6 Oct 1997 Chemical Periodicity 1
Electron Configurations Chemical Periodicity (Ch 8)
• Electron spin & Pauli exclusion principle • configurations
• spectroscopic, orbital box notation• Hund’s rule - electron filling rules • configurations of ATOMS:
• the basis for chemical valence• configurations and properties of IONS• periodic trends in :
• size• ionization energies• electron affinities
Na + Cl NaCl
Mg + O2 MgO
6 Oct 1997 Chemical Periodicity 2
Electrons in atoms are arranged as
SHELLS (n)
SUBSHELLS ()
ORBITALS (m)
Arrangement of Electrons in Atoms
. . . Because there is a 4th quantum number,
the electron spin quantum number, ms.
Each orbital can be assigned
up to 2 electrons!
WHY ?
6 Oct 1997 Chemical Periodicity 3
Electron Spin Quantum Number, ms
• It can be proved experimentally that the electron has a spin. This is QUANTIZED. • The two allowed spin directions are defined by the magnetic spin quantum number, ms
ms = +1/2 and -1/2 ONLY.
6 Oct 1997 Chemical Periodicity 4
Electron Spin Quantum Number
Diamagnetic: NOT attracted to a magnetic field
All electrons are paired N2
MAGNETISM is a macroscopicresult of quantized electron spin
5_magnet.mov
Paramagnetic: attracted to a magnetic field.
Substance has unpaired electrons O2
6 Oct 1997 Chemical Periodicity 5
Pauli Exclusion Principle
• No orbital can have more than 2 electrons
• No two electrons in the same atom can have the same set of 4 quantum
numbers (n, l, ml, ms) OROR
• “Each electron has a unique address.”
• electrons with the same spin keep as far apart as possible• electrons of opposite spin may occupy the same
“region of space” (= orbital)• Consequences:
6 Oct 1997 Chemical Periodicity 6
QUANTUMNUMBERS
n (shell) 1, 2, 3, 4, ...
(subshell) 0, 1, 2, ... n - 1
m (orbital) - ... 0 ... +
ms (electron spin) +1/2, -1/2
6 Oct 1997 Chemical Periodicity 7
Shells, Subshells, Orbitals
n #orbitals #e- Total PERIOD1 0 s 1 2 2 1 (H, He)2 0 s 1 2
1 p 3 6 8 2 (Li…Ne)3 0 s 1 2
1 p 3 6 3 (Na .. Ar)2 d 5 10 18
4 0 s 1 21 p 3 62 d 5 103 f 7 14 32
n 0..(n-1) (2 +1) 2*(2 +1) 2n2 etc, for n = 5, 6
= 0 s = 1 p = 2 d = 3 f
6 Oct 1997 Chemical Periodicity 8
Element Mnemonic Competition
Hey! Here Lies Ben Brown. Could Not Order Fire. Near Nancy Margaret Alice Sits Peggy Sucking Clorets. AreKids Capable ?
6 Oct 1997 Chemical Periodicity 9
Assigning Electrons to Atoms
• Electrons are assigned to orbitals successively in order
of the energy.
• For H atoms, E = - R(1/n2). E depends only on n.
• For many-electron atoms, orbital energy depends on
both n and .
• E(ns) < E(np) < E(nd) ...
6 Oct 1997 Chemical Periodicity 10
Assigning Electrons to Subshells
• In many-electron atom:
a) subshells increase in energy as value of (n + ) increases.
5_manyelE.mov
• In H atom all subshells of same n have same energy.
(n + )= 4
(n + )= 5b) for subshells of same (n +), subshell with lower n is lower in energy.
6 Oct 1997 Chemical Periodicity 11
2s e- spendsmore timeclose to Li3+
nucleus than the 2p e-
Therefore2s is lower in Ethan 3s
Effective Nuclear Charge
• The difference in SUBSHELL energy
e.g. 2s and 2p subshells
is due to effective nuclear charge, Z*.
Charge felt by 2s e- of Li atom
6 Oct 1997 Chemical Periodicity 12
Effective Nuclear Charge, Z*
• Z* is the nuclear charge experienced by an electron. • Z* increases across a period owing to incomplete
shielding by inner electrons.• For VALENCE electrons we estimate Z* as:
• Charge felt by 2s e- in Li Z* = 3 - 2 = 1Be Z* = 4 - 2 = 2B Z* = 5 - 2 = 3
and so on!
Z* = [ Z - (no. of inner electrons) ]
6 Oct 1997 Chemical Periodicity 13
Inner shell or CORE ELECTRONS
VALENCEELECTRONS
Photoelectron Spectroscopy - Measuring IEPhotoelectric effect: h + A A+ + e-
forms basis for DIRECT determination of IE
Kinetic energy of electron = h - IEtherefore: IE = h - KE(e-)
2s 2p
Ne1s 2s 2p
3p 3s
IE (MJ/mol)
Sig
nal
050100
Ar1s
309
6 Oct 1997 Chemical Periodicity 14
Electron Filling Order (Figure 8.7)
6 Oct 1997 Chemical Periodicity 15
Writing Atomic Electron ConfigurationsWriting Atomic Electron Configurations
11 s
value of nvalue of l
no. ofelectrons
SPECTROSCOPIC NOTATIONfor H, atomic number = 1
Two ways of writing configurations. Two ways of writing configurations.
One is called the One is called the spectroscopic notation:
6 Oct 1997 Chemical Periodicity 16
A second way is called the orbital box notation.
Arrowsdepictelectronspin
ORBITAL BOX NOTATIONfor He, atomic number = 2
1s
21 s
One electron has n = 1, = 0, ml = 0, ms = + 1/2
Other electron has n = 1, = 0, ml = 0, ms = - 1/2
Writing Atomic Electron Configurations (2)Writing Atomic Electron Configurations (2)
6 Oct 1997 Chemical Periodicity 17
Electron Configuration tool - see “toolbox”.
6 Oct 1997 Chemical Periodicity 18
LithiumGroup 1A
Z = 3
1s22s1
1s
2s
3s3p
2p
Beryllium Group 2A
Z = 4
1s22s2
1s
2s
3s3p
2p
6 Oct 1997 Chemical Periodicity 19
BoronZ = 5
1s2 2s2 2p1
1s
2s
3s3p
2p
CarbonZ = 6
1s2 2s2 2p2
1s
2s
3s3p
2p
Why not ?
6 Oct 1997 Chemical Periodicity 20
CarbonZ = 6
1s2 2s2 2p2
The configuration of C is an example of HUND’S RULE:
the lowest energy arrangement of electrons in a subshell is that with the MAXIMUM no. of unpaired electrons
1s
2s
3s3p
2pElectrons in a set of orbitals having the same energy, are placed singly as long as possible.
6 Oct 1997 Chemical Periodicity 21
NitrogenNitrogenZ = 7Z = 7
1s1s2 2 2s2s2 2 2p2p33
1s
2s
3s3p
2p
OxygenZ = 8
1s2 2s2 2p4
1s
2s
3s3p
2p
6 Oct 1997 Chemical Periodicity 22
FluorineZ = 9
1s2 2s2 2p5
1s
2s
3s3p
2p
NeonZ = 10
1s2 2s2 2p6
1s
2s
3s3p
2p
Note that we have reached the end of the 2nd period,. . . and the 2nd shell is full!
6 Oct 1997 Chemical Periodicity 23
GROUPS and PERIODS
or “neon core” + 3s1
[Ne] 3s1 (uses rare gas notation)
Na begins a new period.
All Group 1A elements: Li Na K Rb Cs
have [core] ns1 configurations. (n = period #)
SodiumZ = 11
1s2 2s2 2p6 3s1
6 Oct 1997 Chemical Periodicity 24
Periodic Chemical Properties
5_Li.mov
5_Na.mov
5_K.mov
Li
Na
K
Rb
Cs
Alkalis
REACTIVITY SIZE IE (Ionization Energy)
Be
Mg
Ca
Sr
Ba
Alkaline Earths
6 Oct 1997 Chemical Periodicity 25
Alkaline EarthsMetals (ns2) - easily oxidized to M2+
- less reactive than alkalis of same period
reactivity: Be < Mg < Ca < Sr < BaWHY? - • Size INCREASES as group
• VALENCE e- are farther from nucleus• same Z* - Valence e- less tightly held• Therefore valence e- are easier to remove
Typical reactions / compounds
Oxides: M +1/2O2 (g) MO (s) CaO (lime) - #5 Ind. Chem
Halides: M + X2 (g) MXCarbonates: CaCO3 (limestone) CaO + CO2
RECALL: Solubility rules and PRECIPITATION REACTIONS
Sulfates: CaSO4.2H2O (gypsum) CaSO4. 0.5H2O (plaster-of-paris) + 3/2H2O
6 Oct 1997 Chemical Periodicity 26
Relationship of Electron Configuration and Regions of the Periodic Table
f block
s block p blockd block
6 Oct 1997 Chemical Periodicity 27
Transition Metals Transition Metals Table 8.4Table 8.4
• Transition metals (e.g. Sc .. Zn in the 4th period) have the configuration [argon] nsx (n - 1)dy
• also called “d-block” elements.
CopperIronChromium
3d orbitals used for Sc - Zn
6 Oct 1997 Chemical Periodicity 28
To form cations from elements : remove 1 e- (or more) from subshell of highest n [or highest (n + )].
Ion Configurations
P [Ne] 3s2 3p3 - 3e- P3+ [Ne] 3s2 3p0
1s
2s
3s3p
2p
1s
2s
3s3p
2p
6 Oct 1997 Chemical Periodicity 29
Ion Configurations (2)
Transition metals ions:
remove ns electrons and then (n - 1)d electrons.
4s 3d 3d4s
Fe Fe2+
3d4s
Fe3+E4s ~ E3d - exact energyof orbitals depend on whole configuration
Fe [Ar] 4s2 3d6 loses 2 electrons Fe2+ [Ar] 4s0 3d6
6 Oct 1997 Chemical Periodicity 30
Ion Configurations (3)
From the magnetic properties of ions.
Ions (or atoms) with UNPAIRED ELECTRONS are:
PARAMAGNETIC.
Ions (or atoms) without unpaired electrons are:
DIAMAGNETIC.
How do we know the configurations of ions?
6 Oct 1997 Chemical Periodicity 31
General Periodic Trends• Atomic and ionic radii : SIZE• Ionization energy : E(A+) - E(A)• Electron affinity : E(A-) - E(A)
Higher Z*.Electrons heldmore tightly.
Larger orbitals.Electrons held lesstightly.
6 Oct 1997 Chemical Periodicity 32
Atomic Size INCREASESdown a Group
• Size goes UP on going down a GROUP
• Because electrons are added further from the nucleus, there is less attraction.
6 Oct 1997 Chemical Periodicity 33
Atomic Size DECREASES across a period
Size goes DOWN on going across a PERIOD.
Size decreases due to increase in Z*.
Each added electron feels a greater and greater +ve charge.
6 Oct 1997 Chemical Periodicity 34
Atomic RadiiAtomic Radii
6 Oct 1997 Chemical Periodicity 35
Trends in Atomic Size (Figure 8.10)
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
6 Oct 1997 Chemical Periodicity 36
Sizes of Transition Elements(Figure 8.11)
• 3d subshell is inside the 4s subshell.• 4s electrons feel a more or less constant Z*.
• Sizes stay about the same and chemistries are similar!
6 Oct 1997 Chemical Periodicity 37
Ion Sizes - CATIONS
Does the size go up or down when an atom loses an electron to form a cation?
• CATIONS are SMALLER than the parent atoms.• The electron/proton attraction goes UP so size DECREASES.
Forming Forming a cationa cation
Li, 152 pm3 e-, 3 p
+
Li+, 60 pm2 e-, 3 p
6 Oct 1997 Chemical Periodicity 38
F-, 136 pm10 e-, 9 p
-
Does the size go up or down when gaining an electron to form an anion?
Ion Sizes - ANIONS
F, 64 pm9 e-, 9 p
Forming Forming an anionan anion
• ANIONS are LARGER than the parent atoms.
• electron/proton attraction goes DOWN so size INCREASES.
6 Oct 1997 Chemical Periodicity 39
Trends in Ion SizesANIONSCATIONS
Trends in relative ion sizes are the same as atom sizes.
(59 pm)
(207 pm)
6 Oct 1997 Chemical Periodicity 40
Oxidation-Reduction Reactions
• Why do metals lose electrons in
their reactions?
• Why does Mg form Mg2+ ions
and not Mg3+?
• Why do nonmetals take on
electrons?
- related to IE and EA
6 Oct 1997 Chemical Periodicity 41
Mg (g) + 735 kJ Mg+ (g) + e-[Ne]2s1
Ionization Energy (IE)
Mg (g) atom [Ne]2s
Mg
• Energy ‘cost’ is very high to remove an INNER SHELL e- (shell of n < nVALENCE).
• This is why oxidation. no. = Group no.
Mg2+ (g) + 7733 kJ Mg3+ (g) + e- [He]2s22p5
Mg3+
Mg+ (g) + 1451 kJ Mg2+ (g) + e- [Ne]2s0
Mg2+
Mg+
6 Oct 1997 Chemical Periodicity 42
Trends in First 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
6 Oct 1997 Chemical Periodicity 43
Trends in Ionization Energy (2)
• IE increases across a period because Z* increases.• Metals lose electrons more easily than nonmetals.• Metals are good reducing agents.• Nonmetals lose electrons with difficulty.
• IE decreases down a group• Because size increases, reducing ability generally increases down the periodic table. • E.g. reactions of Li, Na, K
6 Oct 1997 Chemical Periodicity 44
2nd IE / 1st IE
LiLi
NaNa
KK
2nd IE: A+ A++ + e-
6 Oct 1997 Chemical Periodicity 45
Electron Affinity (EA)
• A few elements GAIN electrons to form anions.
• Electron affinity is the energy released when an atom gains an electron.
A(g) + e- A-(g) E.A. = E = E(A-) - E(A)
• If E(A-) < E(A) then the anion is more stable and there is an exothermic reaction
6 Oct 1997 Chemical Periodicity 46
• Affinity for electron increases across a period
(EA becomes more negative).
Atom EA (kJ)B -27 C -122 N 0 O -141 F -328
Trends in Electron Affinity (Table 8.5, Figure 8.14)
• Affinity decreases down a group (EA becomes less negative).
F -328 Cl -349 Br -325 I -295
6 Oct 1997 Chemical Periodicity 47
SUMMARY• Electron spin: diamagnetism vs. paramagnetism• Pauli exclusion principle - allowable quantum numbers• configurations
• spectroscopic notation• orbital box notation
• Hund’s rule - electron filling rules • configurations of ATOMS: the basis for chemical valence
• period 2 ; groups• transition metals
• configurations and properties of IONS• periodic trends in :
• size• ionization energies • electron affinities