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Chapter 12
Electrons in Atoms
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Greek Idea
Democritus and Leucippus
Matter is made up of indivisible particles
Dalton - one type of atom for each element
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Thomson’s ModelDiscovered electrons
Atoms were made of positive stuff
Negative electron floating around
“Plum-Pudding” model
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Rutherford’s ModelDiscovered dense positive piece at the center of the atom
Nucleus
Electrons moved around
Mostly empty space
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Bohr’s Model
Why don’t the electrons fall into the nucleus?
Move like planets around the sun.
In circular orbits at different levels.
Amounts of energy separate one level from another.
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Bohr’s Model
Nucleus
Electron
Orbit
Energy Levels
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Bohr’s ModelIn
crea
sing
ene
rgy
Nucleus
First
Second
Third
Fourth
Fifth
} Further away from the nucleus means more energy. There is no “in between” energy Energy Levels
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The Quantum Mechanical ModelEnergy is quantized. It comes in chunks.
A quanta is the amount of energy needed to move from one energy level to another.
Since the energy of an atom is never “in between” there must be a quantum leap in energy.
Schrödinger derived an equation that described the energy and position of the electrons in an atom
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Things that are very small behave differently from things big enough to see.
The quantum mechanical model is a mathematical solution.
It is not like anything you can see.
The Quantum Mechanical Model
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Has energy levels for electrons.
Orbits are not circular.
It can only tell us the probability of finding an electron at a certain distance from the nucleus.
The Quantum Mechanical Model
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The atom is found inside a blurry “electron cloud”
A area where there is a chance of finding an electron.
Draw a line at 90 %
The Quantum Mechanical Model
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Atomic Orbitals
Principal Quantum Number (n) = the energy level of the electron.
Within each energy level the complex math of Schrödinger's equation describes several shapes.
These are called atomic orbitals
Regions where there is a high probability of finding an electron.
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There is 1 s orbital for every energy level Spherical shaped
Each s orbital can hold 2 electrons.
Called the 1s, 2s, 3s, etc.. orbitals.
S orbitals
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P orbitalsStart at the second energy level
3 different directions 3 different shapes
Each can hold 2 electrons
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P Orbitals
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D orbitals Start at the third energy level 5 different shapes Each can hold 2 electrons
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F orbitals Start at the fourth energy level Have seven different shapes 2 electrons per shape
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By Energy Level First Energy
Level only s orbital only 2 electrons 1s2
Second Energy Level
s and p orbitals are available
2 in s, 6 in p 2s22p6
8 total electrons
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By Energy Level Third energy level s, p, and d
orbitals 2 in s, 6 in p, and
10 in d 3s23p63d10
18 total electrons
Fourth energy level
s,p,d, and f orbitals
2 in s, 6 in p, 10 in d, and 14 in f
4s24p64d104f14
32 total electrons
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By Energy Level Any more than
the fourth and not all the orbitals will fill up.
You simply run out of electrons
The orbitals do not fill up in a neat order.
The energy levels overlap
Lowest energy fill first.
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1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5fIn
crea
sing
Ene
rgy
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Electron Configurations The way electrons are arranged in
atoms. Aufbau principle- electrons enter the
lowest energy first. This causes difficulties because of the
overlap of orbitals of different energies. Pauli Exclusion Principle- at most 2
electrons per orbital - different spins
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Electron Configuration Hund’s Rule- When electrons occupy
orbitals of equal energy they don’t pair up until they have to .
Let’s determine the electron configuration for Phosphorus
Need to account for 15 electrons
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1s2s 2p3s 3p 3d4s 4p 4d 4f5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f
1s2 2s2 2p6 3s2 3p6 4s23d104p4
What is it?
Selenium
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Orbitals fill in order Lowest energy to higher energy. Adding electrons can change the
energy of the orbital. Half filled orbitals have a lower
energy. Makes them more stable. Changes the filling order
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Write these electron configurations Titanium - 22 electrons 1s22s22p63s23p64s23d2
•Vanadium - 23 electrons• 1s22s22p63s23p64s23d3
Chromium - 24 electrons 1s22s22p63s23p64s23d4 is expectedBut this is wrong!! Why??
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Light The study of light led to the
development of the quantum mechanical model.
Light is a kind of electromagnetic radiation.
Electromagnetic radiation includes many kinds of waves
All move at 3.00 x 108 m/s = C
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Parts of a wave
Wavelength
AmplitudeOrigin
Crest
Trough
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Parts of Wave Origin - the base line of the energy. Crest - high point on a wave Trough - Low point on a wave Amplitude - distance from origin to crest Wavelength - distance from crest to crest Wavelength - is abbreviated Greek
letter lambda.
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Frequency The number of waves that pass a
given point per second. Units are cycles/sec or hertz (hz) Abbreviated the Greek letter nu
c =
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Frequency and wavelength Are inversely related As one goes up the other goes down. Different frequencies of light is
different colors of light. There is a wide variety of frequencies The whole range is called a spectrum
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Radio waves
Microwaves
Infrared
Ultra-violet
X-Rays
Gamma Rays
Low energy
High energy
Low Frequency
High Frequency
Long Wavelength
Short WavelengthVisible Light
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Atomic Spectrum
How color tells us about atoms
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Prism White light is
made up of all the colors of the visible spectrum.
Passing it through a prism separates it.
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If the light is not white By heating a gas
with electricity we can get it to give off colors.
Passing this light through a prism does something different.
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Wave-Particle DualityJJ Thomson won the Nobel prize for describing the electron as a particle.
His son, George Thomson won the Nobel prize for describing the wave-like nature of the electron.
The electron
is a particle!
The electron is an energy
wave!
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Confused??? You’ve Got Company!
“No familiar conceptions can be woven around the
electron; something unknown is doing we
don’t know what.”
Physicist Sir Arthur Eddington
The Nature of the Physical World
1934
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The Wave-like Electron
Louis deBroglie
The electron propagates through space as an energy
wave. To understand the atom, one must
understand the behavior of
electromagnetic waves.
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c = C = speed of light, a constant (3.00 x 108 m/s) = frequency, in units of hertz (hz, sec-1)
= wavelength, in meters
Electromagnetic radiation propagates through space as a wave moving at the speed of light.
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Types of electromagnetic radiation:
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E = h
EE = Energy, in units of Joules = Energy, in units of Joules (kg·m(kg·m22/s/s22))hh = Planck’s constant (6.626 x 10 = Planck’s constant (6.626 x 10-34-34 J·s)J·s)
= frequency, in units of hertz (hz, sec= frequency, in units of hertz (hz, sec-1-1))
The energy (E ) of electromagnetic radiation is directly proportional to the frequency () of the radiation.
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Long Wavelength
=Low Frequency
=Low ENERGY
Short Wavelength
=High Frequency
=High ENERGY
Wavelength Table
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…produces all of the colors in a continuous spectrum
Spectroscopic analysis of the visible spectrum…
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…produces a “bright line” spectrum
Spectroscopic analysis of the hydrogen spectrum…
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This produces bandsof light with definitewavelengths.
Electron transitionsinvolve jumps of
definite amounts ofenergy.