th

Pa

U

Np

Pu

Am

Cm

Bk

Cf

Es

Fm

Md

No

103 Lr

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?

Study Guide


45

Name

Date

6

CHAPTER

Section 6.3

Class

STUDY GUIDE

Periodic Trends

In your textbook, read about atomic radius and ionic radius.

Circle the letter of the choice that best completes the statement or answers the question. 1. Atomic radii cannot be measured directly because the electron cloud surrounding the

nucleus 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 periodic

table? c.

Generally increase

d.

Generally decrease

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 negative

charge. a. halogen

b. ion

c. isotope

d. molecule

c. losing an electron.

d. losing a neutron.

5. An atom becomes negatively charged by a. gaining an electron.

b. gaining a proton.

6. Which diagram best represents the relationship between the diameter of a sodium atom

and the diameter of a positive sodium ion? a.

b.

Na

46

Na

c.

Na

Na


Na

Na

Study Guide


Generally decrease

Generally decrease

b.

Generally decrease

Generally increase

Generally increase

Generally increase

a.

Name

CHAPTER

Date

6

Class

STUDY GUIDE

Section 6.3 continued In your textbook, read about ionization energy and electronegativity.

Answer the following questions. 7. What is ionization energy?

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

ion.

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?

Study Guide


47

Name

Date

Class

CHAPTER 6


Assessment



4.

7.

2.

5.

8.

3.

6.

9.

10.

Short Answer

Answer each question with complete sentences. 11. 12. 13.

Extended Response

Answer each question with complete sentences.

15.

SAT Subject Test: Chemistry 16. 17. 18. 19.

54



14.

Table of Contents

Reproducible Pages

Chapter 7 Ionic Compounds and Metals MiniLab . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 56 ChemLab . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 57 Teaching Transparency Masters and Worksheets . . . . . . . . . . . . . . . . . . . . . . . . . . 60 Math Skills Transparency Masters and Worksheets . . . . . . . . . . . . . . . . . . . . . . . . . . 68 Study Guide . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 72 Chapter Assessment. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 78



55

Name

Date

Class

mini LAB 7

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

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

1. Read and complete the lab safety form. 2. Examine a property of spring steel by trying to bend open one of three hairpins. Record your observations. 3. Next hold each end of the hairpin with a pair of forceps. Place the curved central loop portion of the hairpin in the top of the blue flame from a laboratory burner. When the metal turns red, pull the hairpin open to form a straight piece of metal. Allow it to cool as you record your observations. Repeat Step 3 for the remaining two hairpins. WARNING: Do not touch the hot metal. Do not hold your hand above the flame of the laboratory burner. 4. To make softened steel, use a pair of forceps to hold all three hairpins vertically in the flame of the laboratory burner until the hairpins glow red all over. Slowly raise the three hairpins straight up and out of the flame so they cool slowly. Slow cooling results in the formation of large crystals.

5. After cooling, bend each of the three hairpins into the shape of the letter J. Record how the metal feels as you bend it. 6. To harden the steel, use tongs to hold two of the bent hairpins in the flame until they are glowing red all over. Quickly plunge the hot metals into a 250-mL beaker containing approximately 200 mL of cold water. Quick-cooling causes 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 hold the remaining hardened metal bend above the flame for a brief period of time. Slowly move the metal back and forth just above the flame until the gray metal turns to an iridescent blue-gray color. Do not allow the metal to become hot enough to glow red. Slowly cool the metal and then try to unbend it using the end of your finger. Record your observations.

Analysis 1. Analyze your results, and identify the two types of steel that appear to have their properties combined 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 be wear-resistant and retain a sharpened edge? Explain your reasoning.

56




Procedure

Name

Date

CHEMLAB

Class

7

Synthesize an Ionic Compound

E

lements combine to form compounds. If energy is released as the compound is formed, the resulting product is more stable than the reacting elements. In this investigation, you will react elements to form two 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 the products you formed are ionic compounds.

Problem

Objectives

Materials

Can the physical properties of a compound indicate that they have ionic bonds?

• Observe evidence of a chemical reaction. • Acquire and analyze information that will enable you to decide if a compound has an ionic bond. • Classify the products as ionic or not ionic.

magnesium ribbon (25 cm) crucible ring stand and ring clay triangle Bunsen burner stirring rod

crucible tongs centigram balance 100-mL beaker distilled water conductivity tester


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.

Pre-Lab 1.

Read the entire CHEMLAB. Identify the variable. List any conditions that must be kept constant.

b. Write the electron configuration of the

magnesium ion.

c. The magnesium ion has an electron

configuration like that of which noble gas? 2.

Write the electron configuration of the magnesium atom.

3.

Repeat question 2 for oxygen and nitrogen.

a. Based on this configuration, will magnesium

lose or gain electrons to become a magnesium ion?



57

Name

Date

CHEMLAB

7

Use the data table in the next column. 5. In your data table, which mass values will be measured directly? Which mass values will be calculated?

4.

6.

Class

Explain what must be done to calculate each mass value that is not measured directly.

4. Measure

the mass of the clean, dry crucible. 5. Roll 25 cm of magnesium ribbon into a loose ball. Place it in the crucible. Measure the mass of the magnesium and crucible together. 6. Place the crucible on the triangle, and heat it with a hot flame (flame tip should be near the crucible). 7. Turn off the burner as soon as the magnesium ignites and begins to burn with a bright white light. Allow it to cool, and measure the mass of the magnesium product and the crucible. 8. Place the dry, solid product in the beaker. 9. Add 10 mL of distilled water to the beaker, and stir. Check the mixture with a conductivity tester. 10. Cleanup and Disposal Dispose of the product as directed by your teacher. Wash out the crucible with water. Return all lab equipment to its proper place.

Mass Data Material(s)

Mass (g)

Procedure Read and complete the lab safety form. 2. Record all measurements in your data table. 3. Position the ring on the ring stand about 7 cm above the top of the Bunsen burner. Place the clay triangle on the ring. 1.

Crucible and Mg ribbon before heating Magnesium ribbon Crucible and magnesium products after heating Magnesium products


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?

58




Empty crucible

Name


CHEMLAB

Date

Class

7

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 of the 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 these results verify that the compounds are ionic?

7.

Error Analysis If the results show that the magnesium lost mass instead of gaining mass, cite possible sources 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 to find out.



59

60


O ion (O2)

Energy 8 protons (8)

+

10 electrons (10)

Formation of Ions

8 protons (8 )

Ca ion (Ca2)

20 protons (20)

2e Two electrons

+

18 electrons (18)

Date




O atom 1s2 2s2 2p4

2e Two electrons

+

8 electrons (8)

Ca atom 1s2 2s2 2p6 3s2 3p6 4s2

Energy

20 protons (20)

+

20 electrons (20)

Name Class

22


Name

Date

Class

22


Formation of Ions


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 the

Ca atom shown?

4. What kind of charge does the Ca atom take on as a result of the reaction? What is the

name 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 the

name 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? How

can you tell?

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

can 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?



61

Name

Date

Class

23


Ionic Bonds

+

62

+

P + Mg Teaching Transparency Masters


?


Mg

e

S

Cl Mg

+

e

+

Cl

?

+ Mg

2

Cl

Cl


Name

Date

Class

23


Ionic Bonds


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 such

an 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 formula

unit of magnesium chloride? Why? What is the formula of magnesium chloride?


6. What kind of compound is magnesium chloride? What happens to electrons during the

formation 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.



63

Date

Class


3s 3p 3s 3p 3s 3p 3s 3p 3s 3p 3s 3p


3s 3p 3s 3p

Periodic Table Row 3

Na

Mg

Al

Si

P

S

Cl

Ar

Formulas for Ionic Compounds

24


Name

64



Name

Date

Class

24


Formulas for Ionic Compounds


1. How many valence electrons are there in an atom of sodium (Na)? What would a sodium

atom 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 reacts

with 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 tend

to 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 reacts

with 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 chlorine

atom 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 aluminum

reacts 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 aluminum

reacts with sulfur? What are the oxidation numbers of each of the ions present?



65

66


2 2 2 2 2 2 2 2 2 2 2

Group 2 Metal Atoms

Date


Metallic Bonding


Group 1 Metal Atoms

Name Class

25



Name

Date

Class

25


Metallic Bonding


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 word

is 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 compare

to 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 compare

to 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 bonding

account for that property?

10. Do metals tend to be brittle, or are they malleable and ductile? How does the model of

metallic bonding account for that property?



67

Name

Date

Class

7


14 Si


15 P

16 S

17 Cl

18 Ar


12 Mg 11 Na

2

3

13 Al

10 Ne 8 O 4 Be 3 Li

Groups 1 2

5 B

13

6 C

14

7 N

Groups 15 16

9 F

17

18

Calculating Numbers of Electrons and Predicting Ionic Change

Periods 68



Name

Date

Class

7


Calculating Numbers of Electrons and Predicting Ionic Change


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)



69

Name

Date

Class

8


Determining Numbers of Ions


Common Polyatomic Ions Name

Ion

Name

NH4

ammonium

IO4

periodate

NO2

nitrite

C2H3O2

acetate

NO3

nitrate

H2PO4

dihydrogen phosphate

HSO4

hydrogen sulfate

CO32

carbonate

OH

hydroxide

SO32

sulfite

CN

cyanide

SO42

sulfate

MnO4

permanganate

S2O32

thiosulfate

HCO3

hydrogen carbonate

O22

peroxide

ClO

hypochlorite

CrO42

chromate

ClO2

chlorite

Cr2O72

dichromate

ClO3

chlorate

HPO42

hydrogen phosphate

ClO4

perchlorate

PO43

phosphate

BrO3

bromate

AsO43

arsenate

IO3

iodate

70




Ion

Name

Date

Class

8


Determining Numbers of Ions


1. Write a simple word equation that illustrates what must be true of total positive charge

and total negative charge in an ionic compound.

2. How many potassium ions (group 1) would be related to balance the charge of each of

the 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 in

a 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



71

Name

Date

7

CHAPTER

Class

STUDY GUIDE

Ionic Compounds and Metals Section 7.1

Ion Formation

In your textbook, read about chemical bonds and formation of ions.

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

electrons

energy level

ions

nucleus

octet

pseudo-noble gas formations

noble gases valence

.

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 electrons, which are the electrons in the outermost

(5)

. The (7)

(6)

are a family of elements that

electrons, which is called a stable (8)


have very little tendency to react. Most of these elements have a set of eight outermost . The relatively stable electron

structures developed by loss of electrons in certain elements of groups 3, 4, 13, and 14 are called (9)

.

For each statement below, write true or false. 10. A positively charged ion is called an anion. 11. Elements in group 1 lose their one valence electron, forming an ion with a

1 charge. 12. Elements tend to react so that they acquire the electron structure of a

halogen. 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 electrons

and becoming anions. 72


Study Guide

Name

Date

7

CHAPTER

Section 7.2

Class

STUDY GUIDE

Ionic Bonds and Ionic Compounds

In your textbook, read about forming ionic bonds and the characteristics of ionic compounds.

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

Study Guide


73

Name

Date

CHAPTER

Class

7

STUDY GUIDE

Section 7.2 continued 9. What is the relationship between lattice energy and the strength of the attractive force

holding 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.

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


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

For each statement below, write true or false. 19. The crystal lattice of ionic compounds affects their melting and boiling

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

compound. 21. The energy of an ionic compound is higher than that of the separate

elements that formed it. 22. Large ions tend to produce a more negative value for lattice energy than

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

energy than ions with smaller charges do. 74


Study Guide

Name

Date

CHAPTER

Section 7.3

Class

7

STUDY GUIDE

Names and Formulas for Ionic Compounds

In your textbook, read about communicating what is in a compound and naming ions and ionic compounds.

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

-ate

cation

electrons

zero

lower right

monatomic

one

oxidation number

-ite

oxyanion

polyatomic

subscript

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

(5)

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) Copyright © Glencoe/McGraw-Hill, a division of The McGraw-Hill Companies, Inc.

ends with (8)

. The name for the ion with fewer oxygen atoms .

In the chemical formula for any ionic compound, the chemical symbol for the is written first, followed by the chemical symbol for the

(9)

. A(n) (11)

(10)

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

a. chlorate

15. ClO4

b. hypochlorite

16. ClO

c. chloride

17. Cl

d. perchlorate

18. ClO3

e. chlorite

Study Guide


75

Name

Date

CHAPTER

7

Class

STUDY GUIDE

Section 7.3 continued 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

76


Study Guide

Name

CHAPTER

Section 7.4

Date

7

Class

STUDY GUIDE

Metallic Bonds and the Properties of Metals

In 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 is not. If the property is a characteristic of metals, explain how metallic bonding accounts for 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 Study Guide


77

Name

Date

Class

CHAPTER 7


Assessment



4.

7.

2.

5.

8.

3.

6.

9.

Short Answer

Answer each question with complete sentences. 10. 11. 12. Extended Response


14.

SAT Subject Test: Chemistry 15.

17.

19.

16.

18.

20.

84



13.

Table of Contents

Reproducible Pages

Chapter 8 Covalent Bonding MiniLab . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 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

85

Name

Date

Class

mini LAB 8

Compare Melting Points How can you determine the relationship between bond type and melting point? The properties of a compound depend on whether the bonds in the compound are ionic or covalent.

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

1. Read and complete the lab safety form. 2. Create a data table for the experiment. 3. Using a permanent marker, draw three lines on the inside bottom of a disposable, 9inch 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 the labeled 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 the compounds 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 and the 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 solid as low, medium, high, or very high.

3. Infer Which compounds are bonded with ionic bonds? Which are bonded with covalent bonds?

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

86




Procedure

Name

CHEMLAB

Date

Class

8

Model Molecular Shapes

C

ovalent bonding occurs when atoms share valence electrons. In the Valence Shell Electron Pair Repulsion (VSEPR) theory, the way in which valence electrons of bonding atoms are positioned, is the basis for predicting a molecule’s shape. This method of visualizing shape is also based on the molecule’s Lewis structure.

Problem

Objectives

Materials

How do the Lewis structure and the positions of valence electrons affect the shape of the covalent compound?

• Predict the shape of molecular compounds • Construct molecular models • Evaluate the strength of covalent bonds • Identify resonance structures • Distinguish polar and nonpolar molecules

molecular-model kit

Pre-Lab do you need to determine whether or not a molecule is polar?

Read the entire CHEMLAB. 2. Review the VSEPR model. What do the initials VSEPR stand for? What is the basis for the predictions of molecular shape made by the VSEPR model?


1.

4.

3.

What is a resonance structure? If a molecule has resonance structures, what can you infer about the types of bonds in that molecule?

How do you determine whether or not a bond is a polar covalent bond? What additional information



87

Name

Date

CHEMLAB

8

Procedure 1. 2. 3.

4. 5. 6.

Class

Read and complete the lab safety form. Create a table to record your data. Note the color used to represent each of these atoms in the molecular modele kit: H, O, P, C, F, S, and N. Draw the Lewis structures of the H2, O2, and N2 molecules. Obtain two hydrogen atoms and one connector and assemble a hydrogen (H2) molecule. Obtain two oxygen atoms and two connectors and assemble an oxygen (O2) molecule. Observe that your model represents a double-bonded diatomic oxygen molecule.

Obtain two nitrogen atoms and three connectors and assemble a nitrogen (N2) molecule. Observe that your model represents a triple-bonded diatomic nitrogen molecule. 8. Recognize that these diatomic molecules are linear in shape because they have only two atoms. 9. Draw the Lewis structure of water (H2O) and construct its molecule. 10.Classify the shape of the H2O molecule using information in Table 8.6. 11. Repeat steps 9 and 10 for PH3, CF4, CO2, SO3, HCN, and CO molecules. 7.

1.

Think Critically Based on the molecular models you built and observed in this lab, rank single, 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 molecule that has three resonance structures, draw the structures, and explain why resonance occurs.

88




Analyze and Conclude

Name

CHEMLAB 4.

Date

Class

8

Recognize Cause and Effect Use the electronegativity difference to determine the polarity of the molecules in Steps 9-11. Based on their calculated bond polarities and the models 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 resonance structures by interchanging a lone pair for a covalent bond.



89

90


SiO2

N2H4

Formula

Step 5

Step 6

O Si O

2

Unnecessary (octets H N N H H N N H complete) H H H H

Step 4

Date


Lewis Structures



O Si O

7

Step Step 2 3

H N N H 14 HH

Step 1

Name Class

26


Name

Date

Class

26


Lewis Structures


1. Step 1 in drawing the Lewis structure for a molecule is to decide which atoms of the

molecule 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 valence

electrons in the atoms in the molecule. Explain why the total number of valence electrons in N2H4 is 14.

3. Step 3 in drawing a Lewis structure requires finding the number of bonding pairs. What

must be done to the result of step 2 to find the number of bonding pairs? Verify that this is 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 of


bonded atoms in N2H4. How many such bonding pairs are shown in step 4, and what symbol 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. Lone pairs are then placed around each terminal atom to achieve a full outer level, and any remaining pairs are assigned to the central atom(s). Explain the drawing that has resulted for N2H4.

6. In step 6, if any central atom drawn in step 5 does not have an octet, lone pairs from the

terminal 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?



91

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Date

Class

27


VSEPR Model and Molecular Shape

H Teaching Transparency Masters


H

H H

H

C

H B

H

H H Chemistry: Matter and Change • Chapter 8

H

H H

N

H N H H H H C H H H B 92

H

O

O H H

H

F

F H

H


Name

Date

Class

27


VSEPR Model and Molecular Shape


1. The shapes of the molecules shown have been determined by means of the VSEPR

model. 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?



93

94


Be 1.57

B 2.04

C 2.55

N 3.04

O 3.44

Date

Electronegativity and Polarity

At 2.2

I 2.66

Br 2.96

Cl 3.16

F 3.98

Group Group Group Group Group 3 4 5 6 7


Li Period 0.98 2

Group Group 1 2

Electronegativities

Name Class


28



Name

Date

Class

28




1. What is electronegativity?

2. The electronegativities of the elements of period 2 and of group 17 of the periodic table are

shown. 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? (Assume

that a bond is nonpolar covalent if the electronegativity difference is 0, polar covalent if the difference is greater than 0 but not more than 1.70, and ionic if the difference is more than 1.70.)

5. Given your answer to question 4 and your knowledge of molecular shapes, is a carbon Copyright © Glencoe/McGraw-Hill, a division of The McGraw-Hill Companies, Inc.

tetrachloride (CCl4) molecule polar or nonpolar? Explain.

6. What kind of bond exists between a nitrogen (N) atom and a fluorine (F) atom? Is a

nitrogen trifluoride (NF3) molecule polar or nonpolar? Explain.

7. What kind of bond exists between a beryllium (Be) atom and a bromine (Br) atom? Is a

beryllium 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 boron

triiodide (BI3) molecule polar or nonpolar? Explain.



95

96

I


(a)

I

N

C

N

(c)

N O

O

6 7 8 9 10

F

F

(d)

F

Si

F

F

F

F

hexaheptaoctanonadeca-

Prefix

(e)

O F

Determining the Names of Binary Compounds and Their Numbers of Atoms

(b)

S

O

O

(never for first element in the formula)

Number of Atoms

Date




S

monoditritetrapenta-

1 2 3 4 5

I

Prefix

Number of Atoms

Name Class

9


Name

Date

Class

9


Determining the Names of Binary Compounds and Their Numbers of Atoms


1. The table shows the prefixes used in naming binary compounds. Notice that each prefix

corresponds to a certain number of atoms. These prefixes are also used in naming many geometric figures with which you probably are familiar. For example, a triangle has three sides. 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, tell

how many atoms of each kind are present. Then write the name of the compound, using the 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 a

molecule of the compound and the total number of atoms per molecule. Then name the compound. a. P4O10

b. S2O7

c. Si3H8

4. What is the formula for each compound listed below? a. tetrasulfur dinitride b. dichlorine monoxide



97


0

2.0

2.19

P 2.58

S

3.0

Polar covalent

1.90

Si 3.16

Cl

Determining Electronegativity Difference and Percent Ionic Character

Electronegativity difference

1.0

1.61

Al

Covalent

1.31

Mg

Date




0

25

50

75

Ionic

0.93

Electronegativity

100

Na

Percent ionic character

98

Period 3 Element

Name Class

10


Name

Date

Class

10


Determining Electronegativity Difference and Percent Ionic Character


1. The electronegativities of the period-3 elements are listed on the transparency. Calculate

the 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 is

considered ionic. For which of the bonded-atoms combinations from question 1 would the bonds be considered ionic according to this rule?

3. Bonds are rarely completely covalent or completely ionic. Rather, they have a percent


ionic character. The graph illustrates that fact. What two variables are plotted on the graph, 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 bonded

atoms listed in question 1. a. Na and Cl

d. Si and Cl

b. Mg and S

e. Si and S

c. Al and P 5. Given what you know about percentages in general, what must be the relationship

between 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. a. Na and Cl

d. Si and Cl

b. Mg and S

e. Si and S

c. Al and P



99

Name

Date

CHAPTER

8

Class

STUDY GUIDE

Covalent Bonding Section 8.1

The Covalent Bond

In your textbook, read about the nature of covalent bonds.

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

molecule

sigma bond

exothermic

pi bond

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

(3)

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. Copyright © Glencoe/McGraw-Hill, a division of The McGraw-Hill Companies, Inc.

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

100


Study Guide

Name

Date

CHAPTER

Section 8.2

Class

8

STUDY GUIDE

Naming Molecules

In your textbook, read about how binary compounds and acids are named from their formulas.

For each statement below, write true or false. 1. Binary molecular compounds are generally composed of a metal and a

nonmetal. 2. The second element in the formula of a binary compound is named using

the 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 the

prefix hydro-. 7. An oxyacid contains only two elements. 8. If the name of the anion of an oxyacid ends in -ate, the acid name

contains the suffix -ous.


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

For each item in Column A, write the letter of the matching item in Column B. Column A 9. CO

Column B a. hydrobromic acid

10. CO2

b. dinitrogen tetroxide

11. H2CO3

c. carbon monoxide

12. NH3

d. nitrous acid

13. N2O4

e. ammonia

14. HNO2

f.

15. HNO3

g. carbonic acid

16. HBr

h. bromic acid

17. HBrO3

i.

Study Guide

nitric acid

carbon dioxide


101

Name

Date

CHAPTER

Section 8.3

Class

8

STUDY GUIDE

Molecular Structures

In your textbook, read about Lewis structures.

For each statement below, write true or false. 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 electron

affinity. 3. In molecules, hydrogen is always a terminal atom. 4. The number of bonding pairs in a molecule is equal to the number of

electrons. 5. To find the total number of electrons available for bonding in a positive

ion, you should add the ion charge to the total number of valence electrons of the atoms present. 6. The electrons in a coordinate covalent bond are donated by both the

bonded atoms. 7. Resonance occurs when more than one valid Lewis structure can be

written for a molecule. 8. Nitrate is an example of an ion that forms resonance structures. Copyright © Glencoe/McGraw-Hill, a division of The McGraw-Hill Companies, Inc.

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 of

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

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

102

Column B

13. Odd number of valence electrons

a. O3

14. Fewer than 8 electrons around an atom

b. BF3

15. More than 8 electrons around central atom

c. NO

16. More than one valid Lewis structure

d. SF6


Study Guide

Name

CHAPTER

Section 8.4

Date

Class

8

STUDY GUIDE

Molecular Shape

In 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 a. tetrahedral.

b. trigonal planar.


d. linear.

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

central atom is a. tetrahedral.

b. trigonal planar.


d. linear.

In your textbook, read about hybridization.

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

hybridization

sp3

identical

methane

The formation of new orbitals from a combination or rearrangement of valence electrons is called (7) (8)

. The orbitals that are produced in this way are 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)

Study Guide

.


103

Name

Date

CHAPTER

Section 8.5

Class

8

STUDY GUIDE


In your textbook, read about electronegativity.

Use the table of electronegativities below to answer the following questions. Electronegativities of Some Elements

1 H 2.20 3 Li 0.98 11 Na 0.93 19 K 0.82 37 Rb 0.82

4 Be 1.57 12 Mg 1.31 20 Ca 1.00 38 Sr 0.95

21 Sc 1.36 39 Y 1.22

22 Ti 1.54 40 Zr 1.33

23 V 1.63 41 Nb 1.6

24 Cr 1.66 42 Mo 2.16

25 Mn 1.55 43 Tc 2.10

26 Fe 1.83 44 Ru 2.2

27 Co 1.88 45 Rh 2.28

28 Ni 1.91 46 Pd 2.20

29 Cu 1.90 47 Ag 1.93

55 Cs 0.79 87 Fr 0.7

56 Ba 0.89 88 Ra 0.9

57 La 1.10 89 Ac 1.1

72 Hf 1.3

73 Ta 1.5

74 W 1.7

75 Re 1.9

76 Os 2.2

77 Ir 2.2

78 Pt 2.2

79 Au 2.4

5 B 2.04

6 C 2.55

7 N 3.04

8 O 3.44

9 F 3.98

30 Zn 1.65 48 Cd 1.69

13 Al 1.61 31 Ga 1.81 49 In 1.78

14 Si 1.90 32 Ge 2.01 50 Sn 1.96

15 P 2.19 33 As 2.18 51 Sb 2.05

16 S 2.58 34 Se 2.55 52 Te 2.1

17 Cl 3.16 35 Br 2.96 53 I 2.66

80 Hg 1.9

81 Tl 1.8

82 Pb 1.8

83 Bi 1.9

84 Po 2.0

85 At 2.2

Metal Metalloid Nonmetal

2. Which element has the highest electronegativity? What is the numerical value? What are the

name 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 the

name 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 exists

between two atoms?

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

For each statement below, write true or false. 6. Ionic compounds are usually soluble in polar substances. 7. In a covalent molecular compound, the attraction between molecules tends

to be strong. 104


Study Guide


1. What is the meaning of the term electronegativity?

Name

CHAPTER

Section 8.5

Date

8

Class

STUDY GUIDE

continued

In your textbook, read about bond polarity.

Using the table of electronegativities on the preceding page, circle the letter of the choice that 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 covalent

and 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 is

polar 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. Study Guide


105

Name

Date

Class

CHAPTER 8


Assessment



4.

7.

2.

5.

8.

3.

6.

9.

10.

Short Answer

Answer each question with complete sentences. 11. 12. 13. Extended Response


SAT Subject Test: Chemistry 15.

17.

16.

18.

112


19.


14.

All Chemistry Worksheets Chapter 5-8

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