1. Chapter 7: Electron Configuration and the Periodic Table
RVCC Fall 2008
CHEM 103-1&2 – General Chemistry I
Chapter 7: Electron Configuration and the Periodic Table
Chemistry: The Molecular Science, 3rd Ed.
by Moore, Stanitski, and Jurs

2. Electromagnetic Radiation
Theories about the arrangement and energy of electrons in atoms are based on experimental studies of the interaction of matter with electromagnetic radiation.
Electromagnetic radiation - consists of oscillating perpendicular electric and magnetic fields that travel through space with the same speed – 3.00 ×108 m/s or 186,000mi/s.
Electric field
Magnetic field
Light, microwaves, x-rays, and TV and radio transmissions are different kinds of electromagnetic waves

3. Electromagnetic Spectrum
energy (and frequency) decrease
wavelength increases
limit your exposure to X-rays (high energy)
radio waves are harmless (low energy)

4. Electromagnetic Radiation - waves
The wavelength, l (lambda), is the distance between any two adjacent identical points of a wave (m, cm, nm).
The frequency, n (nu), of a wave is the number of wavelengths that pass a fixed point in one second (Hz, s-1, or cycles/s).
The amplitude, A, is the height of the crest or the intensity of the radiation.
The phase, φ, is the relative position of the wave. 2

5. Electromagnetic Radiation
The frequency and the wavelength are inversely related to each other. The constant (c) is the speed of light.
Both waves are traveling the same speed.
n = c/λ
cycles/s = m/s / m/cycle
c = 3.00 × 108 m/s
(speed of light in a vacuum
is a constant!) 2

6. Example 1
Calculate the frequency of X-ray that has a wavelength of 8.21 nm.

7. Continuous Spectrum
When a white light from an incandescent lamp passes through a narrow slit and then through a prism it gives a continuous spectrum (containing light of all wavelengths or frequencies).
Prism
CD / Diffraction glasses /slit 2

8. Planck’s Quantum Theory
A heated object gives off radiation of shorter wavelengths with increasing temperatures. (Classical physics cannot explain.)
Plank’s Theory: Matter emits radiation in ‘packets’ that have a minimum (threshold) energy Equantum = h νradiation
Planck’s constant, h = × 10-34Js
Increasing filament T
shorter wavelength
higher energy 2

9. The Photoelectric Effect
(# of ejected e-)
current
increasing E
No current until minimum frequency or energy obtained.
Increasing the intensity (amplitude) of the light, only increases the current. 2

10. The Photoelectric Effect
Einstein postulated that light consists of photons (a particle!) of electromagnetic energy.
E = hν
The only energies an atom can have are 1hn, 2hn, 3hn…

11. Dual Nature of Light
The “wave” and “particle” pictures of light should be regarded as complementary views of the same physical entity.
The equation E = hn displays this duality;
h says that energy only occurs in discrete steps or quanta as a “packets” or “particle” and
n is the frequency of the associated “wave.”
There is experimental evidence for both:
wave – diffraction, refraction…
particle - photoelectric effect, line spectra… 2

12. Dual Nature of Light - particle
What is the energy of one quantum of laser light that has a frequency of 4.57 × 1014 s-1?
E = hn
= (6.626 × Js ) (4.57 × 1014 s-1)
= 3.03 × 10-19J

13. Dual Nature of Light - wave
What is the wavelength of one quantum of laser light that has an energy of 3.03 × 10-19J?
E = hn
= 6.55 × 10-7 m
= 655 nm

14. Photoelectric Effect Example 2 (Ch. 7, #31, #32)
31. Photons of light with sufficient energy can eject electrons from a gold surface. To do so requires photons with an energy equivalent to a 257nm wavelength. Will photons in the visible region ( nm) of the spectrum dislodge electrons from a gold surface?
NO CALCULATION NECESSARY!
32. To eject electrons from the surface of potassium metal requires a minimum energy of 3.69 x J. When 600nm photons shine on a potassium surface, will they cause the photoelectric effect?
= (6.626 × Js ) (2.998 × 108 ms-1)
600nm ·1m/10-9nm
= 3.31 × 10-19J

15. Atomic Line Spectra
The light emitted by a heated gas, such as hydrogen, results in a line spectrum-a spectrum showing only specific wavelengths of light (only certain energy values).
Hydrogen atoms become energized they
give off energy in the form of light.

16. Atomic Line Emission Spectra
Heated solid objects or molecules emit continuous spectra.
Excited atomic gases emit line spectra.
Each element has
a unique pattern.

17. The Bohr Model of the Hydrogen Atom
An ‘ah-ha!’ moment – quanta of light and structure of atom connected.
In 1913, Niels Bohr using the work of Einstein and Planck, offered an explanation of the line spectra and applied new quantum theory to the simplest atom, hydrogen.
Electrons can only occupy discrete positions about the nucleus. 2

18. The Bohr Model of the Hydrogen Atom
LAB!
An electron can only inhabit specific energy levels in an atom. The levels represent orbits of increasing distance (n) from the nucleus.
As n increases, the potential energy increases.
Epotential = -RH/n2

19. The Bohr Model of the Hydrogen Atom
Electrons are typically in their ‘ground state’ or lowest energy configuration.
When energy is absorbed, an electron moves from a low energy level to a higher energy level (ni=2 to nf=3,4,5,6). E=+
This is called an ‘excited state’.
Energy is emitted when an electron moves from a higher energy level back to a lower energy level (i.e. n=3 to n=2). E=-

20. Example
Calculate the wavelength of light that will be emitted when an electron in the hydrogen atom moves from…
n = 6 (E = × J) to n = 2 (E = × J)
Guess what?…
the Bohr model makes accurate predictions only for a one electron atom (Hydrogen).

21. Beyond the Bohr Model: Quantum Mechanics
Schrödinger’s Equations (1926):
Treats an electron as a standing wave. (If light can be a particle why can’t an electron be a wave?!)
The mathematical description of the behavior of electrons as waves is called “wave mechanics” or “quantum mechanics”.
Mathematical functions (wave functions, ) predict the allowed energy states of an electron and the probability of finding that electron in a given region of space.

22. Energy Levels and Orbitals
We can no longer think of an electron as having a precise orbit in an atom.
We can obtain the probability of finding an electron of a given energy and momentum at a given point around the nucleus.
Orbital is a region in space around the nucleus where
there is a high (90 %)probability of finding an electron
The main differences between orbitals and Bohr’s orbits are:
- orbital is three-dimesional – we talk about electron clouds
- we can only determine the probability of finding an electron
in an orbital.

23. Quantum Mechanical Model
An electron density (probability) map plots 2 for each point in space.
The H-atom ground-state orbital
An orbital depicts the space where an electron is most likely (99% of time) located. The quantum number n represents the most probable distance of the electron form the nucleus.

24. Energy Levels and Orbitals
A collection of orbitals within the same probable distance from the nucleus is called an electron shell or energy level
Each shell has one or more subshells within it.
Each subshell has one or more orbitals within it.

25. Quantum Numbers
There are four quantum numbers in the wave equation that describes the position of an electron.
n = principal quantum number
(the orbitals distance from the nucleus)
l = angular momentum quantum number
(the shape of the orbital)
ml = magnetic quantum number
(the orientation of the orbital)
ms = spin quantum number
(limits two electrons per orbital)
Every electron has a unique set (n, l, ml, ms) of quantum numbers.

26. Quantum Numbers
n=4
n=3
n=2
n=1
The principle quantum number, n, is a positive integer that indicates the most probable distance of an orbital from the nucleus.
Rule: n = 1, 2, 3, …
Epot.

27. Quantum Numbers
The angular momentum quantum number, l, represents subshells in the principal levels, and defines the shape of the orbital.
l =
orbital type = s p d f
“stupid pigeons don’t fly”
RULE: l = 0, 1, … n-1
If n=1, l=0 (one subshell)
If n=2, l=0 or 1 (two subshells)
If n=3, l=0 or 1 or 2 (three subshells)
If n=4, l=0 or 1 or 2 or 3 (four subshells)

28. Quantum Numbers
Principle
Shell Subshell(s)
(s)
(p)
(d)
(f)

29. Quantum Numbers The magnetic quantum number, ml
defines the orientation of an orbital in the space around the nucleus of an atom.
RULE: ml = -l …0... +l
If l=0, (s subshell), ml = 0 (one orientation)
If l=1 (p subshell), ml = -1, 0, 1 (three orientations)
If l=2 (d subshell), ml = -2, -1, 0, 1, 2 (five orientations)
If l=3 (f subshell), ml = -3, -2, -1, 0, 1, 2, 3 (seven orientations)

30. The s orbitals (spherical shape)
l = 0 (s orbital)
ml = 0
(one orientation)
Node
Node – a region in space where the probability of finding
an electron is zero.

31. The p orbitals (dumbbell shape)
l = 1 (p orbital)
ml = -1, 0, 1
(three orientations)
Node

32. The d orbitals (cloverleaf shape)
If l = 2 (d orbital)
ml = -2, -1, 0, 1, 2 (five orientations)

33. defines the two possible spin orientations.
Quantum Numbers
The spin magnetic quantum number, ms
ms = +1/2 , -1/2.
defines the two possible spin orientations.
Each orbital (s, px, dxz…) can hold two electrons,
preferred if occupying the same orbital (i.e. pz, )
or preferred if occupying different orbitals at the same energy or sublevel (px and py or dx2y2 and dz2)

34. Quantum Numbers - Summary
distance shape orientation
to nucleus
=n2 1 4 9 16
(s)
(p)
(d)
(f)

35. Trends in Quantum Numbering
nth shell has n subshells
total # of orbitals in the nth shell is n2
The number of orbitals in each subshell equals (2l +1)

36. Practice
How many subshells are there in the electron shell with the principal quantum number n=4?
4 subshells
If n=4, l=0, 1, 2, or 3
How many orbitals are in the n=4 shell?
n2 = 16 orbitals
(1 in l=0, 3 in l=1, 5 in l=2, and 7 in l=4)

37. Practice
Which of the following combination of quantum numbers is allowed? If not, why?
n = 2, l = 1, ml = 2, ms=+½
n = 3, l = 2, ml = 0, ms=-½
n = 1, l = 0, ml = 0, ms=1
n = 3, l = 3, ml = 2, ms=-½
n = 2, l = 0, ml = 0, ms=+½

38. Houses on a Hill – the e- neighborhood about the nucleus
House floor room
n=3 l=2 (d) ml=-2,-1,0,1,2
n=3 l=1 (p) ml=-1,0,1
n=3 l=0 (s) ml= 0
n=4
l=2
ml=-1
ms=+½
n=2 l=1 (p) ml=-1,0,1
n=2 l=0 (s) ml= 0
bed
ms=+½
ms=-½
n=1 l=0 (s) ml=0
nucleus

39. Electron Configurations
Before we start putting electrons in their houses (ground state configurations), there are some rules….

40. Pauli’s Exclusion Principle
No two electrons in an atom may have the same set of four quantum numbers.
(n, l, ml, ms)
Or…only two electrons can
occupy an orbital and they must have opposite spins.

41. Aufbau Principle
Electrons fill the orbitals in order, from lowest energy to highest.
px = py = pz
Principle Shell order: 1s<2s<3s…
Subshell order: 3s<3p<3d
n=2, l=0, ml=0
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p,
5s, 4d, 5p, 6s, 4f, 5d, …

42. Electron Configuration and the Periodic Table
s block
main group elements
d block
transition elements
p block
main group elements
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, …

43. Hund’s Rule C: 1s22s22p2 Orbital Notation 1s 2s 2p or
The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons in a particular set of degenerate (same energy) orbitals.
Orbital Notation
1s s p or
C: 1s22s22p2

44. Electron Configuration of Main Group Elements
condensed notation, orbital notation H
1s1 1s He
1s2 1s 2

45. Electron Configuration of Main Group Elements
core electrons (inner shells)
valence electrons (outer shells, those in highest principal quantum number, n.) Li
1s22s1 1s 2s Be
1s22s2 1s 2s 2

46. B 1s22s22p1 1s 2s 2px 2py 2pz C 1s22s22p2 1s 2s 2px 2py 2pz N
list the 4 quantum numbers
for each pair of electrons
list the 4 quantum numbers
for each pair of electrons
list the 4 quantum numbers
for each pair of electrons N
1s22s22p3 1s 2s
2px
2py
2pz 2

47. Neon has a closed shell configuration. Neon is stable and inert.
1s22s22p4 (=8) 1s 2s
2px
2py
2pz F
1s22s22p5 (=9) 1s 2s
2px
2py
2pz Ne
1s22s22p6 (=10) 1s 2s
2px
2py
2pz
Neon has a closed shell configuration. Neon is stable and inert. 2

48. Noble Gas Notation
P[Ne]3s23p3
P 1s22s22p63s23p3 Ne

49. Electron Configurations of the First Ten Elements
Table 7-5, p.295

50. Electron Configurations of next eight Elements
DO NOW: Ca
(Appendix D) e e e

51. Atom Electron Configurations
Aufbau Diagram and the Periodic Table
2 s 2p
20Ca
3 s 3p
57La
89Ac
4 s 3d 3d 4p
5 s 4d 4d 5p
6 s 5d
4 f 5d 6p
7 s 6d
5 f 6d 7p
Main group
s block
Lanthanides and actinides
f block
Transition elements
d block
Main group
p block
Block identities show where successive e- add.
Note: d “steps down”, f “steps down” again.

52. Electron Configuration of Transition Elements:
(n – 1)d orbitals are filled after ns orbitals and before np prbitals
Z = 21 Sc
1s22s22p63s23p64s23d1 or [Ar] 4s23d1 4s
3dxy 3dyz 3dxz 2

53. Electron Configuration of Transition Elements
Z = 22 Ti
1s22s22p63s23p64s23d2 or [Ar] 4s23d2
Hund’s Rule 4s
3dxy 3dyz 3dxz 2

54. From spectroscopy we know that:
Cr [Ar] 4s23d4 4s
3dxy 3dyz 3dxz
Expected:
Cr [Ar] 4s13d5 4s
3dxy 3dyz 3dxz
From spectroscopy we know that:
Half-filled d subshell, and half-filled s, is more energetically stable. 2

55. Electron Configuration of Transition Elements
1s22s22p63s23p64s23d6 or [Ar] 4s23d6
Z = 26 Fe 4s
3dxy 3dyz 3dxz 2

56. Filled d subshell, and half-filled s, is more energetically stable.
3dxy 3dyz 3dxz
Cu [Ar] 4s23d9 4s
3dxy 3dyz 3dxz
Cu [Ar] 4s13d10
Filled d subshell, and half-filled s, is more energetically stable. 2

57. Electron Configurations of Transition Metals Sc
3d1 4s2 Y
4d1 5s2 La
5d1 6s2 Ti
3d2 4s2 Zr
4d2 5s2 Hf
5d2 6s2 V
3d3 4s2 Nb
4d4 5s1 Ta
5d3 6s2 Cr
3d5 4s1 Mo
4d5 5s1 W
5d4 6s2 Mn
3d5 4s2 Tc
4d5 5s2 Re
5d5 6s2 Fe
3d6 4s2 Ru
4d7 5s1 Os
5d6 6s2 Co
3d7 4s2 Rh
4d8 5s1 Ir
5d7 6s2 Ni
3d8 4s2 Pd
4d10 Pt
5d9 6s1 Cu
3d10 4s1 Ag
4d10 5s1 Au
5d10 4s1 Zn
3d10 4s2 Cd
4d10 5s2 Hg
5d10 6s2
Note: ½ filled and filled shells have extra stability

58. Electron Configurations
Br [Ar]
4s23d104p5 4s
3dxy 3dyz 3dxz
4px
4py
4pz 2

59. Electron Configurations
After Z=57, the f-block fills 4f 5f 2

60. Valence electrons F [He]2s22p5
Valence electrons reside in the outermost shell of an atom. For the main group atoms, this includes ns+np electrons. Where n is the highest principal number.
For transition metals, this includes ns+np+(n-1)d electrons.
Valence electrons are primarily involved in chemical reactions.
Elements within a given group have the same “valence shell configuration.”
This accounts for the similarity of the chemical properties among groups of elements. 2

61. Valence electrons F [He]2s22p5 Cl [Ne]3s23p5 Br [Ar]4s23d104p5
Electron configuration of halogens – group 7A
F [He]2s22p5
Cl [Ne]3s23p5
Br [Ar]4s23d104p5
I [Kr] 5s24d105p5
Filled d-orbitals
are NOT valence e-.
Each halogen has 7 valence electrons 2

62. Lewis Electron-Dot Symbols
A Lewis electron-dot symbol is a symbol in which the electrons in the valence shell of an atom or ion are represented by dots placed around the letter symbol of the element.
[Ne]3s [Ne]3s23p [Ne]3s23p3 [Ne]3s23p5
[Ne]3s [Ne]3s23p [Ne]3s23p [Ne]3s23p6
Group I
Group II
Group III
Group IV
Group V
Group VI
Group VII
Group VIII Na . Si : P S Mg Al Cl Ar Na . Si : P S Mg Al Cl Ar Na . Si : P S Mg Al Cl Ar
Note that the group number indicates the number
of valence electrons. 2

63. Ions and Electron Configuration
Chapter 3
Group VIII Ne : Na . Si : P S Mg Al Cl Ar
Group I
Group II
Group VII
Group VIII
Group VI
Group IV
Group V
Group III X
lose electrons to be like [Ne]
cations (+)
gain electrons to be like [Ar]
anions (-)
Valence electrons are primarily involved in chemical reactions. 2

64. Ion Electron Configurations
Isoelectronic = same number and configuration of electrons. The most stable ion is isoelectronic with the nearest noble gas. [Ne] O-2 F-1 Na+1 Mg+2 Al+3 [Ar] S-2 Cl-1 K+1 Ca+2

65. Ion Electron Configurations
Negative ion: add one e- for each “-”
Positive ion: remove one e- for each “+”
S2-
(16 + 2) = 18 e-
S [Ne] 3s2 3p4
S2- [Ne] 3s2 3p6 or [Ar]
Al+3
(13 - 3) = 10 e-
Al [Ne] 3s23p1
Al+3 [Ne]

66. Electron Configuration of Ions
Cations:
` Li 1s22s1
[He]2s1
Mg 1s22s22p63s2
[Ne]3s2 X
- 1e-
Li+ 1s2
[He] X X
- 2e-
Mg s22s22p6
[Ne] X
Everybody wants to be a noble
gas…or, at least, isoelectronic
with one.

67. Electron configuration of ions
Anions:
O 1s22s22p4
[He]2s22p4
F 1s22s22p5
[He]2s22p5
P 1s22s22p63s23p3
[Ne]3s23p3
+2e-
O s22s22p6
[Ne]
+1e-
F s22s22p6
[Ne]
+3e-
P s22s22p63s23p6
[Ar]

68. Transition Metal Ions (n-1)d e- are added last,
BUT… ns e- are lost first.
Fe [Ar] 4s2 3d6
→ Fe [Ar] 3d6
→ Fe3+ [Ar] 3d5
Mn [Ar] 4s2 3d5 → Mn2+ [Ar] 3d5
→ Mn4+ [Ar] 3d3
→ Mn7+ [Ar]

69. Practice Write the electron configuration for Te.
PPUnit11.ppt
Write the electron configuration for Te.
Write the electron configuration for Ag and Ag+1.
Write the electron configuration for the chlorine ion.
Write the electron configuration for Ba2+.
Write the electron configuration for Os.