1. Electrons in Atoms Rutherford’s model has some limitations
It did not explain the chemical properties of the elements
It did not address the electrons

2. Electrons in Atoms
According to the wave model, light consists of electromagnetic waves.
Low energy
( = 700 nm)
High energy
( = 380 nm)
Frequency  (s-1)
3 x 106
3 x 1012
3 x 1022
102
10-8
10-14

3. Electrons in Atoms
Visible light of different wavelengths can be separated into a spectrum of colors.
In the visible spectrum, red light has the lowest energy.
Violet light has the highest energy.

4. Electrons in Atoms In 1913, Niels Bohr develops a new atomic model
Bohr stated that the electrons orbit the nucleus like the planets orbit the sun.

5. Electrons in Atoms
Each possible electron orbit in Bohr’s model has a fixed energy.
The fixed energies an electron can have are called energy levels.
Each energy level further from the nucleus is of greater energy

6. Electrons in Atoms
The rungs on this ladder are like the energy levels in Bohr’s model.
A person on a ladder cannot stand between the rungs.
Similarly, the electrons in an atom cannot exist between energy levels.

7. Electrons in Atoms
The rungs on this ladder are like the energy levels in Bohr’s model.
The energy levels in atoms are unequally spaced, like the rungs in this unusual ladder.
The higher energy levels are closer together.

8. Electrons in Atoms
The Rutherford model could not explain why elements that have been heated to higher temperatures give off different colors of light.
The Bohr model explains how the energy levels of electrons in an atom change when the atom emits light.

9. Electrons in Atoms
When an electron occupies the lowest possible energy level – it is said to be in its ground state
An electron can absorb energy from an external source
Sun
Fire
Electricity

10. Electrons in Atoms
When the electron absorbs energy it can jump up to higher energy levels – this is called the excited state
However, the electron cannot stay in an excited state
When it returns to its ground state it must release the energy it has absorbed

11. Electrons in Atoms It releases the energy in the form of light
This is called the emission line spectrum
No two elements have the same line emission spectrum

12. Electrons in Atoms
Bohr’s model was based on the emission line spectrum

13. Electrons in Atoms
Unfortunately, Bohr’s model did not apply to other atoms
That led scientists to question his model
They wondered why the electron had to be located in a precise orbit

14. Electrons in Atoms That led to the Heisenburg uncertainty principle
states that it’s impossible to determine both the position and velocity of an electron

15. Electrons in Atoms
Further developments led to Schrodinger developing the quantum mechanical model
This model describes mathematically the position of electrons in an atom
It is based on the allowed energies an electron can have
It shows how likely it is to find an electron in a particular location around the nucleus of an atom.

16. Electrons in Atoms
The probability of finding an electron within a certain volume of space surrounding the nucleus can be represented as a fuzzy cloudlike region called an orbital.
The cloud is more dense where the probability of finding the electron is high.
Electron cloud

17. Electrons in Atoms
To specify the probable location of an electron in an atom, chemists use quantum numbers
Principle quantum number
Angular momentum quantum number
Magnetic quantum number
Spin quantum number

18. Electrons in Atoms
The principle quantum number indicates the main energy level occupied by the electron
Often referred to as the energy level
The energy level corresponds to the periods of the periodic table
1st energy level = period 1

19. Electrons in Atoms
The angular momentum quantum number indicates the shape of the orbital
l = s, p, d, f
The s orbitals are spherical.
The p orbitals are dumbbell shaped.

20. Electrons in Atoms
The magnetic quantum number indicates the orientation of the orbital
m = x, y, z

21. Electrons in Atoms
Parts of the periodic table corresponds to each orbital shape
Groups 1 & 2 – s block
Groups – p block
Groups 3-12 – d block
Bottom two rows – f block

22. Electrons in Atoms
The spin quantum number indicates the spin of the electron
It may be thought of as clockwise or counterclockwise.
A vertical arrow indicates an electron and its direction of spin ( or ).

23. Electrons in Atoms
The numbers and types of atomic orbitals depend on the principal energy level. (each orbital can hold up to 2 electrons)
Summary of Principal Energy Levels and Sublevels
Principal energy level
Number of sublevels
Type of sublevel
Maximum number of electrons
n = 1 1
1s (1 orbital)
n = 2 2
2s (1 orbital), 2p (3 orbitals)
n = 3 3
3s (1 orbital), 3p (3 orbitals),
3d (5 orbitals)
n = 4 4
4s (1 orbital), 4p (3 orbitals),
4d (5 orbitals), 4f (7 orbitals)

24. Electrons in Atoms
To describe the arrangement of the electrons in an atom we use electron configuration
To describe spin we use orbital notation

25. Electrons in Atoms
To determine electron configuration, follow three simple rules
The Aufbau principle states that an electron occupies the lowest energy orbital that can receive it
We fill lowest to highest energy

26. Electrons in Atoms
Increasing energy 6s 5s 4s 3s 2s 1s 6p 5p 5d 4p 4d 4f 3p 3d 2p
The aufbau diagram shows the relative energy levels of the various atomic orbitals. Orbitals of greater energy are higher on the diagram.

27. Electrons in Atoms
To determine electron configuration, follow three simple rules
The Pauli Exclusion principle states that no two electrons in the same atom can have the same set of 4 quantum numbers
Hunds Rule states that orbitals of the same energy must be occupied by one electron before it can be occupied by a second electron.

28. Electrons in Atoms

29. Electrons in Atoms
For larger atoms, the electron configuration can be tedious
We use a shorter method called the noble gas electron configuration