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Many electron atoms and the Periodic Table

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Objectives Explain the scientific basis for the Periodic Table Apply the Aufbau principle, the Pauli Exclusion

principle and Hund’s rule to electrons in an atom Explain the concept of energy levels in an atom and

the order of filling these levels Write the electronic configuration of the first 20

elements Draw and explain a block diagram of the Periodic Table Explain the meaning and position of the transition

elements Explain the periodic variations of atomic size, ionisation

energy, electron affinity and electronegativity

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The Periodic Law

The physical and chemical properties of the elements are a function of the electronic configuration of the their atoms which vary with increasing atomic number in a periodic manner

hence the Periodic Table elements are grouped according to their electronic

structure provides information on chemical properties

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Many Electron Atoms

Describe in terms of hydrogen orbitals Same quantum numbers and shapes energies different

For Hydrogen: Depend on n

3s, 3p, 3d all the same energy

For many e- atom: Different subshells at different energies

2s < 2p depend on n and l

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Orbital Energies

Ene

rgy

1s 1s

n=1

n=3

n=2

2s 2p

3s 3p 3d

0

n=1

n=3

n=2

2s2p

3s

3p

3d

0

Ene

rgy

n=4

4s

Hydrogen Many e- Atoms

Need to consider Effect of increased nuclear charge

Repulsions between electrons

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Effective Nuclear Charge

Z

e- of interest

3 e- found in sphere

Eff. Nuclear charge = 5 - 3 = 2

Assume 5

Net positive charge from nucleus attracting an electron electron shielded by inner e-

effective nuclear chargeZeff = Z - SS = No. of e- between atom and nucleus

hence outer shell e- experience less + charge

effect of “screening” depends on e- distribution need to consider orbital shape

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Orbital shape and Energy

s-orbital e- can be close to nucleus p-orbital further away than an “s” e-

d-orbital further from nucleus than “p” e-

therefore, “s” e- has least screening by other e-

so a larger effective nuclear charge and is more tightly bound(lower energy than “p” ie. more stable)

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Order of Energy Levels for many e- Atom

In general, 1s < 2s < 2p < 3s < 3p…...

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An easy way to remember…...

1s

3d

5p

4p

3p

2p

5s

4s

3s

6d

5f

4f

5d

4d

2s

6s 6p

Increasing Energy

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Energy

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need to consider another property of electrons to determine how electrons populate orbitals

envisage electron as spinning on own axis quantized only 2 spin states

distinguished by the spin -magnetic quantum number ms

Electron Spin

+12

12

-

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Electron Spin

Stern-Gerlach experiment - when a beam of ground state H atoms (1s) is passed through a magnetic field, the beam splits into two beams

REMIND!

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Pauli and Hund

Pauli Exclusion Principle no two electrons in an atom can have the same

4 quantum numbers there are only 2 values of ms

hence, an orbital can only hold 2 electrons and they must have opposite spins

Hund’s Rule If orbitals have the same energy, add electrons

singly with spins parallel first

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The Aufbau Principle

Building up

fill available orbitals with available electrons starting with lowest energy orbitals (most stable)

this gives ground state

Note

don’t forget Pauli and Hund!!

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Building up atoms

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Aufbau Principle The aufbau

principle shows how orbitals are filled: in the order to the left. Two extra rules are needed.

shell 1 shell 2 shell 4shell 3

3d

1s

2s

2p

3s

3p4s

4p

4d

Increasingenergy

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Aufbau Principle-2

Hund’s rule states that when filling a set of orbitals at the same energy (sunshell), one electron is placed in each orbital before pairing occurs.

Pauli’s principle tells us that when placing a second electron in an orbital, its spin must be opposite to the electron already in the orbital. (Spin is usually represented by an arrow in an orbital box.)

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Aufbau Principle-3a “Box” or “orbital diagrams for electron configurations.

H (1e)1S

He (2e)1S

The 1s orbital is filled.Second electron is paired. (Pauli)

Start at 1s

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Aufbau Principle-3b

Li (3e)1S 2S 1s2, 2s1

C (6e)1S 2S 2p

Hund’s rule applies to p subshell

1s2, 2s2, 2p2

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Aufbau Principle-3c

In 2p subshell, Hund’s rule! Next electron follows Pauli principle

O (8e)1S 2S 2p 1s2, 2s2, 2p4

Ti (22e)(22e)

4S 3d

1s2, 2s2, 2p6, 3s2,3p6, 4s2, 3d2

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Electrons and the Periodic Table

Electron fit logically into the periodic table. The s block elements (see next slide) start filling at level 1, the p block at level 2, and the d block at level 3.

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Excited State Atoms

Occur when energy has been supplied to raise e- energy

Ne

1s22s22p6 Ground state

1s22s22p5….5s1 High energy excited state

1s22s22p5….3p1 Low energy excited state

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Transition metals

Consider the elements in the 4th period(after Ar 1s22s22p63s23p6)

after 3p natural sequence would be 3d but 4s has (slightly) lower energy than3d according to Aufbau must fill 4s before 3d

ground state for K is [Ar]4s1 and Ca is [Ar]4s2

as charge increases the energy of 3d decreases in Sc 3d < 4s Sc+ [Ar]3d14s1 Sc2+[Ar]3d1

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Transition metals

energy drop in 3d continues through to Zn

consequences for elements Cu Zn oxidation state 1+ or 2+

beyond Zn 3d electrons have no chemical role

elements from Sc to Zn called d-block elements filling up the d orbital

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Transition metals

anomalies Cr [Ar]3d54s1 expect [Ar]3d44s2

Cu [Ar]3d104s1 expect [Ar]3d94s2

similar occur in fifth period also Ru [Kr]4d75s1 expect [Kr]4d65s2

sixth period filling is erratic energies of 4f, 5d & 6s comparable

seventh period all are radioactive

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The Periodic TableOrganisation of the elements

electronic configurations related to position of element

elements grouped according to type of orbital theouter shell electrons are in

BLOCK: Named for last subshell occupied

GROUP: the columns all elements have same outer orbital e- configuration similar chemical properties

PERIODS: rows all elements same shell

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The Periodic Table Subshell

orbitals with same energy eg. 2p Shell

orbitals with similar energy eg. 2s, 2p Valence Electrons

occupy outermost shell Core Electrons

occupy filled inner shells Cl 1s22s22p6 3s23p5

Ne core valence Closed Shell Atoms

full outer shell - very stable - noble gases

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Periodic Properties

Predicted by considering e- configurations

Sizes of atoms and ions

Ionisation energies

Electron affinities

Electronegativities

Polarising powers and polarisabilities

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Sizes Of Atoms and Ions

Atoms do not have sharply defined boundaries

Hence, need to define atomic size

Atomic size depends on chemical environment

ie. Bonding etc

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This shielding means that each valence electron in effect only “feels” a +1 charge form the nucleus; this occurs for an highly excited valence electron. Otherwise the shielding makes the “seen” charge is higher than +1

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Defining Atomic and Ionic Size

2r

estimating size atomic radius = half thedistance between nearest atomsin element (in condensed phases)

for ions, base ionic radii on interatomicdistance in ionic crystals. (depends on charge...)

Cu Cu Cu+ Cu+2

atom covalent bonding

1.28 1.17 0.96 0.69 Ao

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Sizes of Atoms and Ions

Decrease

Increase

Why?Consider:1. Principle Quantum number (shell)2. Effective nuclear charge

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Increase nuclear charge

but no. of core electrons stay the same

so effective nuclear charge increases while shell remains the same

hence electrons drawn closer to nucleus

hence decrease atomic size

eg. Na 1s22s22p63s1 1.91Ao

Mg 1s22s22p63s2 1.60 Ao

Across a Period…….

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Down A Group….

More distant electron shell occupied while effective nuclear charge the same

hence atomic size increases

eg. Li 1s22s1 1.57 A0

Na 1s22s22p63s1 1.91 Ao

37Atomic radii

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Radius of Ions

Cation < Atomeg. Na+ < Na

0.96 Ao 1.91 Ao

1s22s22p6 1s22s22p63s1

lost an e-

core electrons exposed

more tightly bound

Decreases across a period eg. Na+ > Mg2+

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Radius of Ions

Atom < Anioneg. Cl < Cl

0.99 Ao 1.81 Ao

[Ne]3s23p5 [Ne]3s23p6

gained an e-

electron cloud greater

decreases nuclear pull by each electron

Decreases across a period e.g. S2- > Cl-

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Ionisation Energy

The energy required to remove an electron from a ground state atom

X(g) X+(g)

+ e- E = IE1

Measure of stability of outer shell electron configuration Depends on

size of the atom effective nuclear charge screening effect of inner electrons type of electron

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Ionisation Energy

Increase

Decrease Why?Consider1. Effective nuclear charge2. Distance of e- from nucleus

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Across a Period….

Increase in effective nuclear chargeDecrease in radius

hence increase attraction between e- and nucleus

hence increase IE

Exceptions: “p” less stable than “s” (B < Be)orbitals “singly occupied” more stable

than “doubly occupied” (O < N)

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Down a Group…...

Increase radius whileeffective nuclear charge the same

hence Decrease attraction between e- and nucleus

hence decrease IE

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Ionization energy

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Electron Affinity

The energy released when an e- added to atom to form anion

eg. F(g) + e- F-(g)

EA = 328 kJ/mol

a small EA means e- must be forced to stick measure of ability of atom to accept e-

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Electron Affinities

Same as IEWhy?Consider1. Size2. Effective Nuclear charge

Increase

Decrease

Low for Noble gases

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Electron Affinities

1. F + e- F- EA = 328.0 kJ mol-1

1s2 2s2 2p6

- stable closed shell = Ne

2. Ne + e- Ne- EA = negative

1s2 2s2 2p6 3s1

- new shell, further from nucleus- almost totally screened from nuclear charge

- so unstable

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Electron Affinities

Be 1s22s2 Low EA

Filled s subshellNext e- higher energy level

so need energy to add e-

N 1s22s22p3 Low EA

Half filled “p”Adding another e- will cause e- repulsion

hence unfavourable

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Electronegativity

The ability of an atom to draw e- to itself in a chemical bond

useful for predicting extent of chargetransfer between atoms

eg. “Covalent” “Ionic”

H—HC—HN—HNaCl

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Electronegativity

Related to EA and IE

Cs and F IE1 EA Cs low small F high large

(Cs gives up e- easily, while F accepts e- easily.)

Electron acceptorElectronegative

Electron donorElectropositive

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Electronegativity

Increase

Decrease

( size, nuclear charge )

Size, same effective nuclear charge

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Electronegativity and Bond Type

Numerical scale of electronegativities developed

Paulings electronegativity scale caesium =0.79

fluorine =3.98

For two bonded atoms

(), is a measure of the bond polarity

with the more electronegative atom having more of the electron density

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Electronegativity and Bond Type

H+Cl-

ExamplesNa Cl(0.93) (3.16) () = 2.23, ionic, Na+Cl-

H Cl (2.20) (3.16) () = 0.96, polar covalent

Cl Cl (3.16) (3.16) () = 0, covalent, Cl-Cl

>2.0

<0.4