2. Purpose
• The properties of atoms are due to the properties of the electrons in the
atom. For example, why does sodium like to lose one electron to make an
ion while chlorine likes to gain one electron to make an ion. To
understand this and much more chemical behavior, one needs to
understand the electrons better. In this lesson, we will learn to assign
electrons to specific energy levels and sublevels. At first, it may appear
difficult, but toward the end of the lesson a simple trick using the periodic
table will help you.
3. Electrons occupy different energy levels
• The electrons in an atom exist at different energy levels.
• Each energy level holds up to 2n2 electrons where n represents the energy
level:
– Energy level = 1, 2(1)2 = 2 electrons
– Energy level = 2, 2(2)2 = 8 electrons
– Energy level = 3, 2(3)2 = 18 electrons
– And so on…technically to infinity
• Example: Si with 14 electrons
– 2 electrons in energy level 1
– 8 electrons in energy level 2
– 4 electrons in energy level 3
• Note: While 18 will fit in energy level 3, silicon only has 14 electrons
and 10 were already assigned leaving only 4 to partially fill the third
energy level.
Energy
1
2
3
4
5
∞
4. Pause and Practice
• How many electrons can fit into the 6th energy level for an atom?
5. Pause and Practice
• How many electrons can fit into the 6th energy level for an atom?
– 2n2 electrons fit in an energy level
– 2(6)2 = 72 electrons
– Therefore, up to 72 electrons can fit into energy level 6 for any atom.
6. Energy levels can be broken down into separate
sublevels and orbitals
• From the previous example with energy level 6, you may ask if all of the 72
electrons that can fit into energy level 6 occupy the same space around
the atom. The answer is no. Each energy level can be further broken
down into sublevels and orbitals. We will deal with sublevels now and save
orbitals for the next lesson.
7. Every energy level (n) can be divided into n
sublevels.
• Energy level 1 – 1 sublevel – called 1s
• Energy level 2 – 2 sublevels called 2s and 2p
• Energy level 3 – 3 sublevels called 3s, 3p, and 3d
• Energy level 4 – 4 sublevels called 4s, 4p, 4d, and 4f
• Energy level 5 … and so on technically to infinity.
Energy
1
2
3
4
1s
2s
2p
3s
3p
3d 4s
4p
4d
4f
8. Stop and notice some details
• A few things to notice
– Each energy level (n) has n sublevels
– The sublevels are assigned letters in a certain order
• s – the lowest energy sublevel for each level
• p – second sublevel
• d – third
• f – fourth
• Technically they go on, but we will stop at f.
– There is no 1p sublevel since energy level 1 can only have a single
sublevel and the first one is labeled s. Similarly, there is no 2d or 3f.
– The energy of the 4s is slightly lower than the 3d. As you get to higher
energy sublevels, there is some overlap between energy levels. We
will learn how to use the periodic table to predict this.
9. Putting Electrons into sublevels
• We already learned that an energy level (n) can hold 2n2 electrons. We now
know that energy levels are broken down into sublevels. We now need to
know how many electrons can fit into each sublevel.
• Example: How many electrons fit in energy level 3 and how many in each
sublevel?
– 2(3)2 = 18 total electrons can fit in energy level 3
• 2 in 3s
• 6 in 3p
• 10 in 3d
Sublevel Maximum Number of e-
s 2
p 6
d 10
f 14
10. Summary
Energy Level Types of sublevels/
orbitals
Designations Electrons that
they can hold
1 s 1s 2
2 s
P
2s
2p
2
6
3 s
P
d
3s
3p
3d
2
6
10
4 s
p
d
f
4s
4p
4d
4f
2
6
10
14
11. Assigning electrons to orbitals according to the
Aufbau principle
• The aufbau principle states that we fill lower energy sublevels before we
place electrons into higher energy sublevels.
– Earlier in the lesson we had an example with silicon and assigned 2
electrons to energy level 1, 8 to energy level 2, and 4 to energy level 3.
We can now place the electrons into sublevels.
Energy
1
2
3
1s
2s
2p
3s
3p
3d
2 e-
8 e-
4 e-
2 e-
2 e-
2 e-
6 e-
2 e-
12. The Electron configuration shows how many
electrons are in each sublevel
• We can now write what is called the electron configuration for silicon:
– 1s22s22p63s23p2
Energy
1
2
3
1s
2s
2p
3s
3p
3d
2 e-
8 e-
4 e-
2 e-
2 e-
2 e-
6 e-
2 e-
1s2
Sublevel
Number of electrons in the sublevel
13. Stop and reflect
• By now you might be getting a little lost. Do not worry about assigning
electrons to sublevels too much just yet as it will get much easier in a few
slides. Just understand the big ideas.
– Electrons do not all occupy the same space in an atom
– Electrons occupy specific energy sublevels
14. The problem and the solution
• Problem:
– We will run into a problem if we keep trying to assign electrons like we
did with silicon when we get to a certain number of electrons. The
issue is that the 4s sublevel was actually a little lower in energy than
the 3d sublevel. If we go further up in energy, we find that the 5s is
similarly lower than the 4d and the 4f. In fact, 5s, 5p, and 6s are all
lower in energy than the 4f. This would be difficult to try to
remember.
• Solution:
– We can use the periodic table to determine the order that the
sublevels fill.
15. Labeling the periodic table
• We will use the periodic table to remember the order that sublevels fill.
Take a periodic table and do the following…
16. Label the first two columns in the first row starting
with 1s as shown. Then the first two columns in the
second row with 2s and so on.
6
7
18
8A
17
7A
Period
1
2
3
1
1A
2
2A
4
4B
13
3A
12
2B
9
8B
10
8B
11
1B
4
5
15
5A
16
6A
14
4A
5
5B
3
3B
6
6B
7
7B
8
8B
1s
2s
3s
4s
5s
6s
7s
19. Label the two rows at the bottom of the periodic
table starting with 4f as shown.
6
7
18
8A
17
7A
Period
1
2
3
1
1A
2
2A
4
4B
13
3A
12
2B
9
8B
10
8B
11
1B
4
5
15
5A
16
6A
14
4A
5
5B
3
3B
6
6B
7
7B
8
8B
1s
2s
3s
4s
5s
6s
7s
2p
3p
4p
5p
6p
7p
3d
4d
5d
6d
4f
5f
20. Summary
• Start the sublevel labeling in the following manner
– 1s
– 2p
– 3d
– 4f
• Notice how the number of columns in each block is the number of
electrons that can fit in the sublevel.
– Each s sublevel covers 2 columns, so they hold 2 electrons
– Each p sublevel covers 6 columns s, so they hold 6 electrons
– Each d sublevel covers 10 columns, so they hold 10 electrons
– Each f sublevel covers 14 columns, so they hold 14 electrons
21. We can now use the periodic table to determine
the filling order
• Go left to right in order of the elements:
– 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p
6
7
18
8A
17
7A
Period
1
2
3
1
1A
2
2A
4
4B
13
3A
12
2B
9
8B
10
8B
11
1B
4
5
15
5A
16
6A
14
4A
5
5B
3
3B
6
6B
7
7B
8
8B
1s
2s
3s
4s
5s
6s
7s
2p
3p
4p
5p
6p
7p
3d
4d
5d
6d
4f
5f
22. Let’s look at an example
• What orbitals are occupied by the electrons in a bromine
atom?
– We have to get to atomic number 35. To get there, we go through:
• 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p
18
8A
17
7A
Period
1
2
3
1
1A
2
2A
4
4B
13
3A
12
2B
9
8B
10
8B
11
1B
4 Br
15
5A
16
6A
14
4A
5
5B
3
3B
6
6B
7
7B
8
8B
1s
2s
3s
4s
2p
3p
4p3d
23. Example Continued
• How many electrons are in each sublevel:
– Count the number of elements that each sublevel passes through.
Stop at the designated element.
– 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5
– Note: Even though 4p can hold up to 6 electrons, bromine is the fifth
element in the 4p block. Also, it only has 35 electrons and 30 of them
occupied the other sublevels. This leaves 5 for the 4p.
18
8A
17
7A
Period
1
2
3
1
1A
2
2A
4
4B
13
3A
12
2B
9
8B
10
8B
11
1B
4 Br
15
5A
16
6A
14
4A
5
5B
3
3B
6
6B
7
7B
8
8B
1s
2s
3s
4s
2p
3p
4p3d
24. Pause and Practice
• Take a periodic table and label it 1s, 2s….
• What is the electron configuration of 1s2 2s2… of the following atoms?
– S
– Zn
– Pb
25. Pause and Practice
• Take a periodic table and label it 1s, 2s….
• What is the electron configuration of 1s2 2s2… of the following atoms?
– S
• 1s22s22p63s23p4
– Zn
• 1s22s22p63s23p64s23d10
– Pb
• 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p2
• Note: Notice how each sublevel is filled to its capacity except the last one
which may or may not be filled.
