1. Chemistry Lesson 5 - Electrons in Atoms Mrs. Eline
Bordentown Regional High School

2. 5.1 Electromagnetic radiation and Spectra

3. 5.1 Electromagnetic radiation and Spectra
The wave nature of light c =λ√
velocity of light c =3.0 x 108 m/s
λ = wavelength (m)
√ = frequency (/s)

4. 5.1 Electromagnetic radiation and Spectra

5. 5.1 Electromagnetic radiation and Spectra
Considering light as a wave explains much of its everyday behavior, however, some important aspects are not explained.
The particle nature of light elucidated from heated metal- Max Planck studied the light emitted by heated objects(ex.iron in fire). His study led him to the conclusion that matter can only gain or lose energy in small specific amounts called quanta.

6. 5.1 Electromagnetic radiation and Spectra

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58. 5.1 Electromagnetic radiation and Spectra
Max Planck

59. 5.1 Electromagnetic radiation and Spectra
The photoelectric effect is studied in part because it can be an introduction to wave-particle duality and quantum mechanics.When a surface is exposed to sufficiently energetic electromagnetic energy, light will be absorbed and electrons will be emitted. The threshold frequency is different for different materials. It is visible light for alkali metals, near-ultraviolet light for other metals, and extreme-ultraviolet radiation for nonmetals. The photoelectric effect occurs with photons having energies from a few electronvolts to over 1 MeV. At the high photon energies comparable to the electron rest energy of 511 keV, Compton scattering may occur pair production may take place at energies over MeV.

60. 5.1 Electromagnetic radiation and Spectra
Light’s Dual Nature-In 1905, Albert Einstein proposed that light has a dual nature. As well as having wave properties, it can be thought of as a beam of bundles of energy called photons. A photon is a massless particle that carries a quantum of energy.
energy of photon = energy needed to remove an electron + kinetic energy of the emitted electron
h√ = W + E
h is Planck's constant = × 10-34 m2 joule·s √ is the frequency of the incident photon W is the work function, which is the minimum energy required to remove an electron from the surface of a given metal: h √ 0 E is the maximum kinetic energy of ejected electrons: 1/2 mv2

61. Hydrogen’s Absorption and Emission Spectra
5.1 Electromagnetic radiation and Spectra
Hydrogen’s Absorption and Emission Spectra

62. Helium’s Absorption and Emission Spectra
5.1 Electromagnetic radiation and Spectra
Helium’s Absorption and Emission Spectra

63. Neon’s Absorption and Emission Spectra
5.1 Electromagnetic radiation and Spectra
Neon’s Absorption and Emission Spectra

64. 5.1 Electromagnetic radiation and Spectra
Each element’s atomic emission spectrum is unique and can be used to identify an unknown element.

65. 5.1 Electromagnetic radiation and Spectra
Connection to Astronomy
When looking at stars with telescopes scientists are able to determine the composition of the outer layers of stars by their emission spectra.  Astronomers can determine not only the element, but the temperature and density of that element in the star. The spectral line also can tell us about any magnetic field of the star. The width of the line can tell us how fast the material is moving. 

66. 5.2 The Quantum Mechanical Model of the Atom
Bohr assigned a quantum number to each orbit. He also calculated the radius of each orbit. In its ground state the atom does not radiate energy. When energy is addeeed the electron moves to a higher-energy orbit. When an electron drops from a higher energy orbit to a lower energy orbit an photon is emitted according to the energy difference between the two levels.

67. 5.2 The Quantum Mechanical Model of the Atom
DeBrogile had been thinking that Bohr’s quantized electron orbits had characteristics of waves similar to harmonics on strings. Electrons on circular orbits can only have odd numbers of wavelengths If waves can have pariclelike behavior can particles like electrons behave like waves?
Louis deBrogile

68. 5.2 The Quantum Mechanical Model of the Atom
The DeBrogile equation predicts that moving particles have wave characteristics.

69. 5.2 The Quantum Mechanical Model of the Atom
The Heisenberg Uncertainty Principle
It is impossible to take any measurement of an object without disturbing the object! Therefore you can not know an electron’s location –the act of observing ab electron produces a significant unavoidable uncertainty in the position and motion of the electron. The only quantify that can be known is the probability for an electron to occupy a certain region around the nucleus.

70. 5.2 The Quantum Mechanical Model of the Atom
The Schrodinger Wave Equation and the quantum mechanical model of the atom. Unlike Bohr Schrodinger made no attempt to describe the electron’s path

71. 5.2 The Quantum Mechanical Model of the Atom
Hydrogen’s Atomic Orbitals-4 quantum numbers assigned to atomic orbitals
Principal quantum number n = 1…7

72. 5.2 The Quantum Mechanical Model of the Atom
Shapes of Orbitals

73. 5.3 Electron Configuration
Aufbau principle- Electrons in an atom tend to assume the arrangement that gives the atom the lowest energy possible.
1.All orbitals related to an energy sublevel are of equal energy.
2. In a multi-electron atom, the energy sublevels within a principal energy level have different energies.
3. In order of increasing energy, the sequence of energy sublevels within a principal energy level are s,p,d, and f.
4. Orbitals related to energy sublevels within one principal energy level can overlap orbitals related to energy sublevels within another principal level.

74. 5.4 Electron Configuration

75. 5.4 Electron Configuration
The Pauli Exclusion Principal is the quantum mechanical principle that states that two identical fermions (particles with half-integer spin) cannot occupy the same quantum state simultaneously.
Hund's rule: every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied, and all electrons in singly occupied orbitals have the same spin.

76. 5.4 Electron Configuration
First 2 methods for representing an atom’s electron configuration
1. Orbital Diagram
2. Electron Configuration Notation

77. 5.4 Electron Configuration
3. Noble Gas Notation
Note - when Noble Gas Notation is used the Noble Gas is in Brackets. example [Ne]
Noble gas notation for Sodium Na [Ne]3s1
Noble gas notation for Strontium Sr [Kr] 5s2

78. 5.4 Electron Configuration
4. Lewis Dot Structure Notation
Valence Electrons- these are the electrons that determine the chemical properties of an element-generally those orbitals which occupy the atom’s highest principal energy levels. These are the electrons involved in forming chemical bonds.
Noble gas notation for Sulfur S [Ne]3s23p4
Noble gas notation for Cesium Cs [Xe] 6s1

79. 5.4 Electron Configuration
4. Lewis Dot Structure Notation