1. Arrangement of Electrons in Atoms (Chapter 4) Notes
Part 1 Electromagnetic Radiation

2. I. Properties of Light-Different types of electromagnetic radiation (x-rays, radio waves, microwaves, etc…) SEEM to be very different from one another. Yet they share certain fundamental characteristics. All types of electromagnetic radiation, also called radiant energy, move through a vacuum at a speed of 3.00 x l08 meters per second.

3. A. Wavelength – distance between identical points on successive waves; may be measured in any length unit but is usually dependent on how long the wave is (X-rays are usually measured in nanometers or Angstroms while the very long radio waves might be measured in meter. The Greek letter lambda, , is used to depict wavelength (pg 92)

4. B. Frequency – the number of complete wave cycles that pass a given point in one second: the unit is cycles/second but is written as sec-1, or Hertz. The Greek letter nu,, is used to depict frequency.

5. If the frequency and wavelength are known then the product of the two (wavelength x frequency) is always equal to the same speed. It is known as the speed of light or c.
c = speed of light = 3.00 x l08 m/s
c =
 = wavelength (in m)
 = frequency (in Hz)

6. What is the wavelength of radiation whose frequency is 6
What is the wavelength of radiation whose frequency is 6.24 x l013 sec-1? A: 4.81X10-6 m
2. What is the frequency of radiation whose wavelength is 2.20 x l0-6 nm? (1 m = 1,000,000,000 nm) A: 1.36X1023 s-1 or Hz

7. II. The Photoelectric Effect (pg 93) – refers to the emission of electrons from a metal when light shines on the metal.

8. The wave theory of light (early 1900) could not explain this phenomenon. The mystery of the photoelectric effect involved the frequency of the light striking the metal. For a given metal, no electrons were emitted if the light’s frequency was below a certain minimum – regardless of how long the light was shone. Light was known to be a form of energy, capable of knocking loose an electron from a metal. But the wave theory of light predicted that light of any frequency could supply enough energy to eject an electron. Scientists couldn’t explain why the light had to be of a minimum FREQUENCY in order for the photoelectric effect to occur.
2.0 electron volts = × joules

9. Let’s Review Energy – Radiation of different wavelengths affect matter differently – certain wavelengths (near infrared) may burn your skin with a heat burn, overexposure to X radiation causes tissue damage. These diverse effects are due to differences in the energy of the radiation. Radiation of high frequency and short wavelength are more energetic than radiation of lower frequency and longer wavelength.

10. THE QUANTITATIVE RELATIONSHIP BETWEEN FREQUENCY AND ENERGY WAS DEVELOPED THROUGH THE QUANTUM THEORY OF MAX PLANCK.
The explanation of the photoelectric effect dates back to 1900 when Max Planck revised classical ideas of light by proposing that light, which before was thought of as a collection of waves, consisted of BUNDLES OF ENERGY called QUANTA. A quantum is the minimum quantity of energy that can be lost or gained by an atom.
2.0 electron volts = × joules
Max Planck

11. Planck proposed the following relationship between a quantum of energy and the frequency of radiation:
E = h
h = Planck’s constant = 6.63 x l0-34 Joules  sec
E = energy (in Joules)
 = frequency (in Hz)
2.0 electron volts = × joules

12. Examples:
If a certain light has 7.18 x l0-19 J of energy what is the frequency of this light?
A: 1.08X1015 s-1or Hz
What is the wavelength, in nm, of this light?
A: 278 nm
2. If the frequency of a certain light is 3.8 x l014 Hz, what is the energy of this light? A: 2.5X10-19 J

13. Albert Einstein expanded on Planck’s theory by explaining that electromagnetic radiation has a dual wave-particle nature. While light exhibits many wavelike properties, it can also be thought of as a stream of particles. Each particle of light carries a quantum of energy. Einstein called these particles PHOTONS. A photon is a particle of electromagnetic radiation having zero mass and carrying a quantum of energy.
Albert Einstein

14. Einstein explained the photoelectric effect by proposing that electromagnetic radiation is absorbed by matter only in whole numbers of photons. In order for an electron to be ejected from a metal surface, the electron must be struck by a single photon possessing at least the minimum energy (Ephoton = hv) required to knock the electron loose, this minimum energy corresponds to a minimum frequency. If a photon’s frequency is below the minimum, then the electron remains bound to the metal surface. Electrons in different metals are bound more or less tightly, so different metals require different minimum frequencies to exhibit the photoelectric effect.

15. Example from problem 4: An atom or molecule emitting or absorbing radiation whose wavelength is 589 nm cannot lose or gain energy by radiation except in MULTIPLES OF 3.37x l0-19 J. It cannot, for example, gain 5.00 x l0-19 J from this radiation because this amount is not a multiple of 3.37 x l0-19.

16. In astronomy, it is often necessary to be able to detect just a few photons because the light signals from distant stars are so weak. A photon detector receives a signal of total energy 4.05 x l0-18 J from radiation of 540 nm wavelength. How many photons have been detected? 
A: 11 photons
Excited chromium atoms strongly emit radiation of 427 nm. What is the energy in kilojoules per photon?
A: 4.66X10-22 kJ
7. Light hitting certain chemical substances may cause rupture of a chemical bond. If a minimum energy of 332 kJ is required to break a carbon-chlorine bond in a plastic material, what is the longest wavelength of radiation that possesses the necessary energy? A: 5.99X10-31 m

17. III. The Hydrogen-Atom Line-Emission Spectrum
When investigators passed an electric current through a vacuum tube containing hydrogen gas at low pressure, they observed the emission of a characteristic pinkish glow. When a narrow beam of the emitted light was shined through a prism, it was separated into a series of specific frequencies (and therefore specific wavelengths, c =) of visible light. The bands of light were part of what is known as hydrogen’s LINE-EMISSION SPECTRUM. (page 95)
The lowest energy state of an atom is its ground state.
A state in which an atom has a higher amount of energy is an excited state. When an excited atom returns to its ground state, it gives off energy.

19. IV. Bohr’s Model of Hydrogen – Neils Bohr incorporated Planck’s quantum theory to explain line-emission spectra. Bohr said the absorptions and emissions of light by hydrogen corresponded to energy changes within the atom. The fact that only certain frequencies are absorbed or emitted by an atom tells us that only certain energy changes are possible.
Bohr’s model incorporated (l) Rutherford’s Experiment, which established a nucleus and (2) Einstein’s theory that used Planck’s quantum theory to determine that light is discrete bundles of energy.

20. V. Bohr’s Theory of the Atom:
Electrons cannot have just any energy; only orbits of certain radii having CERTAIN energies are permitted.
Thus, when an electron absorbs quanta of energy, it will cause them to jump away from the nucleus to a higher orbit (energy level or n) and when the electron falls from a high orbit to a lower one, a photon of a particular wavelength is released, and a particular color will be given off. Bohr was able to calculate a set of allowed energies. Each of these allowed energies corresponds to a circular path of a different radius.
Thus the larger the value of n, the farther the electron is from the nucleus and the higher energy it possesses.
The success of Bohr’s model of the hydrogen atom in explaining observed spectral lines led many scientist to conclude that a similar model could be applied to all atoms. It was soon recognized, however, that Bohr’s approach did not explain the spectra of atoms with more than one electron. Nor did Bohr’s theory explain the chemical behavior of atoms.

21. V. Bohr’s Theory of the Atom:

22. Electrons in Atoms (Chapter 4) Notes
Part 2 Quantum Model of the Atom

23. So where are the electrons of an atom located?
A. Various Models of the Atom
Dalton’s Model
Thomson’s Plum Pudding Model
Rutherford’s Model
Bohr’s ‘Solar System’ Model – electrons rotate around the nucleus
Quantum Mechanics Model – modern description of the electron in atoms, derived from a mathematical equation (Schrodinger’s wave equation)

24. B. In 1926, the Austrian physicist Erwin Schrodinger used the hypothesis that electrons have a dual wave/particle nature (developed by Louis de Broglie in 1924) to develop an equation that treated electrons in atoms as waves.
Erwin Schrodinger

25. Electrons as Waves Louis de Broglie (1924)
~1924
Louis de Broglie (1924)
Applied wave-particle theory to electrons
electrons exhibit wave properties
QUANTIZED WAVELENGTHS
Fundamental mode
Second Harmonic or First Overtone
Standing Wave
200
150
100 50
- 50
-100
-150
-200
200
150
100 50
- 50
-100
-150
-200
200
150
100 50
- 50
-100
-150
-200
Louis de Broglie wondered if the converse was true — could particles exhibit the properties of waves?
• de Broglie proposed that a particle such as an electron could be described by a wave whose wavelength is given by
 = h , m
where h is Planck’s constant, m is the mass of the particle, and  is the velocity of the particle.
• de Broglie proposal was confirmed by Davisson and Germer, who showed that beams of electrons, regarded as particles, were diffracted by a sodium chloride crystal in the same manner as X -rays, which were regarded as waves.
• de Broglie also investigated why only certain orbits were allowed in Bohr’s model of the hydrogen atom.
• de Broglie hypothesized that the electron behaves like a standing wave, a wave that does not travel in space.
• Standing waves are used in music: the lowest-energy standing wave is the fundamental vibration, and higher-energy vibrations are overtones and have successively more nodes, points where the amplitude of the wave is zero.
• de Broglie stated that Bohr’s allowed orbits could be understood if the electron behaved like a standing circular wave. The standing wave could exist only if the circumference of the circle was an integral multiple of the wavelength causing constructive interference. Otherwise, the wave would be out of phase with itself on successive orbits and would cancel out, causing destructive interference.
Adapted from work by Christy Johannesson 25

26. Electrons as Waves QUANTIZED WAVELENGTHS n = 4 n = 5
When we refer to a series of measurements being quantized, we are referring to the fact that they are showing up in jumps and not as a smooth, continuous function.  It would be as if an accelerating car were seen as going 5 mph, then 10 mph, then 15 mph, and so on, but not at any speeds in between. 
n = 6
Forbidden
n = 3.3
Courtesy Christy Johannesson 26

27. Schrodinger’s equation results in a series of so called wave functions, represented by the letter  (psi). Although  has no actual physical meaning, the value of 2 describes the probability distribution of an electron. (Same concept covered in Algebra II when dealing with linear regressions and finding best fit lines.)
90% probability of
finding the electron
Orbital
Electron Probability vs. Distance 40 30
Electron Probability (%) 20 10 50
100
150
200
250
Distance from the Nucleus (pm)
Courtesy Christy Johannesson 27

28. Waves are confined to a space and can only have certain frequencies.
We cannot know both the location and velocity of an electron (Heisenberg’s uncertainty principle), thus Schrodinger’s equation does not tell us the exact location of the electron, rather it describes the probability that an electron will be at a certain location in the atom. Here is an overview of electron properties:
Waves are confined to a space and can only have certain frequencies.
Electrons are considered waves confined to the space around an atomic nucleus. Electrons can only exist at specific frequencies. And according to E=hv (Planck’s hypothesis), these frequencies correspond to specific energies (or quantified amounts of energy.)
Electrons, like light waves, can be bent or diffracted.
Orbital
90% probability of
finding the electron

29. C. Heisenberg’s Uncertainty Principle says that there is a fundamental limitation on just how precisely we can hope to know both the location and the momentum of a particle. It turns out that when the radiation used to locate a particle hits that particle, it changes its momentum. Therefore, the position and momentum cannot both be measured exactly. As one is measured more precisely, the other is known less precisely.
Today we say that the electrons are located in a region outside the nucleus called the electron cloud.
Werner Heisenberg

30. Heisenberg Uncertainty Principle
Impossible to know both the velocity and position of an electron at the same time
Werner Heisenberg
~1926 g
Microscope
Werner Heisenberg ( )
The uncertainty principle: a free electron moves into the focus of a hypothetical microscope and is struck by a photon of light;
the photon transfers momentum to the electron. The reflected photon is seen in the microscope, but the electron has moved
out of focus. The electron is not where it appears to be.
A wave is a disturbance that travels in space and has no fixed position.
The Heisenberg uncertainty principle states that the uncertainty in the position of a particle (Δx) multiplied by the uncertainty in its momentum [Δ(m)] is greater than or equal to Planck’s constant divided by 4:
(Δx) [Δ(m)]  h 4
• It is impossible to describe precisely both the location and the speed of particles that exhibit wavelike behavior.
Electron 30

31. I. Electron Cloud – Energy Levels
Electrons are found in various energy levels around the nucleus. The energy levels are analogous to the rungs of a ladder. The lowest rung of the ladder corresponds to the lowest energy level. A person can climb up or down a ladder by going from rung to rung. Similarly, an electron can jump from one energy level to another. A person on a ladder cannot stand between the rungs; similarly, the electrons in an atom cannot exist between energy levels.
For convenience, instead of drawing circular orbits in which the energy is high for each successive orbit, we draw an "Energy Level Diagram" as shown on the right below. The energy levels of an atom are now represented by horizontal lines. Energy increases as you move upward in the diagram.
An analogy to the energy levels in the atom is the "Quantum Stepladder" where the rungs on the ladder correspond to energy levels in the atom.

32. Electron Absorbing Energy (Photon)
A. Quantum: To move from one rung to another, a person climbing a ladder must move just the right distance. To move from one energy level to another, an electron must gain or lose just the right amount of energy. The exact amount of energy required to move from one energy level to another is called a quantum of energy.
B. Photon: When electrons move from one energy level to another energy level we see light – going from one energy level to another energy level gives off an exact amount of light (called a photon).
Electron Absorbing Energy (Photon)
Electron will move from a ground state to an excited state.
Electron Emitting Energy (Photon)
Electron will move from an excited state to a ground state.

33. II. Quantum Mechanics Model of the Atom and Quantum Numbers
Periodic Table with predicted ending electron configurations.
Quantum Numbers – a series of numbers which describe several properties of an energy level (or orbit)

34. Quantum Numbers Four Quantum Numbers:
Specify the “address” of each electron in an atom
UPPER LEVEL
Courtesy Christy Johannesson 34

35. Quantum Numbers Principal Quantum Number ( n )
Angular Momentum Quantum # ( l )
Magnetic Quantum Number ( ml )
Spin Quantum Number ( ms )
Schrödinger used three quantum numbers
(n, l, and ml) to specify any wave functions.
• Quantum numbers provide information
about the spatial distribution of the electron. 35

36. A. Principal Quantum Number, “n” (Energy Levels): energy levels (represented by the letter n) are assigned values in order of increasing energy: n=1,2,3,4, and so forth…. which correspond to the periods in the periodic table. The principle q. n. is related to the size and energy of the orbital. n=1, n=2, n=3, n=4, n=5, etc… Which energy level is furthest away from the nucleus and has electrons with the highest energy - 1, 2,3, or 4?

37. Relative Sizes 1s and 2s 1s 2s
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 334

38. Quantum Numbers Principal Quantum Number ( n ) Energy level
Size of the orbital
n2 = # of orbitals in the energy level 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 38

39. B. Angular Momentum or Azimuthal Quantum Number, “l” (Sublevels): Within each energy level, the electrons are located in various sublevels – there are 4 different sublevels s, p, d, and f. “l” defines the shape of the orbital (s, p, d, & f). The possible values of “l” are limited by the value for “n”. If n = 3, “l” can be 0, 1, or 2, but not 3 or higher. This q.n. is related to the shape of the orbital.

40. 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 40

41. p-Orbitals px pz py
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 335

42. d-orbitals
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 336

43. Atomic Orbitals

44. Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

45. 1s 2s l = 0, is referring to the s sublevel
l = 1, is referring to p sublevel
l = 2, is referring to d sublevel
l = 3, is referring to f sublevel 1s 3p 2p 2s

46. Quantum Numbers f d s p Angular Momentum Quantum # ( l )
Energy sublevel
Shape of the orbital f d s p
Courtesy Christy Johannesson 46

47. Quantum Numbers 2s 2px 2py 2pz
Orbitals combine to form a spherical shape. 2s
2pz
2py
2px
Courtesy Christy Johannesson 47

48. Quantum Numbers Magnetic Quantum Number ( ml ) Orientation of orbital
Specifies the exact orbital within each sublevel
Courtesy Christy Johannesson 48

49. Feeling overwhelmed? Read chapter 4.2!
Chemistry
"Teacher, may I be excused? My brain is full."
Courtesy Christy Johannesson 49

50. C. Orbitals: Where are the electrons in the various sublevels located in relation to the nucleus? Electrons are NOT confined to a fixed circular path, they are, however, found in definite regions of the atoms – these regions are called atomic orbital’s! Each orbital can only hold 2 electrons at a time (Pauli exclusion principle).

51. Within the s sublevel (l=0) there is only 1 orbital (which is spherical) it is called the s orbital.
Within the p sublevel (l=1) there are 3 orbital’s (which are dumbbell shaped) called the px, py, pz orbital’s.
Within the d sublevel (l=2) there are 5 orbital’s (4 of which are cloverleaf shaped) called the dxy, dxz, dyz, dx2-y2, dz2 orbital’s.
Within the f sublevel (l=3) there are 7 orbital’s - which are too complex to draw

52. The magnetic quantum number, ml, refers to the position of the orbital in space relative to other orbital’s. It may have integral numbers ranging from 0 in the s sublevel, 1 to –1 in the p sublevel, 2 to –2 in the d sublevel and 3 to –3 in the f sublevel.

53. ml = 0, is referring to the s orbital
ml = -1, 0, +1, are referring to the three p orbital’s (px, py, and pz)
ml = -2, -1, 0, +1, +2, are referring to the five d orbitals
ml = -3, -2, -1, 0, +1, +2, +3, are referring to the seven f orbitals

54. D. How many electrons can go into each energy level?
Each orbital can hold two electrons. (2n2 = number of electrons per energy level)
The 1st energy level (n=1) only has 1 sublevel called 1s. s only has 1 orbital called the s orbital, so only 2 electrons will be found in the 1st energy level. (2n2 = 2)

55. The 2nd energy level (n=2) has 2 sublevels called 2s and 2p
The 2nd energy level (n=2) has 2 sublevels called 2s and 2p. s only has 1 orbital called the s orbital, p has 3 orbital’s called px, py, and pz orbitals, so 8 electrons will be found in the 2nd energy level. (2n2 = 8)

56. The 3rd energy level (n=3) has 3 sublevels called 3s, 3p, and 3d
The 3rd energy level (n=3) has 3 sublevels called 3s, 3p, and 3d. s only has 1 orbital called the s orbital, p has 3 orbital’s called px, py, and pz orbitals, and d has 5 orbital’s, so 18 electrons will be found in the 3rd energy level. (2n2 = 18)

57. How about the 4th energy level?
It has 4 sublevels called 4s, 4p, 4d, and 4f. s only has 1 orbital, p has 3 orbital’s, d has 5 orbital’s, and f has 7 orbitals, so 32 electrons will be found in the 4th energy level. (2n2 = 32)

58. E. Lets put it all together:
Example of neon atom:

59. Fourth Quantum Number, ms, refers to the magnetic spin of an electron within an orbital. Each orbital can hold two electrons, both with different spins. Clockwise spin is represented with a value of +1/2 and counterclockwise spin is represented with a value of –1/2. Electrons fill the orbital’s one at a time with the same spin (+1/2), then fill up the orbital(s) with electrons of the opposite spin (-1/2).
ms = +1/2 or –1/2

60. Quantum Numbers 4. Spin Quantum Number ( ms ) Electron spin  +½ or -½
An orbital can hold 2 electrons that spin in opposite directions.
Analyzing the emission and absorption spectra of the elements, it was found that for elements having more than one electron, nearly all the lines in the spectra were pairs of very closely spaced lines.
Each line represents an energy level available to electrons in the atom so there are twice as many energy levels available than predicted by the quantum numbers n, l, and ml.
Applying a magnetic field causes the lines in the pairs to split apart.
Uhlenbeck and Goudsmit proposed that the splittings were caused by an electron spinning about its axis.
Courtesy Christy Johannesson 60

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

62. Quantum Numbers Analogy
Energy Levels (n) or Principal Q.N.
n=1 (Weir) n=2 (Liberty Hill) n=3 (Georgetown) n=4 (Austin)
Sublevels (l) or Azimuthal Q.N.
l=0 – s shape 1 bedroom
l=1 – p shape 3 bedroom
l=2 – d shape 5 bedroom
l=3 – f shape 7 bedroom
Orbitals (ml) or Magnetic Q.N.
If l=0 then ml=0 (Represents the 1 bed/orbital in the s sublevel)
If l=1 then ml= -1, 0, 1 (Represents the 3 bed’s/orbital’s in the p sublevel)
If l=2 then ml= -2, -1, 0, 1, 2 (Represents the 5 bed’s/orbital’s in the p sublevel)
If l=3 then ml= -3, -2, -1, 0, 1, 2, 3 (Represents the 7 bed’s/orbital’s in the p sublevel)
Magnetic Spin – Fourth Q.N. (ms)
ms = +1/2 - 1st electron in orbital
ms = -1/2 – 2nd electron in orbital

63. Allowed Sets of Quantum Numbers for Electrons in Atoms
Level n
Sublevel l
Orbital ml
Spin ms 1 -1 2 -2
= +1/2
= -1/2
Allowed Sets of Quantum Numbers for Electrons in Atoms

64. 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

65. Quantum Numbers n shell l subshell ml orbital ms electron spin
1, 2, 3, 4, ...
l subshell
0, 1, 2, ... n - 1
ml orbital
- l l
ms electron spin
+1/2 and - 1/2

66. Electrons In Atoms Notes (Chapter 4) Part 3 Electron Configurations

67. I. Electron Configuration: It should be obvious to you now that it is very difficult to draw a representation or model of atom showing where the electrons are located, so what we do instead is write electron configurations for elements.
Definition of electron configuration: An electron configuration is a written representation of the arrangement of electrons in an atom.

68. When constructing orbital diagrams and electron configurations, keep the following in mind:
Aufbau Principle – electrons fill in order from lowest to highest energy.
The Pauli Exclusion Principle – An orbital can only hold two electrons.
Two electrons in the same orbital must have opposite spins.
You must know how many electrons can be held by each angular momentum number, l. (ie; s can hold 2, 6 for p, l0 for d, 14 for f)
Hund’s rule – the lowest energy configuration for an atom is the one having the maximum number of unpaired electrons for a set of degenerate orbitals. By convention, all unpaired electrons are represented as having parallel spins with spin “up”.

69. Filling Rules for Electron Orbitals
Aufbau Principle: Electrons are added one at a time to the lowest
energy orbitals available until all the electrons of the atom
have been accounted for.
Pauli Exclusion Principle: An orbital can hold a maximum of two electrons.
To occupy the same orbital, two electrons must spin in opposite
directions.
Hund’s Rule: Electrons occupy equal-energy orbitals so that a maximum
number of unpaired electrons results.
*Aufbau is German for “building up”

70. Quantum Numbers Pauli Exclusion Principle
Wolfgang Pauli
Pauli Exclusion Principle
No two electrons in an atom can have the same 4 quantum numbers.
Each electron has a unique “address”:
1. Principal # 
2. Ang. Mom. # 
3. Magnetic # 
4. Spin # 
energy level
sublevel (s,p,d,f)
orbital
electron
Wolfgang Pauli determined that each orbital can contain no more than two electrons.
Pauli exclusion principle: No two electrons in an atom can have the same value of all four quantum numbers (n, l, ml , ms).
By giving the values of n, l, and ml, we specify a particular orbit.
Because ms has only two values (+½ or -½), two electrons (and only two electrons) can occupy any given orbital, one with spin up and one with spin down.
Courtesy Christy Johannesson 70

71. What? How do we write an electron configuration?
1st rule - electrons occupy orbitals that require the least amount of energy for the electron to stay there. So always follow the vertical rule (Aufbau Principle):
You notice, for example, that the 4s sublevel requires less energy than the 3d sublevel; therefore, the 4s orbital is filled with electrons before any electrons enter the 3d orbital!!!! (So just follow the chart and you can’t go wrong!!!!)

72. What? How do we write an electron configuration?
1st rule - electrons occupy orbitals that require the least amount of energy for the electron to stay there. So always follow the vertical rule (Aufbau Principle):

73. You notice, for example, that the 4s sublevel requires less energy than the 3d sublevel; therefore, the 4s orbital is filled with electrons before any electrons enter the 3d orbital!!!! (So just follow the chart and you can’t go wrong!!!!)

74. B. 2nd rule – only 2 electrons can go into any orbital, however, you must place one electron into each orbital in a sublevel before a 2nd electron can occupy an orbital. Orbital’s with only 1 electron in the orbital are said to have an unpaired electron in them.

75. III. Writing Electron Configurations (3 ways):
A. Orbital Notation: an unoccupied orbital is represented by a line______, with the orbitals name written underneath the line. An orbital containing one electron is written as _____, an orbital with two electrons is written as ____. The lines are labeled with the principal quantum number and the sublevel letter.

76. Examples: (Remember that you must place one electron into each orbital before a second electron in placed into an orbital.)
Hydrogen ____ Helium __
1s s
Lithium ___ ____
1s 2s
Carbon ____ ____ ____ ____ _____
1s s px py pz
You try to write the notation for Titanium

77. H = 1s1 He = 1s2 Li = 1s2 2s1 Be = 1s2 2s2 C = 1s2 2s2 2p2 S
THIS SLIDE IS ANIMATED
IN FILLING ORDER 2.PPT H
= 1s1 1s He
= 1s2 1s Li
= 1s2 2s1 1s 2s Be
= 1s2 2s2 1s 2s C
= 1s2 2s2 2p2 1s 2s
2px
2py
2pz S
= 1s2 2s2 2p4 1s 2s
2px
2py
2pz 3s
3px
3py
3pz

78. H
= 1s1 1s e- +1 He
= 1s2 1s e- +2 e-
Coulombic attraction holds valence electrons to atom. Valence electrons are shielded by the kernel electrons.
Therefore the valence electrons are not held as tightly in Be than in He. Be
= 1s2 2s2 1s 2s e- e- +4 e- e-

79. B. Electron Configuration Notation: eliminates the lines and arrows of orbital notation. Instead, the number of electrons in a sublevel is shown by adding a superscript to the sublevel designation. The superscript indicates the number of electrons present in that sublevel.

80. Examples: Hydrogen: 1s1 Helium: 1s2 Lithium: 1s22s1 Carbon: 1s22s22p2
You try to write the notation for Titanium

81. Fe = 1s1 2s22p63s23p64s23d6 26 Iron has ___ electrons. Arbitrary
2px
2py
2pz 3s
3px
3py
3pz 4s 3d 3d 3d 3d 3d
Arbitrary
Energy Scale 18 32 8 2 1s
2s p
3s p
4s p d
5s p d
6s p d f
NUCLEUS e- e- e- e- e- e- e- e- e- e- e- e- e-
+26 e- e- e- e- e- e- e- e- e- e- e- e- e-

82. Electron Configurations
Orbital Filling
Element 1s s px 2py 2pz s Configuration
Orbital Filling
Element 1s s px 2py 2pz s Configuration
Electron
Electron H He Li C N O F Ne Na H He Li C N O F Ne Na
1s1
1s1
1s2
1s2
NOT CORRECT
Violates Hund’s Rule
1s22s1
1s22s1
1s22s22p2
1s22s22p2
1s22s22p3
1s22s22p3
The aufbau principle
1. For hydrogen, the single electron is placed in the 1s orbital, the orbital lowest in energy, and electron configuration is written as 1s1. The orbital diagram is
H: 2p _ _ _
2s _
1s 
2. A neutral helium atom, with an atomic number of 2 (Z = 2), contains two electrons. Place one electron in the lowest-energy orbital, the 1s orbital. Place the second electron in the same orbital as the first but pointing down, so the electrons are paired. This is written as 1s2.
He: 2p _ _ _
1s 
3. Lithium, with Z = 3, has three electrons in the neutral atom. The electron configuration is written as 1s22s1. Place two electrons in the 1s orbital and place one in the next lowest-energy orbital, 2s. The orbital diagram is
Li: 2p _ _ _
2s 
4. Beryllium, with Z = 4, has four electrons. Fill both the 1s and 2s orbitals to achieve 1s22s2:
Be: 2p _ _ _
2s 
1s 
5. Boron, with Z = 5, has five electrons. Place the fifth electron in one of the 2p orbitals. The electron configuration is 1s22s22p1
B: 2p  _ _
2s 
1s 
6. Carbon, with Z = 6, has six electrons. One is faced with a choice — should the sixth electron be placed in the same 2p orbital that contains an electron or should it go in one of the empty 2p orbitals? And if it goes in an empty 2p orbital, will the sixth electron have its spin aligned with or be opposite to the spin of the fifth?
7. It is more favorable energetically for an electron to be in an unoccupied orbital rather than one that is already occupied due to electron-electron repulsions. According to Hund’s rule, the lowest-energy electron configuration for an atom is the one that has the maximum number of electrons with parallel spins in degenerate orbitals. Electron configuration for carbon is 1s22s22p2 and the orbital diagram is
C: 2p   _
8. Nitrogen (Z = 7) has seven electrons. Electron configuration is 1s22s22p3. Hund’s rule gives the lowest-energy arrangement with unpaired electrons as
N: 2p   
9. Oxygen, with Z = 8, has eight electrons. One electron is paired with another in one of the 2p orbitals. The electron configuration is 1s22s22p4:
O: 2p   
2s 
10. Fluorine, with Z = 9, has nine electrons with the electron configuration 1s22s22p5:
F: 2p   
11. Neon, with Z = 10, has 10 electrons filling the 2p subshell. The electron configuration is 1s22s22p6
Ne: 2p   
1s22s22p4
1s22s22p4
1s22s22p5
1s22s22p5
1s22s22p6
1s22s22p6
1s22s22p63s1
1s22s22p63s1 82

83. Short Hand or Noble Gas Notation: Use the noble gases that have complete inner energy levels and an outer energy level with complete s and p orbital’s. Use the noble gas that just precedes the element you are working with.

84. Boron is ls22s22p1
The noble gas preceding Boron is He, so the short way is [He]2s22p1.
Sulfur is ls22s22p63s23p4
Short way: [Ne]3s23p4
Example: Titanium

85. More Practice Problems:
Write electron configurations for each of the following atoms:
1. boron
2. sulfur
3. vanadium
4. iodine
Draw orbital diagrams for these:
5. sodium
6. phosphorus
7. chlorine
Write shorthand electron configuration for the following:
8. Sr
9. Mo
10. Ge

86. Irregular Electron configurations – sometimes the electron configuration is NOT what we would predict it to be. Sometimes electrons are moved because (l) it will result in greater stability for that atom or (2) for some unknown reason??

87. It is very important to define “stable” here. STABLE means:
1. all degenerate (equal energy) orbital’s are FULL
2. all degenerate orbital’s are half-full
3. all degenerate orbital’s are totally empty.

88. Examples – draw the orbital’s (lines or boxes) and fill each orbital with the predicted number of electrons. Predict the electron configuration for Cr #24: [Ar]4s23d6
However, the real E. C. is [Ar]4s13d5. The 4s1 electron has been moved to achieve greater stability.
ALWAYS USE THE ACTUAL E. C. AND NOT THE PREDICTED ONE. YOU WILL HAVE THESE ATOMS WITH IRREGULAR E. C. HIGHLIGHTED OR MARKED ON YOUR PERIODIC TABLE.

89. Electron configurations for Ions-First, determine if the element will lose or gain electrons. Secondly, what number of electrons will be gained or lost? It is recommended that you write the e.c. for the atom and then determine what will happen.

90. For cations (positive ions) – look at the element and decide how many electrons will be lost when it ionizes and keep that in mind when writing the E. C. The last number in the E. C. will now be LESS than what is written on your periodic table.
Ex. Write the electron configuration for magnesium ion: [Ne]3s2 is for the atom. Mg is a metal and will lose its valence (outer) electrons, so the e.c. for Mg2+ is 1s22s22p6
Practice:
1. #3
2. #12
3. #19
4. #13

91. For anions (negative ions) – look at the element and decide how many electrons that element will GAIN when it ionizes. The last number in the E. C. will be MORE than what is written on the periodic table.
Ex. Sulfide ion: Sulfur atom is 1s22s22p4. Sulfur is a nonmetal with 6 valence electrons (2s2 and 2p4) and will gain 2 electrons: 1s22s22p6 is for the sulfide ion.
Practice:
#17 #7
#16
#30