2. Atomic Structure
 There are three subatomic particles which make up
atoms
 1) Protons: 1 AMU, Positive charge
 2) Neutrons: 1 AMU, Neutral charge
 3) Electrons: (virtually) 0 mass, negative Charge
Atomic Mass Unit (AMU): 1
/12
th
the mass of a single carbon atom…
actually equal to: 1.660538782(83) × 10 −24
… This is a unit used to
measure the mass of a single proton or neutron

3. YES: Quarks are fundamental matter particles that
are constituents of neutrons and protons and
other hadrons (Hadrons are particles made from quarks and/or
gluons- carrier particle of the bond). There are six different
types of quarks. Each quark type is called a flavor
(up, down, charm,strange, top, bottom).
Oh…thaaaat’s
a quark

4. What is this? This is a model of a plot of
the probability of where electrons would
be found in a hydrogen atom: Based on
the Bohr experiments

5. QuickTime™ and a
TIFF (Uncompressed) decompressor
are needed to see this picture.

6. 1) All matter is made of atoms. Atoms are
indivisible and indestructible. (True/False)
2) All atoms of a given element are identical in
mass and properties (True/False)
3) Compounds are formed by a combination of
two or more different kinds of atoms.
(True/False)
4) A chemical reaction is a rearrangement of
atoms. (True/False)

7. e- e-
Hey Baby…you get
me spinning…
where you live?
I live at 1s2
, 2s2
, 2p5
Pick me up at 6:30
Where is that? Maybe you
should draw me a map

8. Quantum numbers
N: Shells- average distance from the nucleus as well as its
energy: 1,2,3,4,5,6,7,8… more energy will mean less
stability
L: Subshells- angular momentum, which describes the
shape of the electrons orbital: 0, 1,2…
Ml: Orbitals- magnetic quantum number or orbital
describes the orientation of the orbital in space (ie: x,y,z
axis)
Ms: positive or negative spin direction when paired up in
the same orbital
AP

9. This is how far the electron is from the
nucleus.
1st energy level: ≤ 2 electrons
2nd energy level:≤ 8 electrons
3rd energy level: ≤ 18 electrons
4th energy level: ≤ 32 electrons
5th energy level: ≤ 32 electrons
Etc…
I wonder how many
electrons go into
each energy level
AP

10. L: Subshells
 Subshells are the shapes of the orbits that electrons take
around the nucleus of an electron: S, P, D, F, (G, H)
 1st
shell (N=1): 1 subshell, s (l = 0)
 2nd
shell (=2): 2 subshells s (l = 0), P (l =1)
 3rd
shell (N=3): 3 subshells, s (l = 0), p (l =1), d (l = 2)
 4th
Shell (N=4): 4 subshells, s (l = 0), p (l =1), d (l = 2), f (I =3)
 5th
Shell (N=5): 5 subshells, s (l = 0), p (l =1), d (l = 2), f (I =3), g (I = 4)
AP

11. Orbitals

12. Subshell shapes
S (l = 0)-sphere (1 per energy level)
AP

13. P (l = 1) subshell (3 per energy level
they are found in) AP

14. D (l = 2) subshell (5 per energy
level they are found in)
 5 orientations
AP

15. F (l = 3) subshell (7 per energy level
they are found in)
7 orientations
AP

16. There is also G, H, I
These are so convoluted that there is no real defined
shape… only that they exist

17. How does the l relate to Ml? It
is all about orientation

18. Orbitals: Ml
The spatial orientation… sort of X,Y, Z axis
S subshell: (l = 0), has one orbital, Ml = 0
P subshell (l = 1), has 3 orbials, Ml= -1, 0, or +1
D subshell (l = 2), has 5 orbitals, Ml = -2, -1, 0, +1, or +2
F subshell (l = 3), has 5 orbitals, Ml =-3, -2, -1, 0, +1, +2, or + 3
AP

19. Holy Guacamole…That’s pretty
complicated!!
 No worries. Just remember that each orbital of each
subshell at each energy level can hold no more than 2
electrons
 The total S orbitals for each energy level (1) X 2 = 2
 The total P orbitals for each energy level (3) X 2 = 6
 The total D orbitals for each energy level (5) X 2 = 10
 The total F orbitals for each energy level (7) X 2 = 14
That’s really all that matters when it comes to orbitals
AP

20. Ms Spin direction
When electrons share the same orbital with in their
subshells, they pair up because the electrons start
spinning opposite directions… To signify this we
indicate it as + ½, and – 1/2
AP

21. Putting it together
Quantum numbers
Write the quantum numbers for the atom of the
following:
Na: (3,0,0, ½)
Rb: ( 5, 0, 0, ½)
AP

22. Now What???
So now I know what energy levels are
and what orbitals are aaand how many
electrons I can put in each. What is my
next move???

23. Follow the Map The diagonal chart
Electrons fill in
ORDER from lowest
to highest energy
level… mostly
Think you need a GPS for this address?

24. Here’s an easy way:
The diagonal Chart
1s
2s
3s 2p
4s 3p
5s 4p 3d
6s 5p 4d
7s 6p 5d 4f
8s 7p 6d 5f
Remember:
Each s orbital is 1 x 2 = 2 e-
Each p orbital is 3 x 2 = 6 e-
Each d orbital is 5 x 2 = 10 e-
Each f orbital is 7 x 2 = 14 e-
All electrons will fill in this order at GROUND STATE:
The state of least possible energy in a physical system,
as of elementary particles. Also called ground level
What time is it?

25. Time to Hammer it home
 H: 1s1
 N: 1s2
,2s2
, 2p4
 Ne: 1s2
,2s2
, 2p6

26. s block d block p blockf block

27. Hunds Rule When an electron is added to a subshell it will always
occupy an empty orbital if there is one available (why
doesn’t this apply to S orbitals?)
 Diamagnetic: all electrons paired up in all orbitals
 Paramagnetic: 1 or more electrons not paired in any
orbitals
AP

28. What are the Ground State Electron
Configurations of the following:
Mg:
Ca:
C:
He:
Ar:

29. What observations can you make about Mg
and Ca?
Mg:1s2
, 2s2
, 2p6
, 3s2
How many Valence shell e- ? 2
Ca: 1s2
, 2s2
, 2p6
, 3s2
, 3p6
, 4s2
How many Valence shell e- ? 2
C:1s2
, 2s2
, 2p2
How many Valence shell e- ? 4
What observations can you make about He
and Ar?
He: 1s2
How many Valence shell e- ? 2
Ar: 1s2
, 2s2
, 2p6
, 3s2
, 3p6
How many Valence shell e- ? 8
What observations can you make about Mg
and Ca?
Mg:1s2
, 2s2
, 2p6
, 3s2
How many Valence shell e- ? 2
Ca: 1s2
, 2s2
, 2p6
, 3s2
, 3p6
, 4s2
How many Valence shell e- ? 2
C:1s2
, 2s2
, 2p2
How many Valence shell e- ? 4
What observations can you make about He
and Ar?
He: 1s2
How many Valence shell e- ? 2
Ar: 1s2
, 2s2
, 2p6
, 3s2
, 3p6
How many Valence shell e- ? 8
Answer

30. Orbital Diagrams
A way to represent each electron of an atom/ion
numerically and visually.
Also used to determine if an atom/ion is
paramagnetic or diamagnetic
AP

31. Example AP

32. Procedure:
Always fill one electron in each orbital and one in
each orientation (Ml) before pairing with in the same
orbital
Always fill from lowest energy level on up- if
electrons are higher than that then the atom is in an
excited state
If all electrons are paired in all orbitals, then it is
Diamagnetic, if one electron is unpaired in an orbit,
than it is paramagnetic
AP

33. Example: which is the right way to write the orbital diagram for
the electron configuration: 1S2
, 2S2
, 2 P4
? And is this Diamagnetic
or paramagnetic?
AP

34. Answer
B
and it is paramagnetic
AP

35. The Aufbau principle
In electron configurations the electrons are placed in
order of increasing energy

36. The Pauli exclusion principle
With in an atom, no two electrons can have the same
set of quantum numbers

37. Electrons and Energy
As electrons are farther from the nucleus, their
potential energy increases in the same way that yours
would as you climbed a ladder
Electron energy is Quantized, meaning that the
energy of the electron is related to the specific energy
level it is at and the energy required to keep it there….
So the higher the energy level, the higher the energy.
The electrons do not exist between energy level but at
very discrete levels (N)
AP

38. Calculating the energy of
an electron
En = -2.178 X10-18
Joules
En= the energy of the electron
N = principle quantum number of the electron
N2
AP… just know
setup… no
calculation

39. Energy and Electromagnetic
radiation
E =
hc
Where :
E = Energy released
H = Planck’s constant: 6.63 x 10-34
c = constant speed of light: 3.00 x 108
m/s
λ = wavelength
λ
AP… just know
setup… no
calculation

40. We concentrate mostly on certain electrons called Valence
electrons:
 THE VALENCE SHELL ELECTRONS:
 These are the electrons found in the outer most shell of
the atom: the S and the P orbitals of the outer most energy
level
 Valence shell electrons are determined by the roman
numeral above the groups on the periodic table
 These determine the atomic radii

41. Atomic Radius:
 Distance from the nucleus to the edge of the valence
shell cloud usually in Angstroms or picometers; all
based on probability since electrons are always
moving

42. Relative Atomic Radii
QuickTime™ and a
TIFF (Uncompressed) decompressor
are needed to see this picture.

43. Electronegativity
 The ability to take and hold electrons to itself-
electrostatic (magnetic) attraction (opposites attract)

44. Relative Electronegativity
QuickTime™ and a
TIFF (Uncompressed) decompressor
are needed to see this picture.

45. Compare Electronegativity to
atomic radius
 IN general as atomic radius increases
electronegativity…

46. Answer
 DECREASES

47. Atomic Radius and
Electronegativity
 Use the periodic table to determine the following.
 N or Cs
 Which has the highest electronegativity?
 Which has the largest radius?

48. Answer
 Highest Electronegativity: N
 Largest Radius: Cs

49. Ionization energy
The energy required to remove an electron from an atom
is called the first ionization energy … Energy required to
remove the second electron from an atom is called the
second ionization energy (requires more energy than 1st
)…
so on
Energy is required because the electron is attracted to the
nucleus electrostatically
Ionization energy tends to increase going left to right,
however there is a shielding effect an also as e- pair up in
the same P orbital, it decreases slightly
IE decreases going down the PT in the same group
because of distance and electron shielding effect
AP

50. Ionization energy sequence
Each time that you remove an electron from an atom
there is more required energy, so the more e-
removed the more the IE increases… Once the final
electron is removed which eliminates the last valence
electron there will be a large jump in energy because
the ion is now stable… example: Si
780kJ, 1575 kJ, 3220 kJ, 4350 kJ, 16100 kJ
Notice that there are 4 IEs and the last one is very
large compared to the others

51. IE periodic table
AP

52. Electron affinity
A measure of the change in energy of an atom when
an electron is added to it
If the addition makes it more stable, then energy is
given off
On rare occasions (alkaline earth metals and Nobel
gases) the addition of an e- makes the atom less stable
and therefore requires energy
AP

53. EA periodic table
AP

54. EN, EA, IE
What are the periodic trends when comparing these 3
electron characters?
They are similar
Weird!!!!!
AP

55. A historical perspective of atomic
structure

56. Review
John Dalton: Atomic theory, named elements,
conservation of mass
Dmitre Mendeleev (Lothar Meyer)- arranged
elements into periodic table

57. J.J. Thomson
Noted deflections of the cathode ray and
determined that atoms were both positive
and negative

58. The Plum pudding model: JJ
Thomson

59. Robert Millikan
Examined the behavior of charged oil drops in an
electric field
Calculated the charge on an electron

60. Ernest Rutherford
Noticed how alpha
particles scattered when
passed through gold foil
4
He2+
From this he determined that
the center of the atom
(nucleus) must be the area
that has the positive and that
the atom was mostly empty
space
2

61. Quantum Theory
Max Planck: EM energy is quantized based on a unit
called a quantum or photon (E = hv) So EM energy
does not change smoothly but in distinct steps
This would make a photon a UNIT of energy or a
particle, yet it has wave properties

62. The Bohr model

63. Werner Heisenberg
Heisenberg uncertainty principle
Stated that it is impossible to know both the position
(particle) and the momentum (wave) of an electron at
a particular instant

64. Can only calculate the probability of electron
positions

65. The de Broglie (Pr: De Broy) hypothesis
Said that all matter has wave characteristics and there
is a mathematic relationship between electron waves
and particles
λ = h
λ is wavelength, h is plancks constant, mv (mass x
velocity) momentum of the particle
mv