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Transcript of 1 © 2006 Brooks/Cole - Thomson ATOMIC ELECTRON CONFIGURATIONS AND ORBITAL SHAPES.
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© 2006 Brooks/Cole - Thomson
ATOMIC ELECTRON ATOMIC ELECTRON CONFIGURATIONS AND ORBITAL CONFIGURATIONS AND ORBITAL
SHAPESSHAPES
ATOMIC ELECTRON ATOMIC ELECTRON CONFIGURATIONS AND ORBITAL CONFIGURATIONS AND ORBITAL
SHAPESSHAPES
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© 2006 Brooks/Cole - Thomson
Arrangement of Arrangement of Electrons in AtomsElectrons in Atoms
Arrangement of Arrangement of Electrons in AtomsElectrons in Atoms
Electrons in atoms are arranged asElectrons in atoms are arranged as
SHELLSSHELLS (n) (n)
SUBSHELLSSUBSHELLS (l) (l)
ORBITALSORBITALS (m (mll))
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Each orbital can be assigned no Each orbital can be assigned no
more than 2 electrons!more than 2 electrons!
This is tied to the existence of a 4th This is tied to the existence of a 4th
quantum number, the quantum number, the electron electron
spin quantum number, mspin quantum number, mss..
Arrangement of Arrangement of Electrons in AtomsElectrons in Atoms
Arrangement of Arrangement of Electrons in AtomsElectrons in Atoms
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Electron Electron Spin Spin
Quantum Quantum Number, Number,
mmss
Can be proved experimentally that electronCan be proved experimentally that electronhas a spin. Two spin directions are given byhas a spin. Two spin directions are given bymmss where m where mss = +1/2 and -1/2. = +1/2 and -1/2.
Can be proved experimentally that electronCan be proved experimentally that electronhas a spin. Two spin directions are given byhas a spin. Two spin directions are given bymmss where m where mss = +1/2 and -1/2. = +1/2 and -1/2.
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n ---> shelln ---> shell 1, 2, 3, 4, ...1, 2, 3, 4, ...
l ---> subshelll ---> subshell 0, 1, 2, ... n - 10, 1, 2, ... n - 1
mmll ---> orbital ---> orbital -l ... 0 ... +l-l ... 0 ... +l
mmss ---> electron spin ---> electron spin +1/2 and -1/2+1/2 and -1/2
n ---> shelln ---> shell 1, 2, 3, 4, ...1, 2, 3, 4, ...
l ---> subshelll ---> subshell 0, 1, 2, ... n - 10, 1, 2, ... n - 1
mmll ---> orbital ---> orbital -l ... 0 ... +l-l ... 0 ... +l
mmss ---> electron spin ---> electron spin +1/2 and -1/2+1/2 and -1/2
QUANTUM NUMBERSQUANTUM NUMBERS
Now there are four!Now there are four!
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Pauli Exclusion Pauli Exclusion PrinciplePrinciple
No two electrons in the No two electrons in the same atom can have same atom can have the same set of 4 the same set of 4 quantum numbers.quantum numbers.
That is, each electron has a That is, each electron has a unique address.unique address.
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© 2006 Brooks/Cole - Thomson
Electrons in AtomsElectrons in AtomsElectrons in AtomsElectrons in Atoms
When n = 1, then l = 0When n = 1, then l = 0
this shell has a single orbital (1s) to this shell has a single orbital (1s) to
which 2e- can be assigned.which 2e- can be assigned.
When n = 2, then l = 0, 1When n = 2, then l = 0, 1
2s orbital 2s orbital 2e-2e-
three 2p orbitalsthree 2p orbitals 6e-6e-
TOTAL = TOTAL = 8e-8e-
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© 2006 Brooks/Cole - Thomson
Electrons in AtomsElectrons in AtomsElectrons in AtomsElectrons in Atoms
When n = 3, then l = 0, 1, 2When n = 3, then l = 0, 1, 2
3s orbital 3s orbital 2e-2e-
three 3p orbitalsthree 3p orbitals 6e-6e-
five 3d orbitalsfive 3d orbitals 10e-10e-
TOTAL = TOTAL = 18e-18e-
When n = 3, then l = 0, 1, 2When n = 3, then l = 0, 1, 2
3s orbital 3s orbital 2e-2e-
three 3p orbitalsthree 3p orbitals 6e-6e-
five 3d orbitalsfive 3d orbitals 10e-10e-
TOTAL = TOTAL = 18e-18e-
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© 2006 Brooks/Cole - Thomson
Electrons in AtomsElectrons in AtomsElectrons in AtomsElectrons in Atoms
When n = 4, then l = 0, 1, 2, 3When n = 4, then l = 0, 1, 2, 3
4s orbital 4s orbital 2e-2e-
three 4p orbitalsthree 4p orbitals 6e-6e-
five 4d orbitalsfive 4d orbitals 10e-10e-
seven 4f orbitalsseven 4f orbitals 14e-14e-
TOTAL = TOTAL = 32e-32e-
And many more!And many more!And many more!And many more!
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© 2006 Brooks/Cole - Thomson
Assigning Electrons to Assigning Electrons to AtomsAtoms
Assigning Electrons to Assigning Electrons to AtomsAtoms
• Electrons generally assigned to orbitals of Electrons generally assigned to orbitals of
successively higher energy.successively higher energy.
• For For H atomsH atoms, E = - C(1/n, E = - C(1/n22). E depends only ). E depends only
on n.on n.
• For For many-electron atomsmany-electron atoms, energy depends , energy depends
on both n and l.on both n and l.
• See Active Figure 8.4, Figure 8.5, and Screen 8. 7.See Active Figure 8.4, Figure 8.5, and Screen 8. 7.
• Electrons generally assigned to orbitals of Electrons generally assigned to orbitals of
successively higher energy.successively higher energy.
• For For H atomsH atoms, E = - C(1/n, E = - C(1/n22). E depends only ). E depends only
on n.on n.
• For For many-electron atomsmany-electron atoms, energy depends , energy depends
on both n and l.on both n and l.
• See Active Figure 8.4, Figure 8.5, and Screen 8. 7.See Active Figure 8.4, Figure 8.5, and Screen 8. 7.
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Assigning Electrons to Assigning Electrons to SubshellsSubshells
• In H atom all subshells In H atom all subshells of same n have same of same n have same energy.energy.
• In many-electron atom:In many-electron atom:
a) subshells increase in a) subshells increase in energy as value of n + l energy as value of n + l increases.increases.
b) for subshells of same b) for subshells of same n + l, subshell with n + l, subshell with lower n is lower in lower n is lower in energy.energy.
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Effective Nuclear Charge, Effective Nuclear Charge, Z*Z*
Effective Nuclear Charge, Effective Nuclear Charge, Z*Z*
• Z* is the nuclear charge experienced by Z* is the nuclear charge experienced by the outermost electrons.the outermost electrons. See Figure 8.6 and and See Figure 8.6 and and Screen 8.6.Screen 8.6.
• Explains why E(2s) < E(2p)Explains why E(2s) < E(2p)
• Z* increases across a period owing to Z* increases across a period owing to incomplete shielding by inner electrons.incomplete shielding by inner electrons.
• Estimate Z* by --> [ Estimate Z* by --> [ Z - (no. inner electrons) Z - (no. inner electrons) ]]
• Charge felt by 2s e- in Li Charge felt by 2s e- in Li Z* = 3 - 2 = 1 Z* = 3 - 2 = 1
• Be Be Z* = 4 - 2 = 2Z* = 4 - 2 = 2
• B B Z* = 5 - 2 = 3Z* = 5 - 2 = 3 and so on!and so on!
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Effective Effective Nuclear Nuclear ChargeCharge
Figure 8.6
Electron cloud for 1s electrons
Z* is the nuclear Z* is the nuclear charge experienced charge experienced by the outermost by the outermost electrons.electrons.
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What does shielding do?What does shielding do?What does shielding do?What does shielding do?
• Well… It causes the subshells to have Well… It causes the subshells to have unequal energy. unequal energy.
• Therefore, the energy levels fill in a different Therefore, the energy levels fill in a different order.order.
• Shielding accounts for many periodic Shielding accounts for many periodic properties.properties.
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Electron Electron Filling Filling OrderOrder
Figure 8.5Figure 8.5
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Orbital EnergiesOrbital Energies
CD-ROM Screens 8.9 - 8.13, CD-ROM Screens 8.9 - 8.13, SimulationsSimulations
Orbital energies “drop” as Z* increasesOrbital energies “drop” as Z* increases
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Writing Atomic Electron Writing Atomic Electron ConfigurationsConfigurations
Writing Atomic Electron Writing Atomic Electron ConfigurationsConfigurations
11 s
value of nvalue of l
no. ofelectrons
spdf notationfor H, atomic number = 1
Two ways of Two ways of writing configs. writing configs. One is called One is called the the spdf spdf notation.notation.
Two ways of Two ways of writing configs. writing configs. One is called One is called the the spdf spdf notation.notation.
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Writing Atomic Electron Writing Atomic Electron ConfigurationsConfigurations
Writing Atomic Electron Writing Atomic Electron ConfigurationsConfigurations
Two ways of Two ways of writing writing configs. Other configs. Other is called the is called the orbital box orbital box notation.notation.
Two ways of Two ways of writing writing configs. Other configs. Other is called the is called the orbital box orbital box notation.notation.
Arrowsdepictelectronspin
ORBITAL BOX NOTATIONfor He, atomic number = 2
1s
21 s
Arrowsdepictelectronspin
ORBITAL BOX NOTATIONfor He, atomic number = 2
1s
21 s
One electron has n = 1, l = 0, mOne electron has n = 1, l = 0, m ll = 0, m = 0, mss = + 1/2 = + 1/2
Other electron has n = 1, l = 0, mOther electron has n = 1, l = 0, m ll = 0, m = 0, mss = - 1/2 = - 1/2
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© 2006 Brooks/Cole - Thomson
See “Toolbox” on CD for Electron Configuration tool.See “Toolbox” on CD for Electron Configuration tool.
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Electron Configurations Electron Configurations and the Periodic Tableand the Periodic Table
Active Figure 8.7
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LithiumLithiumLithiumLithium
Group 1AGroup 1A
Atomic number = 3Atomic number = 3
1s1s222s2s11 ---> 3 total electrons ---> 3 total electrons
1s
2s
3s3p
2p
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BerylliumBerylliumBerylliumBeryllium
Group 2AGroup 2A
Atomic number = 4Atomic number = 4
1s1s222s2s22 ---> 4 total ---> 4 total electronselectrons
1s
2s
3s3p
2p
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BoronBoronBoronBoron
Group 3AGroup 3A
Atomic number = 5Atomic number = 5
1s1s2 2 2s2s2 2 2p2p11 ---> --->
5 total electrons5 total electrons
1s
2s
3s3p
2p
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© 2006 Brooks/Cole - Thomson
CarbonCarbonCarbonCarbon
Group 4AGroup 4A
Atomic number = 6Atomic number = 6
1s1s2 2 2s2s2 2 2p2p22 ---> --->
6 total electrons6 total electrons
Here we see for the first time Here we see for the first time
HUND’S RULEHUND’S RULE. When . When placing electrons in a set of placing electrons in a set of orbitals having the same orbitals having the same energy, we place them singly energy, we place them singly as long as possible.as long as possible.1s
2s
3s3p
2p
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© 2006 Brooks/Cole - Thomson
NitrogenNitrogenNitrogenNitrogen
Group 5AGroup 5A
Atomic number = 7Atomic number = 7
1s1s2 2 2s2s2 2 2p2p33 ---> --->
7 total electrons7 total electrons
1s
2s
3s3p
2p
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© 2006 Brooks/Cole - Thomson
OxygenOxygenOxygenOxygen
Group 6AGroup 6A
Atomic number = 8Atomic number = 8
1s1s2 2 2s2s2 2 2p2p44 ---> --->
8 total electrons8 total electrons
1s
2s
3s3p
2p
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FluorineFluorineFluorineFluorine
Group 7AGroup 7A
Atomic number = 9Atomic number = 9
1s1s2 2 2s2s2 2 2p2p55 ---> --->
9 total electrons9 total electrons
1s
2s
3s3p
2p
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NeonNeonNeonNeon
Group 8AGroup 8A
Atomic number = 10Atomic number = 10
1s1s2 2 2s2s2 2 2p2p66 ---> --->
10 total electrons10 total electrons
1s
2s
3s3p
2p
Note that we Note that we have reached the have reached the end of the 2nd end of the 2nd period, and the period, and the 2nd shell is full!2nd shell is full!
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© 2006 Brooks/Cole - Thomson
Electron Configurations Electron Configurations of p-Block Elementsof p-Block Elements
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SodiumSodiumSodiumSodium
Group 1AGroup 1A
Atomic number = 11Atomic number = 11
1s1s2 2 2s2s2 2 2p2p6 6 3s3s11 or or
“ “neon core” + 3sneon core” + 3s11
[Ne] 3s[Ne] 3s1 1 (uses rare gas notation)(uses rare gas notation)
Note that we have begun a new period.Note that we have begun a new period.
All Group 1A elements have All Group 1A elements have [core]ns[core]ns11 configurations. configurations.
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AluminumAluminumAluminumAluminum
Group 3AGroup 3A
Atomic number = 13Atomic number = 13
1s1s2 2 2s2s2 2 2p2p6 6 3s3s2 2 3p3p11
[Ne] 3s[Ne] 3s2 2 3p3p1 1
All Group 3A All Group 3A elements have [core] elements have [core] nsns2 2 npnp1 1 configurations configurations where n is the period where n is the period number.number.
1s
2s
3s3p
2p
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PhosphorusPhosphorusPhosphorusPhosphorus
All Group 5A All Group 5A elements have elements have [core ] ns[core ] ns2 2 npnp3 3
configurations configurations where n is the where n is the period number.period number.
Group 5AGroup 5A
Atomic number = 15Atomic number = 15
1s1s2 2 2s2s2 2 2p2p6 6 3s3s2 2 3p3p33
[Ne] 3s[Ne] 3s2 2 3p3p33
1s
2s
3s3p
2p
Yellow P
Red P
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CalciumCalciumCalciumCalcium
Group 2AGroup 2A
Atomic number = 20Atomic number = 20
1s1s2 2 2s2s2 2 2p2p6 6 3s3s2 2 3p3p66 4s 4s22
[Ar] 4s[Ar] 4s2 2
All Group 2A elements have All Group 2A elements have [core]ns[core]ns2 2 configurations where n configurations where n is the period number.is the period number.
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© 2006 Brooks/Cole - Thomson
Electron Configurations Electron Configurations and the Periodic Tableand the Periodic Table
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© 2006 Brooks/Cole - Thomson
Transition MetalsTransition MetalsTable 8.4Table 8.4
Transition MetalsTransition MetalsTable 8.4Table 8.4
All 4th period elements have the All 4th period elements have the configuration configuration [argon] ns[argon] nsxx (n - 1)d (n - 1)dyy and so are and so are d-blockd-block elements. elements.
CopperCopperIronIronChromiumChromium
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Transition Element Transition Element ConfigurationsConfigurations
3d orbitals used for Sc-Zn (Table 8.4)
3d orbitals used for Sc-Zn (Table 8.4)
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Lanthanides and ActinidesLanthanides and ActinidesLanthanides and ActinidesLanthanides and Actinides
All these elements have the configuration All these elements have the configuration [core] ns[core] nsxx (n - 1)d (n - 1)dy y (n - 2)f(n - 2)fzz and so are and so are f-blockf-block elements. elements.
CeriumCerium[Xe] 6s[Xe] 6s22 5d 5d11 4f 4f11
UraniumUranium[Rn] 7s[Rn] 7s22 6d 6d11 5f 5f33
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Lanthanide Element Lanthanide Element ConfigurationsConfigurations
4f orbitals used for Ce - Lu and 5f for Th - Lr (Table 8.2)
4f orbitals used for Ce - Lu and 5f for Th - Lr (Table 8.2)
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Ion ConfigurationsIon ConfigurationsIon ConfigurationsIon Configurations
To form cations from elements remove To form cations from elements remove 1 or more e- from subshell of highest 1 or more e- from subshell of highest n [or highest (n + l)].n [or highest (n + l)].
P [Ne] 3sP [Ne] 3s22 3p 3p33 - 3e- ---> P - 3e- ---> P3+3+ [Ne] 3s [Ne] 3s22 3p 3p00
1s
2s
3s3p
2p
1s
2s
3s3p
2p
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Ion ConfigurationsIon ConfigurationsIon ConfigurationsIon Configurations
For transition metals, remove ns electrons and For transition metals, remove ns electrons and then (n - 1) electrons.then (n - 1) electrons.
Fe [Ar] 4sFe [Ar] 4s22 3d 3d66
loses 2 electrons ---> Feloses 2 electrons ---> Fe2+2+ [Ar] 4s [Ar] 4s00 3d 3d66
4s 3d 3d4s
Fe Fe2+
3d4s
Fe3+To form cations, always remove electrons of highest n value first!
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Ion ConfigurationsIon ConfigurationsIon ConfigurationsIon Configurations
How do we know the configurations of ions? How do we know the configurations of ions?
Determine the Determine the magnetic propertiesmagnetic properties of ions. of ions.
Sample of Fe2O3
Sample of Fe2O3
with strong magnet
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© 2006 Brooks/Cole - Thomson
Ion ConfigurationsIon ConfigurationsIon ConfigurationsIon Configurations
How do we know the configurations of ions? How do we know the configurations of ions?
Determine the Determine the magnetic propertiesmagnetic properties of ions. of ions.
Ions with Ions with UNPAIRED ELECTRONSUNPAIRED ELECTRONS are are PARAMAGNETICPARAMAGNETIC..
Without unpaired electrons Without unpaired electrons DIAMAGNETICDIAMAGNETIC..
Fe3+ ions in Fe2O3 have 5 unpaired electrons and make the sample paramagnetic.
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PERIODIPERIODIC C
TRENDSTRENDS
PERIODIPERIODIC C
TRENDSTRENDS
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General Periodic General Periodic TrendsTrends
• Atomic and ionic sizeAtomic and ionic size
• Ionization energyIonization energy
• Electron affinityElectron affinity
Higher effective nuclear chargeElectrons held more tightly
Larger orbitals.Electrons held lesstightly.
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Effective Nuclear Charge, Effective Nuclear Charge, Z*Z*
Effective Nuclear Charge, Effective Nuclear Charge, Z*Z*
• Z* is the nuclear charge experienced by the Z* is the nuclear charge experienced by the outermost electrons.outermost electrons. See Figure 8.6 and and Screen 8.6.See Figure 8.6 and and Screen 8.6.
• Explains why E(2s) < E(2p)Explains why E(2s) < E(2p)
• Z* increases across a period owing to incomplete Z* increases across a period owing to incomplete shielding by inner electrons.shielding by inner electrons.
• Estimate Z* by --> [ Estimate Z* by --> [ Z - (no. inner electrons) Z - (no. inner electrons) ]]
• Charge felt by 2s e- in Li Charge felt by 2s e- in Li Z* = 3 - 2 = 1 Z* = 3 - 2 = 1
• Be Be Z* = 4 - 2 = 2Z* = 4 - 2 = 2
• B B Z* = 5 - 2 = 3Z* = 5 - 2 = 3 and so on!and so on!
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Effective Effective Nuclear Nuclear ChargeCharge
Figure 8.6
Electron cloud for 1s electrons
Z* is the nuclear Z* is the nuclear charge experienced charge experienced by the outermost by the outermost electrons.electrons.
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Effective Nuclear Effective Nuclear ChargeCharge
Z*Z*The 2s electron PENETRATES the region The 2s electron PENETRATES the region
occupied by the 1s electron. occupied by the 1s electron.
2s electron experiences a higher positive 2s electron experiences a higher positive charge than expected. charge than expected.
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EffectiveEffective Nuclear Charge, Z* Nuclear Charge, Z*
• Atom Z* Experienced by Electrons in Valence Orbitals
• Li +1.28
• Be -------
• B +2.58
• C +3.22
• N +3.85
• O +4.49
• F +5.13
Increase in Increase in Z* across a Z* across a periodperiod
[Values calculated using Slater’s Rules][Values calculated using Slater’s Rules]
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General Periodic General Periodic TrendsTrends
• Atomic and ionic sizeAtomic and ionic size
• Ionization energyIonization energy
• Electron affinityElectron affinity
Higher effective nuclear chargeElectrons held more tightly
Larger orbitals.Electrons held lesstightly.
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Atomic Atomic RadiiRadiiAtomic Atomic RadiiRadii
Active Figure 8.11Active Figure 8.11
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Atomic Atomic SizeSize
Atomic Atomic SizeSize
• Size goes UPSize goes UP on going on going down a group. down a group. See Figure 8.9.See Figure 8.9.
• Because electrons are Because electrons are added further from the added further from the nucleus, there is less nucleus, there is less attraction.attraction.
• Size goes DOWNSize goes DOWN on going on going across a period.across a period.
• Size goes UPSize goes UP on going on going down a group. down a group. See Figure 8.9.See Figure 8.9.
• Because electrons are Because electrons are added further from the added further from the nucleus, there is less nucleus, there is less attraction.attraction.
• Size goes DOWNSize goes DOWN on going on going across a period.across a period.
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Atomic SizeAtomic SizeAtomic SizeAtomic Size
Size Size decreasesdecreases across a period owing across a period owing to increase in Z*. Each added electron to increase in Z*. Each added electron feels a greater and greater + charge.feels a greater and greater + charge.
LargeLarge SmallSmall
Increase in Z*Increase in Z*
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Trends in Atomic SizeTrends in Atomic SizeSee Active Figure 8.11See Active Figure 8.11
0
50
100
150
200
250
0 5 10 15 20 25 30 35 40
Li
Na
K
Kr
He
NeAr
2nd period
3rd period 1st transitionseries
Radius (pm)
Atomic Number
0
50
100
150
200
250
0 5 10 15 20 25 30 35 40
Li
Na
K
Kr
He
NeAr
2nd period
3rd period 1st transitionseries
Radius (pm)
Atomic Number
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Sizes of Transition ElementsSizes of Transition ElementsSee Figure 8.12See Figure 8.12
Sizes of Transition ElementsSizes of Transition ElementsSee Figure 8.12See Figure 8.12
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Sizes of Transition ElementsSizes of Transition ElementsSee Figure 8.12See Figure 8.12
Sizes of Transition ElementsSizes of Transition ElementsSee Figure 8.12See Figure 8.12
• 3d subshell is inside the 4s 3d subshell is inside the 4s subshell.subshell.
• 4s electrons feel a more or less 4s electrons feel a more or less constant Z*.constant Z*.
• Sizes stay about the same and Sizes stay about the same and chemistries are similar!chemistries are similar!
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Density of Transition Density of Transition MetalsMetals
0
5
10
15
20
25
3B 4B 5B 6B 7B 8B 1B 2B
Group
Den
sit
y (
g/m
L)
6th period6th period
5th period5th period
4th period4th period
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Ion SizesIon SizesIon SizesIon Sizes
Li,152 pm3e and 3p
Li+, 60 pm2e and 3 p
+Does the size goDoes the size goup or down up or down when losing an when losing an electron to form electron to form a cation?a cation?
Does the size goDoes the size goup or down up or down when losing an when losing an electron to form electron to form a cation?a cation?
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Ion SizesIon SizesIon SizesIon Sizes
• CATIONSCATIONS are are SMALLERSMALLER than the than the atoms from which they come.atoms from which they come.
• The electron/proton attraction has The electron/proton attraction has gone UP and so size gone UP and so size DECREASESDECREASES..
Li,152 pm3e and 3p
Li+, 78 pm2e and 3 p
+Forming Forming a cation.a cation.Forming Forming a cation.a cation.
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Ion SizesIon SizesIon SizesIon Sizes
F,64 pm9e and 9p
F- , 136 pm10 e and 9 p
-Does the size go up or Does the size go up or down when gaining an down when gaining an electron to form an electron to form an anion?anion?
Does the size go up or Does the size go up or down when gaining an down when gaining an electron to form an electron to form an anion?anion?
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Ion SizesIon SizesIon SizesIon Sizes
• ANIONSANIONS are are LARGERLARGER than the atoms from than the atoms from which they come.which they come.
• The electron/proton attraction has gone The electron/proton attraction has gone DOWN and so size DOWN and so size INCREASESINCREASES..
• Trends in ion sizes are the same as atom Trends in ion sizes are the same as atom sizes. sizes.
Forming Forming an anion.an anion.Forming Forming an anion.an anion.F, 71 pm
9e and 9pF-, 133 pm10 e and 9 p
-
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Trends in Ion SizesTrends in Ion SizesTrends in Ion SizesTrends in Ion Sizes
Active Figure 8.15Active Figure 8.15
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Redox Reactions
Redox Reactions
Why do metals lose Why do metals lose
electrons in their electrons in their
reactions? reactions?
Why does Mg form MgWhy does Mg form Mg2+2+
ions and not Mgions and not Mg3+3+??
Why do nonmetals take Why do nonmetals take
on electrons?on electrons?
Why do metals lose Why do metals lose
electrons in their electrons in their
reactions? reactions?
Why does Mg form MgWhy does Mg form Mg2+2+
ions and not Mgions and not Mg3+3+??
Why do nonmetals take Why do nonmetals take
on electrons?on electrons?
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Ionization EnergyIonization EnergySee CD Screen 8.12See CD Screen 8.12
Ionization EnergyIonization EnergySee CD Screen 8.12See CD Screen 8.12
IE = energy required to remove an electron IE = energy required to remove an electron from an atom in the gas phase.from an atom in the gas phase.
Mg (g) + 738 kJ ---> MgMg (g) + 738 kJ ---> Mg++ (g) + e- (g) + e-
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Mg (g) + 738 kJ ---> MgMg (g) + 738 kJ ---> Mg++ (g) + e- (g) + e-
MgMg+ + (g) + 1451 kJ ---> Mg(g) + 1451 kJ ---> Mg2+2+ (g) + e- (g) + e-
MgMg++ has 12 protons and only 11 has 12 protons and only 11 electrons. Therefore, IE for Mgelectrons. Therefore, IE for Mg++ > Mg. > Mg.
IE = energy required to remove an electron IE = energy required to remove an electron from an atom in the gas phase.from an atom in the gas phase.
Ionization EnergyIonization EnergySee Screen 8.12See Screen 8.12
Ionization EnergyIonization EnergySee Screen 8.12See Screen 8.12
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Mg (g) + 735 kJ ---> MgMg (g) + 735 kJ ---> Mg++ (g) + e- (g) + e-
MgMg+ + (g) + 1451 kJ ---> Mg(g) + 1451 kJ ---> Mg2+2+ (g) + e- (g) + e-
MgMg2+2+ (g) + 7733 kJ ---> Mg (g) + 7733 kJ ---> Mg3+3+ (g) + e- (g) + e-
Energy cost is very high to dip into a Energy cost is very high to dip into a shell of lower n. shell of lower n. This is why ox. no. = Group no.This is why ox. no. = Group no.
Ionization EnergyIonization EnergySee Screen 8.12See Screen 8.12
Ionization EnergyIonization EnergySee Screen 8.12See Screen 8.12
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Trends in Ionization EnergyTrends in Ionization EnergyTrends in Ionization EnergyTrends in Ionization Energy
Active Figure 8.13Active Figure 8.13
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Trends in Ionization EnergyTrends in Ionization EnergyTrends in Ionization EnergyTrends in Ionization Energy
1 3 5 7 9 11 13 15 17 19 21 23 25 27 29 31 33 350
500
1000
1500
2000
2500
1st Ionization energy (kJ/mol)
Atomic NumberH Li Na K
HeNe
ArKr
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Orbital EnergiesOrbital Energies
CD-ROM Screens 8.9 - 8.13, CD-ROM Screens 8.9 - 8.13, SimulationsSimulations
As Z* increases, orbital energies As Z* increases, orbital energies “drop” and IE increases. “drop” and IE increases.
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Trends in Ionization Trends in Ionization EnergyEnergy
Trends in Ionization Trends in Ionization EnergyEnergy
• IE increases across a period IE increases across a period because Z* increases.because Z* increases.
• Metals lose electrons more Metals lose electrons more easily than nonmetals.easily than nonmetals.
• Metals are good reducing Metals are good reducing agents.agents.
• Nonmetals lose electrons with Nonmetals lose electrons with difficulty.difficulty.
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Trends in Ionization Trends in Ionization EnergyEnergy
Trends in Ionization Trends in Ionization EnergyEnergy
• IE decreases down a group IE decreases down a group
• Because size increases.Because size increases.
• Reducing ability generally Reducing ability generally increases down the periodic increases down the periodic table. table.
• See reactions of Li, Na, KSee reactions of Li, Na, K
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Periodic Trend Periodic Trend in the in the
Reactivity of Reactivity of Alkali Metals Alkali Metals with Waterwith Water
Periodic Trend Periodic Trend in the in the
Reactivity of Reactivity of Alkali Metals Alkali Metals with Waterwith Water
LithiumLithium
SodiumSodium PotassiumPotassium
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Electron AffinityElectron Affinity
A few elements A few elements GAINGAIN electrons electrons to form to form anionsanions..
Electron affinity is the energy Electron affinity is the energy involved when an atom gains involved when an atom gains an electron to form an anion.an electron to form an anion.
A(g) + e- ---> AA(g) + e- ---> A--(g) E.A. = ∆E(g) E.A. = ∆E
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Electron Affinity of OxygenElectron Affinity of Oxygen
∆∆E is E is EXOEXOthermic thermic because O has because O has an affinity for an an affinity for an e-.e-.
[He] O atom
EA = - 141 kJ
+ electron
O [He] - ion
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Electron Affinity of Electron Affinity of NitrogenNitrogen
∆∆E is E is zero zero for Nfor N- -
due to electron-due to electron-electron electron repulsions.repulsions.
EA = 0 kJ
[He] N atom
[He] N- ion
+ electron
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Trends in Electron AffinityTrends in Electron Affinity
Active Figure 8.14Active Figure 8.14
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• See Figure 8.14 and See Figure 8.14 and Appendix FAppendix F
• Affinity for electron Affinity for electron increases across a increases across a period (EA becomes period (EA becomes more positive).more positive).
• Affinity decreases down Affinity decreases down a group (EA becomes a group (EA becomes less positive).less positive).
Atom EAAtom EAFF +328 kJ+328 kJClCl +349 kJ+349 kJBrBr +325 kJ+325 kJII +295 kJ+295 kJ
Atom EAAtom EAFF +328 kJ+328 kJClCl +349 kJ+349 kJBrBr +325 kJ+325 kJII +295 kJ+295 kJ
Trends in Electron AffinityTrends in Electron Affinity
Note effect of atom Note effect of atom size on F vs. Clsize on F vs. Cl