1. Chapter 7
Electronic Structure
of Atoms

2. Recall:
• The concept of atoms proposed by Dalton
explained many important observations:
- such as why compounds always have the
same composition and how chemical rxns.
occur
• Once chemists were convinced of the
existence of atoms…..they naturally began to
ask what atoms looked like?

3. The studies of Thomson, Rutherford and
Chadwick
• Lead to our picture of the atom:
- which includes a solid dense nucleus
containing protons and neutrons
about which we find the negatively charged
electrons

4. In this Chapter we will look at the atomic
structure in more detail.
• In particular we will develop a picture of the electron
arrangements in atoms
• This picture will allow us to account for the chemistry of
the elements
• With this knowledge we will revisit the arrangements of
the elements in the periodic table and see why there
are striking differences in the chemical properties of
elements in the groups

5. Electromagnetic Radiation and Energy
• The sun is a central
source of energy
• Energy from the sun
travels through space in
the form of
electromagnetic
radiation

6. Forms of Electromagnetic Radiation
• Visible light is a form of electromagnetic radiation
• Microwaves are another form of electromagnetic
radiation
• X-rays are yet another form of electromagnetic
radiation
***all of these forms of energy are different, but they are
all forms of electromagnetic radiation

7. Electromagnetic Radiation
• All forms of electromagnetic radiation exhibit
wave-like behavior and travel at the speed of
light in a vacuum (airless space)
speed of light = 3.8 x 1010
cm/s or 186,000 mi/s

8. Waves
• Waves have 3 primary characteristics:
1. Wavelength- λ (Greek letter lambda) is the distance between 2
consecutive peaks or troughs in a wave
2. Frequency- υ (Greek letter nu) it indicates how many waves
pass a given point per second
3. Speed- indicates how fast a given peak moves through space

9. • Electromagnetic radiation from the sun is divided into various
classes (forms) according to λ (wavelength)
• Energy from the sun reaches the earth in the form of visible and
ultra-violet light
• Hot coals emit infrared radiation
• Microwave ovens use microwave radiation to heat food
**** all have different wavelengths

10. Visible Light
• Also known as white light
• When passed through a prism
it separates into a continuous
range of colors with one
gradually blending into
another called the continuous
spectrum(rainbow)
• Separation is due to different
speeds in the prisim
• Each color of light has a
specific wavelength

11. Continuous Spectrum
• Violet has the shortest wavelength
• Red has the longest wavelength

12. Wavelength and energy
• Wavelength and energy have an inverse relationship, as
shown below
• h is Planck’s constant (6.626 × 10-34
J·s)
• c is the speed of light
λλ
hc
EE =∝
1

13. Energy and Wavelength
• Red light with the longest wavelength has the lowest
energy
• Purple light with the shortest wavelength has the
highest energy

14. The Nature of Energy
• The wave nature of light
does not explain how an
object can glow when its
temperature increases.
• Max Planck explained it
by assuming that energy
comes in packets called
quanta.

15. Photoelectric Effect
A freshly polished,
negatively charged
zinc plate looses its
charge if it is
exposed to
ultraviolet light.
This phenomenon
is called the
photoelectric
effect.

16. The Nature of Energy
• Einstein used this
assumption to explain the
photoelectric effect.
• He concluded that energy is
proportional to frequency:
E = hν
where h is Planck’s constant,
6.63 × 10−34
J-s.

17. The Nature of Energy
• Therefore, if one knows the
wavelength of light, one can
calculate the energy in one
photon, or packet, of that
light:
c = λν
E = hν

18. The Nature of Energy
Another mystery
involved the emission
spectra observed from
energy emitted by
atoms and molecules.

19. Heating an Element
• When an element is heated strongly to the
point that the element changes phase to a
gas, the gaseous atoms emit light like the sun
• One might think that the light produced would
result in a continuous spectrum like light
emitted from the sun…………instead…….

20. When an element is heated..
• Only definite or discrete colors are produced
• Since colors are discrete (definite) and the colors
correspond to energies
• The energy being emitted by the atoms of the
element is also discrete

21. An Elements Spectrum
• Is unique to that element
• Are called Atomic Emission Spectra or Line
Spectra
• Are the basis of a fireworks display

22. The Nature of Energy
• Niels Bohr adopted Planck’s
assumption and explained
these phenomena in this way:
1. Electrons in an atom can only
occupy certain orbits
(corresponding to certain
energies).

23. The Nature of Energy
• Niels Bohr adopted Planck’s
assumption and explained
these phenomena in this way:
2. Electrons in permitted orbits
have specific, “allowed”
energies; these energies will not
be radiated from the atom.

24. The Nature of Energy
• Niels Bohr adopted Planck’s
assumption and explained
these phenomena in this way:
3. Energy is only absorbed or
emitted in such a way as to
move an electron from one
“allowed” energy state to
another; the energy is defined
by
E = hν

25. The Nature of Energy
The energy absorbed or emitted from
the process of electron promotion or
demotion can be calculated by the
equation:
∆E = −RH ( )1
nf
2
1
ni
2
-
where RH is the Rydberg constant,
2.18 × 10−18
J, and ni and nf are the
initial and final energy levels of the
electron.

26. Electronic Structure
Good Points
• Electrons in Quantized
Energy Levels
• Maximum # electrons in
each n is 2n2
• Sublevels (s,p,d,f) and #
electrons they hold
Bad Points
• Electrons are placed in
orbits about nucleus
• Only explains emission
spectra of H2
• Does not address all
interactions
• Treats electron as
particle

27. The Wave Nature of Matter
• Louis de Broglie postulated that if light can
have material properties, matter should
exhibit wave properties.
• He demonstrated that the relationship
between mass and wavelength was
λ =
h
mv

28. The Uncertainty Principle
• Heisenberg showed that the more precisely the
momentum of a particle is known, the less
precisely is its position known:
• In many cases, our uncertainty of the
whereabouts of an electron is greater than the
size of the atom itself!
(∆x) (∆mv) ≥
h
4π

29. Dual Nature of Electron
Previous Concept;
A Substance is Either Matter or Energy
• Matter; Definite Mass and Position
Made of Particles
• Energy; Massless and Delocalized
Position not Specificed
Wave-like

30. Dual Nature of Electron
• Electron is both “particle-like” and “wave-like”
at the same time.
• Previous model only considered “particle-like”
nature of the electron

32. Orbitals Replace Orbits
• Bohr Orbits- Both electron position and
energy known with certainty
• Orbitals – Regions of space where an
electrons of a given energy will most likely be
found

33. Quantum Theory
Orbitals Replace Orbits
Orbits Orbitals

34. Quantum Mechanics
• Erwin Schrödinger
developed a mathematical
treatment into which both
the wave and particle
nature of matter could be
incorporated.
• It is known as quantum
mechanics.

35. Schrodinger Wave Equation (Ψ)
Describes size/shape/orientation of orbitals
7.5
• Wave Equation is based on…
1. Dual Nature of Electron (Electron both
particle and wave-like at the same time.)
2. Heisenberg Uncertainty Principle
(Orbitals describe a region in space
an electron will most likely be.)

36. Wave Equation (Ψ)
• Wave Equation describes the size, shape,
and orientation of the orbital the electron
(of a given energy) is in. There are four
variables in the function
-n; Energy and size of orbital
– l; Shape of orbital
– ml; Orientation of orbital
– ms; Electron Spin
(n, l, ml, ms)

37. Quantum Mechanics
• The square of the wave
equation, ψ2
, gives a
probability density map of
where an electron has a
certain statistical likelihood of
being at any given instant in
time.

38. 1. Each electron has a unique set of 4
Quantum Numbers
2. Each orbital described by the Quantum
Numbers can hold a maximum of 2
electrons.

39. Principal Quantum Number, n
• The principal quantum number, n,
describes the energy level on which the
orbital resides.
• The values of n are integers ≥ 0.
• n= 1, 2, 3, 4, ….
distance of e-
from the
nucleus

40. Azimuthal Quantum Number, l
• This quantum number defines the shape of
the orbital.
• Allowed values of l are integers ranging
from 0 to n − 1.
• We use letter designations to communicate
the different values of l and, therefore, the
shapes and types of orbitals.

41. Azimuthal Quantum Number, l
Value of l 0 1 2 3
Type of orbital s p d f

42. Magnetic Quantum Number, ml
• Describes the three-dimensional
orientation of the orbital in space.
• Values are integers ranging from -l to l:
−l ≤ ml ≤ l.
• Therefore, on any given energy level, there
can be up to 1 s orbital, 3 p orbitals, 5 d
orbitals, 7 f orbitals, etc.

43. Magnetic Quantum Number, ml
• Orbitals with the same value of n form a shell.
• Different orbital types within a shell are subshells.

44. s Orbitals
• Value of l = 0.
• Spherical in shape.
• Radius of sphere
increases with increasing
value of n.

45. s Orbitals
Observing a graph of
probabilities of finding an
electron versus distance
from the nucleus, we see
that s orbitals possess
n−1 nodes, or regions
where there is 0
probability of finding an
electron.

46. p Orbitals
• Value of l = 1.
• Have two lobes with a node between them.

47. d Orbitals
• Value of l is 2.
• Four of the five
orbitals have 4
lobes; the other
resembles a p
orbital with a
doughnut
around the
center.

48. f-orbitals

49. Orbital Shapes
Orbital Type Shape Name
s Spherical
p Dumbbell
d Complex
f More complex

50. Energies of Orbitals
• For a one-electron
hydrogen atom, orbitals
on the same energy
level have the same
energy.
• That is, they are
degenerate.

51. Energies of Orbitals
• As the number of
electrons increases,
though, so does the
repulsion between
them.
• Therefore, in many-
electron atoms, orbitals
on the same energy
level are no longer
degenerate.

52. Spin Quantum Number, ms
• In the 1920s, it was
discovered that two
electrons in the same
orbital do not have exactly
the same energy.
• The “spin” of an electron
describes its magnetic
field, which affects its
energy.

53. Spin Quantum Number, ms
• This led to a fourth
quantum number, the spin
quantum number, ms.
• The spin quantum number
has only 2 allowed values:
+1/2 and −1/2.

54. Pauli Exclusion Principle
• No two electrons in the
same atom can have
exactly the same energy.
• For example, no two
electrons in the same atom
can have identical sets of
quantum numbers.

55. How many 2p orbitals are there in an atom?
2p
n=2
l = 1
If l = 1, then ml = -1, 0, or +1
3 orbitals
How many electrons can be placed in the 3d
subshell?
3d
n=3
l = 2
If l = 2, then ml = -2, -1, 0, +1, or +2
5 orbitals which can hold a total of 10 e-
7.6

57. Three Manners to Convey How
Electrons are Arranged
1. Electron Configuration ; List Orbitals and
Number of Electrons in Each
(1s2
2s2
2p6
3s2
…)
2. Quantum Numbers (2,0,0,+1/2)
3. Orbital Diagrams; List Orbitals and show
location of electrons and their spin
1s 2s 2p

58. Electron Configurations
• Distribution of all
electrons in an atom
• Consist of
– Number denoting the
energy level

59. Electron Configurations
• Distribution of all
electrons in an atom
• Consist of
– Number denoting the
energy level
– Letter denoting the type of
orbital

60. Electron Configurations
• Distribution of all
electrons in an atom.
• Consist of
– Number denoting the
energy level.
– Letter denoting the type of
orbital.
– Superscript denoting the
number of electrons in
those orbitals.

61. Writing Atomic Electron
Configurations
Two ways of writing configurations:
• One is called the spdf notation.
spdf notation for H, atomic number = 1
1 s1
no. of
electrons
value of lvalue of n

62. Writing Atomic Electron
Configurations
• The spdf notation can be shortened using
the noble gas notation.
spdf notation for K, atomic number = 19
1s2
2s2
2p6
3s2
3p6
4s1
OR
[Ar]4s1
core e-
valence e-

63. Orbital Diagrams
• Each box represents one
orbital.
• Half-arrows represent the
electrons.
• The direction of the arrow
represents the spin of the
electron.

64. Hund’s Rule
“For degenerate
orbitals, the lowest
energy is attained when
the number of electrons
with the same spin is
maximized.”

65. 7.7
Orbital Diagrams
1s 2s 2p
Carbon; 6 electrons
Electron Configuration; 1s2
2s2
2p2
Orbital Diagram

66. Paramagnetic
unpaired electrons
2p
Diamagnetic
all electrons paired
2p
7.8

67. Periodic Table
• We fill orbitals in
increasing order of
energy.
• Different blocks on the
periodic table, then
correspond to different
types of orbitals.

68. Order of orbitals (filling) in multi-electron atom
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s
7.7

69. Some Anomalies
• This occurs because
the 4s and 3d orbitals
are very close in
energy.
• Cu29
is also an anomaly
with an expected
config. of
[Ar] 4s2
3d9
and an
actual config. of
[Ar] 4s1
3d10
• These anomalies occur
in f-block atoms, as
well.

70. Ion Configurations
• To form cations from elements remove e-’s
from the subshell with the highest n.
P [Ne] 3s2
3p3
- 3e-
→ P3+
[Ne] 3s2
3p0
• For transition metals, remove ns electrons
and then (n - 1) electrons.
Fe [Ar] 4s2
3d6
loses 2 electrons → Fe2+
[Ar] 4s0
3d6

71. Na+
: [Ne] Al3+
: [Ne] F-
: 1s2
2s2
2p6
or [Ne]
O2-
: 1s2
2s2
2p6
or [Ne] N3-
: 1s2
2s2
2p6
or [Ne]
Na+
, Al3+
, F-
, O2-
, and N3-
are all isoelectronic with Ne
What neutral atom is isoelectronic with H-
?
H-
: 1s2
same electron configuration as He
8.2