6 Oct 1997Chemical Periodicity1 Electron Configurations Chemical Periodicity (Ch 8) Electron spin & Pauli exclusion principle configurations spectroscopic,

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


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 ?


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.


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


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:


QUANTUMNUMBERS

n (shell) 1, 2, 3, 4, ...

(subshell) 0, 1, 2, ... n - 1

m (orbital) - ... 0 ... +

ms (electron spin) +1/2, -1/2


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


Element Mnemonic Competition

Hey! Here Lies Ben Brown. Could Not Order Fire. Near Nancy Margaret Alice Sits Peggy Sucking Clorets. AreKids Capable ?


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) ...


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.


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


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) ]


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


Electron Filling Order (Figure 8.7)


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:


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)


Electron Configuration tool - see “toolbox”.


LithiumGroup 1A

Z = 3

1s22s1

1s

2s

3s3p

2p

Beryllium Group 2A

Z = 4

1s22s2

1s

2s

3s3p

2p


BoronZ = 5

1s2 2s2 2p1

1s

2s

3s3p

2p

CarbonZ = 6

1s2 2s2 2p2

1s

2s

3s3p

2p

Why not ?


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.


NitrogenNitrogenZ = 7Z = 7

1s1s2 2 2s2s2 2 2p2p33

1s

2s

3s3p

2p

OxygenZ = 8

1s2 2s2 2p4

1s

2s

3s3p

2p


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!


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


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


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


Relationship of Electron Configuration and Regions of the Periodic Table

f block

s block p blockd block


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


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


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


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?


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.


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.


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.


Atomic RadiiAtomic Radii


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


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!


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


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.


Trends in Ion SizesANIONSCATIONS

Trends in relative ion sizes are the same as atom sizes.

(59 pm)

(207 pm)


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


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+


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


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


2nd IE / 1st IE

LiLi

NaNa

KK

2nd IE: A+ A++ + e-


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


• 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


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

6 Oct 1997Chemical Periodicity1 Electron Configurations Chemical Periodicity (Ch 8) Electron spin & Pauli exclusion principle configurations spectroscopic,

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