Chapter 11a Modern Atomic Theory. Chapter 11 Table of Contents 2 11.1 Rutherford’s Atom 11.2...
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Transcript of Chapter 11a Modern Atomic Theory. Chapter 11 Table of Contents 2 11.1 Rutherford’s Atom 11.2...
Chapter 11a
Modern Atomic Theory
Chapter 11
Table of Contents
2
11.1 Rutherford’s Atom
11.2 Electromagnetic Radiation
11.3 Emission of Energy by Atoms
11.4 The Energy Levels of Hydrogen
11.5 The Bohr Model of the Atom
Section 11.1
Rutherford’s Atom
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3
Nuclear Model of the Atom
• The atom has a small dense nucleus which is positively charged. contains protons (+1 charge). contains neutrons (no charge).
• The remainder of the atom is mostly empty space. contains electrons (–1
charge).
Section 11.1
Rutherford’s Atom
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Copyright © Cengage Learning. All rights reserved 4
• The nuclear charge (n+) is balanced by the presence of n electrons moving in some way around the nucleus.
• What are the electrons doing?• How are the electrons arranged and how do they move?
Section 11.2
Electromagnetic Radiation
Return to TOC
5
Characteristics
• Wavelength ( ) – distance between two peaks or troughs in a wave.
Section 11.2
Electromagnetic Radiation
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6
Different Wavelengths Carry Different Amounts of Energy
Section 11.2
Electromagnetic Radiation
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7
One of the ways that energy travels through space.
Section 11.2
Electromagnetic Radiation
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8
Characteristics
• Frequency ( ) – number of waves (cycles) per second that pass a given point in space
• Speed (c) – speed of light (2.9979×108 m/s)
= c 186,000 miles/s
Section 11.2
Electromagnetic Radiation
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9
Dual Nature of Light
• Wave• Photon – packet of energy
Section 11.2
Electromagnetic Radiation
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10
Characteristics
Ephoton = (hc) / λ
• Energy of a photon of light = Planck’s constant (h) (6.626 x 10-34Js) times the speed of light (c) (2.9979×108 m/s) divided by the wavelength in meters (λ) or the wavelength in nanometers (nm) times ten to the -9 power (550nm = 550 x 10-9m).
OR:
Section 11.2
Electromagnetic Radiation
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11
Let’s Practice!
What is the energy of light with a wavelength of 535 nm?
E=hc/λ
=6.626 x 10-34Js(3.00 x 108m/s)/(535 x 10-9m)
=3.70 x 10-19 J
Section 11.2
Electromagnetic Radiation
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Seeing the Light-A New Model of the Atom
Maxwell Planck-Black Body Radiation
Found that blackbody radiation was quantized.
1900—Nobel Prize in 1918
Section 11.2
Electromagnetic Radiation
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13
Quantized Energy Levels
• The energy levels of all atoms are quantized.
Section 11.2
Electromagnetic Radiation
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Einstein’s Photoelectric Effect (1905--Nobel Prize in 1921)
Only light from a certain color (energy) could eject electrons. Intensity of the light had no effect. Energy is absorbed only at quantized energies!
Section 11.3
Emission of Energy by Atoms
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Copyright © Cengage Learning. All rights reserved 15
Atoms can give off light. They first must receive energy and become excited. The energy is released in the form of a photon. The energy of the photon corresponds exactly to the energy change experienced by the emitting atom.
Section 11.4
The Energy Levels of Hydrogen
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Copyright © Cengage Learning. All rights reserved 16
• Atomic states Excited state – atom with excess energy Ground state – atom in the lowest possible state
• When an H atom absorbs energy from an outside source it enters an excited state.
Section 11.4
The Energy Levels of Hydrogen
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Copyright © Cengage Learning. All rights reserved 17
Energy Level Diagram
• Energy in the photon corresponds to the energy used by the atom to get to the excited state.
Section 11.4
The Energy Levels of Hydrogen
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Stokes Shift-Absorb high energy (UV) and emit low energy (visible).
18
Section 11.4
The Energy Levels of Hydrogen
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19
• Only certain types of photons are produced when H atoms release energy. Why?
Section 11.4
The Energy Levels of Hydrogen
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Line Spectra
http://jersey.uoregon.edu/vlab/elements/Elements.html
Section 11.4
The Energy Levels of Hydrogen
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21
The word laser comes fromlight amplification by stimulated emission of radiation.
Section 11.4
The Energy Levels of Hydrogen
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22
Lasers used to remove blood clots.Laser light transmitted in fiber optics.Cataract RemovalLight Shows
Lasers
Section 11.4
The Energy Levels of Hydrogen
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23
Holograms3D pictures made by Lasers using the interference pattern between reflected laser light from the surface of an object and the undisturbed laser light reflected from a mirror. The Interference pattern is recorded on film. The developed film can then be used by a laser to recreate the image in 3D.
http://www.youtube.com/watch?v=wrxUYzWASvE
http://www.youtube.com/watch?v=E4A_u67EKnU&feature=fvw
https://www.youtube.com/watch?v=AXhGfkGh4vM
http://www.youtube.com/watch?v=cAX8uSc8Fnk&NR=1
Section 11.4
The Energy Levels of Hydrogen
Return to TOC
24
Holograms are made from laser light without using an image forming device. Tupac holographic concert and a holographic fashion display.
http://www.youtube.com/watch?v=mcSYpZchFpI
http://www.youtube.com/watch?v=Zf_eXDPElh0
https://www.youtube.com/watch?v=89KxxpmMhi4
Holograms
Section 11.4
The Energy Levels of Hydrogen
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The Doppler Effect
The doppler effect is the apparent change in frequency of a wave due to the relative motion of the listener and the source of the sound.
The doppler effect also occurs in light waves and is used by astronomers to calculate the speed at which stars are approaching or receding.
Section 11.4
The Energy Levels of Hydrogen
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26
Bohr Model
Line Spectra in Stars and the red shift indicating movement away or towards us.
7 -
Section 11.4
The Energy Levels of Hydrogen
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Copyright © Cengage Learning. All rights reserved 27
Quantized Energy Levels
• Since only certain energy changes occur the H atom must contain discrete energy levels.
Section 11.4
The Energy Levels of Hydrogen
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28
Concept Check
Why is it significant that the color emitted from the hydrogen emission spectrum is not white?
How does the emission spectrum support the idea of quantized energy levels?
Section 11.4
The Energy Levels of Hydrogen
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Copyright © Cengage Learning. All rights reserved 29
Concept Check
When an electron is excited in an atom or ion
a) only specific quantities of energy are released in order for the electron to return to its ground state.
b) white light is never observed when the electron returns to its ground state.
c) the electron is only excited to certain energy levels.
d) All of the above statements are true when an electron is excited.
Section 11.5
The Bohr Model of the Atom
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30
• Quantized energy levels• Electron moves in a circular
orbit.• Electron jumps between
levels by absorbing or emitting a photon of a particular wavelength.
• Actually electrons do not move in a circular orbit.
Niels Bohr1913—Nobel Prize in 1922
Niels Bohr hypothesized that electrons orbit the nucleus just as the planets orbit the sun (planetary model).
Section 11.5
The Bohr Model of the Atom
Return to TOC
31
D
Chapter 11b
Modern Atomic Theory
Section 11.5
The Bohr Model of the Atom
Return to TOC
32
11.6 The Wave Mechanical Model of the Atom
11.7 The Hydrogen Orbitals
11.8 The Wave Mechanical Model: Further Development
11.9 Electron Arrangements in the First Eighteen Atoms on the Periodic Table
11.10 Electron Configurations and the Periodic Table
11.11 Atomic Properties and the Periodic Table
Chapter 11
Modern Model of the Atom
Section 11.6
The Wave Mechanical Model of the Atom
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33
Orbitals
• Nothing like orbits• Probability of finding the electron within a certain space• This model gives no information about when the electron
occupies a certain point in space or how it moves.
Section 11.7
The Hydrogen Orbitals
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Copyright © Cengage Learning. All rights reserved 34
Orbitals
• Orbitals do not have sharp boundaries and are represented by probability distributions or where the electron is likely to be found without regards to movement of the electrons.
• Chemists arbitrarily define an orbital’s size as the sphere that contains 90% of the total electron probability.
Section 11.6
The Wave Mechanical Model of the Atom
Return to TOC
35
Scanning Tunneling Microscope
Kx9GeA7Gfalse
Section 11.6
The Wave Mechanical Model of the Atom
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Louis DeBroglie
• He found that matter (electrons) moved in waves. Just as light behaved like particles and waves, so did matter.
• An 18-wheeler moving down Hwy 99 at 60mph has a wavelength smaller than an atom.
• However, an electron (very light) moves much faster and its wavelength is much larger than its size.
1924 – Nobel Prize in 1929
Section 11.6
The Wave Mechanical Model of the Atom
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Erwin Schrödinger1926 -Nobel Prize in 1933
Found the probability of finding an electron in an atom.
Section 11.6
The Wave Mechanical Model of the Atom
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1926 -Nobel Prize in 1933
Found the probability of finding an electron in an atom.
Erwin Schrödinger
Section 11.7
The Hydrogen Orbitals
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39
Hydrogen Energy Levels
• Hydrogen has discrete energy levels. Called principal energy
levels Labeled with whole
numbers
Section 11.7
The Hydrogen Orbitals
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Copyright © Cengage Learning. All rights reserved 40
Hydrogen Energy Levels
• Each principal energy level is divided into sublevels. Labeled with numbers and letters Indicate the shape of the orbital
Section 11.7
The Hydrogen Orbitals
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Copyright © Cengage Learning. All rights reserved 41
Hydrogen Energy Levels
• The s and p types of sublevel
Section 11.7
The Hydrogen Orbitals
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42
d Orbitals
7 -
Section 11.7
The Hydrogen Orbitals
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43
f-orbtals
http://www.d.umn.edu/~pkiprof/ChemWebV2/AOs/ao4.html
Section 11.7
The Hydrogen Orbitals
Return to TOC
44
Why the different shapes?
1s 3s2s2px
3pz
3d2py
2pz
3py
3px
Section 11.7
The Hydrogen Orbitals
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45
Orbital Labels
1. The number tells the principal energy level.
2. The letter tells the shape. The letter s means a spherical orbital or
shape of the probability distribution of the electron.
The letter p means the orientation. The x, y, or z subscript on a p orbital label tells along which of the coordinate axes the two lobes lie.
Section 11.7
The Hydrogen Orbitals
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46
Hydrogen Orbitals
• Why does an H atom have so many orbitals and only 1 electron? An orbital is a potential space for an electron. Atoms can have many potential orbitals.
Section 11.8
The Wave Mechanical Model: Further Development
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47
Atoms Beyond Hydrogen
• The Bohr model was discarded because it does not apply to all atoms. It did not consider the different energy sublevels or suborbitals within each orbital.
• Atoms beyond hydrogen have multiple electrons that distorts the energy levels due to electron-electron interactions.
• Need one more property to determine how the electrons are arranged: Spin – electrons spin like a top causing a magnetic field. Opposite magnetic fields can attract allowing electrons to occur in pairs if their spin or magnetic field is opposite.
Section 11.8
The Wave Mechanical Model: Further Development
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48
Atoms Beyond Hydrogen
• Pauli Exclusion Principle – an atomic orbital can hold a maximum of 2 electrons and those 2 electrons must have opposite spins.
Section 11.8
The Wave Mechanical Model: Further Development
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Copyright © Cengage Learning. All rights reserved 49
Principal Components of the Wave Mechanical Model of the Atom
1. Atoms have a series of energy levels called principal energy levels (n = 1, 2, 3, etc.).
2. The energy of the level increases as the value of n increases.
3. Each principal energy level contains one or more types of orbitals, called sublevels.
4. The number of sublevels present in a given principal energy level equals n.
Section 11.8
The Wave Mechanical Model: Further Development
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Copyright © Cengage Learning. All rights reserved 50
Principal Components of the Wave Mechanical Model of the Atom
5. The n value is always used to label the orbitals of a given principal level and is followed by a letter that indicates the type (shape) of the orbital (1s, 3p, etc.).
6. An orbital can be empty or it can contain one or two electrons, but never more than two. If two electrons occupy the same orbital, they must have opposite spins.
Section 11.8
The Wave Mechanical Model: Further Development
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51
Principal Components of the Wave Mechanical Model of the Atom
7. The shape of an orbital does not indicate the details of electron movement. It indicates the probability distribution for an electron residing in that orbital.
Section 11.8
The Wave Mechanical Model: Further Development
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Copyright © Cengage Learning. All rights reserved 52
Concept Check
Which of the following statements best describes the movement of electrons in a p orbital?
a) The electron movement cannot be exactly determined.
b) The electrons move within the two lobes of the p orbital, but never beyond the outside surface of
the orbital.
c) The electrons are concentrated at the center (node) of the two lobes.
d) The electrons move along the outer surface of the p orbital, similar to a “figure 8” type of movement.
Section 11.8
The Wave Mechanical Model: Further Development
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53
Energy Level Diagram for
Carbon
Section 11.9
Electron Arrangements in the First Eighteen Atoms on the Periodic Table
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Copyright © Cengage Learning. All rights reserved 54
H Atom
• Electron configuration – electron arrangement 1s1
• Orbital diagram – orbital is a box grouped by sublevel containing arrow(s) to represent electrons
Section 11.9
Electron Arrangements in the First Eighteen Atoms on the Periodic Table
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Copyright © Cengage Learning. All rights reserved 55
Li Atom
• Electron configuration
1s2 2s1
• Orbital diagram
Section 11.9
Electron Arrangements in the First Eighteen Atoms on the Periodic Table
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Copyright © Cengage Learning. All rights reserved 56
O Atom
• The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons in a particular set of degenerate (same energy) orbitals.
Oxygen: 1s 2s 2p
Section 11.9
Electron Arrangements in the First Eighteen Atoms on the Periodic Table
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Copyright © Cengage Learning. All rights reserved 57
• The electron configurations in the sublevel last occupied for the first eighteen elements.
Section 11.9
Electron Arrangements in the First Eighteen Atoms on the Periodic Table
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58
Classifying Electrons
• Core electrons – inner electrons• Valence electrons – electrons in the outermost (highest)
principal energy level of an atom 1s22s22p6 (valence electrons = 8) The elements in the same group on the periodic table
have the same valence electron configuration. Elements with the same valence electron
arrangement show very similar chemical behavior.
Section 11.9
Electron Arrangements in the First Eighteen Atoms on the Periodic Table
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Copyright © Cengage Learning. All rights reserved 59
Concept Check
How many unpaired electrons does the element cobalt (Co) have in its lowest energy state?
a) 0
b) 2
c) 3
d) 7
3d suborbitals
Section 11.9
Electron Arrangements in the First Eighteen Atoms on the Periodic Table
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Copyright © Cengage Learning. All rights reserved 60
Concept Check
Can an electron in a phosphorus atom ever be in a 3d orbital? Choose the best answer.
a) Yes. An electron can be excited into a 3d orbital.
b) Yes. A ground-state electron in phosphorus is located in a 3d orbital.
c) No. Only transition metal atoms can have electrons located in the d orbitals.
d) No. This would not correspond to phosphorus’ electron arrangement in its ground state.
Section 11.9
Electron Arrangements in the First Eighteen Atoms on the Periodic Table
Return to TOC
61
Quantum #’s are like an Address.
What do you need to know to find out where you live?
State City Street House
Principle Quantum # (n)
Angular Quantum # (l)
Magnetic Quantum # (ml)
Spin Quantum # (ms)
Section 11.10
Electron Configurations and the Periodic Table
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Copyright © Cengage Learning. All rights reserved 62
• Look at electron configurations for K through Kr.
Section 11.10
Electron Configurations and the Periodic Table
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Copyright © Cengage Learning. All rights reserved 63
Orbital Filling and the Periodic Table
Section 11.10
Electron Configurations and the Periodic Table
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Copyright © Cengage Learning. All rights reserved 64
Orbital Filling
1. In a principal energy level that has d orbitals, the s orbital from the next level fills before the d orbitals in the current level.
2. After lanthanum, which has the electron configuration [Xe]6s25d1, a group of fourteen elements called the lanthanide series, or the lanthanides, occurs. This series of elements corresponds to the filling of the seven 4f orbitals.
Section 11.10
Electron Configurations and the Periodic Table
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65
Orbital Filling
3. After actinum, which has the configuration [Rn]7s26d1,a group of fourteen elements called the actinide series, or actinides, occurs. This series corresponds to the filling of the seven 5f orbitals.
Section 11.10
Electron Configurations and the Periodic Table
Return to TOC
66
Orbital Filling
4. Except for helium, the group numbers indicate the sum of electrons in the ns and np orbitals in the highest principal energy level that contains electrons (where n is the number that indicates a particular principal energy level). These electrons are the valence electrons.
Section 11.10
Electron Configurations and the Periodic Table
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Copyright © Cengage Learning. All rights reserved 67
Exercise
Determine the expected electron configurations for each of the following.
a) S
1s22s22p63s23p4 or [Ne]3s23p4
b) Ba
[Xe]6s2
c) Eu
[Xe]6s24f7
Section 11.10
Electron Configurations and the Periodic Table
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Copyright © Cengage Learning. All rights reserved 68
• Write electron configurations for the following:
1. Al
2. Sc
3. K
4. Br
5. Zn
6. Hg
7 -
1s22s22p63s23p1
1s22s22p63s23p64s23d1
1s22s22p63s23p64s1
1s22s22p63s23p64s23d104p5
1s22s22p63s23p64s23d10
1s22s22p63s23p64s23d104p65s24d105p66s2 4f145d10
Section 11.10
Electron Configurations and the Periodic Table
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Copyright © Cengage Learning. All rights reserved 69
• Write the abbreviated electron configuration for the following:
1. Magnesium
2. Carbon
3. Boron
4. Chlorine
5. Selenium
7 -
– [Ne] 3s23p5
– [Ne] 3s2
– [He] 2s22p2
– [He] 2s22p1
– [Ar] 4s23d104p4
Section 11.10
Electron Configurations and the Periodic Table
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Copyright © Cengage Learning. All rights reserved 70
• Write the electron configuration in long and abbreviated notation for the following ions.1. Br-
2. N3-
3. K+
4. Sr2+
5. S2-
6. Ni2+
7 -
[Kr] isoelectronic with Kr
[Ne] isoelectronic with Ne
[Ar] isoelectronic with Ar[Kr] isoelectronic with Kr[Ar] isoelectronic with Ar
[Ar]4s23d6 isoelectronic with Fe
Section 11.10
Electron Configurations and the Periodic Table
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71
Section 11.11
Atomic Properties and the Periodic Table
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Copyright © Cengage Learning. All rights reserved 72
Metals and Nonmetals
• Metals tend to lose electrons to form positive ions.• Nonmetals tend to gain electrons to form negative ions.
Section 11.11
Atomic Properties and the Periodic Table
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Copyright © Cengage Learning. All rights reserved 73
Ionization Energy
• Energy required to remove an electron from a gaseous atom or ion. X(g) → X+(g) + e–
Mg → Mg+ + e– I1 = 735 kJ/mol (1st IE)
Mg+ → Mg2+ + e– I2 = 1445 kJ/mol (2nd IE)
Mg2+ → Mg3+ + e– I3 = 7730 kJ/mol *(3rd IE)
*Core electrons are bound much more tightly than valence electrons.
Section 11.11
Atomic Properties and the Periodic Table
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Copyright © Cengage Learning. All rights reserved 74
Ionization Energy
• In general, as we go across a period from left to right, the first ionization energy increases.
• Why? Electrons added in the same principal
quantum level do not completely shield the increasing nuclear charge caused by the added protons.
Electrons in the same principal quantum level are generally more strongly bound from left to right on the periodic table.
Section 11.11
Atomic Properties and the Periodic Table
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75
Ionization Energy
• In general, as we go across a period the ionization energy increases.
• As we go up a group from top to bottom, the first ionization energy increases.
Section 11.11
Atomic Properties and the Periodic Table
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Copyright © Cengage Learning. All rights reserved 76
Concept Check
Which atom would require more energy to remove an electron? Why?
Na Cl
Section 11.11
Atomic Properties and the Periodic Table
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Copyright © Cengage Learning. All rights reserved 77
Concept Check
Which atom would require more energy to remove an electron? Why?
Li Cs
Section 11.11
Atomic Properties and the Periodic Table
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Copyright © Cengage Learning. All rights reserved 78
Atomic Size
• In general as we go across a period from left to right, the atomic radius decreases. Effective nuclear charge increases, therefore
the valence electrons are drawn closer to the nucleus, decreasing the size of the atom.
• In general atomic radius increases in going down a group. Orbital sizes increase in successive principal
quantum levels.
Section 11.11
Atomic Properties and the Periodic Table
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Copyright © Cengage Learning. All rights reserved 79
Relative Atomic Sizes for Selected Atoms (Fig. 11-36)
Section 11.11
Atomic Properties and the Periodic Table
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80
Concept Check
Which should be the larger atom? Why?
Na Cl
Section 11.11
Atomic Properties and the Periodic Table
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Copyright © Cengage Learning. All rights reserved 81
Concept Check
Which should be the larger atom? Why?
Li Cs
Section 11.11
Atomic Properties and the Periodic Table
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Copyright © Cengage Learning. All rights reserved 82
Concept Check
Which is larger?• The hydrogen 1s orbital• The lithium 1s orbital
Which is lower in energy?•The hydrogen 1s orbital•The lithium 1s orbital
Section 11.11
Atomic Properties and the Periodic Table
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Copyright © Cengage Learning. All rights reserved 83
Exercise
Arrange the elements oxygen, fluorine, and sulfur according to increasing:
Ionization energy
S, O, F
Atomic size
F, O, S