1. Chapter 5 Section 1 History of the Periodic Table
The Periodic Law
Chapter 5
Section 1 History of the Periodic Table
Section 2 Electron Configuration and the Periodic Table
Section 3 Electron Configuration and Periodic
Properties

2. 5.1 Mendeleev and Chemical Periodicity
Section 1 History of the Periodic Table
Chapter 5
5.1 Mendeleev and Chemical Periodicity
Mendeleev noticed that when the elements were arranged in order of increasing atomic mass, certain similarities in their chemical properties appeared at regular intervals.
Repeating patterns are referred to as
periodic.
Mendeleev created a table in which
elements with similar properties were
grouped together—a periodic table of the
elements.

4. Mendeleev and Chemical Periodicity, continued
Section 1 History of the Periodic Table
Chapter 5
Mendeleev and Chemical Periodicity, continued
After Mendeleev placed all the known elements in his periodic table, several empty spaces were left.
In 1871 Mendeleev predicted the existence and properties of elements that would fill three of the spaces.
By 1886, all three of these elements had been discovered.

5. Properties of Some Elements Predicted By Mendeleev
Section 1 History of the Periodic Table
Chapter 5
Properties of Some Elements Predicted By Mendeleev

6. Moseley and the Periodic Law
Section 1 History of the Periodic Table
Chapter 5
Moseley and the Periodic Law
In 1911, the English scientist Henry Moseley discovered that the elements fit into patterns better when they were arranged according to atomic number, rather than atomic weight.
The Periodic Law states that the physical and chemical properties of the elements are periodic functions of their atomic numbers.

7. The Modern Periodic Table
Section 1 History of the Periodic Table
Chapter 5
The Modern Periodic Table
an arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column, or group.
The periodic table is used to determine the electron configuration, atomic number, symbol, name, and atomic mass
Elements of the same family have similar properties because their electron configurations are similar.

8. Periodicity of Atomic Numbers
Section 1 History of the Periodic Table
Chapter 5
Periodicity of Atomic Numbers

9. The Modern Periodic Table
Section 1 History of the Periodic Table
Chapter 5
The Modern Periodic Table
Periodicity: regularly repeating pattern
The difference between atomic numbers down any group of the p.t. changes at regular intervals (Ex. Group 1 & 18)

10. Periods and Blocks of the Periodic Table
Section 2 Electron Configuration and the Periodic Table
Chapter 5
Periods and Blocks of the Periodic Table
Elements are arranged vertically in the periodic table in groups that share similar chemical properties.
Elements are also organized horizontally in rows, or periods.
The length of each period is determined by the number of electrons that can occupy the sublevels being filled in that period.

11. Relationship Between Periodicity and Electron Configurations
Section 2 Electron Configuration and the Periodic Table
Chapter 5
Relationship Between Periodicity and Electron Configurations

12. Section 2 Electron Configuration and the Periodic Table
Chapter 5
The elements of Group 18 of the periodic table are known as the Noble Gases.
Helium, neon, argon, xenon, etc.
Unreactive, colorless, odorless gases at room temperature

13. Section 2 Electron Configuration and the Periodic Table
Chapter 5
Group 1 elements of the periodic table are known as the alkali metals.
Very reactive (not found in pure form in nature), soft, silver colored, shiny; low density
Group 2 elements called the alkaline-earth metals.
Very reactive (but less reactive than alkali metals); silver colored, more dense than previous

14. Chapter 5 Transition metals Less reactive than alkaline-earth metals
Section 2 Electron Configuration and the Periodic Table
Chapter 5
Transition metals
Less reactive than alkaline-earth metals
Shiny, good conductors of thermal and electric energy; higher densities and melting points

15. Section 2 Electron Configuration and the Periodic Table
Chapter 5
In the periodic table, the f-block elements are wedged between Groups 3 and 4 in the sixth and seventh periods.
Their position reflects the fact that they involve the filling of the 4f sublevel.
The first row of the f block, the lanthanides, are shiny metals similar in reactivity to the Group 2 alkaline-earth metals.
The second row of the f block, the actinides, are between actinium and rutherfordium. The actinides are all radioactive.

17. Section 2 Electron Configuration and the Periodic Table
Chapter 5
The p-block elements together with the s-block elements are called the main-group (representative) elements.
The properties of elements of the p block vary greatly.
Includes metals, metalloids (semi-metals), nonmetals

18. Section 2 Electron Configuration and the Periodic Table
Chapter 5
The metalloids, or semi-metals, are located between nonmetals and metals in the p block.
The elements of Group 17 are known as the halogens.
The halogens are the most reactive nonmetals.
They react vigorously with most metals to form examples of the type of compound known as salts.
“Other metals” metals of the p block are generally harder & denser than s-block alkaline-earth metals, but softer & less dense than d-block metals.

19. Section 2 Electron Configuration and the Periodic Table
Chapter 5

20. Periods and Blocks of the Periodic Table
Section 2 Electron Configuration and the Periodic Table
Chapter 5
Periods and Blocks of the Periodic Table
Hydrogen has an electron configuration of 1s1, but despite the ns1 configuration, it does not share the same properties as the elements of Group 1.
Hydrogen is a unique element.
Like the Group 2 elements, helium has an ns2 group configuration. Yet it is part of Group 18.
Because its highest occupied energy level is filled by two electrons, helium possesses special chemical stability.

21. Periods and Blocks of the Periodic Table, continued
Section 2 Electron Configuration and the Periodic Table
Chapter 5
Periods and Blocks of the Periodic Table, continued
Sample Problem A
a. Without looking at the periodic table, identify the group, period, and block in which the element that has the electron configuration [Xe]6s2 is located.
b. Without looking at the periodic table, write the electron configuration for the Group 1 element in the third period. Is this element likely to be more reactive or less reactive than the element described in (a)?

22. Periods and Blocks of the Periodic Table, continued
Section 2 Electron Configuration and the Periodic Table
Chapter 5
Periods and Blocks of the Periodic Table, continued
Sample Problem A Solution
The element is in Group 2, as indicated by the group configuration of ns2.
It is in the sixth period, as indicated by the highest principal quantum number in its configuration, 6.
The element is in the s block.

23. Periods and Blocks of the Periodic Table, continued
Section 2 Electron Configuration and the Periodic Table
Chapter 5
Periods and Blocks of the Periodic Table, continued
Sample Problem A Solution, continued
In a third-period element, the highest occupied energy level is the third main energy level, n = 3. The 1s, 2s, and 2p sublevels are completely filled.
This element has the following configuration:
1s22s22p63s1 or [Ne]3s1
Because it is in Group 1, this element is likely to be more reactive than the element described in (a), which is in Group 2.

24. Periods and Blocks of the Periodic Table, continued
Section 2 Electron Configuration and the Periodic Table
Chapter 5
Periods and Blocks of the Periodic Table, continued
Sample Problem B
An element has the electron configuration [Kr]4d55s1. Without looking at the periodic table, identify the period, block, and group in which this element is located. Then, consult the periodic table to identify this element and the others in its group.

25. Periods and Blocks of the Periodic Table, continued
Section 2 Electron Configuration and the Periodic Table
Chapter 5
Periods and Blocks of the Periodic Table, continued
Sample Problem B Solution
The number of the highest occupied energy level is 5, so the element is in the fifth period.
There are five electrons in the d sublevel, which means that it is incompletely filled. The d sublevel can hold 10 electrons. Therefore, the element is in the d block.
For d-block elements, the number of electrons in the ns sublevel (1) plus the number of electrons in the (n − 1)d sublevel (5) equals the group number, 6.
This Group 6 element is molybdenum. The others in Group 6 are chromium, tungsten, and seaborgium.

26. Periods and Blocks of the Periodic Table, continued
Section 2 Electron Configuration and the Periodic Table
Chapter 5
Periods and Blocks of the Periodic Table, continued
Sample Problem C
Without looking at the periodic table, write the outer electron configuration for the Group 14 element in the second period. Then, name the element, and identify it as a metal, nonmetal, or metalloid.

27. Periods and Blocks of the Periodic Table, continued
Section 2 Electron Configuration and the Periodic Table
Chapter 5
Periods and Blocks of the Periodic Table, continued
Sample Problem C Solution
The group number is higher than 12, so the element is in the p block.
The total number of electrons in the highest occupied s and p sublevels is therefore equal to the group number minus 10 (14 − 10 = 4).
Two electrons are in the s sublevel, so two electrons must also be present in the 2p sublevel.
The outer electron configuration is 2s22p2.
The element is carbon, C, which is a nonmetal.

28. Periods and Blocks of the Periodic Table, continued
Section 2 Electron Configuration and the Periodic Table
Chapter 5
Periods and Blocks of the Periodic Table, continued
Sample Problem D
Name the block and group in which each of the following elements is located in the periodic table. Then, use the periodic table to name each element. Identify each element as a metal, nonmetal, or metalloid. Finally, describe whether each element has high reactivity or low reactivity.
a. [Xe]4f145d96s1 c. [Ne]3s23p6
b. [Ne]3s23p5 d. [Xe]4f66s2

29. Periods and Blocks of the Periodic Table, continued
Section 2 Electron Configuration and the Periodic Table
Chapter 5
Periods and Blocks of the Periodic Table, continued
Sample Problem D Solution
The 4f sublevel is filled with 14 electrons. The 5d sublevel is partially filled with nine electrons. Therefore, this element is in the d block.
The element is the transition metal platinum, Pt, which is in Group 10 and has a low reactivity.
b. The incompletely filled p sublevel shows that this element is in the p block.
A total of seven electrons are in the ns and np sublevels, so this element is in Group 17, the halogens.
The element is chlorine, Cl, and is highly reactive.

30. Periods and Blocks of the Periodic Table, continued
Section 2 Electron Configuration and the Periodic Table
Chapter 5
Periods and Blocks of the Periodic Table, continued
Sample Problem D Solution, continued
c. This element has a noble-gas configuration and thus is in Group 18 in the p block.
The element is argon, Ar, which is an unreactive nonmetal and a noble gas.
d. The incomplete 4f sublevel shows that the element is in the f block and is a lanthanide.
Group numbers are not assigned to the f block.
The element is samarium, Sm. All of the lanthanides are reactive metals.

31. Relating Period Length and Sublevels Filled
Visual Concepts
Chapter 5
Relating Period Length and Sublevels Filled

32. Section 3 Electron Configuration and Periodic Properties
Chapter 5
Objectives
Define atomic and ionic radii, ionization energy, electron affinity, and electronegativity.
Compare the periodic trends of atomic radii, ionization energy, and electronegativity, and state the reasons for these variations.
Define valence electrons, and state how many are present in atoms of each main-group element.
Compare the atomic radii, ionization energies, and electronegativities of the d-block elements with those of the main-group elements.

33. Section 3 Electron Configuration and Periodic Properties
Chapter 5
Review:
Elements on the periodic table are arranged according to their atomic numbers
As a result, there is a relationship between periodic law and electron configurations

34. 7 Periodic Properties of the Periodic Table
Section 3 Electron Configuration and Periodic Properties
Chapter 5
7 Periodic Properties of the Periodic Table
Valence Electrons
Atomic Radii
Ionic Radii
Ionization Energy
Electron Affinity
Electronegativity
Metallic Behavior

35. Trend 1: Valence Electrons
Section 3 Electron Configuration and Periodic Properties
Chapter 5
Trend 1: Valence Electrons
The electrons available to be lost, gained, or shared in the formation of chemical compounds
Main group elements: valence electrons are located in s & p sublevels
Inner-shell electrons are not involved in compound formation
Held too tightly by the nucleus to be involved in forming bonds

36. Trend 1: Valence Electrons
Section 3 Electron Configuration and Periodic Properties
Chapter 5
Trend 1: Valence Electrons
General Electron Configurations
Group 1:
Group 2:
Group 13:
Group 14:
Group 15:
Group 16:
Group 17:
Group 18:
Typical Ion Charges +1 +2 +3
+4/-4 -3 -2 -1
No charge
ns1
ns2
ns2np1
ns2np2
ns2np3
ns2np4
ns2np5
ns2np6

37. Chapter 5 Trend 2: Atomic Radii
Section 3 Electron Configuration and Periodic Properties
Chapter 5
Trend 2: Atomic Radii
one-half the distance between the nuclei of identical atoms that are bonded together.
Determined by taking 2 atoms in gas phase and measuring distance from nucleus to nucleus / 2
Units of picometers (pm)

38. Shielding Effect: Chapter 5
Section 3 Electron Configuration and Periodic Properties
Chapter 5
Shielding Effect:
Inner-shell electrons (those not in highest occupied energy level) shield the valence electrons from the pull of the positive nucleus
“Pull of nucleus” known as effective nucleus charge
Remember: as atomic numbers increase, not only are you adding an additional electron, the nucleus is also gaining a proton.

39. Chapter 5 Across a Period
Section 3 Electron Configuration and Periodic Properties
Chapter 5
Across a Period
valence e-’s increase, but no new inner-shell e-’s
Increasing positive charge of nucleus (effective nuclear charge) pulls electrons toward the nucleus.
Atoms tend to become smaller
Down a Group
New energy level added, previous valence e-’s add to number of inner-shell electrons
size of electron cloud increases because more effective shielding occurs
Atoms tend to become larger

40. Section 3 Electron Configuration and Periodic Properties
Chapter 5

41. Chapter 5 Trend 3: Ionic Radii
Section 3 Electron Configuration and Periodic Properties
Chapter 5
Trend 3: Ionic Radii
positive ion: Cation (loss of electrons)
negative ion: Anion (gain of electrons)
Neutral  Cation: decrease in atomic radius
effective nuclear charge increases, electron cloud shrinks
Neutral  Anion: increase in atomic radius
Size of electron cloud increases

42. Chapter 5 Across a Period
Section 3 Electron Configuration and Periodic Properties
Chapter 5
Across a Period
effective nuclear charge pulls electrons toward the nucleus
Ionic Radii gradually tend to become smaller
Down a Group
size of electron cloud increases because more effective shielding occurss
Ionic Radii gradually end to become larger

43. Section 3 Electron Configuration and Periodic Properties
Chapter 5

44. Periodic Trends of Radii
Section 3 Electron Configuration and Periodic Properties
Chapter 5
Periodic Trends of Radii

45. A + energy  A+ + e- Chapter 5 Trend 4: Ionization Energy
Section 3 Electron Configuration and Periodic Properties
Chapter 5
Trend 4: Ionization Energy
The energy required to remove one electron from a neutral atom of an element
Remove 1st e-: first ionization energy, IE1
Remove 2nd e-: second ionization energy, IE2
Removing an electron always requires energy input
Measured in KJ/mol
A + energy  A+ + e-

46. Chapter 5 Across a Period
Section 3 Electron Configuration and Periodic Properties
Chapter 5
Across a Period
Effective nuclear charge increases across a period & makes it harder to remove electrons
Ionization Energies generally increase
Down a Group
Electrons removed from higher energy levels are farther from the nucleus and easier to remove
Ionization Energies generally decrease

47. Trend 5: Electron Affinity
Section 3 Electron Configuration and Periodic Properties
Chapter 5
Trend 5: Electron Affinity
Energy change (gained or lost) that occurs when an electron is acquired by a neutral atom
Across a Period
Increasing effective nuclear charge along the same sublevel attracts electrons more strongly
Electron affinity generally increases
Down a Group
Larger electron cloud, e-’s further from nucleus
Electron affinity generally decreases

48. Trend 6: Electronegativity
Section 3 Electron Configuration and Periodic Properties
Chapter 5
Trend 6: Electronegativity
measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound.
Valence electrons hold atoms together in chemical compounds.
In many compounds, the negative charge of the valence electrons is concentrated closer to one atom than to another.

49. Across a Period
Increasing effective nuclear charge creates more pulls for electrons of other ion
Electronegativities tend to increase
Down a Group
Increased size of electron cloud decreases ability of nucleus to pull electrons of other ion
Electronegativities tend to decrease or remain about the same

50. Trend 7: Metallic Behavior
Section 3 Electron Configuration and Periodic Properties
Chapter 5
Trend 7: Metallic Behavior
Typical metallic properties that are periodic
Ability to be reduced (become positive ions)
Formation of basic oxide & ionic chlorides
Ability to displace hydrogen from other compounds
Metallic behavior tends to decrease across period, and increase down a group
Most metallic element - Fr
Least metallic element - F

51. In Summary…
Increasing effective nuclear charge across a period (poor shielding)
Decrease atomic radii
Decrease ionic radii
Increase ionization energy
Increase electron affinity
Increase electronegativity

52. In Summary…
Increasing size of electron cloud down a group (shielding effect increasing)
Increase in atomic radii
Increase in ionic radii
Decrease ionization energy
Decrease electron affinity
Decrease electronegativity

53. Electronegativity
Electronegativity

54. End of Chapter 5 Show

55. Chapter 5 Multiple Choice
Standardized Test Preparation
Chapter 5
Multiple Choice
1. In the modern periodic table, elements are arranged according to
A. decreasing atomic mass.
B. Mendeleev’s original model.
C. increasing atomic number.
D. when they were discovered.

56. Chapter 5 Multiple Choice
Standardized Test Preparation
Chapter 5
Multiple Choice
1. In the modern periodic table, elements are arranged according to
A. decreasing atomic mass.
B. Mendeleev’s original model.
C. increasing atomic number.
D. when they were discovered.

57. Chapter 5 Multiple Choice
Standardized Test Preparation
Chapter 5
Multiple Choice
2. Group 17 elements, the halogens, are the most reactive of the nonmetal elements because they
A. require only one electron to fill their outer energy level.
B. have the highest ionization energies.
C. have the largest atomic radii.
D. are the farthest to the right in the periodic table.

58. Chapter 5 Multiple Choice
Standardized Test Preparation
Chapter 5
Multiple Choice
2. Group 17 elements, the halogens, are the most reactive of the nonmetal elements because they
A. require only one electron to fill their outer energy level.
B. have the highest ionization energies.
C. have the largest atomic radii.
D. are the farthest to the right in the periodic table.

59. Chapter 5 Multiple Choice 3. The periodic law states that
Standardized Test Preparation
Chapter 5
Multiple Choice
3. The periodic law states that
A. the chemical properties of elements can be grouped according to periodicity.
B. the properties of the elements are functions of atomic mass.
C. all elements in the same group have the same number of valence electrons.
D. all elements with the same number of occupied energy levels must be in the same group.

60. Chapter 5 Multiple Choice 3. The periodic law states that
Standardized Test Preparation
Chapter 5
Multiple Choice
3. The periodic law states that
A. the chemical properties of elements can be grouped according to periodicity.
B. the properties of the elements are functions of atomic mass.
C. all elements in the same group have the same number of valence electrons.
D. all elements with the same number of occupied energy levels must be in the same group.

61. Chapter 5 Multiple Choice
Standardized Test Preparation
Chapter 5
Multiple Choice
4. As you move left to right across Period 3 from Mg to Cl, the energy needed to remove an electron from an atom
A. generally increases.
B. generally decreases.
C. does not change.
D. varies unpredictably.

62. Chapter 5 Multiple Choice
Standardized Test Preparation
Chapter 5
Multiple Choice
4. As you move left to right across Period 3 from Mg to Cl, the energy needed to remove an electron from an atom
A. generally increases.
B. generally decreases.
C. does not change.
D. varies unpredictably.

63. Chapter 5 Multiple Choice
Standardized Test Preparation
Chapter 5
Multiple Choice
5. Which of the following elements has the highest electronegativity?
A. oxygen.
B. hydrogen.
C. fluorine.
D. carbon.

64. Chapter 5 Multiple Choice
Standardized Test Preparation
Chapter 5
Multiple Choice
5. Which of the following elements has the highest electronegativity?
A. oxygen.
B. hydrogen.
C. fluorine.
D. carbon.

65. Chapter 5 Multiple Choice 6. The noble gases have
Standardized Test Preparation
Chapter 5
Multiple Choice
6. The noble gases have
A. high ionization energies.
B. high electron affinities.
C. large atomic radii.
D. a tendency to form both cations and anions.

66. Chapter 5 Multiple Choice 6. The noble gases have
Standardized Test Preparation
Chapter 5
Multiple Choice
6. The noble gases have
A. high ionization energies.
B. high electron affinities.
C. large atomic radii.
D. a tendency to form both cations and anions.

67. Chapter 5 Multiple Choice
Standardized Test Preparation
Chapter 5
Multiple Choice
7. Which electron configuration is not correct?
A. O2– [He]2s22p6
B. Mg2+ [He]2s22p6
C. V3+ [Ar]3d2
D. Al3+ [Ar]2s22p6

68. Chapter 5 Multiple Choice
Standardized Test Preparation
Chapter 5
Multiple Choice
7. Which electron configuration is not correct?
A. O2– [He]2s22p6
B. Mg2+ [He]2s22p6
C. V3+ [Ar]3d2
D. Al3+ [Ar]2s22p6

69. Chapter 5 Multiple Choice
Standardized Test Preparation
Chapter 5
Multiple Choice
8. Which two elements are more likely to have the same charge on their ions?
A. Se and As
B. Sn and Si
C. Ca and Rb
D. I and Xe

70. Chapter 5 Multiple Choice
Standardized Test Preparation
Chapter 5
Multiple Choice
8. Which two elements are more likely to have the same charge on their ions?
A. Se and As
B. Sn and Si
C. Ca and Rb
D. I and Xe

71. Chapter 5 Multiple Choice
Standardized Test Preparation
Chapter 5
Multiple Choice
9. Which list ranks the elements Sr, Te, Kr, Ru, and Cs in order of increasing electron affinity?
A. Sr < Te < Ru < Cs < Kr
B. Te < Ru < Sr < Cs < Kr
C. Cs < Sr < Ru < Te < Kr
D. Kr < Cs < Sr < Ru < Te

72. Chapter 5 Multiple Choice
Standardized Test Preparation
Chapter 5
Multiple Choice
9. Which list ranks the elements Sr, Te, Kr, Ru, and Cs in order of increasing electron affinity?
A. Sr < Te < Ru < Cs < Kr
B. Te < Ru < Sr < Cs < Kr
C. Cs < Sr < Ru < Te < Kr
D. Kr < Cs < Sr < Ru < Te

73. Chapter 5 Short Answer Standardized Test Preparation
10. The second ionization energies for the elements S–Ti are listed in a scrambled order below. Assign the correct IE2 value to each element (Hint: S has IE2 = 2251 kJ/mol, and Ti has IE2 = 1310 kJ/mol.) Explain your reasoning.
IE2 values (kJ/mol): 2666, 2297, 3051, 1235, 2251, 1310, and 1145

74. Chapter 5 Short Answer Standardized Test Preparation
10. The second ionization energies for the elements S–Ti are listed in a scrambled order below. Assign the correct IE2 value to each element (Hint: S has IE2 = 2251 kJ/mol, and Ti has IE2 = 1310 kJ/mol.) Explain your reasoning.
IE2 values (kJ/mol): 2666, 2297, 3051, 1235, 2251, 1310, and 1145
Answer: S: 2551 kJ/mol; Cl: 2297 kJ/mol; Ar: 2666 kJ/mol; K: 3051 kJ/mol; Ca: 1145 kJ/mol; Sc: 1235 kJ/mol; Ti: 1310 kJ/mol. For the second ionization, the general trend is for increasing IE2 across the period in Groups 218 with Group 1 having the highest IE2. IE2 decreases going down a group.

75. Standardized Test Preparation
Chapter 5
Short Answer
11. What group most commonly forms 2– ions? Explain your reasoning.

76. Standardized Test Preparation
Chapter 5
Short Answer
11. What group most commonly forms 2– ions? Explain your reasoning.
Answer: Group 16 most commonly forms 2– ions, because these elements require only two more electrons to fill their shell (obtain a noble-gas configuration).

77. Chapter 5 Extended Response Standardized Test Preparation
12. An ordered list of atomic radii for 14 consecutive elements is shown below. Without using the graph in your book, make a graph of these atomic radii versus the element’s atomic number. Explain your reasoning.
Atomic radii (pm): 75, 73, 72, 71, 186, 160, 143, 118, 110, 103, 100, 98, 227, and 197.

78. Chapter 5 Extended Response Standardized Test Preparation
12. An ordered list of atomic radii for 14 consecutive elements is shown below. Without using the graph in your book, make a graph of these atomic radii versus the element’s atomic number. Explain your reasoning.
Atomic radii (pm): 75, 73, 72, 71, 186, 160, 143, 118, 110, 103, 100, 98, 227, and 197.
Answer: There are eight numbers between the first and second minimum; therefore, these eight elements must either be the second or third period. There are four numbers before the first maximum; hence, these four numbers must be elements in the second row because the first row contains only two elements (H and He). Therefore, the eight elements must be in the third period. The first atomic radius listed must correspond to atomic number 7, which is N. (continued on next slide)

79. Chapter 5 Extended Response
Standardized Test Preparation
Chapter 5
Extended Response
12. Answer, continued: Your graph should look like the following: