chapter 13 “electrons in atoms” credits: stephen l. cotton charles page high school mr. daniel...
TRANSCRIPT
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Chapter 13
“Electrons in Atoms”
Credits: Stephen L. CottonCharles Page High School
Mr. DanielOlympic High School
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Section 13.3Physics and the Quantum Mechanical Model
OBJECTIVES:• Describe the relationship between the
wavelength and frequency • Distinguish between quantum mechanics
and classical mechanics. of light.Identify the source of atomic emission spectra.•Explain how the frequencies of emitted light are related to changes in electron energies.
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Light Visible light is a type of electromagnetic
radiation. Electromagnetic radiation is a form of energy
and includes many types: gamma rays, x-rays, radio waves, visible light…
Speed of light (c) = 3.00 x 108 m/s All electromagnetic radiation travels at this
same rate when measured in a vacuum
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Parts of a Wave
Wavelength
AmplitudeNode
Crest
Trough
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- Page 373
“R O Y G B I V”
Frequency Increases
Wavelength Shorter
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Long Wavelength
Low Frequency
Low ENERGY
Short Wavelength
High Frequency
High ENERGY
Wavelength Table
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Wavelength and Frequency: Are inversely related
• As one increases the other decreases. Different frequencies of visible light are different
colors. There is a wide variety of frequencies Spectrum: A whole range of electromagnetic
wavelengths. (e.g. the visible light spectrum)
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Radiowaves
Microwaves
Infrared .
Ultra-violet
X-Rays
GammaRays
Low Frequency
High Frequency
Long Wavelength
Short Wavelength
Visible Light
Low Energy
High Energy
The Electromagnetic Spectrum
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So what is Energy? All energy is quantized A quantum is a “packet” of energy.
Not all quanta (plural) are the same size.
(eggs are not all the same size either, but all are eggs)
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So what is Light Energy? Light is a form of energy. Therefore, light must be quantized A quantum of light energy is called a photon.
Einstein determined that light is not only a wave, but is also a particle!
He demonstrated it in an experiment that showed the photoelectric effect
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Photoelectric Effect
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Photoelectric EffectPhotoelectric EffectExperiment demonstrates the particle nature of light.Experiment demonstrates the particle nature of light.
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So what is Light Energy? (con’t)
Therefore, light has what is called wave-particle duality. It has characteristics of both waves and particles.
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Wave-Particle Duality (again)J.J. 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 Physics of the Very Small
Quantum mechanics explains how very small particles behave• Quantum mechanics is an explanation for
subatomic particles and atoms as waves Classical mechanics describes the motions of
bodies much larger than atoms
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Section 13.1Models of the Atom
OBJECTIVES:
• Identify the inadequacies in the Rutherford atomic model.
• Identify the new proposal in the Bohr model of the atom.
• Describe the energies and positions of electrons according to the quantum mechanical model.•Describe how the shapes of orbitals related to different sublevels differ.
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Ernest Rutherford’s Model Discovered dense positive
piece at the center of the atom- “nucleus”
Electrons would surround and move around the nucleus
Atom is mostly empty space It did not explain the chemical
properties of the elements – a better description of the electron behavior was needed
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Niels Bohr’s Model Why don’t the electrons fall into the nucleus?
He agreed with Rutherford that electrons move around the nucleus, But:
• In specific circular paths, or orbits (like planets around the sun), at specific energy levels.
• An amount of fixed energy separates one electron energy level from another.
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The Bohr Model of the Atom
Niels Bohr
I pictured the electrons orbiting the nucleus much like planets orbiting the sun.
However, electrons are found in specific energy levels around the nucleus, and can jump from one level to another.
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Bohr’s Model Electrons occupy specific energy levels
• analogous to the rungs of a ladder The electron cannot exist between energy
levels, just like you can’t stand between rungs on a ladder
A quantum of energy is the amount of energy required to move an electron from one energy level to another (plural: quanta)
Since the energy of an atom is never “in between” there must be a quantum leap in energy.
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Changing the energy Let’s look at a hydrogen atom, with only one
electron, and in the first energy level.
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Changing the energy Heat, electricity, or light can move the
electron up to different energy levels. The electron is now said to be in an “excited state”
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Changing the energy The electron is unstable at the higher energy
level and as it falls back to the ground state, it gives the energy back in the form of light
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They fall down in specific steps Each step has a different energy (quantum)
and results in a different color of light.
Changing the energy
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Origin of Line SpectraOrigin of Line Spectra
Balmer seriesBalmer series
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The further electrons fall, the more energy is released and the higher the frequency of light emitted.
This is a simplified explanation! Remember, the orbitals also have different
sublevels within the principle energy levels
Ultraviolet Visible Infrared
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Atomic Spectra
White light is made up of all the colors of the visible spectrum.
Passing it through a prism separates it.
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But not all light is 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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Atomic Spectrum Each element gives
off its own characteristic colors of light.
The colors can be used to identify the atom.
This is how we know what elements stars are made of.
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• This is called a bright line spectrum
• Unique to each element, like fingerprints!
• Very useful for identifying elements
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Explanation of Atomic Spectra
ground state - the lowest energy level of the electron.
In summary. When an electron at ground state receives a quantum of energy it jumps directly to a higher energy level. The electron is unstable and immediately drops to a lower energy level. As it drops it gives off the same amount of
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The Quantum Mechanical Model
Problems with Bohr’s theory :Problems with Bohr’s theory :
It was only successful for H- no other It was only successful for H- no other elements followed his predictions.elements followed his predictions.
It introduced the quantum idea artificially.It introduced the quantum idea artificially.
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Heisenberg Uncertainty Principle
You can find out where the electron is, but not its energy
OR…
You can know how much energy it has, but not where it is!
“One cannot simultaneously determine both the position and momentum of an electron.”
Werner Heisenberg
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Heisenberg Uncertainty Principle
It is impossible to know exactly the location and velocity of a particle simultaneously.
The better we know one, the less we know the other.
Measuring one property, changes the other.
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Moving Electron
Photon
Before
Electron velocity changes
Photon wavelengthchanges
After
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In 1926, Erwin Schrodinger derived an equation that described the energy and probable position of the electrons in an atom
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Schrodinger’s Wave Equation22
2 2
8dh EV
m dx
His equation determined the probabilityprobability of finding a single electron along a single axis (x-axis)
Erwin SchrodingerErwin Schrodinger
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Particles which are very small and travel very quickly (like electrons) behave very differently from objects big enough to observe.
The quantum mechanical model is a mathematical solution describing how those particles act.
The Quantum Mechanical Model
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Describes energy levels for electrons. Electrons move in an unpredictable manner We can only determine the probability of
finding an electron a certain distance from the nucleus.
The Quantum Mechanical Model
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The electrons are probably located inside a blurry “electron cloud”
The area where there is the greatest chance of finding an electron.
The Quantum Mechanical Model
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Atomic Orbitals Principal Quantum Number (n) = the energy
level of the electron: 1, 2, 3, etc. Within each energy level, there are sub-
levels (like theater seats arranged in sections): letters s, p, d, and f
The complex math of Schrodinger’s equation describes several shapes
These are the atomic orbitals - regions where there is a 90% probability of finding an electron.
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Principal Quantum NumberThe Principle Quantum Number (n) denotes the shell (energy level) in which the electron is located.
The maximum number of electrons that fit into an energy level can be calculated:
2n2
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Summary
s
p
d
f
# of orbitals
Max. electrons
Starts at energy level
1 2 1
3 6 2
5 10 3
7 14 4
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Types of OrbitalsTypes of Orbitals
s orbitals orbital p orbitalp orbital d orbitald orbital
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Types of
Atomic Orbitals
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By Energy Level First Energy Level Has only an s
sublevel only 2 electrons 1s2
Second Energy Level
Has s and p sublevels
2 e- in s, 6 e- in p 2s22p6
8 total electrons
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By Energy Level Third energy level Has s, p, and d
sublevels 2 e- in s, 6 e- in p,
and 10 e- in d 3s23p63d10
18 total electrons
Fourth energy level Has s, p, d, and f
sublevels 2 e- in s, 6 e- in p, 10
e- in d, and 14 e- in f 4s24p64d104f14
32 total electrons
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By Energy Level Beyond the fourth
energy level, not all sublevels fill up.
You simply run out of electrons
So only the s, p, d and f sublevels are used
Because the energy levels overlap the orbitals do not fill up in a consistent pattern
However, the lowest energy orbitals fill first.
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Section 13.2Electron Arrangement in Atoms
OBJECTIVES:• Describe how to write the electron
configuration for an atom.• Explain why the actual electron configurations for some elements differ from those predicted by the aufbau principle.
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Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
aufbau diagram
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Electron Configurations… …are the way electrons are arranged in
various orbitals around the nuclei of atoms. Three rules tell us how:
1) Aufbau principle - electrons enter the lowest energy sublevels first.
• This becomes complex because of the overlap of orbitals of different energies – follow the diagram!
2) Pauli Exclusion Principle – there are at most 2 electrons per orbital - with opposite spins
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Pauli Exclusion Principle
No two electrons in an atom can have the same four quantum numbers.
Wolfgang Pauli
To show the different direction of spin, a pair in the same orbital is written as:
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Electron Configurations3) Hund’s Rule- When electrons occupy
orbitals of equal energy, they don’t pair up until each orbital has one electron.
Let’s write the electron configuration for Phosphorus
We need to account for all 15 electrons in phosphorus
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The first two electrons go into the 1s orbital
Notice the opposite direction of the spins
only 13 more to go...Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
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The next electrons go into the 2s orbital
only 11 more...
Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
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• The next electrons go into the 2p orbital
• only 5 more...
Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
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• The next electrons go into the 3s orbital
• only 3 more...Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
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Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
• The last three electrons go into the 3p sublevel.
They each go into separate orbitals (Hund’s)
• 3 unpaired electrons
• = 1s22s22p63s23p3
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Orbitals fill in an order: Lowest energy to higher energy.
Adding electrons can change the energy of the orbital. Full sublevels are the most stable arrangement.
Half filled sublevels have a lower energy than partially filled sublevels, and are next most stable.
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Write the electron configurations for these elements:
Zirconium - 40 electrons [Kr] 5s2 4d2
Tantalum - 73 electrons [Xe] 6s2 4f14 5d3
Chromium - 24 electrons [Ar] 4s2 3d4 (expected)But this is not what happens with Chromium
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Chromium is actually: [Ar]4s13d5
Why? This gives us two half filled orbitals (the others
are all still full) Half full is slightly lower in energy. The same principal applies to copper…
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Copper’s electron configuration
Copper has 29 electrons so we expect: [Ar] 4s2 3d9
But the actual configuration is: [Ar]4s13d10
This change gives one more filled orbital and one that is half filled.
Remember these exceptions: Groups ending in d4 and d9
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Irregular configurations of Cr and Cu
Chromium steals a 4s electron to make its 3d sublevel HALF FULL
Copper steals a 4s electron to FILL its 3d sublevel
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