1. Chapter 4 Electronic Structure of Atoms
Chemistry, The Central Science, 11th edition
Theodore L. Brown; H. Eugene LeMay, Jr.;
and Bruce E. Bursten
Chapter 4 Electronic Structure of Atoms
John D. Bookstaver
St. Charles Community College
Cottleville, MO
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2. Waves
To understand the electronic structure of atoms, one must understand the nature of electromagnetic radiation.
The distance between corresponding points on adjacent waves is the wavelength ().
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3. Waves
The number of waves passing a given point per unit of time is the frequency ().
For waves traveling at the same velocity, the longer the wavelength, the smaller the frequency.
Wavelength: It is the distance between two consecutive peaks or troughs in a wave.
Frequency: It indicates how many waves pass a given point per second.
Speed: It indicates how fast a given peak is moving through the space.
Speed of light ,c=ln ,where l=wave length and n =frequency
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4. EMR
Electromagnetic Radiation: Electromagnetic radiation is one of the ways in which energy travels through space. All forms of EMR compose the electromagnetic radiation spectrum, which includes sun rays, microwaves, X- rays, visible spectrum, UV rays and IR rays.
Some characteristics of EMR are:
All electromagnetic radiation moves at a constant speed of about 3.0 X 108 m/s.
All EMR exhibit wave like behavior. Waves have three primary characteristics:
Wavelength: It is the distance between two consecutive peaks or troughs in a wave.
Frequency: It indicates how many waves pass a given point per second.
Speed: It indicates how fast a given peak is moving through the space.
Speed of light ,c=ln ,where l=wave length and n =frequency

5. Electromagnetic Radiation
All electromagnetic radiation travels at the same velocity: the speed of light (c),
3.00  108 m/s.
Therefore,
c = 
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6. Refer to the following EMR spectrum. (visible spectrum Roy G. BiV)
Imagine you have invented a machine that allows you to see all types of EMR. Make a list of type of EMR you might see in your home.
What will happen if you change the red light in the dark room for photo processing with the yellow light and why?

7. Wave Nature of Light
Before the concept of quantization of energy, the wave like nature of light/energy was widely accepted.
Two properties that exhibit the wave like behavior of light are interference and diffraction.
Animation Diffraction:
Animation for Interference:

8. The Nature of Energy
The wave nature of light does not explain how an object can glow when its temperature increases.
Max Planck explained it by assuming that energy comes in packets called quanta.
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9. Planck’s Theory of Quantization of Energy
To explain this, Planck suggested that the energy transfer or exchange is not a continuous process, but is done in small packets of energy called by him as quantum.(Word quantum means fixed amount.) So, he introduced the concept of quantization of energy.
According to Planck’s theory, E= hv,
where E= energy of radiation
h= Planck’s constant
v= frequency of radiation
Planck’s theory: Before Planck’s theory, the wave model of the light was widely accepted. But it was unable to explain some phenomenon for example change in the radiation (wave length) emitted by an object with the change in temperature.

10. Quantization of Energy: Max Plank
www
Where else do you see quantization in real life?

11. The Nature of Energy
Therefore, if one knows the wavelength of light, one can calculate the energy in one photon, or packet, of that light:
c = 
E = h
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12. The Nature of Energy
Another mystery in the early 20th century involved the emission spectra observed from energy emitted by atoms and molecules.
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13. The Nature of Energy
Einstein used this assumption to explain the photoelectric effect.
He concluded that energy is proportional to frequency:
E = h
where h is Planck’s constant,  10−34 J-s.
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14. Photoelectric Effect: It refers to the emission of electrons from a metal, when the light shines on the metal. For each metal the frequency of light needed to release the electrons is different. But the wave theory of light could not explain it. The photoelectric effect led scientists to think about the dual nature of light i.e. as a wave and a particle both.
Animation 1:
Animation 2:

15. Dual Wave- Particle Nature of Light
Dual nature (Wave- Particle nature) of light: Albert Einstein in 1905, expanded on Planck’s theory by introducing the idea of the dual nature of the radiation. Albert Einstein called the quantum of energy in light as ‘photons’.
Light as wave Light as particle DUAL NATURE OF LIGHT

16. http://www. nobel. se/physics/educational/tools/quantum/w-p-images/wp
Photoelectric Effect
Max Planck suggested that the hot object emits energy in small, specific amounts called quanta.
1905, Einstein said electromagnetic radiation has a dual wave- particle nature.

17. The Nature of Energy
For atoms and molecules one does not observe a continuous spectrum, as one gets from a white light source.
Only a line spectrum of discrete wavelengths is observed.
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18. Continuous and Line Spectra (Absorption and Emission Spectrum- Line Spectra)(
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um%26svnum%3D10%26hl%3Den%26lr%3D%26sa%3DN)

19. Bohr Model Of the Hydrogen Atom
The energy levels of Hydrogen ( As explained by Bohr’s Model) :
An excited atom can release some or all of its excess energy by emitting a photon, thus moving to a lower energy state.
The lowest possible energy state of an atom is called the ‘ground state’.
Different wavelengths of light carry different amount of energy per photon. Ex. A beam of red light has a lower energy photons than beam of blue light.
Ephoton=hv

20. The Nature of Energy
Niels Bohr adopted Planck’s assumption and explained these phenomena in this way:
Electrons in an atom can only occupy certain orbits (corresponding to certain energies).
Animation :
Animation:
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21. The Nature of Energy
Niels Bohr adopted Planck’s assumption and explained these phenomena in this way:
Electrons in permitted orbits have specific, “allowed” energies; these energies will not be radiated from the atom.
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22. The Nature of Energy
Niels Bohr adopted Planck’s assumption and explained these phenomena in this way:
Energy is only absorbed or emitted in such a way as to move an electron from one “allowed” energy state to another; the energy is defined by
E = h
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23. Light Equations E=h n E is energy (in joules)
c= l n
c is the speed of light (3.0 x 108 m/s)
lambda is the wave length (in m, cm, or nm)
n represents the frequency (waves/s or hertz)
E=h n
E is energy (in joules)
h is Planck’s Constant (h= x 10-34J s)
n is frequency
E=h(c/ l)

24. Dual Wave-Particle nature by De Broglie
Quantum Theory: Describes mathematically the wave properties of electrons or other very small particles treating e as waves and using Heisenberg’s and De Broglie’s principles.
. Heisenberg Uncertainty Principle: States that it is impossible to determine simultaneously both the position and velocity of an electron or any other particle.
Dual Wave-Particle nature by De Broglie

25. Quantum Mechanical Model of Atom (Schrodinger’s Model)
Erwin Schrödinger, in developing a quantum-mechanical model for the atom, began with a classical equation for the properties of waves. He modified this equation to take into account the mass of a particle and the de Broglie relationship between mass and wavelength. The important consequences of the quantum-mechanical view of atoms are the following: (
The energy of electrons in atoms is quantized.
The number of possible energy levels for electrons in atoms of different elements is a direct consequence of wave-like properties of electrons.
The position and momentum of an electron cannot both be determined simultaneously.
The region in space around the nucleus in which an electron is most probably located is what can be predicted for each electron in an atom. Electrons of different energies are likely to be found in different regions. The region in which an electron with a specific energy will most probably be located is called an atomic orbital.

26. Quantum Mechanics
Erwin Schrödinger developed a mathematical treatment into which both the wave and particle nature of matter could be incorporated.
It is known as quantum mechanics.
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27. Quantum Mechanics
The wave equation is designated with a lower case Greek psi ().
The square of the wave equation, 2, gives a probability density map of where an electron has a certain statistical likelihood of being at any given instant in time.
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28. Quantum Numbers
Schrodinger arrived at three functions while solving the wave equation for electrons and came up with three quantum numbers, each corresponding to a property of atomic orbital. The fourth quantum number was later introduced to clarify the spin of electrons. use(
Solving the wave equation gives a set of wave functions, or orbitals, and their corresponding energies.
Each orbital describes a spatial distribution of electron density.
An orbital is described by a set of three quantum numbers.
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29. Principal Quantum Number (n)
The principal quantum number, n, describes the energy level on which the orbital resides.
The values of n are integers ≥ 1.
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30. Angular Momentum Quantum Number (l)
This quantum number defines the shape of the orbital.
Allowed values of l are integers ranging from 0 to n − 1.
We use letter designations to communicate the different values of l and, therefore, the shapes and types of orbitals.
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31. Angular Momentum Quantum Number (l)
Value of l 1 2 3
Type of orbital s p d f
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32. Mosby-Year Book Inc.

33. Magnetic Quantum Number (ml)
The magnetic quantum number describes the three-dimensional orientation of the orbital.
Allowed values of ml are integers ranging from -l to l:
−l ≤ ml ≤ l.
Therefore, on any given energy level, there can be up to 1 s orbital, 3 p orbitals, 5 d orbitals, 7 f orbitals, etc.
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34. Magnetic Quantum Number (ml)
Orbitals with the same value of n form a shell.
Different orbital types within a shell are subshells.
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35. Principle energy level (n) Type of sublevel
Orbitals and Electron Capacity of the First Four Principle Energy Levels
Principle energy level (n)
Type of sublevel
Number of orbitals per type
Number of orbitals per level(n2)
Maximum number of electrons (2n2) 1 s 2 4 8 p 3 9 18 d 5 16 32 f 7

36. s Orbitals The value of l for s orbitals is 0.
They are spherical in shape.
The radius of the sphere increases with the value of n.
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37. s Orbitals
Observing a graph of probabilities of finding an electron versus distance from the nucleus, we see that s orbitals possess n−1 nodes, or regions where there is 0 probability of finding an electron.
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38. p Orbitals The value of l for p orbitals is 1.
They have two lobes with a node between them.
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39. d Orbitals The value of l for a d orbital is 2.
Four of the five d orbitals have 4 lobes; the other resembles a p orbital with a doughnut around the center.
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40. Energies of Orbitals
For a one-electron hydrogen atom, orbitals on the same energy level have the same energy.
That is, they are degenerate.
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41. Energies of Orbitals
As the number of electrons increases, though, so does the repulsion between them.
Therefore, in many-electron atoms, orbitals on the same energy level are no longer degenerate.
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42. Spin Quantum Number, ms
In the 1920s, it was discovered that two electrons in the same orbital do not have exactly the same energy.
The “spin” of an electron describes its magnetic field, which affects its energy.
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43. Spin Quantum Number, ms
This led to a fourth quantum number, the spin quantum number, ms.
The spin quantum number has only 2 allowed values: +1/2 and −1/2.
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44. Pauli Exclusion Principle
No two electrons in the same atom can have exactly the same energy.
Therefore, no two electrons in the same atom can have identical sets of quantum numbers.
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45. Electron Configurations
This shows the distribution of all electrons in an atom.
Each component consists of
A number denoting the energy level,
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46. Electron Configurations
This shows the distribution of all electrons in an atom
Each component consists of
A number denoting the energy level,
A letter denoting the type of orbital,
Animation:
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47. Electron Configurations
This shows the distribution of all electrons in an atom.
Each component consists of
A number denoting the energy level,
A letter denoting the type of orbital,
A superscript denoting the number of electrons in those orbitals.
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48. Orbital Diagrams Each box in the diagram represents one orbital.
Half-arrows represent the electrons.
The direction of the arrow represents the relative spin of the electron.
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49. Hund’s Rule
“For degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized.”
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50. Periodic Table We fill orbitals in increasing order of energy.
Different blocks on the periodic table (shaded in different colors in this chart) correspond to different types of orbitals.
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51. Some Anomalies
Some irregularities occur when there are enough electrons to half-fill s and d orbitals on a given row.
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52. Some Anomalies For instance, the electron configuration for copper is
[Ar] 4s1 3d5
rather than the expected
[Ar] 4s2 3d4.
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53. Some Anomalies
This occurs because the 4s and 3d orbitals are very close in energy.
These anomalies occur in f-block atoms, as well.
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54. Quantum Number Terms Ground State: Lowest energy state of an atom
Excited State: A state in which an atom has a higher potential energy then it has in its ground state
Orbital: A 3D region around the nucleus that indicates the probable location of an electron
Quantum Numbers: Specify the properties of atomic orbitals and the properties of electrons in orbitals
Principle Quantum Number: (n) Indicates the main energy level occupied by the electron.
Angular Momentum Quantum Number: (l) Indicates the shape of the orbital.
Magnetic Quantum Number: (m) Indicates the orientation of an orbital around the nucleus.
Spin Quantum Number: (+1/2, -1/2) Indicates the two fundamental spin states of an electron in an orbital

55. Electron Configuration Theories
Electron Configuration Theories
Hund’s Rule: Orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron, and all electrons in singly occupied orbitals must have the same spin.
Aufbau’s Principle: An electron occupies the lowest-energy orbital that can receive it.
Pauli’s Exclusion Principle: No two electrons in the same atom can have the same set of four quantum numbers.

56. Ways to Represent Electron Configuration
Expanded Electron Configuration
Condensed Electron Configurations
Orbital Notation
Electron Dot Structure
Write the above four electron configurations for Zinc, Zinc ion and Cu ion.