1. Arrangement of Electrons in an Atom
Chapter 4
Arrangement of Electrons in an Atom

2. 4.1 Refinements of the atomic model
Models of the atom so far:
Dalton – atoms are like little “bb’s” - then the electron gets discovered
Thomson – atom is like a charged “bb”
Rutherford - Gold foil experiment – hollow charged “bb”
Bohr model of the atom (1913) – Neils Bohr – Danish Physicist
The Bohr model of the atom comes from the idea that light is waves of energy

3. The Bohr Atom (1913) All the positive charge was in the nucleus
Electrons orbited the nucleus much like planets orbit the sun (at fixed distances)
The closer the electrons to the nucleus, the less energy it has.
The farther the electron is from the nucleus, the more energy it has.

4. The Electromagnetic Spectrum
Visible light, x-rays, ultraviolet radiation, infrared radiation, microwaves and radio waves are all part of the electromagnetic spectrum

5. The Electromagnetic Spectrum
The spectrum consists of electromagnetic radiation – energy that travels like a wave
Waves can be described by the wave equation which includes velocity (c = speed of light), wavelength (λ) and frequency (ν).
Wavelength (definition) = the distance between peaks of a wave
Light through prism leads to high energy (violet) low energy (red)

6. The Electromagnetic Spectrum
ROYGBIV - colors of the visible spectrum
Bright Line Spectrum (BLS) – caused by e- emitting energy as they return to lower energy levels energy level.
heat sodium - yellow light 2 c
heat lithium - red light
elements can appear to give off the same color light, but each will have its own BLS
BLS - used to determine identity of an element
BLS - validates Bohr’s idea that electrons jump to different energy levels and give off different wavelengths of light

7. The Electromagnetic Spectrum
Light from the sun (white light) appears as a continuous spectrum of light.
Continuous Spectrum of Light (definition) = There are no discrete, individual wavelengths of light but rather all wavelengths appear, one after the other in a continuous fashion
Spectroscopy (definition) = the study of substances from the light they emit.
We will use spectroscopes (An instrument that splits light into its component colors) and flame tests to study elements because each element emits a different spectrum of light when exited .

8. Birght Line Spectrum
Bohr proposed that the energy possessed by an e- in a H- atom and the radius of the orbit are quantized (bls)
Quantized (definition): a specific value (of energy)
Like a set of stairs, the energy states of an electron is quantized – i.e. electrons are only found on a specific step
The ramp is an example of a continuous situation in which any energy state is possible up the ramp

9. Bohr’s Energy Absorption Process
Light or energy excites an e- from a lower energy level (e- shell) to a higher energy level
These energy levels are “ quantized “ (the e- cannot be in between levels), the e- disappears from one shell and reappears in another
This absorption or excitation process is called a quantum leap or quantum jump

10. Bohr’s Energy Absorption Process
Ground State Analogy = a spring and two balls
Both the atom and e- now have higher energy
This is an energy emission process and what we observe in the hydrogen line spectrum
The e- absorbs energy in the ground state and is excited to a higher level

11. Bohr’s Energy Absorption Process
When energy is added, the electron is found in the “excited state.”
The Excited State (definition) = an unstable, higher energy state of an atom
An illustration of Bohr’s Hydrogen atom (from ground to excited state):

12. Bohr’s Energy Absorption Process
The atomic line spectral lines - when an e- in an excited state decays back to the ground state
The electron loses energy, light (colors) is emitted and the e- returns to the ground state
This is another illustration of bls.

13. The Bohr Model - Summary
When an atom absorbs energy, its electrons are promoted to a higher energy level. When the electron drops back down, energy is given off in the form of light.
Each distance fallen back is a specific energy, and therefore, a specific color.
3. Since electrons can fall from level 5 to 4, 5 to 3, etc., many colors are produced.

14. The Bohr Model - Summary
Bohr also predicted that since electrons would occupy specific energy levels and each level holds a specific number of electrons
The maximum capacity of the first (or innermost) electron shell is two e-.
Any element with more than two e-, the extra e- reside in additional electron shells.

15. The Bohr Model - Summary
Electron Configurations for Selected Elements
Group IA
Lithium
VIA
Oxygen
VIIA
Fluorine
VIIIA
Neon IA
Sodium
The number of e- per shell = 2n2 (where n is the shell number)

16. Bohn Models
Draw Bohr Diagrams for the elements Save room to draw them short hand also

17. Short-Hand e- Configuration
Short Hand Bohr Model
Write the symbol of the element
Use a ) to represent each shell
Write the # of e- in each shell
Ex.
Element
Short-Hand e- Configuration
Hydrogen
H )1e-
Lithium
Li )2e- )1e-
Fluorine
       F  )2e- )7e-
Sodium
Na )2e- )8e- )1e-

18. The Truth About Bohr Models
At atomic # 19 (z = 19), there is a a break in the pattern. One would expect that energy level #3 would continue to fill up. However, the next two electrons go into the next energy level. Look at K and Ca: IA
VIIIA H
+ ) 1
IIA
IIIA
IVA VA
VIA
VIIA He 2 Li
+ ) )
2 1 Be
2 2 B
2 3 C
2 4 N
2 5 O
2 6 F
2 7 Ne
2 8 Na
+ ) ) )
2 8 1 Mg
2 8 2 Al
2 8 3 Si
2 8 4 P
2 8 5 S
2 8 6 Cl
2 8 7 Ar
2 8 8 K
+ ) ) ) ) Ca
+ ) ) ) )

19. The Truth Continued……
So, there is a relationship between the main column # and the number of outershell electrons.
Column # = the number of valence electrons
And, there is a relationship between the row # and the number of energy levels.
Row # = the number of shells
The Bohr model truly works well for the H atom only
for elements larger than H the model does not work.

20. Do ws# 1, question 1- Use short-hand configuration
Bohr Summed Up
Bohr made 2 huge contributions to the development of modern atom theory
He explained the atomic line spectra in terms of electron energies
He introduced the idea of quantized electron energy levels in the atom
The Bohr atom lasted for about 13 years and was quickly replaced by the quantum mechanical model of the atom. The Bohr model is a good starting point for understanding the quantum mechanical model of the atom
Do ws# 1, question 1- Use short-hand configuration

21. 4.2 Quantum Numbers and Atomic Orbitals
4.3 Electron Configuration

22. Quantum Numbers & Atomic Orbitals
The Bohr model describes the atom as having definite orbitals occupied by electron particles.
Schrödinger (1926) introduced wave mechanics to describe electrons – proved Bohr’s Model to be a lie
Based his idea that electrons behaved like light (photons).
Electrons show diffraction (interference) properties like light.
Treats electrons as waves that are found in orbitals.
Orbitals (definition) = clouds that show region of probable location of a particular electron.

23. Wave Mechanical Model
The Bohr model really is the wave mechanical model
There are many types of orbitals – we can see them on the periodic table

25. Subatomic Orbitals S P D Type # sublevels Total # e Shape s 1 2 sphere3 6
peanut d 5 10
dumbbell f 7 14
flower S P D

26. Quantum Numbers An electron’s address
principle (n): what shell, level, the e- is in n = 1,2,3...7
azimuthal (l): energy sub level - s, p, d, f
magnetic – orientation of orbital about the nucleus (s has only 1, p has 3, etc.)
spin - clockwise or counterclockwise (+1/2 or -1/2)

27. Label Your Periodic Tabel
On your periodic table, shade azimuthal s,p,d,f blocks different colors
Label the principal quantum numbers…1-7
Label the valence electrons across the top

28. Electron Configuration
Electron Configuration - a representation of the arrangement of electrons in an atom

29. Electron Configuration
Examples of electron Configuration
1. Li 1s22s1
2. C 1s22s22p6
# of e- in that shell
principle
azimuthal

30. Electron Configuration
Take note that after 4s is filled, 3d is than filled before 4p.
…… 6s than 4f than 5d than 6p
When writing out the electron configuration, always write your numbers in numerical order
Y 1s22s22p63s23p64s23d104p65s24d1 – NO!
Y 1s22s22p63s23p63d104s24p64d15s2

31. Electron Configuration
Examples: Be O Ca Mn

32. Electron Configuration
Examples Pb Os

33. Electron Configuration
Short Hand
Write the name of the last noble gas
Write the electron config. that follows
Ex. Fe [Ar]3d64s2
Exceptions
Cr [Ar] 3d54s1
Cu, Ag, Au- all s’s donate 1 e- to make the d orbital full
Cu [Ar] 4s13d10

34. Orbital Notation
Electrons enter orbitals in a set pattern. For the most part, they follow these rules:
1) The Aufbau Principle - electrons must fill lower energy levels before entering higher levels.

35. Orbital Notation
Orbitals are like "rooms" within which electrons "reside".
The s subshell has one s-orbital. The p subshell has three p-orbitals. The d subshell has 5 and f has 7.
Each orbital can hold at most 2 electrons

36. Orbital Notation 2. Hund’s Rule (better known as the Bus Rule)
Before any second electron can be placed in a sub level, all the orbitals of that sub level must contain at least one electron – spread out the e- before pairing them up.
3. Pauli Exclusion Principle - electrons occupying the same orbital must have opposite spin.
See a good online illustration at

37. Orbital Notation H 1s F 1s 2s 2p

38. Orbital Notation
Examples:
Li F
Na Sc

39. Orbital Notation Ca N Fe
We can also do shorthand orbital notation (outer shell only)
Ca N Fe Ag
[Kr] 4d105s1 Ag
[Kr] 4d 5s

40. Significance of Electron Configurations
Valence shell electrons - outermost electrons involved with bonding
no atom has more than 8 valence electrons
Noble gases - 8 valence electrons – least reactive of all elements
Lewis Dot structures: NSEW (cheating) also show correct way, count to 8
Lewis Dot Structures