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AP Chemistry

The Atom

2015-08-25 www.njctl.org

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Table of Contents: The Atom (Pt. B)

· Periodic Table · Periodic Trends

Click on the topic to go to that section

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Periodic Table

Return to Table

• f Contents

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

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

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3 To which group on the periodic table does Neon

belong?

A

Alkali Metals

B

Transition Metals

C

Noble Gases

D

Alkaline Earth Metals

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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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5 To which group on the periodic table does

Beryllium belong?

A

Alkali Metals

B

Transition Metals

C

Halogens

D

Alkaline Earth Metals

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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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1A 2A 8A 1 2 18 3A 4A 5A 6A 7A 13 14 15 16 17 8B 3B 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

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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 block B d block C p block D f block

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8 The highlighted elements below are in the

A s block B d block C p block D f block

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9 The highlighted elements below are in the

A s block B d block C p block D f block

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

A Alkali Metals B Alkaline Earth Metals C Halogens D Noble Gases

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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 20 / 113 Shorthand Configurations

Noble 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

• rbital of the row that the element is in until the

configuration is complete.

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

A Alkali Metals B Alkaline Earth Metals C Transition Metals D Halogens

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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 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 Stability

Elements 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 ( d

5, f7)

1 2 3 4 5 6 7 6 7

Slide 27 / 113 Stability

Next in order of stability are elements with full sublevels. ~Full Energy Level ~Full Sublevel (s, p, d, f) ~Half Full Sublevel ( d

5, f7)

1 2 3 4 5 6 7 6 7

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Finally, 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) 1 2 3 4 5 6 7 6 7

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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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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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18 The elements in the periodic table which lack

• ne electron from a filled shell are referred to

as: A

noble gases.

B

halogens.

C

alkali metals.

D

transition elements.

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. 1 2 3 4 5 6 7 6 7

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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

1 2 3 4 5 6 7 6 7

Cr

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1 2 3 4 5 6 7

1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s

4f 5d 6p 5f

7s 6d 7p

6f 7d 7f

Energy

Energies of Orbitals

Because of how close the f and d orbitals are to the s

• rbitals an electron very little

energy is required to move an electron from the s orbital (leaving it half full) to the f or d

• rbital, 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.)

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Copper Expected: [Ar] 4s2 3d9 Actual: [Ar] 4s1 3d10 Copper gains stability when an electron from the 4s

• rbital fills the 3d orbital.

Electron Configuration Exceptions

1 2 3 4 5 6 7 6 7

Cu

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19 The electron configuration for Copper (Cu) is

A [Ar] 4s24d9 B [Ar] 4s14d9 C [Ar] 4s23d9 D [Ar] 4s13d10

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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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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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Periodic Trends

Return to Table

• f Contents

Slide 40 / 113 Periodic Trends

There are four main trends in the periodic table: Size or radius of atoms/ions Electronegativity Ionization energy Metallic 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+

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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.

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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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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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24 In which subshell does an electron in a calcium

atom experience the greatest effective nuclear charge?

A 1s B

2s C 3s D 3p

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

atomic number radius (pm) 200 100 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 H Li Na He Ne 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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25 What is Zeff for Boron (B)?

5+

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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

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27 What is Zeff for Carbon (C)?

6+

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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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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

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

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31 Across a period from left to right

Zeff _____________.

A increases B decreases C remains the same

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32 Down a group from top to bottom

Zeff _____________.

A increases B decreases C remains the same

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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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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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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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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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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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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 62 / 113 Ionic Radii

When 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

• ver a larger number of

electrons, resulting in a weaker coulombic attraction. Answer Answer

Slide 63 / 113 Ionic Radii

Let'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

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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 Radii

Below is an example of an isoelectronic series. In an isoelectronic series the atoms/ions have the same number of electrons.

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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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40 Which ion below has the largest radius? A O2- B Li+ C I- D N3-

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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

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

E # 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

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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 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?

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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

• rbitals (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

• rbital will frequently have atomic radii that do not match up with

the normal trend.

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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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45 Which of the following correctly ranks the elements

from highest to lowest electronegativity?

A Cl > S > P B Br > Cl > F C K > Na > Li D N > O > F

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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

• f 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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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 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 on Ionization Energy

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Discontinuity #1

The 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

• rbital 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 #2

The 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.

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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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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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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

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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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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

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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/mol 3rd: 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.

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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/mol 3rd: X2+ + IE --> X3+ + e- 14,850 kJ/mol

Answer

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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 Si D F 1st IE = 345 kJ 2nd IE = 456 kJ 3rd IE = 3,400 kJ 4th IE = 3,876 kJ

Slide 92 / 113 Ionization Energy and PES

Ionization 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 Practice

Let'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 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?

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55 The following PES spectrum shows the valence "p"

• rbital 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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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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57 Below is an actual PES spectrum of palladium (Pd). Which of the following would be TRUE? (Note: the

• uter 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

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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 Intensity 10 10 100 28.6 1.09 1.72 Binding Energy Intensity 10 10 100 39.6 1.40 2.45

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 99 / 113 Ionization Energy and Metallic Character

Metals 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-metal 1st Ionization Energy (kJ/mol) Na metal 496 O non-metal 1314

Slide 100 / 113 Ionization Energy and Metallic Character

We can predict, based on ionization energies, where the metals and non-metals are on the periodic table. semi-metals

• r metalloids

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

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59 Which of the following has the elements correctly

• rdered by increasing metallic character?

A Li < Be < B B Ca < K < Ga C Ga < Ca < K D Rb < Cs < As

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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

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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

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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

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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)

• f valence

electrons Electron Configuration Lose/ Gain electrons Ionic Charge H 1 1s1 gain 1 lose 1

• 1

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

• 1

Ne 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.

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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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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

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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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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.

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

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

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