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. The Hydrogen-Atom Line-Emission Spectrum
When current is passed through a gas, it goes from the ground state to the excited state.
It emits light known as the emission-line spectrum.
When an excited hydrogen atom falls to its ground state, it emits a photon of radiation.

7. 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.

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

9. 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.

10. 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.

11. 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

12. 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.

13. 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

14. Shapes of s, p, and d-Orbitals
s orbital
p orbitals
• p orbitals
– Orbitals with l = 1 are p orbitals and contain a nodal plane that includes the nucleus, giving rise to a “dumbbell shape.”
– The size and complexity of the p orbitals for any atom increase as the principal quantum number n increases.
• d orbitals
– Orbitals with l = 2 are d orbitals and have more complex shapes with at least two nodal surfaces.
• f orbitals
– Orbitals with l = 3 are f orbitals, and each f orbital has three nodal surfaces, so their shapes are complex.
d orbitals

15. Maximum Number of Electrons In Each Sublevel
Sublevel Number of Orbitals of Electrons s p d f
LeMay Jr, Beall, Robblee, Brower, Chemistry Connections to Our Changing World , 1996, page 146

16. Copyright © 2006 Pearson Benjamin Cummings. All rights reserved.

17. 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

18. 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.

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

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

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

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

23. 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

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

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

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

27. 35

28. 1s1 Periodic Patterns 1st Period s-block # element in block
Example - Hydrogen
1s1
# element in block
1st Period
s-block
Courtesy Christy Johannesson

29. 1s2 2s2 2p3 N 1s 2s 2p 7e- Notation N Orbital Notation 7
Notation
Orbital Notation 1s 2s 2p N
7e-
Electron Configuration
1s2 2s2 2p3
Courtesy Christy Johannesson

30. 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

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

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

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