Energy and. Electron configuration

1. Energy and Electrons

2. Electrons
 What are electrons?

3. Electron Configuration
 Electron configuration is the
arrangement of electrons in an atom,
molecule or other body.
 The electrons occupy specific
probability regions (known as orbitals),
whose shapes and electron capacity
vary.

4. Electron States
 The lowest energy level that an
electron can exist in is the ground
state (n=1,2,…7).
 Higher energy levels are called
excited states (n>1).
 In a given orbital, an electron that
never radiates or absorbs energy is
said to be in a stationary state.

5. Energy as Light
 If an atom absorbs exactly the right
amount of energy, the electron will
rise to the next orbital or energy
level.
 If an atom released an exact amount
of energy, it would fall to a lower
orbital or energy level.
 The energy released appeared as a
photon of light.

6. Where are the electrons?
 According to the Bohr model, electrons
are found in energy levels around the
nucleus.
 There are 7 different energy levels
 The farther the energy level is from the
nucleus, the greater the amount of
energy it holds.

7. To Summarize ...
Principal
Energy Level
Type(s) of
Sublevel
Number of
Orbitals
Maximum #
of Electrons
1 s 1
1 total
2
2 total
2 s
p
1
3
4 total
2
6
8 total
3 s
p
d
1
3
5
9 total
2
6
10
18 total
4 s
p
d
f
1
3
5
7
16 total
2
6
10
14
32 total
Level n n types n2
orbitals 2n2
electrons

8. For example ...
 Calculate the number and types of
sublevels, the total number of
orbitals, and the maximum number
of electrons with an energy level of 6
 n = 6
 types of orbitals = n = 6
 there are n2 orbitals on level 6 = 36
 there are 2n2 electrons = maximum 72
electrons

9. Which state do you live in?
 For each element, the atomic number
= number of protons = number of
electrons
 The electron configuration must equal
the number of protons in the atom and
advance by 2n2
 Advance by 2 electrons, 4, 9, 16, etc.

10. Electron Behavior
 An electron can orbit in specified
energy levels (orbitals). The further
the orbital from the nucleus, the
greater the energy level.
 Orbital: a region in an atom where there
is a high probability of finding electrons

11. How many can fit?
 Each level can hold a specific number
of electrons
 n = the energy level
 2n2 = the maximum number of electrons
that energy level can hold
 So, the first energy level can hold
 2 x 12 = 2 electrons

12. Draw a Bohr Atom: Ground
State Configuration
 Put all of the protons in the nucleus
 p+
 Determine the number of electrons in
the atom
 e- = p+
 Add your electrons to each energy
level completely filling one level before
moving to the next highest level.
 Remember the max is 2n2 electrons in
each level

13. Bohr-ing Electrons!
 Atomic Number = 13
13p+ 2e- 8e- 3e-
For a total of 13 electrons

14. You practice a few
 Draw the ground-state electron
configuration for
 Oxygen
 Potassium

15. Oxygen
 Atomic Number (# electrons) = 8
 Maximum # electrons = 8 = 2n2
 Level & Types of orbitals = n = 2
 Number of orbitals = n2 = 4
 Fill the electron orbitals like this:
 2 + 6

16. Potassium
 Atomic Number (# electrons) = 19
 n = 3.08 = 3
 Fill the orbitals 2 + 8 + 9

17. Classwork
 Draw the Bohr electron diagrams for
the 1st 20 elements

18. Quantum Numbers
 To define the region in which electrons
can be found, there are 4 quantum
numbers assigned
 Quantum number: a number that
specifies the properties of electrons

19. The principal of it all
 The principal quantum number, n,
indicates the main energy level
occupied by the electron
 The values for n are in positive whole
integers (1, 2, 3, 4)
 As n increases, the distance from the
nucleus increases

20. Next comes the l
 The angular momentum quantum
number, l, indicates the shape of the
orbital
 If l = 0, then there is an s orbital
 If l = 1, p orbital
 If l = 2, d orbital
 If l = 3, f orbital

21. Okay, now with the m
 Next comes the magnetic quantum
number, m, which is a subset of the l
quantum number
 This number indicates the numbers
and orientations of the orbitals
 The number of orbitals include 1s, 3p,
5d and 7f orbitals.

22. Last, but not least, 
 The spin quantum number, +1/2 or -1/2
(), indicates the orientation of the
electron’s magnetic field relative to an
outside magnetic field

23. Orbitals
 Each orbital is associated with a different
letter: s, p, d, f, g, . . .
 As chemists, we will only look at the s, p, d
and f orbitals
 Each orbital can accommodate only 2
electrons with opposite spins
 Empty, half-filled and filled orbitals contain
0, 1 and 2 electrons, respectively

24. Pauli Exclusion Principle
 The principle that states that two
particles of a certain class cannot be in
the exact same energy state.
 In other words, only 2 electrons can
occupy a single orbital

25. Heisenburg Uncertainty
Principle
 You cannot know simultaneously
 Where an electron is, and
 How fast it is moving

26. S Orbitals
 There is only one type of s orbital and
it is present on every principal energy
level.
 The s orbital is spherical

27. The p Orbital
 There are three types of p orbitals (px,
py and pz)
 p orbitals are on every energy level
except level 1
 All p orbitals have a dumbell shape

28. The d Orbitals
 There are 5 types of d orbitals and
they are located on every energy level
except for level 1 and level 2

29. The f Orbitals
 There are 7 types of f orbitals and they
are located on every energy level
except for levels 1, 2 and 3
 The shapes of the f orbitals are
extremely complex

30. Aufbau Principle
 The principle that states that the
structure of each successive element
is obtained by adding one proton to the
nucleus of the atom and one electron
to the lowest energy orbital.
 In other words, the electrons must fill
the lowest energy level available

31. The Diagonal Rule
 The basic rule for assigning electrons to
atoms is that electrons should occupy
the lowest energy state possible.
 To determine the relative energies or
sublevels, use the diagonal rule.
 Work from left and follow each arrow
from tail to head and work from left to
right.

32. Start-up: October 7
 What is an electron configuration?

33. The Chart
7s 7p
6s 6p 6d
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s

34. The Arrows
7s 7p
6s 6p 6d
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
Start Here

35. Write the electron filling pattern
for the first 5
sublevels.
Element Atomic # Electron Config.
H 1 1s1
He 2 1s2
Li 3 1s22s1
Be 4 1s22s2
B 5 1s22s22p1

36. Shortcut
 In order to conserve space, we use
shorthand
 When you reach a noble gas, the next
element will begin a new principal
energy level (He, Ne, Ar, Kr, Xe, Rn)
 Place the symbol in brackets in place
of the configuration scheme before it

37. Sodium, for example
 Neon Atomic Number = 10
 Electron Configuration:
 1s2 2s2 2p6
 Sodium Atomic Number = 11
 Electron Configuration:
 [Ne] 3s1

38. Classwork
 Write the electron configuration of
Bromine using the diagonal rule and
the “shorthand” notation
 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5
 [Ar] 4s2 3d10 4p5 or [Ar] 3d10 4s2 4p5

39. Fun Facts
 The orbital names s, p, d, and f stand
for names given to groups of lines in
the spectra of the alkali metals. These
line groups are called sharp, principal,
diffuse, and fundamental.

40. Homework
 P. 99 numbers 1 - 11