1. Chapter 4.1 The Development of a New Atomic Model

2. Properties of Light Light as a wave:
Visible light is a type of electromagnetic radiation, along with X-rays, ultraviolet and infrared light, microwaves, and radio waves.
These form the electromagnetic spectrum.
Waves have a repetitive nature and can be measured by wavelength () & frequency ().
Wavelength unit is cm or nm.
Frequency unit is waves/sec or hertz (Hz).

3. c = speed of light or 3.00 x 108 m/s
Frequency and wavelength are related to each other through the following equation:
c = 
c = speed of light or 3.00 x 108 m/s
As wavelength increases, frequency decreases and vice versa.

4. Change cm to m and plug into equation
Example
Determine the frequency () of light whose wavelength () is 6.87 x 10-8 cm.
c = 
c = 3.00 x 108 m/s
= c/
Change cm to m and plug into equation
= 3.00 x 108 m/s
6.87 x m
 = .437 x 1018 ≈ 4.37 x 1017 Hz

5. Light as a particle:
Photoelectric effect is the emission of electrons from metal when light shines on it.
A quantum of energy is the minimum amount of energy that can be lost or gained by an atom.
A photon is a particle of electromagnetic radiation with no mass and carrying a quantum of energy.

6. Ch 4.2 Notes The Quantum Model of the Atom
Objectives:
To describe the quantum mechanical model of the atom.
To describe the relative sizes and shapes of s and p orbitals.

7. Electrons as Waves
Behavior of electrons is similar to the behavior of waves.
Electron waves can only exist at certain frequencies.

8. Heisenberg Uncertainty Principle
Werner Heisenberg had an idea on how to detect the location of electrons.
Heisenberg Uncertainty Principle: it is impossible to determine simultaneously both the position and velocity (speed) of an electron.

9. The Schrödinger Wave Equation
Developed an equation that treated electrons in atoms as waves and quantization of electron energies was an outcome of the equation.
Quantum Theory: mathematically describes the wave properties of electrons and other very small particles.

10. Principal Quantum Number ( n )
Main energy level (shell)
Size of the orbital
PERIOD #
Number of orbitals per main energy level is equal to n2.
Number of electrons = 2n2. 1s 2s
s Orbitals
– Orbitals with l = 0 are s orbitals and are spherically symmetrical, with the greatest probability of finding the electron occurring at the nucleus.
– All orbitals with values of n > 1 and l  0 contain one or more nodes.
– Three things happen to s orbitals as n increases:
1. they become larger, extending farther from the nucleus
2. they contain more nodes
3. for a given atom, the s orbitals become higher in energy as n increases due to the increased distance from the nucleus 3s
Courtesy Christy Johannesson

11. Angular Momentum Quantum Number
Indicates the shape of the orbital, represented by l.
Values of l allowed are 0 and all positive integers less than or equal to n -1.
Ex. If n=2, shapes are l = 0 and l = 1
A sublevel consists of orbitals in a main energy level with the same value of l.

12. Magnetic Quantum Number
Indicates the orientation of an orbital around the nucleus, represented by m.
Values of m are whole numbers, including zero, from –l to +l

13. Electron capacities Electron capacities
Copyright © 2006 Pearson Benjamin Cummings. All rights reserved.

14. Spin Quantum Number
A single orbital has a maximum of two electrons, and those electrons must have opposite spins.
It has two values +1/2 and -1/2

15. Ch 4.3 Electron Configurations
Electron Configurations: the arrangement of electrons in an atom.
Each element has a unique electron configuration.
Ground-state Electron Configuration is the lowest energy arrangement of the electrons for an element.

16. General Rules Aufbau Principle
Electrons fill the lowest energy orbitals first.
“Lazy Tenant Rule”

17. General Rules WRONG RIGHT Hund’s Rule
Within a sublevel, place one electron per orbital before pairing them.
“Empty Bus Seat Rule”
WRONG
RIGHT
Courtesy Christy Johannesson

18. General Rules Pauli Exclusion Principle
Wolfgang Pauli
Pauli Exclusion Principle
Each orbital can hold TWO electrons with opposite spins.
Courtesy Christy Johannesson

19. Electron Filling in Periodic Tables s p 1 2 d 3 4 5 6 7 f

20. Elements in the s - blocks
really should include He, but He has the properties of the noble gases, and has a full outer level of electrons. 31

21. The P-block p1 p2 p3 p4 p6 p5 32

22. Transition Metals - d block
Note the change in configuration.
s1 d5
s1 d10 d1 d2 d3 d5 d6 d7 d8
d10 33

23. F - block f1 f5 f2 f3 f4 f6 f7 f8 f9 f10 f11 f12 f14 f13
Called the “inner transition elements” f1 f5 f2 f3 f4 f6 f7 f8 f9
f10
f11
f12
f14
f13 34

24. 35

25. 1s1 Periodic Patterns 1st Period s-block Example - Hydrogen
# element in block
1st Period
s-block
Courtesy Christy Johannesson

26. Periodic Patterns p s d (n-1) f (n-2)
Courtesy Christy Johannesson

27. 1s2 2s2 2p3 N Notation 1s 2s 2p 7e- N Orbital Diagram 7 7
Notation
Orbital Diagram 1s 2s 2p N
7e-
Electron Configuration
1s2 2s2 2p3
Courtesy Christy Johannesson

28. Electron Filling in Periodic Tables s p 1 2 d 3 K
4s1 Ca
4s2 Sc
3d1 Ti
3d2 V
3d3 Cr
3d4 Cu
3d9 Cr
3d5 Mn
3d5 Fe
3d6 Co
3d7 Ni
3d8 Cu
3d10 Zn
3d10 Ga
4p1 Ge
4p2 As
4p3 Se
4p4 Br
4p5 Kr
4p6 4 Cr
4s13d5 Cu
4s13d10 Cr
4s13d5 4s 3d Cu
4s13d10 4s 3d

29. Order in which subshells are filled with electrons (Fig 19 pg 116)
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d …

30. Shorthand Electron ConfigurationGe
72.61 32
Shorthand Electron Configuration
Example - Germanium
[Ar]
4s2
3d10
4p2
Courtesy Christy Johannesson

31. S 16e- 1s2 2s2 2p6 3s2 3p4 S 16e- [Ne] 3s2 3p4 Notation Core Electrons
32.066 16
Notation
Longhand Configuration
S 16e-
1s2
2s2
2p6
3s2
3p4
Core Electrons
Valence Electrons
Shorthand Configuration
S 16e- [Ne] 3s2 3p4
Courtesy Christy Johannesson