This document discusses quantum numbers and their role in describing the size, shape, and orientation of atomic orbitals. It explains that there are four quantum numbers - principal, angular, magnetic, and spin. The principal quantum number determines the electron shell or energy level, while the angular and magnetic quantum numbers further specify the subshell and orbital within that subshell. The spin quantum number refers to the spin of the electron. Factors that influence ionization energy such as atomic radius, nuclear charge, and electron shielding are also summarized.

1. QUANTUM NUMBERS
CAPE Unit 1

2. SCHROEGINDER WAVE EQUATION
 Wave-particle duality of electrons
 The position of an electron is described in terms of
probability density
 Orbital
 region (volume of space around the nucleus) where
there is a high probability of finding an electron of a
given energy
 Atomic model
 3-D
 3 quantum numbers (principal, angular, magnetic)

3. QUANTUM NUMBERS
 Describe the size, shape and orientation in space of
the orbitals
 Principal Quantum Number (n)
 Energy level of the electron
 Maximum number of electrons at n level is 2n2
Energy Level No. of electrons
n = 1 2
n = 2 8
n = 3 18
n = 4 32

4. QUANTUM NUMBERS
 Angular Quantum Number (l)
 Sublevels in n & Shape of the orbitals
 s, p, d, f
 Each energy level has n sublevels
Energy Level No. of
Sublevels
Sublevels
n = 1 1 1s
n = 2 2 2s, 2p
n = 3 3 3s, 3p, 3d
n = 4 4 4s, 4p, 4d, 4f

5. QUANTUM NUMBERS
 Magnetic Quantum Number
 Number of orbitals within a sublevel
Sublevel No. of Orbitals Max. Electrons
s 1 2
p 3 6
d 5 10
f 7 14

6. QUANTUM NUMBERS
 Fourth Quantum Number????????
 Spin Quantum Number
 Each electron has a magnetic field and a spin
associated with that electron
 Pauli Exclusion Principle
 No more than two (2) electrons can occupy an orbital
 Two (2) electrons in the same orbital must have
opposite spins
 NO TWO ELECTRONS IN AN ATOM HAVE THE SAME
FOUR QUANTUM NUMBERS

7. SHAPES OF ORBITALS
CAPE Unit 1
Dr. Z. Clarke

8. SHAPES OF ORBITALS
 s orbital
 Each energy level has one s orbital
 Maximum number of electrons = 2
 Spherical
 1s and 2s orbitals are similar in shape however electron
density is closer to the nucleus for the 1s orbital

9. SHAPES OF ORBITALS
 p orbitals
 Each energy level has three (3) degenerate p orbitals
 i.e. 3 orbitals of EQUAL ENERGY
 Dumb-bell shape

10. ELECTRONIC CONFIGURATION
CAPE Unit 1
Dr. Z. Clarke

11. ELECTRONIC CONFIGURATIONS
 s orbitals have slightly lower energy than the p
orbitals at the same energy level i.e. 2s < 2p
 s orbital will ALWAYS fill before corresponding p
orbitals
 s orbital have the lowest energy then p, d, f
 s < p < d < f

12. ELECTRONIC CONFIGURATIONS
 Anomaly
 Irregularity in the position of the 3d and 4s orbitals
 3d has slightly more energy than 4s
 4s fills first then 3d orbitals followed by 4p orbitals

13. ELECTRONIC CONFIGURATIONS
 Describes the arrangement of electrons in the
orbitals of an atom
 How are electronic configurations worked out?
 Electrons are added one at a time, starting with the
lowest energy orbital (Aufbau Principle)
 No more than two electrons can occupy an orbital (Pauli
Exclusion Principle)
 Electrons fill degenerate orbitals one at a time with
parallel spin before a second electron is added with
opposite spin (Hund’s Rule)

14. ELECTRONIC CONFIGURATIONS
 How do we write electronic configurations?
 Principal Quantum number (1, 2, 3 etc)
 Symbol for the orbital (s, p, d, f)
 Superscript that shows the number of electrons in the
sublevel
 number of electrons in orbital
energy level1s2
type of orbital

15. ELECTRONIC CONFIGURATIONS
Atomic Number Symbol Electronic
Configuration
1 H 1s1
2 He 1s2 or [He]
3 Li [He] 2s1
4 Be [He] 2s2
5 B [He] 2s2 2p1
6 C [He] 2s2 2p2
7 N [He] 2s2 2p3
8 O [He] 2s2 2p4
9 F [He] 2s2 2p5
10 Ne [He] 2s2 2p6 or [Ne]

16. ELECTRONIC CONFIGURATIONS
Atomic Number Symbol Electronic
Configuration
11 Na [Ne] 3s1
12 Mg [Ne] 3s2
13 Al [Ne] 3s2 3p1
14 Si [Ne] 3s2 3p2
15 P [Ne] 3s2 3p3
16 S [Ne] 3s2 3p4
17 Cl [Ne] 3s2 3p5
18 Ar [Ne] 3s2 3p6 or [Ar]
19 K [Ar] 4s1
20 Ca [Ar] 4s2

17. ELECTRONIC CONFIGURATIONS
Atomic Number Symbol Electronic
Configuration
21 Sc [Ar] 4s2 3d1
22 Ti [Ar] 4s2 3d2
23 V [Ar] 4s2 3d3
24 Cr [Ar] 4s1 3d5
25 Mn [Ar] 4s2 3d5
26 Fe [Ar] 4s2 3d6
27 Co [Ar] 4s2 3d7
28 Ni [Ar] 4s2 3d8
29 Cu [Ar] 4s1 3d10
30 Zn [Ar] 4s2 3d10

18. ELECTRONIC CONFIGURATIONS –
ABBREVIATED
 He, Ne and Ar have electronic configurations with
filled shells of orbitals
 Abbreviated electronic configurations
He = 1s2 or [He]
Ne = 1s2 2s2 2p6 or [Ne]
Ar = 1s2 2s2 2p6 3s2 3p6 or [Ar]

19. ELECTRONIC CONFIGURATIONS - SPECIAL
 After 3p orbitals are filled, 4s orbital is filled before
the 3d orbital
 4s orbital is at a slightly lower energy than the 3d
 K is [Ar] 4s1
 Ca is [Ar] 4s2
 Sc is [Ar] 4s2 3d1

20. ELECTRONIC CONFIGURATIONS - SPECIAL
 After Sc, the 3d orbitals are filled
 Irregularity is seen in the electronic configuration of
Cr and Cu
 Cr is [Ar] 4s1 3d5
 Cu is [Ar] 4s1 3d10

21. ELECTRONIC CONFIGURATIONS - SPECIAL
 One electron has been transferred from the 4s
orbital to the 3d orbital
 Half-filled and filled sublevels of 3d orbitals decreases
 Energy
 Spin pairing of the 4s orbital increases
 Energy

22. IONIZATION ENERGY
CAPE Unit 1
Dr. Z. Clarke

23. IONIZATION ENERGY
 1st Ionization Energy of an element
 Energy needed to convert 1 mole of its gaseous atoms
into gaseous ions with a single positive charge
M(g) M+
(g) + e-
 Energy required to remove each successive
electron is called the 2nd, 3rd, 4th, etc. ionization
energy
 Ionization energies are positive because it requires
energy to remove an electron

24. IONIZATION ENERGY –
INFLUENCING FACTORS
 Magnitude of ionization energy
 how strongly the electron to be lost is attracted to the
nucleus
 Factors that influence ionization energy
 Atomic Radii
 Nuclear Charge
 Shielding (Screening)

25. IONIZATION ENERGY – ATOMIC RADII
 Atomic Radii
 Distance of the outer electron is from the nucleus
 As distance increases ( ), nuclear attraction for the
outer electron decreases ( ), ionization energy
decreases( )

26. IONIZATION ENERGY – ATOMIC RADII
 Successive Ionization Energies of Sodium (Na)
Ionization
Energy
Energy Orbital
Electron Lost
From
1st 496 3s
2nd 4562 2p
3rd 6912 2p
4th 9543 2p
5th 13353 2p
6th 16610 2p
7th 20114 2p

27. IONIZATION ENERGY – NUCLEAR CHARGE
 Nuclear Charge
 As nuclear charge increases, attraction of the nucleus
for the outer electron increases, ionization energy
increases
 Atomic Radii and Electron Shielding (Screening) can
outweigh the effect of nuclear charge
 Cs has a larger nuclear charge than Na, loses electron more
readily than Na

28. IONIZATION ENERGY –
SHIELDING (SCREENING)
 Screening Effect of Inner Electrons
 Electrons experience repulsion by other electrons
 Outer electrons are shielded from the attraction of the
nucleus by repelling effect of inner electrons
 Screening effect of electrons in lower energy levels is
more effective than electrons in higher energy levels

29. IONIZATION ENERGY –
SHIELDING (SCREENING)
 Screening Effect of Inner Electrons
 Electrons in same energy level has negligible screening
effect on each other
 As screening effect becomes more effective, ionization
energy decreases