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Slide 1 / 113 Slide 2 / 113 AP Chemistry The Atom 2015-08-25 www.njctl.org Slide 3 / 113 Table of Contents: The Atom (Pt. B) · Periodic Table · Periodic Trends Click on the topic to go to that section Slide 4 / 113 Periodic Table Return to Table of Contents Slide 5 / 113 Recall that the periodic law states that the physical and chemical properties of the elements tend to recur in a systematic way when arranged by increasing atomic number. The Periodic Law Slide 6 / 113 Special Groups Some groups have distinctive properties and are given special names. Alkali Metals Alkaline Earth Metals Halogens Noble Gases Transition Metals

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Slide 1 / 113 Slide 2 / 113

AP Chemistry

The Atom

2015-08-25

www.njctl.org

Slide 3 / 113

Table of Contents: The Atom (Pt. B)

· Periodic Table

· Periodic Trends

Click on the topic to go to that section

Slide 4 / 113

Periodic Table

Return to Tableof Contents

Slide 5 / 113

Recall that the periodic law states that the physical and chemical properties of the elements tend to recur in a systematic way when

arranged by increasing atomic number.

The Periodic Law

Slide 6 / 113

Special Groups

Some groups have distinctive properties and are given special names.

Alk

ali M

etal

sA

lkal

ine

Eart

h M

etal

s

Hal

ogen

sN

oble

Gas

es

Transition Metals

Slide 7 / 113

1 What is the atomic number for the element in period 3, group 16?

Slide 7 (Answer) / 113

1 What is the atomic number for the element in period 3, group 16?

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16

Slide 8 / 113

2 What is the atomic number for the element in period 5, group 3?

Slide 8 (Answer) / 113

2 What is the atomic number for the element in period 5, group 3?

[This object is a pull tab]A

nsw

er

6

Slide 9 / 113

3 To which group on the periodic table does Neon belong?

A Alkali Metals

B Transition Metals

C Noble GasesD Alkaline Earth Metals

Slide 9 (Answer) / 113

3 To which group on the periodic table does Neon belong?

A Alkali Metals

B Transition Metals

C Noble GasesD Alkaline Earth Metals

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C

Slide 10 / 113

4 To which group on the periodic table does Iron belong?

A Alkali Metals

B Transition Metals

C Halogens

D Alkaline Earth Metals

Slide 10 (Answer) / 113

4 To which group on the periodic table does Iron belong?

A Alkali Metals

B Transition Metals

C Halogens

D Alkaline Earth Metals

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B

Slide 11 / 113

5 To which group on the periodic table does Beryllium belong?

A Alkali Metals

B Transition Metals

C Halogens

D Alkaline Earth Metals

Slide 11 (Answer) / 113

5 To which group on the periodic table does Beryllium belong?

A Alkali Metals

B Transition Metals

C Halogens

D Alkaline Earth Metals

[This object is a pull tab]A

nsw

er

D

Slide 12 / 113

6 Two elements are studied. One with atomic number X and one with atomic number X+1. It is known that element X is a Noble Gas. Which group on the periodic table is X+1 in?

A Transition Metals

B Halogens

C Alkali Metals

D There is no way to tell

Slide 12 (Answer) / 113

6 Two elements are studied. One with atomic number X and one with atomic number X+1. It is known that element X is a Noble Gas. Which group on the periodic table is X+1 in?

A Transition Metals

B Halogens

C Alkali Metals

D There is no way to tell

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C

Slide 13 / 113

1A 2A 8A 1 2 183A 4A 5A 6A 7A

13 14 15 16 17 8B3B 4B 5B 6B 7B 1B 2B 3 4 5 6 7 8 9 10 11 12

}

There are two methods for labeling the groups, the older method shown in black on the top and the newer method shown in blue on the bottom.

Periodic Table & Electron Configuration

Slide 14 / 113

Click here to view an Interactive Periodic Table that shows orbitals for each Element

Click here for an electron orbital game.

Periodic Table & Electron Configuration

Slide 15 / 113

7 The highlighted elements below are in the

A s blockB d blockC p blockD f block

Slide 15 (Answer) / 113

7 The highlighted elements below are in the

A s blockB d blockC p blockD f block

[This object is a pull tab]A

nsw

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B

Slide 16 / 113

8 The highlighted elements below are in the

A s blockB d blockC p blockD f block

Slide 16 (Answer) / 113

8 The highlighted elements below are in the

A s blockB d blockC p blockD f block

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A

Slide 17 / 113

9 The highlighted elements below are in the

A s blockB d blockC p blockD f block

Slide 17 (Answer) / 113

9 The highlighted elements below are in the

A s blockB d blockC p blockD f block

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D

Slide 18 / 113

10 The electron configuration ending ns2p6 belongs in which group of the periodic table

A Alkali MetalsB Alkaline Earth MetalsC HalogensD Noble Gases

Slide 18 (Answer) / 113

10 The electron configuration ending ns2p6 belongs in which group of the periodic table

A Alkali MetalsB Alkaline Earth MetalsC HalogensD Noble Gases

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D

Slide 19 / 113

11 An unknown element has an electron configuration ending in s2. It is most likely in which group?

A Alkaline Earth Metals

B Halogens

C Alkali Metals

D Transition Metals

Slide 19 (Answer) / 113

11 An unknown element has an electron configuration ending in s2. It is most likely in which group?

A Alkaline Earth Metals

B Halogens

C Alkali Metals

D Transition Metals

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A

Slide 20 / 113

Shorthand ConfigurationsNoble Gas elements are used to write shortened electron configurations.

To write a Shorthand Configuration for an element:

(1) Write the Symbol of the Noble Gas element from the row before it in brackets [ ].

(2) Add the remaining electrons by starting at the s orbital of the row that the element is in until the configuration is complete.

Slide 21 / 113

12 What would be the expected "shorthand" electron configuration for Sulfur (S)?A [He]3s23p4

B [Ar]3s24p4

C [Ne]3s23p3

D [Ne]3s23p4

Slide 21 (Answer) / 113

12 What would be the expected "shorthand" electron configuration for Sulfur (S)?A [He]3s23p4

B [Ar]3s24p4

C [Ne]3s23p3

D [Ne]3s23p4

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D

Slide 22 / 113

13 Which of the following represents an electron configuration of a Halogen?

A [He]2s1

B [Ne]3s23p5

C [Ar]4s23d2

D [Kr]5s24d105p4

Slide 22 (Answer) / 113

13 Which of the following represents an electron configuration of a Halogen?

A [He]2s1

B [Ne]3s23p5

C [Ar]4s23d2

D [Kr]5s24d105p4

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B

Slide 23 / 113

14 The electron configuration [Ar]4s23d5 belongs in which group of the periodic table

A Alkali MetalsB Alkaline Earth MetalsC Transition MetalsD Halogens

Slide 23 (Answer) / 113

14 The electron configuration [Ar]4s23d5 belongs in which group of the periodic table

A Alkali MetalsB Alkaline Earth MetalsC Transition MetalsD Halogens

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B

Slide 24 / 113

15 Which of the following represents an electron configuration of an alkaline earth metal?

A [He]2s1

B [Ne]3s23p6

C [Ar]4s23d2

D [Xe]6s2

Slide 24 (Answer) / 113

15 Which of the following represents an electron configuration of an alkaline earth metal?

A [He]2s1

B [Ne]3s23p6

C [Ar]4s23d2

D [Xe]6s2

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D

Slide 25 / 113

Stability

When the elements were studied scientists noticed that some of them do not react in certain situations in which others do. These elements were labeled "stable" because they did not change easily. When these stable elements were grouped together, it was noted that periodically, there were patterns in the occurrence of stable elements.

Today we recognize that this difference in stability is due to electron configurations.

Slide 26 / 113

StabilityElements of varying stability fall into one of 3 categories. The most stable atoms have completely full energy levels.

~Full Energy Level ~Full Sublevel (s, p, d, f) ~Half Full Sublevel ( d5, f7)

1234567

67

Slide 27 / 113

StabilityNext in order of stability are elements with full sublevels.

~Full Energy Level ~Full Sublevel (s, p, d, f) ~Half Full Sublevel ( d5, f7)

1234567

67

Slide 28 / 113

StabilityFinally, the elements with half full sublevels are also stable, but not as stable as elements with fully energy levels or sublevels.

~Full Energy Level ~Full Sublevel (s, p, d, f) ~Half Full Sublevel ( d 5, f7)

1234567

67

Slide 29 / 113

16 The elements in the periodic table that have completely filled shells or subshells are referred to as:

A noble gases.

B halogens.

C alkali metals.

D transition elements.

Slide 29 (Answer) / 113

16 The elements in the periodic table that have completely filled shells or subshells are referred to as:

A noble gases.

B halogens.

C alkali metals.

D transition elements.

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A

Slide 30 / 113

17 Alkaline earth metals are more stable than alkali metals because...

A they have a full shell.

B they have a full subshell.

C they have a half-full subshell.

D they contain no p orbitals.

Slide 30 (Answer) / 113

17 Alkaline earth metals are more stable than alkali metals because...

A they have a full shell.

B they have a full subshell.

C they have a half-full subshell.

D they contain no p orbitals.

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B

Slide 31 / 113

18 The elements in the periodic table which lack one electron from a filled shell are referred to as:

A noble gases.

B halogens.

C alkali metals.

D transition elements.

Slide 31 (Answer) / 113

18 The elements in the periodic table which lack one electron from a filled shell are referred to as:

A noble gases.

B halogens.

C alkali metals.

D transition elements.

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B

Slide 32 / 113

Electron Configuration Exceptions

You should know the basic exceptions in the d- and f-sublevels. These fall in the circled areas on the table below.

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67

Slide 33 / 113

Chromium Expect: [Ar] 4s2 3d4 Actually: [Ar] 4s1 3d5

For some elements, in order to exist in a more stable state, electrons from an s sublevel will move to a d sublevel, thus providing the stability of a half-full sublevel. To see why this can happen we need to examine how "close" d and s sublevels are.

Electron Configuration Exceptions

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67

Cr

Slide 34 / 113

12

3

4

5

6

7

1s

2s

2p

3s

3p4s

3d4p

5s

4d5p

6s 4f5d

6p5f7s

6d7p

6f7d

7f

Ene

rgy

Energies of Orbitals

Because of how close the f and d orbitals are to the s orbitals an electron very little energy is required to move an electron from the s orbital (leaving it half full) to the f or d orbital, causing them to also be half full.

(It's kind of like borrowing a cup of sugar from a neighbor. You'd only borrow it from someone you were close to and only if you needed it.)

Slide 35 / 113

CopperExpected: [Ar] 4s2 3d9 Actual: [Ar] 4s1 3d10

Copper gains stability when an electron from the 4sorbital fills the 3d orbital.

Electron Configuration Exceptions

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67

Cu

Slide 36 / 113

19 The electron configuration for Copper (Cu) is

A [Ar] 4s24d9

B [Ar] 4s14d9

C [Ar] 4s23d9

D [Ar] 4s13d10

Slide 36 (Answer) / 113

19 The electron configuration for Copper (Cu) is

A [Ar] 4s24d9

B [Ar] 4s14d9

C [Ar] 4s23d9

D [Ar] 4s13d10

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D

Slide 37 / 113

20 What would be the expected "shorthand" electron configuration for Silver (Ag)?A [Kr]5s25d9

B [Ar]5s14d10

C [Kr]5s24d9

D [Kr]5s14d10

Slide 37 (Answer) / 113

20 What would be the expected "shorthand" electron configuration for Silver (Ag)?A [Kr]5s25d9

B [Ar]5s14d10

C [Kr]5s24d9

D [Kr]5s14d10

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D

Slide 38 / 113

21 What would be the expected "shorthand" electron configuration for Molybdenum (Mb)?A [Kr]5s25d4

B [Ar]5s24d4

C [Kr]5s14d5

D [Kr]5s24d4

Slide 38 (Answer) / 113

21 What would be the expected "shorthand" electron configuration for Molybdenum (Mb)?A [Kr]5s25d4

B [Ar]5s24d4

C [Kr]5s14d5

D [Kr]5s24d4

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C

Slide 39 / 113

Periodic Trends

Return to Tableof Contents

Slide 40 / 113

Periodic Trends

There are four main trends in the periodic table:

Size or radius of atoms/ionsElectronegativityIonization energyMetallic character

Slide 41 / 113

Periodic Trends

These four periodic trends are all shaped by the interactions between the positive charge of the atomic nucleus and the

negative charges of electrons.

Slide 42 / 113

Effective Nuclear Charge

In a multi-electron atom, electrons are both attracted to the positive nucleus and repelled by other electrons.

The nuclear charge that an electron experiences depends on both factors.

For example, here's sodium.

Valence 3s electron

[Ne] inner shell electrons (10-)

Nucleus (11+)

10-

11+

-

Combined effect = 11-10 = 1+

Slide 43 / 113

This is the effective nuclear charge experience by the valence electron. The inner shell electrons are shielding the outermost electron from experiencing all but +1 e worth of charge!

The effective nuclear charge (Zeff) is found by Zeff = Z − S

Effective Nuclear Charge

Valence 3s electron

S (inner shell electrons) = 10-

Z = 11+

Zeff = 11-10 = 1+

10-

11+

-

Z is the atomic number (number of protons)

S is the shielding constant and represents the number of electrons in the inner shells of an atom.

Slide 44 / 113

22 Two elements are studied. One with atomic number X and one with atomic number X+1. Assuming element X is not a Noble Gas, which element has the larger shielding constant?

A Element X

B Element X+1

C They are both the same

D More information is needed

Slide 44 (Answer) / 113

22 Two elements are studied. One with atomic number X and one with atomic number X+1. Assuming element X is not a Noble Gas, which element has the larger shielding constant?

A Element X

B Element X+1

C They are both the same

D More information is needed

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C

Slide 45 / 113

23 Two elements are studied. One with atomic number X and one with atomic number X+1. It is known that element X is a Noble Gas. Which element has the larger shielding constant?

A Element X

B Element X+1

C They are both the same

D More information is needed

Slide 45 (Answer) / 113

23 Two elements are studied. One with atomic number X and one with atomic number X+1. It is known that element X is a Noble Gas. Which element has the larger shielding constant?

A Element X

B Element X+1

C They are both the same

D More information is needed

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B

Slide 46 / 113

24 In which subshell does an electron in a calcium atom experience the greatest effective nuclear charge? A 1s

B 2sC 3sD 3p

Slide 46 (Answer) / 113

24 In which subshell does an electron in a calcium atom experience the greatest effective nuclear charge? A 1s

B 2sC 3sD 3p

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A

Slide 47 / 113

Let's examine the trend in atomic radii for the first 18 elements.

atomic number

radius (pm)

0

200

100

0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18

H

LiNa

HeNe

Ar

We clearly see two trends!

1. As atomic number increases down a group, the radii increase. Why?

H < Li < Na

2. As atomic number increases across a period, the radii decrease. Why?

Li > Be > B > C > N > O > F > Ne

Atomic Radii

Slide 47 (Answer) / 113

Let's examine the trend in atomic radii for the first 18 elements.

atomic number

radius (pm)

0

200

100

0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18

H

LiNa

HeNe

Ar

We clearly see two trends!

1. As atomic number increases down a group, the radii increase. Why?

H < Li < Na

2. As atomic number increases across a period, the radii decrease. Why?

Li > Be > B > C > N > O > F > Ne

Atomic Radii

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1. An additional energy shell is gained increasing the atomic radii.

2. The effective nuclear charge increase, creating a stronger

Coulombic attraction between the nucleus and the valence electrons.

Slide 48 / 113

25 What is Zeff for Boron (B)?

5+

Slide 48 (Answer) / 113

25 What is Zeff for Boron (B)?

5+

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3

Slide 49 / 113

26 Compare the radial size of boron to lithium and beryllium.

A Li>Be>B

B Li<Be<B

C Li>B>Be

D Be<Li<B

Slide 49 (Answer) / 113

26 Compare the radial size of boron to lithium and beryllium.

A Li>Be>B

B Li<Be<B

C Li>B>Be

D Be<Li<B

[This object is a pull tab]A

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A

Slide 50 / 113

27 What is Zeff for Carbon (C)?

6+

Slide 50 (Answer) / 113

27 What is Zeff for Carbon (C)?

6+

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4

Slide 51 / 113

28 Compare the radial size of carbon to boron and nitrogen.

A C>N>B

B C<N<B

C B>C>N

D B<C<N

Slide 51 (Answer) / 113

28 Compare the radial size of carbon to boron and nitrogen.

A C>N>B

B C<N<B

C B>C>N

D B<C<N

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C

Slide 52 / 113

29 Which of the following equations correctly calculates the Coulombic force between the valence electrons and the nucleus of an oxygen atom?

A F = k(2e)2/r2

B F = k(4e)2/r2

C F = k(6e)2/r2

D F = k(8e)2/r2

Slide 52 (Answer) / 113

29 Which of the following equations correctly calculates the Coulombic force between the valence electrons and the nucleus of an oxygen atom?

A F = k(2e)2/r2

B F = k(4e)2/r2

C F = k(6e)2/r2

D F = k(8e)2/r2[This object is a pull tab]

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B

Slide 53 / 113

30 Given the atomic number of the smallest element in the 2nd period.

Slide 53 (Answer) / 113

30 Given the atomic number of the smallest element in the 2nd period.

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10

Slide 54 / 113

31 Across a period from left to right Zeff _____________.

A increases

B decreasesC remains the same

Slide 54 (Answer) / 113

31 Across a period from left to right Zeff _____________.

A increases

B decreasesC remains the same

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A

Slide 55 / 113

32 Down a group from top to bottom Zeff _____________.

A increases

B decreasesC remains the same

Slide 55 (Answer) / 113

32 Down a group from top to bottom Zeff _____________.

A increases

B decreasesC remains the same

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C

Slide 56 / 113

33 Atomic radius generally increases as we move __________.

A down a group and from right to left across a period

B up a group and from left to right across a period

C down a group and from left to right across a period

D up a group and from right to left across a period

Slide 56 (Answer) / 113

33 Atomic radius generally increases as we move __________.

A down a group and from right to left across a period

B up a group and from left to right across a period

C down a group and from left to right across a period

D up a group and from right to left across a period

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A

Slide 57 / 113

34 The atomic radius of main-group elements generally increases down a group because __________.

A effective nuclear charge increases down a group

B effective nuclear charge decreases down a group

C effective nuclear charge zigzags down a group

D the principal quantum number of the valence orbitals increases

Slide 57 (Answer) / 113

34 The atomic radius of main-group elements generally increases down a group because __________.

A effective nuclear charge increases down a group

B effective nuclear charge decreases down a group

C effective nuclear charge zigzags down a group

D the principal quantum number of the valence orbitals increases

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B

Slide 58 / 113

35 Of the following, which gives the correct order for atomic radius for Mg, Na, P, Si and Ar?

A Mg > Na > P > Si > Ar

B Ar > Si > P > Na > Mg

C Si > P > Ar > Na > Mg

D Na > Mg > Si > P > Ar

Slide 58 (Answer) / 113

35 Of the following, which gives the correct order for atomic radius for Mg, Na, P, Si and Ar?

A Mg > Na > P > Si > Ar

B Ar > Si > P > Na > Mg

C Si > P > Ar > Na > Mg

D Na > Mg > Si > P > Ar

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D

Slide 59 / 113

36 Which of the following atoms would have a smaller atomic radii than Ar and why?

A Fe - It has more core electrons

B Si - It has fewer core electrons

C O - It has fewer core electrons

D Ne - it has a higher nuclear charge (Z)

Slide 59 (Answer) / 113

36 Which of the following atoms would have a smaller atomic radii than Ar and why?

A Fe - It has more core electrons

B Si - It has fewer core electrons

C O - It has fewer core electrons

D Ne - it has a higher nuclear charge (Z)

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C

Slide 60 / 113

37 Two elements are studied. One with atomic number X and one with atomic number X+1. Assuming element X is not a Noble Gas, which element has the larger atomic radius?

A Element X

B Element X+1

C They are both the same

D More information is needed

Slide 60 (Answer) / 113

37 Two elements are studied. One with atomic number X and one with atomic number X+1. Assuming element X is not a Noble Gas, which element has the larger atomic radius?

A Element X

B Element X+1

C They are both the same

D More information is needed

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A

Slide 61 / 113

38 Two elements are studied. One with atomic number X and one with atomic number X+1. It is known that element X is a Noble Gas. Which element has the larger atomic radius?

A Element X

B Element X+1

C They are both the same

D More information is needed

Slide 61 (Answer) / 113

38 Two elements are studied. One with atomic number X and one with atomic number X+1. It is known that element X is a Noble Gas. Which element has the larger atomic radius?

A Element X

B Element X+1

C They are both the same

D More information is needed

[This object is a pull tab]A

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B

Slide 62 / 113

Ionic RadiiWhen electrons are gained or lost, the effect on the radii can be

dramatic or slight but there are some certainties.

If an atom loses electrons, the radii will

decrease. Why?

Ca --> Ca2+ + 2e-

194 pm 99 pm

When electrons are lost, the remaining electrons feel a

stronger coulombic attraction from the nucleus.

If an atom gains electrons, the radii will

increase. Why?

F + e- --> F-

42 pm 136 pm

When electrons are gained, the nuclear charge is spread

over a larger number of electrons, resulting in a

weaker coulombic attraction.

Answer Answer

Slide 63 / 113

Ionic RadiiLet's rank a series of atoms and ions in order of increasing radii.

Al3+ Al Mg Mg2+

Whenever comparing radii, use the following procedure:

1. Determine the energy level of the atom/ion.

2. For atoms in the same energy level, use the nuclear charge (Z) to determine the radii.

Al3+ Al Mg Mg2+

Energy Level 2 3 3 2

"Z" 13 13 12 12

Al3+ < Mg2+ < Al < Mg

radius (pm) 50 < 65 < 118 < 145

Slide 64 / 113

In this case, Na+, Mg2+, Al3+, O2-, and F- are all isoelectronic with Ne. As a result, they all experience the same core shielding.

The ionic radii then decreases with an increasing nuclear charge.Al3+ < Mg2+ < Na+ < F- < O2-

Z = 13 12 11 9 8

Ionic RadiiBelow is an example of an isoelectronic series. In an isoelectronic

series the atoms/ions have the same number of electrons.

Slide 65 / 113

39 Which of the following influences the atomic/ionic radii?

A the number of neutrons

B the amount of core electrons between the nucleus and the valence electrons

C the number of protons

D B and C

Slide 65 (Answer) / 113

39 Which of the following influences the atomic/ionic radii?

A the number of neutrons

B the amount of core electrons between the nucleus and the valence electrons

C the number of protons

D B and C

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D

Slide 66 / 113

40 Which ion below has the largest radius?

A O2-

B Li+

C I-

D N3-

Slide 66 (Answer) / 113

40 Which ion below has the largest radius?

A O2-

B Li+

C I-

D N3-

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C

Slide 67 / 113

41 Which of the following pairs correctly shows the proper relationship between the two atoms/ions in terms of atomic/ionic radii?

A Ca < Ca2+

B F < F-

C V < Mn

D Ca < Be

Slide 67 (Answer) / 113

41 Which of the following pairs correctly shows the proper relationship between the two atoms/ions in terms of atomic/ionic radii?

A Ca < Ca2+

B F < F-

C V < Mn

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B

Slide 68 / 113

42 Which of the following would correctly rank the following in order of decreasing atomic/ionic radii?

A V4+ > V5+ > F > F-

B V4+ > V5+ > F- > F

C V5+ > V4+ > F- > F

D V5+ > V4+ > F > F-

Slide 68 (Answer) / 113

42 Which of the following would correctly rank the following in order of decreasing atomic/ionic radii?

A V4+ > V5+ > F > F-

B V4+ > V5+ > F- > F

C V5+ > V4+ > F- > F

D V5+ > V4+ > F > F-

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B

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43 Isotopes of an element, like C-12 and C-13, are likely to have different atomic radii?

Yes

No

Slide 69 (Answer) / 113

43 Isotopes of an element, like C-12 and C-13, are likely to have different atomic radii?

Yes

No

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NO

Slide 70 / 113

Electronegativity

Since the electromagnetic force decreases rapidly as distance increases, we could say that the distance between charges affects an atoms ability to hold onto electrons more than the actual number of protons in the nucleus.

Applying Coulomb's Law:

FE # 1/ r2 FE # qnucleus

Of course, if the distance doesn't change much and the number of protons does, then the electrons will still be held to the atom tightly.

Slide 71 / 113

Electronegativity Trends

Electronegativity is a measure of the ability of atoms in a molecule to attract electrons to themselves.

On the periodic chart, electronegativity increases as you go…

from left to right across a row

from the bottom to the top of a column

Slide 72 / 113

Slide 73 / 113

Electronegativity

There are two notable exceptions to this.

Electronegativity is very closely related to atomic size.

Generally, as atomic size decreases, electronegativity increases.

Can you explain why there is a direct relationship between atomic size and electronegativity?

Slide 73 (Answer) / 113

Electronegativity

There are two notable exceptions to this.

Electronegativity is very closely related to atomic size.

Generally, as atomic size decreases, electronegativity increases.

Can you explain why there is a direct relationship between atomic size and electronegativity?

[This object is a pull tab]

Ans

werThe smaller the radius, the greater

the electrostatic force on the valence electrons

Slide 74 / 113

Electronegativity Exception #1

The Noble Gases are some of the smallest atoms, but are usually left out of electronegativity trends since they neither want electrons nor want to get rid of electrons.

Using your knowledge of electron configurations, why do you think noble gases are left out of electronegativity trends?

Slide 74 (Answer) / 113

Electronegativity Exception #1

The Noble Gases are some of the smallest atoms, but are usually left out of electronegativity trends since they neither want electrons nor want to get rid of electrons.

Using your knowledge of electron configurations, why do you think noble gases are left out of electronegativity trends?

[This object is a pull tab]

Ans

wer Because noble gases have a full

shell they are stable. Gaining or losing an electron would decrease

the stability of a noble gas atom

Slide 75 / 113

The Transition Metals have some unexpected trends in electronegativity because of their d and sometimes f orbitals.

Electronegativity Exception #2

The electrons located in the 3d orbitals (and all d and f orbitals after that) do not contribute as much to the shielding constants of the elements as electrons in the s and p orbitals.

As such, elements with configurations that end in a d or f orbital will frequently have atomic radii that do not match up with the normal trend.

Slide 76 / 113

44 Electronegativity __________ from left to right within a period and __________ from top to bottom within a group.

A decreases, increases

B increases, increases

C increases, decreases

D decreases, decreases

Slide 76 (Answer) / 113

44 Electronegativity __________ from left to right within a period and __________ from top to bottom within a group.

A decreases, increases

B increases, increases

C increases, decreases

D decreases, decreases

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C

Slide 77 / 113

45 Which of the following correctly ranks the elements from highest to lowest electronegativity?

A Cl > S > P

B Br > Cl > FC K > Na > Li

D N > O > F

Slide 77 (Answer) / 113

45 Which of the following correctly ranks the elements from highest to lowest electronegativity?

A Cl > S > P

B Br > Cl > FC K > Na > Li

D N > O > F

[This object is a pull tab]

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A

Slide 78 / 113

46 Which of the following BEST explains why fluorine has a higher electronegativity than oxygen?

A F has a higher nuclear charge and less shielding than O

B F has a higher nuclear charge and similar shielding of O

C F has the equivalent nuclear charge and less shielding than O

D F has the equivalent nuclear charge and more shielding than O

Slide 78 (Answer) / 113

46 Which of the following BEST explains why fluorine has a higher electronegativity than oxygen?

A F has a higher nuclear charge and less shielding than O

B F has a higher nuclear charge and similar shielding of O

C F has the equivalent nuclear charge and less shielding than O

D F has the equivalent nuclear charge and more shielding than O

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B

Slide 79 / 113

47 An element with a small electronegativity value is likely to have...

A Valence shell PES peaks with high binding energies

B A high nuclear charge and a low amount of shielding

C A low nuclear charge and a high amount of shielding

D Both A and C

Slide 79 (Answer) / 113

47 An element with a small electronegativity value is likely to have...

A Valence shell PES peaks with high binding energies

B A high nuclear charge and a low amount of shielding

C A low nuclear charge and a high amount of shielding

D Both A and C[This object is a pull tab]

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C

Slide 80 / 113

Ionization Energy

Ca Ca+ + e- Ca+ Ca2+ + e-

The ionization energy is the amount of energy required to remove an electron from the ground state of a gaseous atom or ion.

The first ionization energy is the energy required to remove the first electron.

The second ionization energy is the energy required to remove the second electron, etc.

1e-

1e-

Slide 81 / 113

Trends in First Ionization Energies

How is ionization energy related to electronegativity and Z eff?

Click here for an animation onIonization Energy

Slide 81 (Answer) / 113

Trends in First Ionization Energies

How is ionization energy related to electronegativity and Z eff?

Click here for an animation onIonization Energy

[This object is a pull tab]

Ans

wer In general, ionization energy

increases with increasing electronegativity and effective

nuclear charge.

Slide 82 / 113

Discontinuity #1The first is between Groups 2 and 13 (3A).As you can see on the chart to the right, the ionization energy actually decreases from Group 2 to Group 13 elements. The electron removed for Group 13 elements is from a p orbital and removing this electron actually adds stability.

The electron removed is farther from nucleus, there is a small amount of repulsion by the s electrons.

The atom gains stability by having a full s orbital, and an empty p orbital.

Slide 83 / 113

Discontinuity #2The second is between Groups 15 and 16.

Using your knowledge of electron configurations and the stability of atoms explain why the first ionization energy for a Group 16 element would be less than that for a Group 15 element in the same period.

Slide 83 (Answer) / 113

Discontinuity #2The second is between Groups 15 and 16.

Using your knowledge of electron configurations and the stability of atoms explain why the first ionization energy for a Group 16 element would be less than that for a Group 15 element in the same period.

[This object is a pull tab]

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wer Group 15 is more stable because it

has a half-full sublevel

Slide 84 / 113

48 What is the ionization energy?

A Energy change associated with the gain of an electron

B Measure of the attraction of an atom for electrons when in a compound

C Pull of the neutrons on the electrons

D Amount of energy required to remove an electron from an atom or ion

Slide 84 (Answer) / 113

48 What is the ionization energy?

A Energy change associated with the gain of an electron

B Measure of the attraction of an atom for electrons when in a compound

C Pull of the neutrons on the electrons

D Amount of energy required to remove an electron from an atom or ion

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D

Slide 85 / 113

49 Which of the following would NOT influence the ionization energy?

A The shielding from core electrons

B The extent to which an orbital is full

C The nuclear charge

D The number of principal energy levels between the valence electrons and the nucleus

E All of these influence the ionization energy

Slide 85 (Answer) / 113

49 Which of the following would NOT influence the ionization energy?

A The shielding from core electrons

B The extent to which an orbital is full

C The nuclear charge

D The number of principal energy levels between the valence electrons and the nucleus

E All of these influence the ionization energy

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E

Slide 86 / 113

50 Which of the following elements would be expected to have a higher ionization energy than magnesium (Mg)?

A Al

B Ca

C Na

D B

Slide 86 (Answer) / 113

50 Which of the following elements would be expected to have a higher ionization energy than magnesium (Mg)?

A Al

B Ca

C Na

D B

[This object is a pull tab]

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D

Slide 87 / 113

51 Which of the following correctly ranks the elements below in order of decreasing ionization energy?

A Ne > O > N

B Ne > N > O

C H > He > Ne

D Li > Mg > Ga

Slide 87 (Answer) / 113

51 Which of the following correctly ranks the elements below in order of decreasing ionization energy?

A Ne > O > N

B Ne > N > O

C H > He > Ne

D Li > Mg > Ga

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B

Slide 88 / 113

52 Which of the following pairs are correct in terms of relative first ionization energy and why?

A O2- < Ne , due to smaller nuclear charge on oxide ion

B Li > Na , due to increased shielding in the Na atom

C Zn > Cu , due to a higher nuclear charge in zinc

D Cl > S , due to the smaller nuclear charge in sulfur

E All of these

Slide 88 (Answer) / 113

52 Which of the following pairs are correct in terms of relative first ionization energy and why?

A O2- < Ne , due to smaller nuclear charge on oxide ion

B Li > Na , due to increased shielding in the Na atom

C Zn > Cu , due to a higher nuclear charge in zinc

D Cl > S , due to the smaller nuclear charge in sulfur

E All of these [This object is a pull tab]

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E

Slide 89 / 113

The Periodic Law and Ionization Energy

Unless you're hydrogen, you've got multiple electrons that can be lost. As a result we have to distinguish between 1st, 2nd, 3rd, etc.

ionization energies.

Ionization Ionization Energy

1st: Na + IE --> Na+ + e- 496 kJ/mol

2nd: Na+ + IE --> Na2+ + e- 4560 kJ/mol3rd: Na2+ + IE --> Na3+ + e- 6,900 kJ/mol

4th: Na3+ + IE --> Na4+ + e- 9540 kJ/mol

Note the huge jump in ionization energy from the 1st to the 2nd. After sodium loses it's first electron, it is isoelectronic with [Ne], with

an extremely stable full s and p orbital and minimal shielding.

Each successive ionization energy is always higher than the previous. This is due to the higher nuclear charge felt by the

remaining electrons.

Slide 90 / 113

53 Which of the following elements best fits the data provided below?

A Li

B C

C Be

D Ne

Ionization Ionization Energy

1st: X + IE --> X+ + e- 900 kJ/mol

2nd: X+ + IE --> X2+ + e- 1757 kJ/mol3rd: X2+ + IE --> X3+ + e- 14,850 kJ/mol

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Slide 91 / 113

54 An atom has the following values for its first four ionization energies. Which of the following elements would fit this data?

A Na

B Mg

C SiD F

1st IE = 345 kJ2nd IE = 456 kJ3rd IE = 3,400 kJ4th IE = 3,876 kJ

Slide 91 (Answer) / 113

54 An atom has the following values for its first four ionization energies. Which of the following elements would fit this data?

A Na

B Mg

C SiD F

1st IE = 345 kJ2nd IE = 456 kJ3rd IE = 3,400 kJ4th IE = 3,876 kJ

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B

Slide 92 / 113

Ionization Energy and PESIonization energy data can be determined from PES (photoelectron spectroscopy). Recall that PES looks at the energy required to remove electrons from an atom. Each orbital appears as a peak on the spectrum.

The PES spectrum clearly shows that the core electrons require the most energy to remove. It also shows that Be has a higher 1st IE for the removal of the valence electrons than does Li. This is expected as Be has a higher "Z".

Li (1s)

Be (1s)Be (2s)

Li (2s)

Intensity

binding energy

Slide 93 / 113

PES PracticeLet's interpret another PES spectra, this one of nitrogen and oxygen.

Intensity

binding energy

N (2s) N (1s)N (2p)

O (2p)O (2s) O (1s)

Why is the N (2p) peak greater than the O (2p) peak?

Why is the N(2s) peak less than the O (2s) peak?

Slide 93 (Answer) / 113

PES PracticeLet's interpret another PES spectra, this one of nitrogen and oxygen.

Intensity

binding energy

N (2s) N (1s)N (2p)

O (2p)O (2s) O (1s)

Why is the N (2p) peak greater than the O (2p) peak?

Why is the N(2s) peak less than the O (2s) peak? [This object is a pull tab]

Ans

wer N has a half-full "p" orbital

increasing the ionization energy

O has the higher nuclear charge

Slide 94 / 113

Ionization Energy and PES

Click to go to an interactive PES spectra database

and answer the questions.

1. Why is the binding energy of the electrons greater in He than H?

2. Which peak in the Li spectra represents the valence electrons?

3. Why is the valence peak binding energy less in Li than in H?

4. Why is the core peak (1s) binding energy greater in Li than in H?

Slide 94 (Answer) / 113

Ionization Energy and PES

Click to go to an interactive PES spectra database

and answer the questions.

1. Why is the binding energy of the electrons greater in He than H?

2. Which peak in the Li spectra represents the valence electrons?

3. Why is the valence peak binding energy less in Li than in H?

4. Why is the core peak (1s) binding energy greater in Li than in H? [This object is a pull tab]

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1. Similar shielding but greater "Z"

2. Peak with lower binding energy

3. Increased shielding due to core 1s electrons, lessens coulombic force

4. Lithium has a higher nuclear charge "Z" so higher coulombic attractions

Slide 95 / 113

55 The following PES spectrum shows the valence "p" orbital peaks for Si and for C. Which of the following would be TRUE?

A The Si peak is of lower energy due to it's higher nuclear charge

B The Si peak is of higher energy due to the increased shielding from core electrons

C The Si peak is of lower energy due to the increased shielding from core electrons

D The Si peak is of higher energy due to its higher nuclear charge

Intensity

binding energy

Slide 95 (Answer) / 113

55 The following PES spectrum shows the valence "p" orbital peaks for Si and for C. Which of the following would be TRUE?

A The Si peak is of lower energy due to it's higher nuclear charge

B The Si peak is of higher energy due to the increased shielding from core electrons

C The Si peak is of lower energy due to the increased shielding from core electrons

D The Si peak is of higher energy due to its higher nuclear charge

Intensity

binding energy

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C

Slide 96 / 113

56 The 3s peak for magnesium should have a higher binding energy than that of the 4s peak in calcium due to calcium's higher amount of shielding by core electrons?

True

False

Slide 96 (Answer) / 113

56 The 3s peak for magnesium should have a higher binding energy than that of the 4s peak in calcium due to calcium's higher amount of shielding by core electrons?

True

False

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TRUE

Slide 97 / 113

57 Below is an actual PES spectrum of palladium (Pd). Which of the following would be TRUE? (Note: the outer 5s and 4d peaks are not shown)

A Compared to Pd, the 3d peak in Cd would be found to the left of the 3d Pd peak

B Compared to Pd, the 3d peak in Rb would be of a higher binding energy due to lower nuclear charge

C Compared to Pd, the 3p peak in Kr should be found to the left of the 3p peak in Pd

3s 3p

3d

4p 4s

Slide 97 (Answer) / 113

57 Below is an actual PES spectrum of palladium (Pd). Which of the following would be TRUE? (Note: the outer 5s and 4d peaks are not shown)

A Compared to Pd, the 3d peak in Cd would be found to the left of the 3d Pd peak

B Compared to Pd, the 3d peak in Rb would be of a higher binding energy due to lower nuclear charge

C Compared to Pd, the 3p peak in Kr should be found to the left of the 3p peak in Pd

3s 3p

3d

4p 4s

[This object is a pull tab]A

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A

Slide 98 / 113

58 Based on the PES data below, what would be TRUE regarding atoms 1 and 2?

A I only B II and III only C 1 and III only

D II and IV only

Binding Energy

Inte

nsity

0 10

10 100

28.6

1.091.72

Binding Energy

Inte

nsity

0 10

10 100

39.6

1.402.45

1 2

I. Atom 1 has a smaller atomic radii II. Atom 2 has a larger first ionization energy

III. Both atoms are in the same period

IV. Both atoms are in the same group

Slide 98 (Answer) / 113

58 Based on the PES data below, what would be TRUE regarding atoms 1 and 2?

A I only B II and III only C 1 and III only

D II and IV only

Binding Energy

Inte

nsity

0 10

10 100

28.6

1.091.72

Binding Energy

Inte

nsity

0 10

10 100

39.6

1.402.45

1 2

I. Atom 1 has a smaller atomic radii II. Atom 2 has a larger first ionization energy

III. Both atoms are in the same period

IV. Both atoms are in the same group[This object is a pull tab]

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B

Slide 99 / 113

Ionization Energy and Metallic CharacterMetals are generally described as being able to lose

electrons readily which promotes conductivity.

Since metals lose electrons easily, they must have low ionization energies compared to non-metals.

Element Metal or Non-metal1st Ionization

Energy (kJ/mol)Na metal 496

O non-metal 1314

Slide 100 / 113

Ionization Energy and Metallic CharacterWe can predict, based on ionization energies, where the

metals and non-metals are on the periodic table.

semi-metals or metalloids

Notice that an element becomes more metallic as the shielding increases and as the nuclear charge - for a given

level of shielding - decreases.

Slide 101 / 113

59 Which of the following has the elements correctly ordered by increasing metallic character?

A Li < Be < B

B Ca < K < Ga

C Ga < Ca < K

D Rb < Cs < As

Slide 101 (Answer) / 113

59 Which of the following has the elements correctly ordered by increasing metallic character?

A Li < Be < B

B Ca < K < Ga

C Ga < Ca < K

D Rb < Cs < As[This object is a pull tab]

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C

Slide 102 / 113

60 Which of the following ranks the metals in order of increasing reactivity?

A Li < Na < Mg < K

B Mg < Li < Na < K

C K < Li < Na < K

D Li < Fe < Zn < Au

Slide 102 (Answer) / 113

60 Which of the following ranks the metals in order of increasing reactivity?

A Li < Na < Mg < K

B Mg < Li < Na < K

C K < Li < Na < K

D Li < Fe < Zn < Au

[This object is a pull tab]

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B

Slide 103 / 113

61 Which of the following would be TRUE?

A The higher the ionization energy, the less metallic an element will be

B The lower the ionization energy, the less metallic an element will be

C For a given amount of core electron shielding, the higher the nuclear charge, the more metallic an element will be

D Both A and C

Slide 103 (Answer) / 113

61 Which of the following would be TRUE?

A The higher the ionization energy, the less metallic an element will be

B The lower the ionization energy, the less metallic an element will be

C For a given amount of core electron shielding, the higher the nuclear charge, the more metallic an element will be

D Both A and C

[This object is a pull tab]

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A

Slide 104 / 113

62 Why is lead considered a metal and carbon a non-metal despite being in the same group?

A Lead has a greater S value

B Lead has a greater Zeff

C Lead is more electronegative

D All of the above

Slide 104 (Answer) / 113

62 Why is lead considered a metal and carbon a non-metal despite being in the same group?

A Lead has a greater S value

B Lead has a greater Zeff

C Lead is more electronegative

D All of the above[This object is a pull tab]

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A

Slide 105 / 113

The charges of ions is also periodic in nature, as seen on the graph below. The formation of ions depends on ionization energy and electronegativity.

ion charge

+1

+2

+3

-1

-2

-3

atomic number 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18

+4

Ionic Charge

Slide 106 / 113

Ionic Charge This trend in ionic charge can be easily explained if we apply the

quantum model of the atom.

Element

Principal Quantum

Number (N) of valence electrons

Electron Configuration

Lose/Gain

electrons

Ionic Charge

H 1 1s1 gain 1

lose 1

-1

+1He 1 1s2 NA NA Li 2 [He]2s1 lose 1 +1Be 2 [He]2s2 lose 2 +2

B 2 [He]2s22p1 lose 3 +3

C 2 [He]2s22p2 lose 4 +4

N 2 [He]2s22p3 gain 3 -3

O 2 [He]2s22p4 gain 2 -2 F 2 [He]2s22p5 gain 1 -1Ne 2 [He]2s22p6 NA NA

Na 3 [Ne]3s1 lose 1 +1

The pattern recurs with every increase in the

principal quantum number. Atoms lose or gain electrons to obtain a full shell or subshell, thereby increasing their

stability.

Slide 107 / 113

63 Which of the following BEST explains why O and S both form ions with a -2 charge?

A They both have the same atomic number

B They are both in the same period

C They both have the same electron configuration

D They both have the same number of valence electrons

Slide 107 (Answer) / 113

63 Which of the following BEST explains why O and S both form ions with a -2 charge?

A They both have the same atomic number

B They are both in the same period

C They both have the same electron configuration

D They both have the same number of valence electrons

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D

Slide 108 / 113

64 An atom with the electron configuration of [Kr]5s24d2 would be in the same group as _____ and have a likely charge of ____.

A Sc, +1

B Hf, +4

C Ti, +3

D Zn, +2

Slide 108 (Answer) / 113

64 An atom with the electron configuration of [Kr]5s24d2 would be in the same group as _____ and have a likely charge of ____.

A Sc, +1

B Hf, +4

C Ti, +3

D Zn, +2

[This object is a pull tab]A

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B

Slide 109 / 113

65 Atoms on the right side of the chart tend to form negative ions because...

A Their principal energy level is almost empty

B Their principal energy level is almost full

C Their atomic number is less than other elements in that period

D Both B and C

Slide 109 (Answer) / 113

65 Atoms on the right side of the chart tend to form negative ions because...

A Their principal energy level is almost empty

B Their principal energy level is almost full

C Their atomic number is less than other elements in that period

D Both B and C

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B

Slide 110 / 113

Let's use quantum theory to explain the trends we see amongst the charges of the transition elements.

Question 1: Elements within the Fe group can form ions of both +2 and +3 charges. Explain why the +3 charge is more common:

Question 2: Why do the elements in the zinc group tend to only form ions with a +2 charge?

Transition Metal Ions Fe = [Ar]4s23d6

The 4s electrons are readily lost yielding the +2 ion.

A half-full "d" orbital is quite stable so Fe will lose 1 d orbital electron as well to yield the +3 ion.

Slide 110 (Answer) / 113

Let's use quantum theory to explain the trends we see amongst the charges of the transition elements.

Question 1: Elements within the Fe group can form ions of both +2 and +3 charges. Explain why the +3 charge is more common:

Question 2: Why do the elements in the zinc group tend to only form ions with a +2 charge?

Transition Metal Ions

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2[This object is a pull tab]

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1

Fe = [Ar]4s23d6

The 4s electrons are readily lost yielding the +2 ion.

A half-full "d" orbital is quite stable so Fe will lose 1 d orbital electron as well to yield the +3 ion.

Slide 111 / 113

66 What is/are the possible charge(s) on a chromium ion?

Slide 112 / 113

67 What is/are the possible charge(s) on a copper ion?

Slide 113 / 113

68 What is/are the possible charge(s) on a manganese ion?