1. The Arrangement of Electrons in Atoms
Ch. 4 Modern Chemistry

2. A New Atomic Model Cathode Ray Tube Summary
As of 1911, Rutherford’s Gold Foil Experiment had shown scientists the nucleus of an atom was very small, dense, and positively charged.
Why was Rutherford’s model of the atom incomplete?
Rutherford’s model did not explain how electrons were distributed around the nucleus.
Just assumed electrons located around nucleus

3. Wave-like Properties of Light
Over the next decade, scientists began to discover the dual nature of electrons - particles and waves.
Electromagnetic radiation is a form of energy that exhibits wavelike behavior as it travels through space.
Together, all the forms of electromagnetic radiation form the electromagnetic spectrum.
Visible spectrum: visible portion of spectrum
But there are many other forms of energy

4. Electromagnetic Spectrum
Electromagnetic radiation is the emission and transmission of energy in the form of electromagnetic waves.

5. Properties of Light
Wavelength (λ) : distance between corresponding points on adjacent waves
Frequency (ν):
number of waves that pass a given point per unit time time (usually 1 sec)

6. Properties of Light c = λν
Frequency and wavelength are mathematically related to each other:
c = λν
c : speed of light (m/s)
λ : wavelength (m)
ν : frequency of wave (s−1 or 1/s).

7. Photoelectric Effect Particle Description of Light
Electromagnetic radiation strikes the surface of the metal
Electrons are ejected from the metal causing an electric current

8. Photoelectric Effect
The photoelectric effect refers to the emission of electrons from a metal when light shines on the metal.
Not all forms of electromagnetic radiation were “powerful” enough to eject electrons from the surface
A minimum amount of energy was needed
A quantum of energy - the minimum quantity of energy that can be lost or gained by an atom

9. German physicist Max Planck proposed the following relationship:
Photoelectric Effect
German physicist Max Planck proposed the following relationship:
E = hν
E : quantum of energy (J – joules)
ν : frequency of the radiation emitted; s−1
h : fundamental constant called Planck’s constant;
h = × 10−34 J• s

10. Particle Description of Light
photon : particle of electromagnetic radiation having zero mass and carrying a quantum of energy.
The energy of a particular photon depends on the frequency of the radiation.
Ephoton = hν

11. The Hydrogen Line Emission Spectrum

12. The Hydrogen Line Emission Spectrum
When investigators passed 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 four specific colors of the visible spectrum.
The four bands of light were part of what is known as hydrogen’s line-emission spectrum.

13. The Hydrogen Line Emission Spectrum
What had scientists expected to observe?
A continuous range of frequency (all bands of ER)

14. The Hydrogen Line Emission Spectrum
The lowest energy state of an atom is its ground state.
A state in which an atom has a higher potential energy than it has in its ground state is an excited state.

15. Bohr Model of the Atom
By 1922, Niels Bohr proposed a hydrogen-atom model that linked the atom’s electron to photon emission.
electrons can circle the nucleus only in allowed paths, called orbits.
The energy of the electron is higher when it is in an orbit farther from the nucleus.

16. Bohr Model of the Atom
When electrons move from an excited state to a ground state or to a lower energy state, a photon of energy is emitted from an atom
Process called emission (high to low orbit)
absorption (low to high orbit): energy is added to an atom to move an electron from lower energy level to higher energy level
3) The energy of a photon can be directly calculated if you know the frequency of radiation being emitted

17. Photo Absorption & Emission

18. Quantum Model of the Atom
Section 4.2
Quantum Model of the Atom

19. Electrons as Waves
French scientist Louis de Broglie suggested that electrons be considered waves confined to the space around an atomic nucleus.
electron waves could exist only at specific frequencies.
It followed that if each electron has a specific frequency, it also has a corresponding energy.
Recall equation E = hv

20. Heisenberg Uncertainty Principle
German physicist Werner Heisenberg proposed that any attempt to locate a specific electron with a photon knocks the electron off its course.
The Heisenberg uncertainty principle: it is impossible to determine simultaneously both the position and velocity of an electron or any other particle.

21. The Schrödinger Wave Equation
In 1926, Austrian physicist Erwin Schrödinger developed an equation that treated electrons as waves in an atom.
Heisenberg & Schrödinger ideas laid the foundation for modern quantum theory.
Quantum theory describes mathematically the wave properties of electrons and other very small particles.

22. The Schrödinger Wave Equation
Electrons do not travel around the nucleus in neat orbits, as Bohr had postulated, but rather in certain regions called orbitals.
An orbital: three-dimensional region around the nucleus - indicates the probable location of an electron.
Quantum numbers specify the properties of atomic orbitals and the properties (location & spin) of electrons in orbitals.
4 quantum numbers

23. Schrodinger Equation
Quantum Theory mathematically explains wave properties of electrons and similarly sized particles.
The solutions described shapes in space that electrons were highly likely to be moving around in.

24. Atomic Orbitals and Quantum Numbers
An electron’s quantum numbers describe:
primary energy level (n)
How close to the nucleus it is
shape of momentum ( l )
direction of momentum ( ml )
spin direction ( ms )

25. Quantum Numbers
Quantum numbers describe the behavior of an atom’s electrons

26. Quantum Numbers
“n” represents the main energy level an electron occupies
The bigger “n” gets, the further away from the nucleus the electron gets
If more than 1 electron has the same value for “n” they are in the same “shell”
“n” can only be in integer values; n≥1

27. Angular Momentum Quantum Number l
l represents the sublevel (or shape) of the electron cloud
Each number value corresponds to a shape
Value for l
Orbital letter
shape s
Sphere 1 p
Dumbbell 2 d
Butterfly 3 f
Complex

29. Magnetic Quantum Number m
m represents the orientation of the orbital about the nucleus
The values for m can be -l, 0 , +l
The amount of m values correspond to the number of orientations of that shape
Value for m
# orbitals of that shape
Sublevel letter
shape 1 s
Sphere
-1, 0, 1 3 p
Dumbbell
-2, -1, 0, 1, 2 5 d
Butterfly
-3, -2, -1, 0, 1, 2, 3 7 f
Complex

30. Magnetic Quantum Number m.

31. ms = +½ ms = -½ Spin Quantum Number
Like charges repel therefore, electrons repel each other
Electrons want to always move in opposite directions if they have to share an orbital
Spin quantum number can be +1/2 or -1/2
ms = +½
ms = -½
Spin Quantum Number

32. Quantum Numbers

33. Application of quantum numbers Electron configuration
Section 4.3
Application of quantum numbers
Electron configuration
Orbital notation

34. Electron Configuration
Shows the arrangement of electrons in an atom
Each element has a unique electron configuration
An element will have the orbitals of the elements preceding it plus any additional orbitals to account for it’s extra electrons
Ex:
He has a level 1 s orbital,
Li has both a level 1 s orbital and a level 2 s orbital

35. Electron-Configuration Notation
Electron-Configuration Notations uses the principle quantum number, the orbital letter, and the number of electrons in superscript
How to write e-configurations? Read the PT as you would read a book, from left to right. Start at H and “travel” to the desired element writing down principal quantum number, sublevel and number of electrons.
H = 1s He = 1s B = 1s22s22p1

36. Orbital Diagram AKA Orbital notation – another application of quantum numbers
Orbital notation is a visual notation using arrows to represent electrons and lines to represent orbitals
Write the electron configuration for the element
Draw horizontal line to represent every orbital in each sublevel for every energy level
Arrange orbitals from lowest to highest energy from left to right
Group like orbitals together ex: put all p orbitals of one level closer together

37. Orbital Notation
H ___ 1s
He ___
B ___ ___ ___ ___ ___ s 2s p

38. Electron Configuration & Orbital Diagram
Three RULES!
Aufbau principle: an electron occupies the lowest-energy orbital available
If we have 5 electrons, how can we fit them so the lowest energy orbitals fill first?
Energy
2px
2py
2pz 2s 1s

39. Electron Configuration
Pauli Exclusion Principle: no two electrons of one atom can have the same set of 4 quantum numbers
If we have 4 electrons, but they only fill 2 energy levels, how can we arrange them so they are different?
Energy
2px
2py
2pz 2s 1s

40. Electron Configuration
Hund’s Rule: orbitals of equal energy all fill with one electron before a second electron may be added.
If we have 7 electrons, how can we fill the orbitals by energy level?
Energy
2px
2py
2pz 2s 1s

41. Nobel Gas Notation Nobel Gas Notation is a short cut notation
We know a Nobel gas will have all the orbitals up to that Nobel gas entirely filled
Therefore, we can write a Nobel Gas plus any new orbitals corresponding to the new element
Ca = 1s22s22p63s13p64s2
Ar precedes Ca; Ar = 1s22s22p63s13p6
Therefore: Ca = [Ar]4s2

42. Valence Electrons & Inner Electrons
Every element has a set of valence electrons. For the s and p block, there are 8 valence electrons
Electrons in the valence shell are typically in the highest occupied energy level
For Argon (1s2 2s2 2p6 3s2 3p6) this is 3
For Berylium (1s2 2s2) this is 2
All non-valence electrons are inner-shell electrons
If we know an electron configuration of a neutral atom, we can figure out which element it is

43. Exceptions There are several exceptions to the rules we learned.
We will study two of these exceptions; Chromium, Cr and Copper, Cu
The d-block elements are most stable when there are 10 electrons present (2 per orbital)
The second most stable arrangement is with 5 electrons (one per orbital)
To achieve this, an electron can be “borrowed” from the s-block.

44. Exceptions
Write the noble gas configuration for copper following the rules we have learned so far.
[Ar]4s23d9
Now, to make more stable, “borrow” an electron from the s-block and write a new configuration
[Ar]4s13d10
Now write the noble gas configuration for chromium.
[Ar]4s23d4
Apply the new exception  [Ar]4s13d5

45. How should we look at the atom?
Models of the Atom – An Atomic Rant