Chapter 6 Electronic Structure of Atoms. Why is the electron structure important? When atoms react...
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Transcript of Chapter 6 Electronic Structure of Atoms. Why is the electron structure important? When atoms react...
![Page 1: Chapter 6 Electronic Structure of Atoms. Why is the electron structure important? When atoms react it is the electrons of the atom that interact. The.](https://reader031.fdocuments.net/reader031/viewer/2022013012/5697bfa11a28abf838c95e2b/html5/thumbnails/1.jpg)
Chapter 6
Electronic Structure of Atoms
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Why is the electron structure important?
When atoms react it is the electrons of the atom that interact.
The electron structure is the arrangement of electrons within the atom.
Much of our understanding of the electronic structure of atoms comes from analysis of light emitted or absorbed by substances
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Light
Light is a type of electromagnetic radiation. Electromagnetic radiation carries energy
through space (aka radiant energy). Electromagnetic radiation moves at the speed of light.
Many different kinds of electromagnetic radiation.
Goes back to the “wave theory” of electron movement.
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Parts of a wave
Wavelength
Frequency = number of cycles in one secondMeasured in hertz 1 hertz = 1 cycle/second
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Frequency =
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Kinds of EM waves
There are many different and Radio waves, microwaves, x rays
and gamma rays are all examples. Light is only the part our eyes can
detect.GammaRays
Radiowaves
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Visible light
Visible light has a wavelength of 400-750 nm
Violet- highest energy Red- lowest energy The only type of electromagnetic
radiation we can see.
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The speed of light
Because all electromagnetic energy moves at the speed of light, wavelength and frequency must be related (inversely)
C= in a vacuum is 2.998 x 108 m/s c = The many different properties of
electromagnetic radiation come from the different wavelengths.
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Problems
What is the wavelength of light with a frequency 5.89 x 105 Hz?
What is the frequency of blue light with a wavelength of 484 nm?
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In 1900
Matter and energy were seen as different from each other in fundamental ways.
Matter was particles. Energy could come in waves, with
any frequency.
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The study of Energy
When something is heated, it gives off radiation.
When trying to understand the relationship between temperature and intensity of wavelengths.
Current laws could not account for the fact that white hot is hotter than red hot.
(Bunsen Burner)
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Planck Planck found energy can be released or
absorbed only in chunks (he called quantum)
E = h He proposed the energy of a single quantum
equals its frequency times a constant Planck’s constant (h = 6.626 x 10-34 J s) Energy is emitted/ absorbed in multiples of
this (nh (climb a ladder)
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Einstein’s Theory of Light
Used Planck’s quantum theory to explain the photoelectric effect.
Photoelectric Effect- (shining a light on a piece of metal causes the metal to emit electrons, each metal requires a specific frequency to get it started)
Said electromagnetic radiation is quantized in particles called photons.
Energy of light dependent on frequency High frequency (short wavelengths) x-rays- high
energy photons able to cause tissue damage
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A Photon
Each photon has energy = h = hc/ Combine this with E = mc2 You get the apparent mass of a
photon. m = h / (c) This is why you wear a lead shield
when you have an x-ray, but not when you use your microwave!
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Is light a wave, or is it made of particles?
Both!! Dual Nature of light This dual nature also true for
matter
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Matter as a wave
Using the velocity v instead of the wavelength we get.
De Broglie’s equation = h/mv Can calculate the wavelength of an
object.
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Examples
The laser light of a CD is 7.80 x 102 m. What is the frequency of this light?
What is the energy of a photon of this light?
What is the apparent mass of a photon of this light?
What is the energy of a mole of these photons?
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What is the wavelength?
of an electron with a mass of 9.11 x 10-31 kg traveling at
1.0 x 107 m/s? Of a softball with a mass of 0.10 kg
moving at 125 mi/hr?
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Spectrum
The range of frequencies present in light.
White light has a continuous spectrum.
All the colors are possible. A rainbow.
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Hydrogen spectrum
Emission spectrum because these are the colors it gives off or emits.
Called a line spectrum. There are just a few discrete lines showing
410 nm
434 nm
486 nm
656 nm
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What this means
Only certain energies are allowed for the hydrogen atom.
Can only give off certain energies. Use E = h= hc / Energy in the in the atom is
quantized.
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Niels Bohr
Developed the quantum model of the hydrogen atom.
He said the atom was like a solar system.
The electrons were attracted to the nucleus because of opposite charges.
Didn’t fall in to the nucleus because it was moving around.
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The Bohr Ring Atom
He didn’t know why but only certain energies were allowed.
He called these allowed energies energy levels.
Putting Energy into the atom moved the electron away from the nucleus.
From ground state to excited state. When it returns to ground state it gives
off light of a certain energy.
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The Bohr Ring Atom
n = 3n = 4
n = 2n = 1
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The Bohr Model
n is the energy level for each energy level the energy is Z is the nuclear charge, which is +1 for
hydrogen. E = -2.178 x 10-18 J (Z2 / n2 ) n = 1 is called the ground state when the electron is removed, n = E = 0
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We are worried about the change
When the electron moves from one energy level to another.
E = Efinal - Einitial
E = -2.178 x 10-18 J Z2 (1/ nf2 - 1/
ni2)
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Examples
Calculate the energy need to move an electron from its first energy level to the third energy level.
Calculate the energy released when an electron moves from n= 4 to n=2 in a hydrogen atom.
Calculate the energy released when an electron moves from n= 5 to n=3 in a He+1 ion
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When is it true?
Only for hydrogen atoms and other monoelectronic species.
Why the negative sign? To increase the energy of the electron
you make it closer to the nucleus. the maximum energy an electron can
have is zero, at an infinite distance.
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The Bohr Model
Doesn’t work. Only works for hydrogen atoms. Electrons don’t move in circles. The quantization of energy is right,
but not because they are circling like planets.
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The Quantum Mechanical Model
A totally new approach. De Broglie said matter could be
like a wave. De Broglie said they were like
standing waves. The vibrations of a stringed
instrument.
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What’s possible?
There are only certain allowed waves.
In the atom there are certain allowed waves called electrons.
1925 Erwin Schroedinger described the wave function of the electron.
Much math but what is important are the solution.
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Schroedinger’s Equation
Solutions to the equation are called orbitals.
These are not Bohr orbits. Each solution is tied to a certain
energy. These are the energy levels.
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There is a limit to what we can know
We can’t know how the electron is moving or how it gets from one energy level to another.
The Heisenberg Uncertainty Principle.
There is a limit to how well we can know both the position and the momentum of an object.
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Quantum Numbers
There are many solutions to Schroedinger’s equation (n, l, ml)
Principal quantum number (n) size and energy of of an orbital.
Has integer values >0
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Quantum numbers Angular momentum quantum number l . shape of the orbital. integer values from 0 to n-1 l = 0 is called s l = 1 is called p l =2 is called d l =3 is called f l =4 is called g
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S orbitals
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P orbitals
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P Orbitals
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D orbitals
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F orbitals
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F orbitals
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Quantum numbers
Magnetic quantum number (m l)
integer values between - l and + l tells direction in each shape. Electron spin quantum number (m s)
Can have 2 values. either +1/2 or -1/2
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Polyelectronic Atoms More than one electron. three energy contributions. The kinetic energy of moving
electrons. The potential energy of the
attraction between the nucleus and the electrons.
The potential energy from repulsion of electrons.
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Polyelectronic atoms Can’t solve Schroedinger’s equation
exactly. Difficulty is repulsion of other
electrons. Solution is to treat each electron as
if it were effected by the net field of charge from the attraction of the nucleus and the repulsion of the electrons.
Effective nuclear charge
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+10
11 electrons
+10 10 otherelectrons
e-Zeff
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The Periodic Table
Developed independently by German Julius Lothar Meyer and Russian Dmitri Mendeleev (1870”s).
Didn’t know much about atom. Put in columns by similar properties. Predicted properties of missing
elements.
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Aufbau Principle
Aufbau is German for building up. As the protons are added one by
one, the electrons fill up hydrogen-like orbitals.
Fill up in order of energy levels.
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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
He with 2 electrons
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More on Electron Configuration
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Details
Valence electrons- the electrons in the outermost energy levels (not d).
Core electrons- the inner electrons. Hund’s Rule- The lowest energy
configuration for an atom is the one have the maximum number of of unpaired electrons in the orbital.
C 1s2 2s2 2p2
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Fill from the bottom up following the arrows
1s2s 2p3s 3p 3d4s 4p 4d 4f
5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f
• 1s2
• 2• electrons
2s2
• 4
2p6 3s2
• 12
3p6 4s2
• 20
3d10 4p6
5s2
• 38
4d10 5p6 6s2
• 56
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Details
Elements in the same column have the same electron configuration.
Put in columns because of similar properties.
Similar properties because of electron configuration.
Noble gases have filled energy levels. Transition metals are filling the d
orbitals
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Exceptions
Ti = [Ar] 4s2 3d2 V = [Ar] 4s2 3d3
Cr = [Ar] 4s1 3d5
Mn = [Ar] 4s2 3d5
Half filled orbitals. Scientists aren’t sure of why it
happens same for Cu [Ar] 4s1 3d10
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More exceptions
Lanthanum La: [Xe] 6s2 5d1 Cerium Ce: [Xe] 6s2 4f1 5d1
Promethium Pr: [Xe] 6s2 4f3 5d0
Gadolinium Gd: [Xe] 6s2 4f7 5d1
Lutetium Pr: [Xe] 6s2 4f14 5d1
We’ll just pretend that all except Cu and Cr follow the rules.
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More Polyelectronic
We can use Zeff to predict properties, if we determine it’s pattern on the periodic table.
Can use the amount of energy it takes to remove an electron for this.
Ionization Energy- The energy necessary to remove an electron from a gaseous atom.
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Shielding Electrons on the higher energy levels
tend to be farther out. Have to look through the other
electrons to see the nucleus. They are less effected by the nucleus. lower effective nuclear charge If shielding were completely effective,
Zeff = 1 Why isn’t it?
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Penetration
There are levels to the electron distribution for each orbital.
2s
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GraphicallyGraphically
Penetration
2s
Rad
ial P
roba
bili
ty
Distance from nucleus
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GraphicallyR
adia
l Pro
babi
lity
Distance from nucleus
3s
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Rad
ial P
roba
bili
ty
Distance from nucleus
3p
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Rad
ial P
roba
bili
ty
Distance from nucleus
3d
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Rad
ial P
roba
bili
ty
Distance from nucleus
4s
3d
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How orbitals differ
The more positive the nucleus, the smaller the orbital.
A sodium 1s orbital is the same shape as a hydrogen 1s orbital, but it is smaller because the electron is more strongly attracted to the nucleus.
The helium 1s is smaller as well. This provides for better shielding.
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Periodic Trends Ionization energy the energy required to
remove an electron form a gaseous atom Highest energy electron removed first. First ionization energy (I1) is that
required to remove the first electron. Second ionization energy (I2) - the
second electron etc. etc.
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Trends in ionization energy for Mg
• I1 = 735 kJ/mole
• I2 = 1445 kJ/mole
• I3 = 7730 kJ/mole
The effective nuclear charge increases as you remove electrons.
It takes much more energy to remove a core electron than a valence electron because there is less shielding.
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Explain this trend
For Al• I1 = 580 kJ/mole
• I2 = 1815 kJ/mole
• I3 = 2740 kJ/mole
• I4 = 11,600 kJ/mole
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Across a Period
Generally from left to right, I1
increases because there is a greater nuclear charge
with the same shielding. As you go down a group I1
decreases because electrons are farther away.
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It is not that simple
Zeff changes as you go across a
period, so will I1 Half filled and filled orbitals are
harder to remove electrons from. here’s what it looks like.
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Firs
t Ion
izat
ion
ener
gy
Atomic number
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Firs
t Ion
izat
ion
ener
gy
Atomic number
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Firs
t Ion
izat
ion
ener
gy
Atomic number
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Atomic Radius
The atomic radius increases going down a group.
The atomic radius decreases going across a period
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Electron Affinities Non metals have the most negative
electron affinity. (Release the most energy when they gain an electron, thus becoming more stable)
Metals have the least negative (sometimes 0) electron affinity because they don’t become more stable or lose energy when they gain an electron
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Parts of the Periodic Table
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The information it hides
Know the special groups It is the number and type of valence
electrons that determine an atoms chemistry.
You can get the electron configuration from it.
Metals lose electrons have the lowest IE Non metals- gain electrons most
negative electron affinities.
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The Alkali Metals
Doesn’t include hydrogen- it behaves as a non-metal
decrease in IE increase in radius Decrease in density decrease in melting point Behave as reducing agents
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Reducing ability
Lower IE< better reducing agents Cs>Rb>K>Na>Li works for solids, but not in aqueous
solutions. In solution Li>K>Na Why? It’s the water -there is an energy
change associated with dissolving
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The reaction with water Na and K react explosively with water Li doesn’t. Even though the reaction of Li has a
more negative H than that of Na and K Na and K melt H does not tell you speed of reaction More in Chapter 12.