1. ELECTRONS IN ATOMS

2. Bohr’s Model of the Atom
1913 Niels Bohr
Linked electron with photon emission
Electrons circle the nucleus in exact paths
Paths farther from the nucleus are higher in energy
When electrons drop to lower energy levels, the photon emitted is equal to the energy difference between the levels

3. Problems with Bohr’s Model
It only worked for hydrogen, atoms with more electrons did not fit the model
It did not fully explain the chemical behavior of atoms

4. Quantum Model
Electrons as Waves – If light has a dual nature, could electrons also?
DeBroglie – electrons considered waves confined to the space around the nucleus – electrons could exist only at certain frequencies – quantized energy levels
Electrons like waves can be bent or diffracted; show interference

5. Heisenberg Uncertainty
Electrons are detected by their interaction with photons
Uncertainty Principle: It is impossible to determine simultaneously both the position and velocity of an electron or any other particles

6. Schrodinger Wave Equation
1926 – Used dual nature to develop equation that treated electrons as waves
Like Bohr – electrons exist in quantized energy levels
Unlike Bohr – Electrons do not travel in exact pathways, but are located in orbitals of electron density probability

7. Atomic Orbitals & Quantum Numbers
Electron Cloud – surface drawn where electrons are likely to be found (orbital)
Quantum Numbers – Mathematical description of electrons in atoms
90% probability

8. Principal Quantum Number
Main energy level; n= 1,2,3..
As n increases, energy and distance from the nucleus increases
Always equals the number of sublevels within the principle energy level

9. Shape of orbital
Indicates the shape of an orbital; the number of shapes possible = n
Also referred to as sublevles
Its value ranges from 0 to n-1
Shapes – spherical (s,), dumbbell (p,), diffuse (d,), and more complex (f,)

10. Electron Configurations
Defined: Ways in which electrons are arranged around the nuclei of atoms, from lowest to highest energy.

11. Types of Sublevels/Orbitals
s – 1 orbital – 2 electrons
p – 3 orbital – 6 electrons
d – 5 orbitals – 10 electrons
f – 7 orbitals – 14 electrons
Remember each orbital holds up to 2 electrons.

12. Summary of Electron Positions
Principal Energy Level
Number of Sublevels
Types of Sublevels
Number of Electrons 1 1s 2
2s, 2p 8 3
3s, 3p 3d 18 4
4s, 4p, 4d, 4f 32

13. Hydrogen Line-Emission
Ground State
Excited State
When energy is passed through hydrogen gas, an electron is excited to higher energy levels. When the electron falls back to its ground state, the energy is released as electromagnetic radiation. Visible light can be separated into separated bands of color known as a line emission spectrum. These bands are associated with specific frequencies (energy)
DEMONSTRATIONS