1 © 2006 brooks/cole - thomson atomic electron configurations and orbital shapes

1

© 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

2


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))

3


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..



4


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.

5


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!

6


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.

7


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-

8



When n = 3, then l = 0, 1, 2When n = 3, then l = 0, 1, 2



five 3d orbitalsfive 3d orbitals 10e-10e-


When n = 3, then l = 0, 1, 2When n = 3, then l = 0, 1, 2





9



When n = 4, then l = 0, 1, 2, 3When n = 4, then l = 0, 1, 2, 3




seven 4f orbitalsseven 4f orbitals 14e-14e-


And many more!And many more!And many more!And many more!

10


11


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.

12


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.

13


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!

14


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.

15


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.

16


Electron Electron Filling Filling OrderOrder

Figure 8.5Figure 8.5

17


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

18


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.

19




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

20


See “Toolbox” on CD for Electron Configuration tool.See “Toolbox” on CD for Electron Configuration tool.

21


Electron Configurations Electron Configurations and the Periodic Tableand the Periodic Table

Active Figure 8.7

22


LithiumLithiumLithiumLithium

Group 1AGroup 1A

Atomic number = 3Atomic number = 3

1s1s222s2s11 ---> 3 total electrons ---> 3 total electrons

1s

2s

3s3p

2p

23


BerylliumBerylliumBerylliumBeryllium

Group 2AGroup 2A


1s1s222s2s22 ---> 4 total ---> 4 total electronselectrons

1s

2s

3s3p

2p

24


BoronBoronBoronBoron

Group 3AGroup 3A


1s1s2 2 2s2s2 2 2p2p11 ---> --->

5 total electrons5 total electrons

1s

2s

3s3p

2p

25


CarbonCarbonCarbonCarbon

Group 4AGroup 4A


1s1s2 2 2s2s2 2 2p2p22 ---> --->


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

26


NitrogenNitrogenNitrogenNitrogen

Group 5AGroup 5A


1s1s2 2 2s2s2 2 2p2p33 ---> --->


1s

2s

3s3p

2p

27


OxygenOxygenOxygenOxygen

Group 6AGroup 6A


1s1s2 2 2s2s2 2 2p2p44 ---> --->


1s

2s

3s3p

2p

28


FluorineFluorineFluorineFluorine

Group 7AGroup 7A


1s1s2 2 2s2s2 2 2p2p55 ---> --->


1s

2s

3s3p

2p

29


NeonNeonNeonNeon

Group 8AGroup 8A


1s1s2 2 2s2s2 2 2p2p66 ---> --->


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!

30


Electron Configurations Electron Configurations of p-Block Elementsof p-Block Elements

31


SodiumSodiumSodiumSodium

Group 1AGroup 1A


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.

32


AluminumAluminumAluminumAluminum

Group 3AGroup 3A


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

33


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


1s1s2 2 2s2s2 2 2p2p6 6 3s3s2 2 3p3p33

[Ne] 3s[Ne] 3s2 2 3p3p33

1s

2s

3s3p

2p

Yellow P

Red P

34


CalciumCalciumCalciumCalcium

Group 2AGroup 2A


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.

35


Electron Configurations Electron Configurations and the Periodic Tableand the Periodic Table

36


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

37


Transition Element Transition Element ConfigurationsConfigurations

3d orbitals used for Sc-Zn (Table 8.4)

3d orbitals used for Sc-Zn (Table 8.4)

38


39


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

40


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)

41


42


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

43



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!

44



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

45



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.

46


PERIODIPERIODIC C

TRENDSTRENDS

PERIODIPERIODIC C

TRENDSTRENDS

47


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.

48




• 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!

49


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.

50


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.

51


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]

52


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.

53


Atomic Atomic RadiiRadiiAtomic Atomic RadiiRadii

Active Figure 8.11Active Figure 8.11

54


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.

55


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*

56


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

57


Sizes of Transition ElementsSizes of Transition ElementsSee Figure 8.12See Figure 8.12


58




• 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!

59


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



60


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?

61



• 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.

62



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?

63



• 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

-

64


Trends in Ion SizesTrends in Ion SizesTrends in Ion SizesTrends in Ion Sizes


65


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?

66


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-

67



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


68



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.



70


Trends in Ionization EnergyTrends in Ionization EnergyTrends in Ionization EnergyTrends in Ionization Energy


71


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

72


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.

73


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.

74




• 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

75


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

76


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

77


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

78


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

79


Trends in Electron AffinityTrends in Electron Affinity


80


• 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

1 © 2006 brooks/cole - thomson atomic electron configurations and orbital shapes

Documents

e slide

s electrons z

brookscole thomson electrons

brookscole thomson n

atoms electrons

brookscole thomson slide

electron atoms

outermost electrons