1. Electrons

2. Review of Atoms and Atomic Theory
So far we have described the atom as a nucleus of protons and neutrons surrounded by electrons – this works well for a simple explanation, but does not explain certain properties of elements.
For example, why do metals give off a characteristic color when heated in a flame?
Answer: Because the electrons are giving off light energy as they return to a lower energy level. More about this later!!!

3. Dalton’s Atomic Theory, Review:
All elements are composed of tiny indivisible particles called atoms
Atoms of the same element are identical. The atoms of one element are different from atoms of a different element
Atoms of different elements can physically mix together, or chemically combine
in simple whole-number ratios to
form compounds
Chemical reactions occur when
atoms are separated, joined or
rearranged.

4. JJ Thompson
JJ Thompson (who discovered the electron) realized that Dalton’s model did not take into account electrons – so he proposed a revised model often called the plum-pudding model
This had a central positively charged mass onto which electrons were just stuck to the outside.
It did not describe the number of protons and electrons or how they were arranged.

5. Ernest Rutherford
After discovering the nucleus, Ernest Rutherford proposed the nuclear atom, in which electrons surrounded a dense nucleus – leaving lots of empty space
Later experiments showed a nucleus of + charged protons and neutrally charged neutrons, surrounded by – charged electrons.
So why didn’t the – charged electrons fall into the + charged protons of the nucleus?

6. Bohr Model of Electrons
In the early 1900’s, Niels Bohr, who was a student of Rutherford, came up with a new atomic theory
Bohr proposed that electrons were in circular orbits around the nucleus, much like planets orbit the sun (thus it was called the planetary model)
Bohr suggested that electrons in a particular orbit have a fixed energy, and that electrons do not lose energy and cannot fall into the nucleus.

7. Energy Levels of Electrons
The energy level of an electron is the region around the nucleus where the electron is likely to be moving. These are also called electron shells.
These energy levels are like the rungs on a ladder – just like you can not
stand between rungs on a
ladder, electrons cannot be
between energy levels.

8. Quantums of Energy
A quantum of energy is the amount of energy needed to move an electron from its present energy level to the next high level.
The term quantum leap comes from describing this sudden change from
one level to the next.
The higher the energy
level, the farther the
electron typically is
from the nucleus.

9. The Quantum Mechanical Model (The Cloud Model)
Finally, in 1926 Erwin Schrödinger took the atom to its current shape
The modern description of the electrons in an atom is called the quantum mechanical model and is based on the mathematical equations from Schrödinger.

10. The Cloud Model
The Quantum Mechanical Model (also called the Cloud Model) does not describe electrons in an orbit like the planets.
The probability of finding an electron within a certain volume of space surrounding the nucleus is represented as fuzzy cloud, which represents where an electron would be found 90% of the time.

11. Atomic Orbitals
In the Bohr model we had orbits, but since the quantum mechanical model does not have orbits, but rather regions of probability, the regions are called atomic orbitals. (They may also be called
electron clouds.)

12. Principle Quantum Number-the main Energy Level occupied by an electron
The main energy level of an electron (or shell) is designated by means of a principle quantum number (n)
This is assigned in order of increasing energy levels: (shells) n =1,2,3,4,5,. . .
Within each shell,
there are sublevels. The
number of sublevels found
will be the same as the
value of n.

13. Where will the electrons be found?
In each shell, electrons will be found in regions of probability called atomic orbitals.
Each orbital is given a letter symbol: S, P, D, or F: primarily based on shape
Each orbital can hold
two electrons.

14. S, P, D, or F: Shapes of Orbitals
S orbitals are spherical
P orbitals are shaped like dumbbells (or peanuts): there are 3 of these: x, y, and z
D orbitals are like two sets of dumbbells at right angles to each other
F orbitals are too complex to visualize

15. Sublevel 1: 1s
For the first shell (n =1), there is one sublevel called 1s.
In the s atomic orbital, there is an equal probability of finding an
electron in any direction
from the nucleus
S=spherical

16. Sublevel 2: 2s and 2p
For the second shell (n =2), there are two sublevels: 2s and 2p
The 2s orbital is spherical, while the 2p orbitals are dumbbell shape
The 2p orbital is of higher energy level than the 2s orbital
The 2p orbital consists of three p orbitals, each aligned on the x, y & z axis
Thus the second energy level has four orbitals 2s, 2z, 2y, 2z

17. Sublevel 3: 3s, 3p, and 3d
For the third shell (n=3), there are 3 sublevels: 3s, 3p and 3d
As before, there is one 3s orbital and three 3p orbitals – but 3d has 5 orbitals, for a total of 9 orbitals
How many electrons
can Sublevel 3 hold?

18. Sublevel 4: 4s, 4p, 4d, and 4f
For the fourth shell (n=4), there are 4 sublevels: 4s, 4p, 4d and 4f
The 4th principle energy level has 16 orbitals: One 4s, three 4 p, five 4d and seven 4f orbitals
How many electrons
can the 4th sublevel
hold?

19. Electron Configuration
The arrangement of electrons in an atom is usually shown by writing an electron configuration.
Example: Hydrogen (AtomicNumber 1) is 1s1
Helium (Atomic Number 2) is 1s2
Neon (Atomic Number 10) is 1s2 2s2 2p6

20. Rule 1: Aufbau Principle
Aufbau Principle:Electrons always fill orbitals of lower energy first. 1s is filled before 2s, and 2s before 2p.
Electrons are placed in increasing energy:
1s<2s<2p<3s<3p<4s<3d<4p<5s<4d<5p<6s<4f<5d<6p<7s<5f<6d<7p

21. An Easier Way to Use the Aufbau Principle:

22. Rule 2: Pauli Exclusion Principle
This discovery was made in 1925 by German chemist Wolfgang Pauli →
No two electrons within a particular atom can have identical quantum numbers. If two electrons do occupy the same orbital, they must have opposite spin.

23. Rule 3: Hund’s Rule
When an electron joins an atom and has to choose between two or more orbitals of the same energy, the electron will prefer to enter an empty orbital rather than one already occupied.
As more electrons are added
to the atom, electrons will
half-fill orbitals of the same
energy before pairing with
existing electrons to fill orbitals.

24. Practice:Writing out Electron Configuration of Silicon
Step A
Step B
Step C
Step D
Step E
1s22s2p63s23p2

25. Now, You Practice!
Write out the electron configuration of Potassium: