Download CH4 PT1 Arrangement of Electrons

Arrangement of Electrons
in Atoms
Part One
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Section 1 Development of a New
Atomic Model
• Explain the mathematical relationship amount
the speed, wavelength, and frequency of
electromagnetic radiation.
• Discuss the dual wave-particle nature of light.
• Discuss the significance of the photoelecvtric
effect and the line-emission spectrum of H to
the development of the atomic model.
• Describe the Bohr model of the H atom.
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Section 2 Quantum Model
• Discuss De Broglie’s role in the development
of the quantum model of the atom.
• Compare and contrast the Bohr model and the
quantum model of the atom.
• Explain how the Heisenberg uncertainty
principle and the Schrodinger wave equation
led to the idea of atomic orbitals.
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• Electromagnetic Radiation (EMR) or light is a
form of energy that moves as a wave through
space.
• Electromagnetic Spectrum is made up of many
kinds of EMR: visible, X rays, ultraviolet,
infrared, microwaves, and radiowaves.
Section 3 – Electron Configuration
• Aufbrau principle, Hund’s rule, Pauli exclsuion
principle
• Notation types (orbital, e-configuration,
noble-gas)
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Electromagnetic (E-M) Waves (LIGHT!)
Do not require a medium through which to travel
Light travels at 3.0 x 108 m/s in a vacuum or air
Its wavelength and frequency varies according to the type of E-M wave
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Higher frequency 
Greater energy 
More penetration
What type has
a. Greatest frequency?
b. Less frequency than
infrared light?
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c = ln
• c , the speed of light which is 3x108 m/s in a
vacuum or air. Units: m/s
• l, wavelength or distance between
corresponding points on adjacent waves
– Units: m or nm
• n, frequency or number of waves passing a
point in a given amount of time. Units: Hertz,
Hz or 1/s or s-1
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Light Problems: What is the frequency
of light whose wavelength is 600 nm?
nm means 10-9 m
c = ln --> n = c/l
n = c/l =
3x108 m = 5 x 1014 s-1
600x10-9 m s
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Photoelectric Effect
• This is the emission of electrons from a metal
when electromagnetic radiation shines on the
metal.
P. E. shows that energy is
emitted in small, specific
packets called quanta.
A quantum of energy is the
minimum quantity of energy
that can be lost or gained by an
atom.
The photoelectric effect showed that light behaves as particles, too!
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E=hn
*
E, energy of a quantum of radiation in joules, J
h, Planck’s constant is 6.626 x 10-34 Js
n, frequency in s-1
Problem: What is the frequency of a photon whose
energy is 3.4 x 10-19 J?
n = E/h = 3.4 x 10-19 J / 6.626 x 10 -34 Js
n = 5.1 x 1014 s-1
*Wavelength-frequency relationship was
proposed by Planck in 1900.
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• Einstein explained the photoelectric effect was
due to metal absorbing energy in discrete
amounts of photons.
• Ground state – lowest energy state of an atom
• Excited state – state where an atom has a
higher potential energy than ground state.
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Bohr Model
• 1913 – Niels Bohr proposed a hydrogen atom
model where electrons circle the nucleus only
in allowed paths or orbits with a definite
amount of energy.
• If an electron absorbs energy, it can go to a
higher level.
• If in a higher energy level, an electron can
emit a certain amount of energy to move to a
lower level.
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chemweb.ucc.ie
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Quantum Model of the Atom
• Questions were unanswered regarding how
electrons could be particles yet they gave off
waves of light.
• De Broglie suggested that electrons could be
considered waves confined to space around a
nucleus only at specific frequencies.
• Diffraction experiments proved that electron
beams can interfere with each other and produce
areas of low energy and high energy areas as a
result of interference.
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• Heisenberg Uncertainty Principle (1927) – it is
impossible to determine simultaneously both the
position and velocity of an electron or any other
particle.
• Schrodinger Wave Equation (1926) - developed
an equation that treated electrons in atoms as
waves.
• Heisenberg and Schrodinger laid the foundation
for mathematical descriptions of wave properties
of very small particles such as electrons – the
probable location of electrons around the
nucleus.
• AKA: Quantum Theory and Quantum Numbers.
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• End of Section 1 and part of Section 2 of
Chapter 4
• Arrangement of Electrons in Atoms
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