dr. s. m. condren chapter 7 quantum theory of the atom

Dr. S. M. Condren

Chapter 7

Quantum Theory of the Atom

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Electromagnetic Radiation

Electromagnetic wave• A wave of energy having a frequency

within the electromagnetic spectrum and propagated as a periodic disturbance of the electromagnetic field when an electric charge oscillates or accelerates.

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Electromagnetic Radiation

Electromagnetic wave

• wavelength

• frequency

• amplitude

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Electromagnetic Radiation

= c

where=> frequency

=> wavelength

c => speed of light

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Wave Nature of the Electron

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Electromagnetic Spectrum

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Line Emission Spectrum

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Line Spectrum

A spectrum produced by a luminous gas or vapor and appearing as distinct lines characteristic of the various elements constituting the gas.

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Emission Spectrum

The spectrum of bright lines, bands, or continuous radiation characteristic of and determined by a specific emitting substance subjected to a specific kind of excitation.

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Ground State

The state of least possible energy in a physical system, as of elementary particles. Also called ground level.

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Excited State

Being at an energy level higher than the ground state.

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Photoelectric Effect

• the emission of electrons by substances, especially metals, when light falls on their surfaces.

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Photoelectric Effect

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Quantum Mechanics

Quantum theory• the theory of the structure and behavior of

atoms and molecules.

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Black Body Radiation

http://www.cbu.edu/~mcondren/C11599/BBvis.mov

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Electromagnetic Radiation

Ehi - Elo = hc/where E => energy

h => Planck's constant

c => speed of light

=> wavelength

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Photons

The quantum of electromagnetic energy, generally regarded as a discrete particle having zero mass, no electric charge, and an indefinitely long lifetime.

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Dispersion of White Light

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The Atomic Spectrum of Hydrogen and the Bohr Model

Bohr Model for the Hydrogen Atom

mnr = nh/2p

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Bohr Atom

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Bohr Model

E = -B/n2

where n => quantum number

1, 2, 3, 4, 5, 6, 7, etc

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Bohr Model

for hydrogen

ground state: n = 1

excited state: n > 1

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Bohr Model

E = (-2.179 X 10-18 J/part.)

(6.022 X 1023 part./mole)

(1 kJ/103 J)/n2

= (-1312 kJ/mol)(1/n2)

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Electron Transition in a Hydrogen Atom

Lyman series => ultraviolet

n > 1 ==> n = 1

Balmer series => visible light

n > 2 ==> n = 2

Paschen series => infrared

n > 3 ==> n = 3

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Line Spectra

See CHEMWORKS software

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Absorption Spectrum

• Light shinning on a sample causes electrons to be excited from the ground state to an excited state

• wavelengths of that energy are removed from transmitted spectra

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Knowing diamond is transparent, which curve best represents the absorption spectrum of diamond (see below)?

A, B, C

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According to the energy diagram below for the Bohr model of the hydrogen atom, if an electron jumps from E1 to E2, energy is

absorbed

emitted

not involved

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Heisenberg, Werner1901–76, German physicist

1932 Nobel Prize in physics

A founder of QUANTUM MECHANICS, he is famous for his uncertainty principle, which states that it is impossible to determine both the position and momentum of a subatomic particle (such as the electron) with arbitrarily high accuracy.

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Heissenberg Uncertainty Principle

“it is impossible to determine both the position and momentum of a subatomic particle (such as the electron) with arbitrarily high accuracy”

The effect of this principle is to convert the laws of physics into statements about relative, instead of absolute, certainties.

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Orbitals

• region of probability of finding an electron around the nucleus

• 5 types => s p d f

• maximum of 2 electrons per orbital

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Pure Atomic Orbitals

shape # of orbitals / energy level

s spherical 1

p dumbbell 3

d complex 5

f very complex 7

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Atomic Orbitals, s-type

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Atomic Orbitals, p-type

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Atomic Orbitals, d-type

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Atomic Orbitals, f-type

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Shapes of Orbitals

http://www.colby.edu/chemistry/OChem/DEMOS/Orbitals.html