CHAPTER 5
description
Transcript of CHAPTER 5
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CHAPTER 5
The Structure of Atoms
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Chapter Outline
Subatomic Particles
1. Fundamental Particles
2. The Discovery of Electrons
3. Canal Rays and Protons
4. Rutherford and the Nuclear Atom
5. Atomic Number
6. Neutrons
7. Mass Number and Isotopes
8. Mass spectrometry and Isotopic Abundance
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Chapter Goals
9. The Atomic Weight Scale and Atomic Weights
The Electronic Structures of Atoms
10. Electromagnetic radiation
11. The Photoelectric Effect
12. Atomic Spectra and the Bohr Atom
13. The Wave Nature of the Electron
14. The Quantum Mechanical Picture of the Atom
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Chapter Goals
15. Quantum Numbers
16. Atomic Orbitals
17. Electron Configurations
18. Paramagnetism and Diamagnetism
19. The Periodic Table and Electron Configurations
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Fundamental Particles
Particle Mass (amu) Charge
Electron (e-) 0.00054858 -1
Proton (p,p+) 1.0073 +1
Neutron(n,n0) 1.0087 0
Three fundamental particles make up atoms. The following table lists these particles together with their masses and their charges.
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The Discovery of Electrons
Humphrey Davy in the early 1800’s passed electricity through compounds and noted:– that the compounds decomposed into elements.– Concluded that compounds are held together by
electrical forces. Michael Faraday in 1832-1833 realized that the
amount of reaction that occurs during electrolysis is proportional to the electrical current passed through the compounds.
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The Discovery of Electrons
Cathode Ray Tubes experiments performed in the late 1800’s & early 1900’s.
– Consist of two electrodes sealed in a glass tube containing a gas at very low pressure.
– When a voltage is applied to the cathodes a glow discharge is emitted.
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The Discovery of Electrons
These “rays” are emitted from cathode (- end) and travel to anode (+ end).– Cathode Rays must be negatively charged!
J.J. Thomson modified the cathode ray tube experiments in 1897 by adding two adjustable voltage electrodes.– Studied the amount that the cathode ray beam was
deflected by additional electric field.
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The Discovery of Electrons
Modifications to the basic cathode ray tube experiment.
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The Discovery of Electrons
Thomson used his modification to measure the charge to mass ratio of electrons.Charge to mass ratio
e/m = -1.75881 x 108 coulomb/g of e-
Thomson named the cathode rays electrons. Thomson is considered to be the “discoverer of
electrons”. TV sets and computer screens are cathode ray
tubes.
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The Discovery of Electrons
Robert A. Millikan won the 1st American Nobel Prize in 1923 for his famous oil-drop experiment.
In 1909 Millikan determined the charge and mass of the electron.
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The Discovery of Electrons
Millikan determined that the charge on a single electron = -1.60218 x 10-19 coulomb.
Using Thomson’s charge to mass ratio we get that the mass of one electron is 9.11 x 10-28 g.– e/m = -1.75881 x 108 coulomb– e = -1.60218 x 10-19 coulomb– Thus m = 9.10940 x 10-28 g
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Canal Rays and Protons
Eugene Goldstein noted streams of positively charged particles in cathode rays in 1886.
– Particles move in opposite direction of cathode rays. – Called “Canal Rays” because they passed through holes (channels or
canals) drilled through the negative electrode. Canal rays must be positive.
– Goldstein postulated the existence of a positive fundamental particle called the “proton”.
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Rutherford and the Nuclear Atom
Ernest Rutherford directed Hans Geiger and Ernst Marsden’s experiment in 1910. - particle scattering from thin Au foils – Gave us the basic picture of the atom’s structure.
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Rutherford and the Nuclear Atom
In 1912 Rutherford decoded the -particle scattering information.– Explanation involved a nuclear atom with electrons
surrounding the nucleus .
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Rutherford and the Nuclear Atom
Rutherford’s major conclusions from the -particle scattering experiment
1. The atom is mostly empty space.
2. It contains a very small, dense center called the nucleus.
3. Nearly all of the atom’s mass is in the nucleus.
4. The nuclear diameter is 1/10,000 to 1/100,000 times less than atom’s radius.
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Rutherford and the Nuclear Atom
Because the atom’s mass is contained in such a small volume:– The nuclear density is 1015g/mL.– This is equivalent to 3.72 x 109 tons/in3.– Density inside the nucleus is almost the same as a
neutron star’s density.
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Atomic Number
The atomic number is equal to the number of protons in the nucleus.
– Sometimes given the symbol Z.– On the periodic chart Z is the uppermost number in each
element’s box. In 1913 H.G.J. Moseley realized that the atomic
number determines the element .– The elements differ from each other by the number of protons
in the nucleus. – The number of electrons in a neutral atom is also equal to the
atomic number.
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Neutrons
James Chadwick in 1932 analyzed the results of -particle scattering on thin Be films.
Chadwick recognized existence of massive neutral particles which he called neutrons.– Chadwick discovered the neutron.
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Mass Number and Isotopes
Mass number is given the symbol A. A is the sum of the number of protons and neutrons.
– Z = proton number N = neutron number– A = Z + N
A common symbolism used to show mass and proton numbers is
AuCa, C, example for E 19779
4820
126
AZ
Can be shortened to this symbolism.
etc. Ag,Cu, N, 1076314
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Mass Number and Isotopes
Isotopes are atoms of the same element but with different neutron numbers.
– Isotopes have different masses and A values but are the same element.
One example of an isotopic series is the hydrogen isotopes.
1H or protium is the most common hydrogen isotope. one proton and no neutrons
2H or deuterium is the second most abundant hydrogen isotope. one proton and one neutron
3H or tritium is a radioactive hydrogen isotope. one proton and two neutrons
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Mass Number and Isotopes
The stable oxygen isotopes provide another example. 16O is the most abundant stable O isotope.
• How many protons and neutrons are in 16O?neutrons 8 and protons 8
17O is the least abundant stable O isotope. How many protons and neutrons are in 17O?
18O is the second most abundant stable O isotope.How many protons and neutrons in 18O?
neutrons 9 and protons 8
neutrons 10 and protons 8
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Mass Spectrometry andIsotopic Abundances
Francis Aston devised the first mass spectrometer.– Device generates ions that pass down an evacuated path inside
a magnet.– Ions are separated based on their mass.
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Mass Spectrometry andIsotopic Abundances
There are four factors which determine a particle’s path in the mass spectrometer.1 accelerating voltage2 magnetic field strength3 masses of particles4 charge on particles
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Mass Spectrometry andIsotopic Abundances
Mass spectrum of Ne+ ions shown below.– How scientists determine the masses and
abundances of the isotopes of an element.
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The Atomic Weight Scale and Atomic Weights
If we define the mass of 12C as exactly 12 atomic mass units (amu), then it is possible to establish a relative weight scale for atoms.
– 1 amu = (1/12) mass of 12C by definition– What is the mass of an amu in grams?
Example 5-1: Calculate the number of atomic mass units in one gram.
– The mass of one 31P atom has been experimentally determined to be 30.99376 amu.
– 1 mol of 31P atoms has a mass of 30.99376 g.
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The Atomic Weight Scale and Atomic Weights
Pg 30.99376
atoms P 10 6.022g) (1.000
31
3123
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The Atomic Weight Scale and Atomic Weights
– Thus 1.00 g = 6.022 x 1023 amu.– This is always true and provides the conversion factor between grams and amu.
Pamu 10 6.022 atom P
amu 30.99376
Pg 30.99376
atoms P 10 6.022g) (1.000
312331
31
3123
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The Atomic Weight Scale and Atomic Weights
The atomic weight of an element is the weighted average of the masses of its stable isotopes
Example 5-2: Naturally occurring Cu consists of 2 isotopes. It is 69.1% 63Cu with a mass of 62.9 amu, and 30.9% 65Cu, which has a mass of 64.9 amu. Calculate the atomic weight of Cu to one decimal place.
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The Atomic Weight Scale and Atomic Weights
amu) .9(0.691)(62 weight atomic
isotopeCu 63
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The Atomic Weight Scale and Atomic Weights
isotopeCu isotopeCu 6563
amu) .9(0.309)(64 amu) .9(0.691)(62 weight atomic
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The Atomic Weight Scale and Atomic Weights
copperfor amu 63.5 weight atomic
amu) .9(0.309)(64 amu) .9(0.691)(62 weight atomic
isotopeCu isotopeCu 6563
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The Atomic Weight Scale and Atomic Weights
Example 5-3: Naturally occurring chromium consists of four isotopes. It is 4.31% 24
50Cr, mass = 49.946 amu, 83.76% 24
52Cr, mass = 51.941 amu, 9.55% 24
53Cr, mass = 52.941 amu, and 2.38% 24
54Cr, mass = 53.939 amu. Calculate the atomic weight of chromium.
You do it!You do it!
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The Atomic Weight Scale and Atomic Weights
amu 51.998
amu 1.2845.056 43.506 2.153
amu) 53.9390.0238(amu) 52.941(0.0955
amu) 51.941(0.8376amu) 49.946(0.0431 weight atomic
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The Atomic Weight Scale and Atomic Weights
Example 5-4: The atomic weight of boron is 10.811 amu. The masses of the two naturally occurring isotopes 5
10B and 511B, are 10.013
and 11.009 amu, respectively. Calculate the fraction and percentage of each isotope.
You do it!You do it! This problem requires a little algebra.
– A hint for this problem is x + (1-x) = 1
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The Atomic Weight Scale and Atomic Weights
x
x
xx
xx
xx
199.0
-0.9960.198-
amu 11.009 - 10.013amu 11.009-10.811
amu 11.009-11.009 10.013
amu) (11.0091amu) (10.013 amu 10.811
isotope Bisotope B 1110
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The Atomic Weight Scale and Atomic Weights
Note that because x is the multiplier for the 10B isotope, our solution gives us the fraction of natural B that is 10B.
Fraction of 10B = 0.199 and % abundance of 10B = 19.9%.
The multiplier for 11B is (1-x) thus the fraction of 11B is 1-0.199 = 0.801 and the % abundance of 11B is 80.1%.
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The Electronic Structures of AtomsThe Electronic Structures of AtomsElectromagnetic Radiation
The wavelengthwavelength of electromagnetic radiation has the symbol
Wavelength is the distance from the top (crest) of one wave to the top of the next wave.
– Measured in units of distance such as m,cm, Å.– 1 Å = 1 x 10-10 m = 1 x 10-8 cm
The frequencyfrequency of electromagnetic radiation has the symbol
Frequency is the number of crests or troughs that pass a given point per second.
– Measured in units of 1/time - s-1
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Electromagnetic Radiation
The relationship between wavelength and frequency for any wave is velocity =
For electromagnetic radiation the velocity is 3.00 x 108 m/s and has the symbol c.
Thus c = forelectromagnetic radiation
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Electromagnetic Radiation
Molecules interact with electromagnetic radiation.– Molecules can absorb and emit light.
Once a molecule has absorbed light (energy), the molecule can:1. Rotate2. Translate3. Vibrate4. Electronic transition
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Electromagnetic Radiation
For water:– Rotations occur in the microwave portion of spectrum.– Vibrations occur in the infrared portion of spectrum.
– Translation occurs across the spectrum.– Electronic transitions occur in the ultraviolet portion of spectrum.
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Electromagnetic Radiation
Example 5-5: What is the frequency of green light of wavelength 5200 Å?
m10 5.200 Å 1
m 10 x 1 Å) (5200
c c
7-10-
1-14
7-
8
s 10 5.77
m10 5.200
m/s 10 3.00
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Electromagnetic Radiation
In 1900 Max Planck studied black body radiation and realized that to explain the energy spectrum he had to assume that:
1. energy is quantized
2. light has particle character
Planck’s equation is
sJ 10x 6.626 constant s Planck’ h
hc or E h E
34-
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Electromagnetic Radiation
Example 5-6: What is the energy of a photon of green light with wavelength 5200 Å?What is the energy of 1.00 mol of these photons?
photon per J10 3.83 E
) s 10 s)(5.77J10 (6.626 E
h E
s 10 x 5.77 that know we5,-5 Example From
19-
1-1434-
-114
kJ/mol 231 photon)per J10 .83photons)(3 10 (6.022
:photons of mol 1.00For 19-23
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The Photoelectric Effect
Light can strike the surface of some metals causing an electron to be ejected.
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The Photoelectric Effect
What are some practical uses of the photoelectric effect?
You do it!
Electronic door openers Light switches for street lights Exposure meters for cameras Albert Einstein explained the photoelectric effect
– Explanation involved light having particle-like behavior.– Einstein won the 1921 Nobel Prize in Physics for this work.
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Atomic Spectra and the Bohr Atom
An emission spectrum is formed by an electric current passing through a gas in a vacuum tube (at very low pressure) which causes the gas to emit light.– Sometimes called a bright line spectrum.
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Atomic Spectra and the Bohr Atom
An absorption spectrum is formed by shining a beam of white light through a sample of gas.– Absorption spectra indicate the wavelengths of light that
have been absorbed.
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Atomic Spectra and the Bohr Atom
Every element has a unique spectrum. Thus we can use spectra to identify elements.
– This can be done in the lab, stars, fireworks, etc.
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Atomic Spectra and the Bohr Atom
Atomic and molecular spectra are important indicators of the underlying structure of the species.
In the early 20th century several eminent scientists began to understand this underlying structure.– Included in this list are:– Niels Bohr– Erwin Schrodinger – Werner Heisenberg
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Atomic Spectra and the Bohr Atom
Example 5-7: An orange line of wavelength 5890 Å is observed in the emission spectrum of sodium. What is the energy of one photon of this orange light?
You do it!You do it!
J 10375.3
m 10890.5
m/s 1000.3sJ 10626.6
hchE
m 10890.5Å
m 10 1 Å 5890
19
7
834
7-10
m 10890.5Å
m 10 1 Å 5890 7
-10
hchE
m 10890.5Å
m 10 1 Å 5890 7
-10
m 10890.5
m/s 1000.3sJ 10626.6
hchE
m 10890.5Å
m 10 1 Å 5890
7
834
7-10
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Atomic Spectra and the Bohr Atom
The Rydberg equation is an empirical equation that relates the wavelengths of the lines in the hydrogen spectrum.
hydrogen of spectrumemission
in the levelsenergy theof
numbers therefer to sn’
n n
m 10 1.097 R
constant Rydberg theis R
n
1
n
1R
1
21
1-7
22
21
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Atomic Spectra and the Bohr Atom
Example 5-8. What is the wavelength of light emitted when the hydrogen atom’s energy changes from n = 4 to n = 2?
22 1-7
22
21
12
4
1
2
1m 10 1.097
1
n
1
n
1R
1
2n and 4n
22
21
12
n
1
n
1R
1
2n and 4n
16
1
4
1m 10 1.097
1
4
1
2
1m 10 1.097
1
n
1
n
1R
1
2n and 4n
1-7
22 1-7
22
21
12
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Atomic Spectra and the Bohr Atom
m 104.862
m 10 2.057 1
1875.0m 10 1.097 1
0625.0250.0m 10 1.097 1
7-
1-6
1-7
1-7
1-6
1-7
1-7
m 10 2.057 1
1875.0m 10 1.097 1
0625.0250.0m 10 1.097 1
1875.0m 10 1.097 1
0625.0250.0m 10 1.097 1
1-7
1-7
0625.0250.0m 10 1.097
1 1-7
Notice that the wavelength calculated from the Rydberg equation matches the wavelength of the green colored line in the H spectrum.
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Atomic Spectra and the Bohr Atom
In 1913 Neils Bohr incorporated Planck’s quantum theory into the hydrogen spectrum explanation.
Here are the postulates of Bohr’s theory.1. Atom has a number of definite and discrete
energy levels (orbits) in which an electron may exist without emitting or absorbing electromagnetic radiation.
As the orbital radius increases so does the energy
1<2<3<4<5......
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Atomic Spectra and the Bohr Atom
2. An electron may move from one discrete energy level (orbit) to another, but, in so doing, monochromatic radiation is emitted or absorbed in accordance with the following equation.
E E
hc h E E - E
12
1 2
Energy is absorbed when electrons jump to higher orbits.
n = 2 to n = 4 for example
Energy is emitted when electrons fall to lower orbits.
n = 4 to n = 1 for example
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Atomic Spectra and the Bohr Atom
3. An electron moves in a circular orbit about the nucleus and it motion is governed by the ordinary laws of mechanics and electrostatics, with the restriction that the angular momentum of the electron is quantized (can only have certain discrete values).
angular momentum = mvr = nh/2h = Planck’s constant n = 1,2,3,4,...(energy levels)
v = velocity of electron m = mass of electron
r = radius of orbit
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Atomic Spectra and the Bohr Atom
Light of a characteristic wavelength (and frequency) is emitted when electrons move from higher E (orbit, n = 4) to lower E (orbit, n = 1).– This is the origin of emission spectra.
Light of a characteristic wavelength (and frequency) is absorbed when electron jumps from lower E (orbit, n = 2) to higher E (orbit, n= 4)– This is the origin of absorption spectra.
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59
Atomic Spectra and the Bohr Atom
Bohr’s theory correctly explains the H emission spectrum.
The theory fails for all other elements because it is not an adequate theory.
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60
The Wave Nature of the Electron
In 1925 Louis de Broglie published his Ph.D. dissertation.– A crucial element of his dissertation is that electrons
have wave-like properties.– The electron wavelengths are described by the de
Broglie relationship.
particle of velocity v
particle of mass m
constant s Planck’ hmv
h
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61
The Wave Nature of the Electron
De Broglie’s assertion was verified by Davisson & Germer within two years.
Consequently, we now know that electrons (in fact - all particles) have both a particle and a wave like character.– This wave-particle duality is a fundamental property
of submicroscopic particles.
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62
The Wave Nature of the Electron
Example 5-9. Determine the wavelength, in m, of an electron, with mass 9.11 x 10-31 kg, having a velocity of 5.65 x 107 m/s.
– Remember Planck’s constant is 6.626 x 10-34 Js which is also equal to 6.626 x 10-34 kg m2/s2.
m 1029.1
m/s 1065.5kg 109.11
sm kg 10626.6
mv
h
11
731-
2234
m/s 1065.5kg 109.11
sm kg 10626.6
mv
h
731-
2234
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63
The Wave Nature of the Electron
Example 5-10. Determine the wavelength, in m, of a 0.22 caliber bullet, with mass 3.89 x 10-3 kg, having a velocity of 395 m/s, ~ 1300 ft/s.
You do it!You do it!
m 1031.4
m/s 953kg 103.89
sm kg 10626.6
mv
h
34
3-
2234
m/s 953kg 1089.3
sm kg 10626.6
mv
h
3-
2234
• Why is the bullet’s wavelength so small compared to the electron’s wavelength?
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64
The Quantum Mechanical Picture of the Atom
Werner Heisenberg in 1927 developed the concept of the Uncertainty Principle.
It is impossible to determine simultaneously both the position and momentum of an electron (or any other small particle).– Detecting an electron requires the use of
electromagnetic radiation which displaces the electron! Electron microscopes use this phenomenon
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65
The Quantum Mechanical Picture of the Atom
Consequently, we must must speak of the electrons’ position about the atom in terms of probability functions.
These probability functions are represented as orbitals in quantum mechanics.
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66
The Quantum Mechanical Picture of the Atom
Basic Postulates of Quantum Theory
1. Atoms and molecules can exist only in certain energy states. In each energy state, the atom or molecule has a definite energy. When an atom or molecule changes its energy state, it must emit or absorb just enough energy to bring it to the new energy state (the quantum condition).
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67
The Quantum Mechanical Picture of the Atom
2. Atoms or molecules emit or absorb radiation (light) as they change their energies. The frequency of the light emitted or absorbed is related to the energy change by a simple equation.
hc
h E
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68
The Quantum Mechanical Picture of the Atom
3. The allowed energy states of atoms and molecules can be described by sets of numbers called quantum numbers.
Quantum numbers are the solutions of the Schrodinger, Heisenberg & Dirac equations.
Four quantum numbers are necessary to describe energy states of electrons in atoms.
EV8
b
equationdinger oSchr
2
2
2
2
2
2
2
2
..
zyxm
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69
Quantum Numbers
The principal quantum number has the symbol – n.n = 1, 2, 3, 4, ...... “shells”
n = K, L, M, N, ......
The electron’s energy depends principally on n .
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70
Quantum Numbers
The angular momentum quantum number has the symbol . = 0, 1, 2, 3, 4, 5, .......(n-1) = s, p, d, f, g, h, .......(n-1)
tells us the shape of the orbitals. These orbitals are the volume around the atom
that the electrons occupy 90-95% of the time.This is one of the places where Heisenberg’s
Uncertainty principle comes into play.
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71
Quantum Numbers
The symbol for the magnetic quantum number is m.m = - , (- + 1), (- +2), .....0, ......., ( -2), ( -1),
If = 0 (or an s orbital), then m = 0. – Notice that there is only 1 value of m.
This implies that there is one s orbital per n value. n 1
If = 1 (or a p orbital), then m = -1,0,+1.– There are 3 values of m.
Thus there are three p orbitals per n value. n 2
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72
Quantum Numbers
If = 2 (or a d orbital), then m = -2,-1,0,+1,+2.– There are 5 values of m.
Thus there are five d orbitals per n value. n 3
If = 3 (or an f orbital), then m = -3,-2,-1,0,+1,+2, +3. – There are 7 values of m.
Thus there are seven f orbitals per n value, n 4
Theoretically, this series continues on to g,h,i, etc. orbitals.
– Practically speaking atoms that have been discovered or made up to this point in time only have electrons in s, p, d, or f orbitals in their ground state configurations.
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73
Quantum Numbers
The last quantum number is the spin quantum number which has the symbol ms.
The spin quantum number only has two possible values.
– ms = +1/2 or -1/2– ms = ± 1/2
This quantum number tells us the spin and orientation of the magnetic field of the electrons.
Wolfgang Pauli in 1925 discovered the Exclusion Principle.
– No two electrons in an atom can have the same set of 4 quantum numbers.
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74
Atomic Orbitals
Atomic orbitals are regions of space where the probability of finding an electron about an atom is highest.
s orbital properties:– There is one s orbital per n level. = 0 1 value of m
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75
Atomic Orbitals
s orbitals are spherically symmetric.
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76
Atomic Orbitals
p orbital properties:– The first p orbitals appear in the n = 2 shell.
p orbitals are peanut or dumbbell shaped volumes.– They are directed along the axes of a Cartesian coordinate
system.
There are 3 p orbitals per n level. – The three orbitals are named px, py, pz.
– They have an = 1.– m = -1,0,+1 3 values of m
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77
Atomic Orbitals
p orbitals are peanut or dumbbell shaped.
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78
Atomic Orbitals
d orbital properties:– The first d orbitals appear in the n = 3 shell.
The five d orbitals have two different shapes:– 4 are clover leaf shaped.– 1 is peanut shaped with a doughnut around it.– The orbitals lie directly on the Cartesian axes or are rotated
45o from the axes.
222 zy-xxzyzxy d ,d ,d ,d ,d
There are 5 d orbitals per n level.–The five orbitals are named –They have an = 2.–m = -2,-1,0,+1,+2 5 values of m
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79
Atomic Orbitals
d orbital shapes
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80
Atomic Orbitals
f orbital properties:– The first f orbitals appear in the n = 4 shell.
The f orbitals have the most complex shapes. There are seven f orbitals per n level.
– The f orbitals have complicated names.– They have an = 3– m = -3,-2,-1,0,+1,+2, +3 7 values of m
– The f orbitals have important effects in the lanthanide and actinide elements.
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81
Atomic Orbitals
f orbital shapes
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82
Atomic Orbitals
Spin quantum number effects:– Every orbital can hold up to two electrons.
Consequence of the Pauli Exclusion Principle.– The two electrons are designated as having– one spin up and one spin down
Spin describes the direction of the electron’s magnetic fields.
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83
Paramagnetism and Diamagnetism
Unpaired electrons have their spins aligned or – This increases the magnetic field of the atom.
Atoms with unpaired electrons are called paramagnetic .– Paramagnetic atoms are attracted to a magnet.
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84
Paramagnetism and Diamagnetism
Paired electrons have their spins unaligned – Paired electrons have no net magnetic field.
Atoms with paired electrons are called diamagneticdiamagnetic. – Diamagnetic atoms are repelled by a magnet.
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85
Paramagnetism and Diamagnetism
Because two electrons in the same orbital must be paired, it is possible to calculate the number of orbitals and the number of electrons in each n shell.
The number of orbitals per n level is given by n2. The maximum number of electrons per n level is
2n2.– The value is 2n2 because of the two paired electrons.
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86
Paramagnetism and Diamagnetism
Energy Level # of Orbitals Max. # of e-
nn nn22 2n2
1 1 2
2 4 8
You do it!You do it!
3 9 18
4 16 32
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87
The Periodic Table and Electron Configurations
The principle that describes how the periodic chart is a function of electronic configurations is the Aufbau Principle.
The electron that distinguishes an element from the previous element enters the lowest energy atomic orbital available.
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88
The Periodic Table and Electron Configurations
The Aufbau Principle describes the electron filling order in atoms.
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89
The Periodic Table and Electron Configurations
There are two ways to remember the correct filling order for electrons in atoms.1. You can use this mnemonic.
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90
The Periodic Table and Electron Configurations
2. Or you can use the periodic chart .
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91
The Periodic Table and Electron Configurations
Now we will use the Aufbau Principle to determine the electronic configurations of the elements on the periodic chart.
1st row elements.
22
11
1s He
1s H
ionConfigurat 1s
11 1s H
ionConfigurat 1s
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92
The Periodic Table and Electron Configurations
2nd row elements.
•Hund’s rule tells us that the electrons will fill thep orbitals by placing electrons in each orbital singly and with same spin until half-filled. Thenthe electrons will pair to finish the p orbitals.
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93
The Periodic Table and Electron Configurations
3rd row elements 62
18
5217
4216
3215
2214
1213
212
111
3p s3 Ne NeAr
3p s3 Ne Ne Cl
3p s3 Ne Ne S
3p s3 Ne Ne P
3p s3 Ne Ne Si
3p s3 Ne Ne Al
s3 Ne Ne Mg
s3 Ne NeNa
ionConfigurat 3p 3s
52
17
4216
3215
2214
1213
212
111
3p s3 Ne Ne Cl
3p s3 Ne Ne S
3p s3 Ne Ne P
3p s3 Ne Ne Si
3p s3 Ne Ne Al
s3 Ne Ne Mg
s3 Ne Ne Na
ionConfigurat 3p 3s
42
16
3215
2214
1213
212
111
3p s3 Ne Ne S
3p s3 Ne Ne P
3p s3 Ne Ne Si
3p s3 Ne Ne Al
s3 Ne Ne Mg
s3 Ne Ne Na
ionConfigurat 3p 3s
32
15
2214
1213
212
111
3p s3 Ne Ne P
3p s3 Ne Ne Si
3p s3 Ne Ne Al
s3 Ne Ne Mg
s3 Ne Ne Na
ionConfigurat 3p 3s
2
12
111
s3 Ne Ne Mg
s3 Ne Ne Na
ionConfigurat 3p 3s
111 s3 Ne Ne Na
ionConfigurat 3p 3s
22
14
1213
212
111
3p s3 Ne Ne Si
3p s3 Ne Ne Al
s3 Ne Ne Mg
s3 Ne Ne Na
ionConfigurat 3p 3s
12
13
212
111
3p s3 Ne Ne Al
s3 Ne Ne Mg
s3 Ne Ne Na
ionConfigurat 3p 3s
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94
The Periodic Table and Electron Configurations
4th row elements
119 4s Ar ArK
ionConfigurat 4p 4s 3d
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95
The Periodic Table and Electron Configurations
2
20
119
4s Ar ArCa
4s Ar ArK
ionConfigurat 4p 4s 3d
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96
The Periodic Table and Electron Configurations
it! do You Sc
4s Ar ArCa
4s Ar ArK
ionConfigurat 4p 4s 3d
21
220
119
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97
The Periodic Table and Electron Configurations
12
21
220
119
3d 4s Ar Ar Sc
4s Ar ArCa
4s Ar ArK
ionConfigurat 4p 4s 3d
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98
The Periodic Table and Electron Configurations
it! do You Ti
3d 4s Ar Ar Sc
4s Ar ArCa
4s Ar ArK
ionConfigurat 4p 4s 3d
22
1221
220
119
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99
The Periodic Table and Electron Configurations
22
22
1221
220
119
3d 4s Ar Ar Ti
3d 4s Ar Ar Sc
4s Ar ArCa
4s Ar ArK
ionConfigurat 4p 4s 3d
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100
The Periodic Table and Electron Configurations
3223
2222
1221
220
119
3d 4s Ar Ar V
3d 4s Ar Ar Ti
3d 4s Ar Ar Sc
4s Ar ArCa
4s Ar ArK
ionConfigurat 4p 4s 3d
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101
The Periodic Table and Electron Configurations
orbitals. filled completely and filled-half with
associatedstability of measureextra an is There
3d 4s Ar ArCr
3d 4s Ar Ar V
3d 4s Ar Ar Ti
3d 4s Ar Ar Sc
4s Ar ArCa
4s Ar ArK
ionConfigurat 4p 4s 3d
5124
3223
2222
1221
220
119
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102
The Periodic Table and Electron Configurations
5225 3d 4s Ar Ar Mn
ionConfigurat 4p 4s 3d
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103
The Periodic Table and Electron Configurations
it! do You Fe
3d 4s Ar Ar Mn
ionConfigurat 4p 4s 3d
26
5225
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104
The Periodic Table and Electron Configurations
62
26
5225
3d 4s Ar Ar Fe
3d 4s Ar Ar Mn
ionConfigurat 4p 4s 3d
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105
The Periodic Table and Electron Configurations
72
27
6226
5225
3d 4s Ar Ar Co
3d 4s Ar Ar Fe
3d 4s Ar Ar Mn
ionConfigurat 4p 4s 3d
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106
The Periodic Table and Electron Configurations
82
28
7227
6226
5225
3d 4s Ar Ar Ni
3d 4s Ar Ar Co
3d 4s Ar Ar Fe
3d 4s Ar Ar Mn
ionConfigurat 4p 4s 3d
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107
The Periodic Table and Electron Configurations
it! do You Cu
3d 4s Ar Ar Ni
3d 4s Ar Ar Co
3d 4s Ar Ar Fe
3d 4s Ar Ar Mn
ionConfigurat 4p 4s 3d
29
8228
7227
6226
5225
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108
The Periodic Table and Electron Configurations
reason. same y theessentiallfor
andCr like exceptionAnother
3d 4s Ar Ar Cu
3d 4s Ar Ar Ni
3d 4s Ar Ar Co
3d 4s Ar Ar Fe
3d 4s Ar Ar Mn
ionConfigurat 4p 4s 3d
10129
8228
7227
6226
5225
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109
The Periodic Table and Electron Configurations
102
30
10129
8228
7227
6226
5225
3d 4s Ar Ar Zn
3d 4s Ar Ar Cu
3d 4s Ar Ar Ni
3d 4s Ar Ar Co
3d 4s Ar Ar Fe
3d 4s Ar Ar Mn
ionConfigurat 4p 4s 3d
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110
The Periodic Table and Electron Configurations
110231 4p 3d 4s Ar ArGa
ionConfigurat 4p 4s 3d
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111
The Periodic Table and Electron Configurations
it! do You Ge
4p 3d 4s Ar ArGa
ionConfigurat 4p 4s 3d
32
110231
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112
The Periodic Table and Electron Configurations
2102
32
110231
4p 3d 4s Ar Ar Ge
4p 3d 4s Ar ArGa
ionConfigurat 4p 4s 3d
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113
The Periodic Table and Electron Configurations
3102
33
210232
110231
4p 3d 4s Ar Ar As
4p 3d 4s Ar Ar Ge
4p 3d 4s Ar ArGa
ionConfigurat 4p 4s 3d
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114
The Periodic Table and Electron Configurations
it! do You Se
4p 3d 4s Ar Ar As
4p 3d 4s Ar Ar Ge
4p 3d 4s Ar ArGa
ionConfigurat 4p 4s 3d
34
310233
210232
110231
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115
The Periodic Table and Electron Configurations
4102
34
310233
210232
110231
4p 3d 4s Ar Ar Se
4p 3d 4s Ar Ar As
4p 3d 4s Ar Ar Ge
4p 3d 4s Ar ArGa
ionConfigurat 4p 4s 3d
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116
The Periodic Table and Electron Configurations
5102
35
410234
310233
210232
110231
4p 3d 4s Ar ArBr
4p 3d 4s Ar Ar Se
4p 3d 4s Ar Ar As
4p 3d 4s Ar Ar Ge
4p 3d 4s Ar ArGa
ionConfigurat 4p 4s 3d
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117
The Periodic Table and Electron Configurations
6102
36
510235
410234
310233
210232
110231
4p 3d 4s Ar ArKr
4p 3d 4s Ar ArBr
4p 3d 4s Ar Ar Se
4p 3d 4s Ar Ar As
4p 3d 4s Ar Ar Ge
4p 3d 4s Ar ArGa
ionConfigurat 4p 4s 3d
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The Periodic Table and Electron Configurations
Now we can write a complete set of quantum numbers for all of the electrons in these three elements as examples.
– Na– Ca– Fe
First for 11Na.– When completed there must be one set of 4 quantum numbers for
each of the 11 electrons in Na(remember Ne has 10 electrons)
111 s3 Ne NeNa
ionConfigurat 3p 3s
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119
The Periodic Table and Electron Configurations
1/2 0 0 1 e 1
m m n -st
s
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120
The Periodic Table and Electron Configurations
electrons s 11/2 0 0 1 e 2
1/2 0 0 1 e 1
m m n
-nd
-st
s
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The Periodic Table and Electron Configurations
1/2 0 0 2 e 3
electrons s 11/2 0 0 1 e 2
1/2 0 0 1 e 1
m m n
-rd
-nd
-st
s
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122
The Periodic Table and Electron Configurations
electrons s 21/2 0 0 2 e 4
1/2 0 0 2 e 3
electrons s 11/2 0 0 1 e 2
1/2 0 0 1 e 1
m m n
-th
-rd
-nd
-st
s
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123
The Periodic Table and Electron Configurations
1/2 1- 1 2 e 5
electrons s 21/2 0 0 2 e 4
1/2 0 0 2 e 3
electrons s 11/2 0 0 1 e 2
1/2 0 0 1 e 1
m m n
-th
-th
-rd
-nd
-st
s
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124
The Periodic Table and Electron Configurations
1/2 0 1 2 e 6
1/2 1- 1 2 e 5
electrons s 21/2 0 0 2 e 4
1/2 0 0 2 e 3
electrons s 11/2 0 0 1 e 2
1/2 0 0 1 e 1
m m n
-th
-th
-th
-rd
-nd
-st
s
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125
The Periodic Table and Electron Configurations
1/2 1 1 2 e 7
1/2 0 1 2 e 6
1/2 1- 1 2 e 5
electrons s 21/2 0 0 2 e 4
1/2 0 0 2 e 3
electrons s 11/2- 0 0 1 e 2
1/2 0 0 1 e 1
m m n
-th
-th
-th
-th
-rd
-nd
-st
s
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126
The Periodic Table and Electron Configurations
1/2 1 1 2 e 8
1/2 1 1 2 e 7
1/2 0 1 2 e 6
1/2 1- 1 2 e 5
electrons s 21/2 0 0 2 e 4
1/2 0 0 2 e 3
electrons s 11/2 0 0 1 e 2
1/2 0 0 1 e 1
m m n
-th
-th
-th
-th
-th
-rd
-nd
-st
s
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127
The Periodic Table and Electron Configurations
1/2 0 1 2 e 9
1/2 1 1 2 e 8
1/2 1 1 2 e 7
1/2 0 1 2 e 6
1/2 1- 1 2 e 5
electrons s 21/2 0 0 2 e 4
1/2 0 0 2 e 3
electrons s 11/2 0 0 1 e 2
1/2 0 0 1 e 1
m m n
-th
-th
-th
-th
-th
-th
-rd
-nd
-st
s
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128
The Periodic Table and Electron Configurations
electrons p 2
1/2 1 1 2 e 10
1/2 0 1 2 e 9
1/2 1 1 2 e 8
1/2 1 1 2 e 7
1/2 0 1 2 e 6
1/2 1- 1 2 e 5
electrons s 21/2 0 0 2 e 4
1/2 0 0 2 e 3
electrons s 11/2 0 0 1 e 2
1/2 0 0 1 e 1
m m n
-th
-th
-th
-th
-th
-th
-th
-rd
-nd
-st
s
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129
The Periodic Table and Electron Configurations
electron s 31/2 0 0 3 e 11
electrons p 2
1/2 1 1 2 e 10
1/2 0 1 2 e 9
1/2 1 1 2 e 8
1/2 1 1 2 e 7
1/2 0 1 2 e 6
1/2 1- 1 2 e 5
electrons s 21/2 0 0 2 e 4
1/2 0 0 2 e 3
electrons s 11/2 0 0 1 e 2
1/2 0 0 1 e 1
m m n
-th
-th
-th
-th
-th
-th
-th
-th
-rd
-nd
-st
s
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130
The Periodic Table and Electron Configurations
Next we will do the same exercise for 20Ca.– Again, when finished we must have one set of 4 quantum numbers for each of the 20 electrons
in Ca. We represent the first 18 electrons in Ca with the symbol [Ar].
220 4s Ar [Ar]Ca
ionConfigurat 4p 4s 3d
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131
The Periodic Table and Electron Configurations
1/2 0 0 4 e 19 ]Ar[
m m n
-th
s
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132
The Periodic Table and Electron Configurations
electrons s 41/2 0 0 4 e 20
1/2 0 0 4 e 19]Ar[
m m n
-th
-th
s
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133
The Periodic Table and Electron Configurations
Finally, we do the same exercise for 26Fe.– We should have one set of 4 quantum numbers for each of the 26 electrons in Fe.
To save time and space, we use the symbol [Ar] to represent the first 18 electrons in Fe
6226 3d 4s Ar Ar Fe
ionConfigurat 4p 4s 3d
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134
The Periodic Table and Electron Configurations
1/2 0 0 4 e 19 ]Ar[
m m n
-th
s
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135
The Periodic Table and Electron Configurations
electrons s 41/2 0 0 4 e 20
1/2 0 0 4 e 19]Ar[
m m n
-th
-th
s
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136
The Periodic Table and Electron Configurations
1/2 2- 2 3 e 21
electrons s 41/2 0 0 4 e 20
1/2 0 0 4 e 19 ]Ar[
m m n
-st
-th
-th
s
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The Periodic Table and Electron Configurations
it! doYou e 22
1/2 2- 2 3 e 21
electrons s 41/2 0 0 4 e 20
1/2 0 0 4 e 19 ]Ar[
m m n
-nd
-st
-th
-th
s
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138
The Periodic Table and Electron Configurations
1/2 1- 2 3 e 22
1/2 2- 2 3 e 21
electrons s 41/2 0 0 4 e 20
1/2 0 0 4 e 19 [Ar]
m m n
-nd
-st
-th
-th
s
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139
The Periodic Table and Electron Configurations
1/2 0 2 3 e 23
1/2 1- 2 3 e 22
1/2 2- 2 3 e 21
electrons s 41/2 0 0 4 e 20
1/2 0 0 4 e 19 [Ar]
m m n
-rd
-nd
-st
-th
-th
s
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140
The Periodic Table and Electron Configurations
1/2 1 2 3 e 24
1/2 0 2 3 e 23
1/2 1- 2 3 e 22
1/2 2- 2 3 e 21
electrons s 41/2 0 0 4 e 20
1/2 0 0 4 e 19 [Ar]
m m n
-th
-rd
-nd
-st
-th
-th
s
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The Periodic Table and Electron Configurations
shell d filled-half
1/2 2 2 3 e 25
1/2 1 2 3 e 24
1/2 0 2 3 e 23
1/2 1- 2 3 e 22
1/2 2- 2 3 e 21
electrons s 41/2 0 0 4 e 20
1/2 0 0 4 e 19 [Ar]
m m n
-th
-th
-rd
-nd
-st
-th
-th
s
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The Periodic Table and Electron Configurations
it! doYou e 26
1/2 2 2 3 e 25
1/2 1 2 3 e 24
1/2 0 2 3 e 23
1/2 1- 2 3 e 22
1/2 2- 2 3 e 21
electrons s 41/2 0 0 4 e 20
1/2 0 0 4 e 19 [Ar]
m m n
-th
-th
-th
-rd
-nd
-st
-th
-th
s
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143
The Periodic Table and Electron Configurations
1/2 2- 2 3 e 26
1/2 2 2 3 e 25
1/2 1 2 3 e 24
1/2 0 2 3 e 23
1/2 1- 2 3 e 22
1/2 2- 2 3 e 21
electrons s 41/2 0 0 4 e 20
1/2 0 0 4 e 19 [Ar]
m m n
-th
-th
-th
-rd
-nd
-st
-th
-th
s