# CM1001 Lecture 5

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```					Lecture 5
Electronic Configuration of Periodic Table

shells

1s                                                1s K
2s                                       2p          L
3s                                       3p          M
4s       3d            3 d               4p          N
5s       4d            4 d               5p          O
6s       5d            5 d               6p          P
7s       6d   6 d                       p block
s block              d block

4 f
5 f
f block
Electronic Configuration of Periodic Table

• Aufbau Principle: Add one proton to nucleus and one electron
to the lowest-energy orbital available in the electron shell.
• s block: 2 boxes  1 orbital  2 electrons. For s: l= 0; m= 0
i.e. 1 orbital on s subshell.
• p block: 6 boxes  3 orbitals  with 2 electrons each.
For p: l=1; m= -1, m= 0, m= 1 i.e. 3 orbitals on p subshell
• d block: 10 boxes  5 orbitals  with 2 electrons each.
For d: l = 2; m = -2, m = -1, m= 0, m =1, m =2 i.e. 5 orbitals
on d subshell
• f block: 14 boxes  7 orbitals  with 2 electrons each.
For f: l =3; m =-3, m =-2, m =-1, m =0, m =1, m =2, m =3 i.e.
7 orbitals on f subshell
Electron Energy Diagrams

• A visual method of showing the energy of electrons within an
element: 3 steps to build an energy diagram:
•   e.g. Iron [Fe] Z=26  26 protons  26 electrons
•   Step 1: Write down shorthand electronic configuration.
•   Step 2: Build energy diagram with subshells
•   Step 3: Fill in electrons with arrows counting up to Z.

Step 1    1s2 2s2 2p6 3s2 3p6 4s2 3d6 (short hand)
Electron Energy Diagrams

• A visual method of showing the energy of electrons within an
element: 3 steps to build an energy diagram:
•   e.g. Iron [Fe] Z=26  26 protons  26 electrons
•   Step 1: Write down shorthand electronic configuration.
•   Step 2: Build energy diagram with subshells
•   Step 3: Fill in electrons with arrows counting up to Z.

Step 1    1s2 2s2 2p6 3s2 3p6 4s2 3d6 (short hand)
Electron Energy Diagrams

Energy
3d
4s

3p            Step 2
3s

2p            Step 3
2s

1s
Valence Electrons
• Every element has both core electrons and valence electrons,
e.g. Magnesium: Mg Z=12  12 electrons:
1s2 2s2 2p6 3s2
core electrons        valence electrons
• Core electrons are electrons in fully filled shells
• Valence electrons are electrons in the outermost shell that is
not fully filled with the exception of the noble gases that all
have fully filled shells
He: 1s2,                     Ne: {He} 2s2 2p6,
Ar: {Ne} 3s2 3p6,            Kr: {Ar} 4s2 3d10 4p6,
Xe: {Kr} 5s2 4d10 5p6.       Rn: {Xe} 6s2 4f14 5d10 6p6 .
Construction of the Periodic Table

• Electron shells fill in a systematic fashion so that patterns can
be recognised in the electronic configuration.
• Elements listed in horizontal Rows are called Periods. A new
period is started each time the value of the principle quantum
number, n, increases, i.e. each time the valence electrons enter
a new shell.
• Arrange rows so that elements with similar electronic
configuration lie above one another to form vertical columns
called Groups, similar electronic configuration meaning
similar chemistry.
Group Properties

• Group 1 - elements with only one valence electron: These are
called the Alkali-Metal Group

Electronic configuration
3Lithium      Li      {He}2s1
11Sodium
Physical Properties
Na      {Ne}3s1       metals i.e good
19Potassium   K       {Ar}4s1       conductors, soft,
37Rubidium    Rb      {Kr}5s1       low melting point
55Cesium      Cs      {Xe}6s1       and boiling point
87Francium    Fr      {Rn}7s1
Group I Chemical Properties

;
.           .                            -

;
Na + :Cl     →   NaCl (Na+ and :Cl : )

;

;
Alkali + Halogen → Ionic Compound

.
Na + H2O     →         NaOH + H+
Alkali + Water →         Ionic Compound
NaOH →            Na+ + OH-

• cations: positive ions (e.g. Na+, K+, H+)
• anions: negative ions (Cl, OH)
• Me → Me+ + e (Me = Alkali metal)
Group 17 (or 7A) Halogens
• Elements with one electron less than their nearest nobel gas:
The Halogens: (Greek: halogen=salt former)

Electronic Configuration:
9Fluorine   F      {He}2s22p5
17Chlorine Cl      {Ne}3s23p5
35Bromine Br       {Ar}4s23d104p5
53Iodine    I      {Kr}5s24d105p5

• Physical Properties:
Highly coloured - volatile - non-metals - bad conductors -
occur in nature as diatomic molecules X2 (X=Halogen), e.g.
Cl2. F2 and Cl2 are gases, Br2 is a liquid and I2 is a solid.
Chemical Properties of the Halogens

X + H              HX
or better:
X2 + H2            2 HX

X + Me              MeX (Me+ + X) (see group 1)

• Aqueous solutions of HX contain high concentrations of H+,
i.e. are acids:

HX + H2O            H+ + X- + H2O
Group 18 (or 8A) - The Noble Gases
• Special group of elements within the periodic table. They all
have full electron shells and are highly non-reactive.
• Physical properties: Colourless gases (at normal temperature
and pressure) – lowest boiling and melting points of all
elements.
• Chemical properties: The most chemically unreactive of all
elements, up to recently thought to be totally unreactive and
were called Inert Gas Elements.
Summary
• Groups (columns) have similar electronic configuration and
similar chemistry, e.g., Noble gases have full shells, Alkali
metals have one electron more than Noble gases. Halogens
have one electron less than Noble gas configuration.
Summary Continued.

• Properties of Periods (Rows): Periods are characterised by the
gradual filling of valence shells – all atoms in a period have
different electronic configuration.
• Similar chemical behaviour is not expected, however a gradual
change of the metallic properties of Group 1 to the non-
metallic properties of Group 17 is expected.

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