Download Properties of atoms result from electron configuration

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Lecture 6: Sub-­‐atomic & quantum structure Lecture 6 Topics: Brown, chapter 6 & 7 1. Atomic properties from e-­‐ configuration
2. The true nature of the atom? •  Light (and electrons) behave as waves & particles 3. Developing a new physics for atoms
•  A quick tour of quantum mechanics 4. Bohr’s quantum planetary model
5. Applying quantum mechanics to the atom •  Electrons inhabit orbitals 7.1 6. Orbital filling and electron configuration •  AuQau & orbital diagrams 6.8 – 6.9 6.1 6.2 – 6.4 6.2 6.6 – 6.7 Properties of atoms result from electron configuration Elements in each column have the same configuration. 1
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Atomic properties & electron configuration The physical & chemical properties of each element are a result of the arrangement of the electrons of an atom around its nucleus. •  Number of electrons dictates how they are “configured” around the atoms nucleus. •  The configuration of electrons is dictated by their energy levels. • The outermost shell of electrons (the valance shell) bonds to other atoms. Columns: atoms with the same valance shell configuration •  Similar physical & chemical properties. Rows (periods): atoms with the same number of electron shells First pages of Ch 7 The beauty of the periodic table Mendeleev, creator of the table, predicted the nature of elements not yet discovered by relating their properties to those of adjacent and well characterized elements. First pages of Ch 7 2
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p.217-­‐ Examples: periodic table knowledge Which have the most similar physical & chemical characteristics? B, Ca, F, He, Mg, P •  Of the elements listed only Ca & Mg are found in the same column. •  Elements found in the same column have the most similar properties. Find sodium & bromine in the table. a) Give the atomic number of each. b) Label each as a metal, nonmetal, or metalloid. •  Na is in column IA; number 11 & a metal. •  Br is in column VIIA; number 35 & a nonmetal. Metals West Coast -­‐ luster, conduct heat & electricity, solids (except Hg) Metalloids Staircase that separates metals & nons. Properties lie between metals & nons. Nonmetals Great Lakes & East Coast -­‐ Properties are non-­‐metallic. Can be gases, liquids or solids at room temp. First pages of Ch 7 Lecture 6: Sub-­‐atomic & quantum structure Lecture 6 Topics: Brown, chapter 6 & 7 1. Atomic properties from e-­‐ configuration 2. The true nature of the atom? •  Light (and electrons) behave as waves & particles 3. Developing a new physics for atoms
•  A quick tour of quantum mechanics 4. Bohr’s quantum planetary model
5. Applying quantum mechanics to the atom •  Electrons inhabit orbitals 7.1 6. Orbitals filling and electron configuration •  AuQau & orbital diagrams 6.8 – 6.9 6.1 6.2 – 6.4 6.2 6.6 – 6.7 3
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The true nature of the atom? Failure of the planetary model Newtonian physics cannot describe atomic behavior. Electrons, like light, are ‘waveicles’. Back to the atom! We left off at the planetary model of the atom developed by Hantaro Nagaoka in 1903 and was considered in Rutherford’s nuclear model paper. But the planetary model already had problems: •  Electrons orbiting the nucleus would be accelerating and therefore emitting energy like radio signals; & •  As energy is emitted it is lost, and the electron would spiral in and crash. the nuclear
“death spiral”
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Newtonian physics just doesn’t cut it The behavior of our everyday world can be described by classical, Newtonian, physics. However, at the end of the 1800s it was clear that Newtonian physics didn’t accurately describe the behavior of light and matter at the atomic scale. For example: Why atoms don’t collapse? Give that some thought… Three phenomena that could not be understood with Newtonian physics were investigated and explained by radical new ideas that would lead to a new type of physics, the field that would become quantum mechanics. Blackbody radiation
Max Planck
1900 Photoelectric effect
Albert Einsetein 1905 Niels Bohr
1913 Line spectra of elements Blackbody radiation Hot objects emit energy as both heat and color. The color of hot objects correlates with their temperature. • 
In turn of the century Germany, this was of great economic interest, as the newly unified country wanted to dominate the lighting industry. They wanted to find the best filament for light bulbs. The relationship between color & temperature was not continuous. In 1900, Max Planck found that the energy emitted or radiated from the object was in discrete energy bundles that he called quanta. E = nhν Notice that all values of E are multiples of E = energy (J) Planck’s constant and are therefore not n = integer h = Planck’s constant (J-­‐s) continuous, but discrete.
ν = frequency (1/s) 5
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Photoelectric effect In 1839 it was observed that when visible or uv light is aimed at a clean metal surface, electrons are ejected from the metal, causing an electrical current. Unexpectedly, a light energy threshold was required before electrons would be emitted. In 1905, Albert Einstein applied Planck’s idea of quanta to explain the effect. •  Einstein called these quantum bundles of light photons. •  Einstein’s work bolstered the quantum concept. E = hν E = energy (J) h = Planck’s constant (J-­‐s) ν = frequency (1/s) Elemental line spectra When gases are heated to incandescence, they emit light at a series of characteristic wavelengths due to the atomic composition of the gas. These ‘line spectra’ are discrete rather than continuous. Why? In 1913, Niels Bohr built on the work of Planck and Einstein and hypothesized that atomic energies are discrete rather than continuous. •  The lines of elemental spectra correspond to the possible energy states of the atom’s electrons. So Bohr applied the concept of quanta to matter and to the atom in particular. 6
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Line spectra of deep winter p.217-­‐ The aurora borealis occurs when solar wind and magnetospheric plasma lose energy, hit Earth’s atmosphere and excite atmospheric molecules. The energy lost upon return to ground state produces the colors of the aurora. windows2universe.org Yea, but is this quantum stuff really relevant? Does all this quantum stuff really matter outside of the classroom and lab? Yep, we’re about to look at the development of the quantum mechanical behavior of the atom, a theoretical framework that expresses the behavior of matter at the atomic scale. Quantum mechanics now plays critical roles in science, medicine, engineering and technology. This model has correctly predicted the behavior of matter at the atomic level across disciplines. It’s estimated that quantum mechanics is involved in 30% of the US gross domestic product. 7
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Lecture 6: Sub-­‐atomic & quantum structure Lecture 6 Topics: Brown, chapter 6 & 7 1. Atomic properties from e-­‐ configuration 2. The true nature of the atom? •  Light (and electrons) behave as waves & particles 3. Developing a new physics for atoms •  A quick tour of quantum mechanics 4. Bohr’s quantum planetary model
5. Applying quantum mechanics to the atom •  Electrons inhabit orbitals 7.1 6. Orbital filling and electron configuration •  AuQau & orbital diagrams 6.8 – 6.9 6.1 6.2 – 6.4 6.2 6.6 – 6.7 Developing a new physics Particles and waves? Light behaves as waves & particles Electrons do too! Best proof of the dualistic nature of atoms 8
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Developing a new quantum physics In fact, creating a model of the atom consistent with line spectrum data would require a more complete understanding of the implications of quantum energy. Through the 1920s, the work of a number of physicists fleshed out and developed the quantum mechanical model of the atom. In addition to the concept of quantized energy, a number of other ideas were critical to this work: •  Light behaves as a wave. Landé suggests the •  Light behaves as a particle. term ‘waveicle’. •  Matter can behave as a wave. •  Wave functions can describe the energy levels of atoms. •  At the atomic level two properties of a particle cannot be determined exactly at the same time: uncertainty (undeterminant). Let’s start with some clear definitions Particle •  A minute portion of matter. •  The least possible amount. •  Mass •  Size (diameter) Wave •  An oscillation that moves outward from a disturbance. •  A periodic disturbance of the particles of a substance that may be propagated without net movement of the particles, such as in the passage of undulating motion, heat, or sound. •  A variation of an electromagnetic field in the propagation of light or other radiation through a medium or vacuum. •  Wavelength (λ) distance from peak to peak or trough to trough (nm) •  Frequency (ν) the number of peaks that pass a point over time (1/s) •  Amplitude the height of the wave (nm) 9
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Light behaves as a wave In 1801,Thomas Young’s double-­‐slit experiment demonstrated that light is a wave. •  Waves of light passing through the two slits meet and interfere. •  Where waves meet out of phase, addition of waves of opposite signs yields little light. •  Where waves meet in-­‐phase, addition produces strong light. •  Stripes of strong light are separated by darker stripes of little light. Light behaves as particles In 1905, Einstein’s explanation of the photoelectric effect showed that light could behave as a particle, a photon. In this view, photons act as ‘bullets’, hitting the surface of the metal and knocking electrons free. http://www.daviddarling.info/images/photoelectric_effect.jpg
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First image of light as both wave & particle In 2015, physicists at EPFL (École polytechnique fédérale de Lausanne) in Switzerland were the first to ever get photographic evidence of light behaving both as a particle & a wave. The video linked below describes their experiment. Video of experimental design: http://phys.org/news/2015-­‐03-­‐particle.html Electrons behave like waves In 1924, Louis de Broglie showed that electrons (a form of matter) could behave like waves. He described e-­‐ wavelength as: λ = h/mν This work implied that the behavior of electrons in an atom could be described mathematically by a wave function (equation). De Broglie shot a beam of electrons at a metal surface and observed a diffraction pattern typically produced by waves. Examples of electron diffraction patterns from graphene & variants. 11
Wave functions describe atomic energies In 1926, Erwin Schrödinger showed that the energy levels of electrons in atoms – shown by elemental line spectra -­‐ could be described by mathematical wave functions, confirming the wave nature of . electrons. -­‐ δ
ih ψ = Ĥψ δt Schrödinger’s work confirmed the wave behavior of electrons and allowed scientists to calculate, or predict, the likely location of an atom’s electrons. The square of the wave function, ψ2, represents the probability that the e-­‐ will be found at that location: probability density or electron density. This concept of electron density plots led to the development of electron orbitals. Schrödinger is sometimes called the ‘father of quantum mechanics’. Orbital of H atom imaged by a quantum microscope (Stodolna, 2013)
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Particle behavior is also uncertain In 1927, Werner Heisenberg found that pairs of properties of particles cannot have exact values at the same time when the particles are of subatomic scale. He called this ‘indeterminant’ behavior. For example, if you know a subatomic particle’s speed or momentum exactly, you cannot know its location exactly at the same moment. •  So, the location and movement of electrons cannot be known precisely at any particular moment. •  Heisenberg’s work supported that of Schrödinger. We don’t see this uncertainty in our “macro” world because the size of the uncertainty is many, many times smaller than the size of the object. •  Uncertainty ‘works’ in the subatomic world because the size of the uncertainty and the size of the object are similar. 12
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Best proof of the electron’s dualist nature In 1965, Richard Feynman proposed a thought experiment: repeat Young’s double-­‐slit experiment with electrons. A number of groups have now done it! •  Pozzi et al. (1974) •  Tonomura et al. (1982) •  Batelaan et al. (2013) Let’s look at Dr. Quantum’s great video recap! https://www.youtube.com/watch?
v=DfPeprQ7oGc Roger Bach, Damian Pope, Sy-­‐Hwang Liou, Herman Batelaan. Controlled double-­‐slit electron diffraction. New Journal of Physics, 2013; 15 (3): 033018 DOI: 10.1088/1367-­‐2630/15/3/03301 So, why don’t atoms collapse? In our macro world atoms would collapse when the negatively charged electrons were pulled into the nucleus by electrostatic attraction. Unless we constantly added energy to the electrons, their velocity would slow, and they would spiral into the nucleus and collapse the atom. But, in the subatomic quantum world, atoms are subject to quantum mechanics rather than the Newtonian laws of motion. •  Electrons have quantum energy and that level of energy can only be changed transiently. •  Uncertainty says that if an electron did approach the nucleus the uncertainty of its location would be greatly decreased. To compensate, the uncertainty of its speed or momentum would have to increase. That increased speed would push the electron away from the nucleus. 13
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I know where we stand right now, Dr. Heisenberg, but where are we going? Lecture 6: Sub-­‐atomic & quantum structure Lecture 6 Topics: Brown, chapter 6 & 7 1. Atomic properties from e-­‐ configuration 2. The true nature of the atom? •  Light (and electrons) behave as waves & particles 3. Developing a new physics for atoms
•  A quick tour of quantum mechanics 4. Bohr’s quantum planetary model
5. Applying quantum mechanics to the atom •  Electrons inhabit orbitals 7.1 6. Orbital filling and electron configuration •  AuQau & orbital diagrams 6.8 – 6.9 6.1 6.2 – 6.4 6.2 6.6 – 6.7 14
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Bohr’s quantum planetary model Electrons are configured in shells & subshells. Failure… and the need for better application of QM Bohr’s quantum planetary model In 1913, Bohr applied what he’d learned about the quantum nature of elemental line spectra to the atom. Bohr’s three postulates: can occupy only very specific energy levels (quantum levels) & 1. Electrons aren’t found between these levels. 2. Electrons staying in “allowed” quantum energy levels do not radiate energy. (Thus, they don’t lose or change energy levels.) 3. If electrons change energy levels (move from one quantum level to another) they radiate energy in the form of photons. e-­‐ 3 e-­‐ 2 1 e-­‐ p/n p.216-­‐222 15
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Bohr’s success was not complete Bohr’s largest triumph was that his quantum planetary model explained the existence of elemental line spectra. •  Spectral lines corresponded to the planetary rings of the electrons. •  And were caused by those fixed, discrete (aka quantum) energy levels. Alas, this rescue of the planetary model was only temporary because: •  The model worked only for atoms with one electron (H). •  Doesn’t explain why electrons don’t lose energy. •  Essentially, Bohr viewed electrons as particles. Lecture 6: Sub-­‐atomic & quantum structure Lecture 6 Topics: Brown, chapter 6 & 7 1. Atomic properties from e-­‐ configuration 2. The true nature of the atom? •  Light (and electrons) behave as waves & particles 3. Developing a new physics for atoms
•  A quick tour of quantum mechanics 4. Bohr’s quantum planetary model
5. Applying quantum mechanics to the atom •  Electrons inhabit orbitals 7.1 6. Orbital filling and electron configuration •  AuQau & orbital diagrams 6.8 – 6.9 6.1 6.2 – 6.4 6.2 6.6 – 6.7 16
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Applying quantum mechanics to the atom Electrons inhabit orbitals. Orbitals are 3D shapes. Orbitals have energy levels. Orbitals have orientation. Electrons have spin. Shells with orbitals A simple modification of Bohr’s model allows us to see one effect of orbitals on the configuration of electrons around the nucleus. p.228-­‐ 3d Shells Subshells
#e-­‐
e-­‐ pairs 1
s
2
1 3p 2
s
2
1 2p 3s p
6
3 1s p/n 2s 3 
s
p
d
2
6
10
1 3 5 Notice that each subshell (or orbital) can only hold a max of 2 electrons. How many electrons in an atom? Atomic # Which is the valance shell? Outermost; here 3rd shell What causes chemical bonding? Overlapping, or interaction, of electrons in valence shells of 2 atoms. 17
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p.232-­‐ Orbitals: electron density plots Since orbitals are wave functions, they are also probabilities. The energy and shape of an orbital describe the space in which you will most likely find the electron at any one time; 90% of the time. AKA electron density map Imagine drawing a probability map of your whereabouts for a given day of the week. Imagine drawing a probability map of your whereabouts for a given day of the week. Sure you’re going to take unexpected trips, but for the most part it’s pretty simple to predict where you’ll be. Imagine carrying a transmitter all day. If it sent a signal every minute or two you’d have a pretty good dot plot or “probability map”. p.224-­‐ p.225-­‐ Orbitals: density dictates shape Of course, the shape of the space occupied by the electrons is shown by the probability plot as seen here: Here probability is shown on the y-­‐axis and radial distance from the nucleus for s orbitals in shells 1, 2 & 3. Notice that the bulk of electrons are located in the same type of ‘shape’, but further from the nucleus in each successive shell. Total densities are the same, but distributed more broadly in 2 & 3. 1s 2s 3s p.224-­‐32 18
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Orbital shapes: s orbital Remember that each orbital can contain, at most, 2 electrons. So the single s orbital holds 2. But there are 3 p orbitals, that can hold a total of 6 electrons. And the 5 d orbitals can hold a total of 10 electrons. p orbital d orbital p.228-­‐32 Orbital energy levels The energy levels of orbitals increase in a straightforward manner until we reach the junction of the 3rd & 4th shells. This table shows which orbitals are filled for each column of elements. 1) electropositive metals (IA, IIA) 2) transition metals (d orbitals) 3) non-­‐metals (& metalloids) (other As) p.232-­‐4, 240 19
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p.242-­‐ Electron spin Pauli’s exclusion principle states that each orbital can contain no more than two electrons. And if two electrons occupy a single orbital they must have opposite spin. Stern-­‐ Gerlach exp’t This gives rise to the intrinsic spin of some molecules – particularly those with a single valence electron. This allows NMR & MRI.
And so we use “up” & “down” arrows when describing configuration. p.234 Lecture 6: Sub-­‐atomic & quantum structure Lecture 6 Topics: Brown, chapter 6 & 7 1. Atomic properties from e-­‐ configuration 2. The true nature of the atom? •  Light (and electrons) behave as waves & particles 3. Developing a new physics for atoms
•  A quick tour of quantum mechanics 4. Bohr’s quantum planetary model
5. Applying quantum mechanics to the atom •  Electrons inhabit orbitals 7.1 6. Orbital filling and electron configuration •  Aujau & orbital diagrams 6.8 – 6.9 6.1 6.2 – 6.4 6.2 6.6 – 6.7 20
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Orbital filling & electron configuration Hund’s rule: electrons fill one at a time Pauli’s exclusion principle: opposite spin Configuration & the periodic table Orbital diagrams p.235 Notice that the orbital ‘boxes’ fill one electron (e-­‐) at a time. What property of e-­‐ explains their desire to be ‘alone’? Negative charge -­‐> repulsion. Notice that the two e-­‐ occupying an orbital have opposite ‘spin’. Electron spin creates magnetic fields, minimized by opposite spin. Two e-­‐ with opposite spin are ‘paired’. 21
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Orbital filling: Aujau H Li Na K Rb Cs Fr 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10 7p6 He Ne Ar Kr Xe Rn ?? p.233 Examples: orbital filling Draw the orbital diagram for oxygen. O is element 8: 1s22s22p4 Draw the electron configuration of phosphorous. Atomic number 15: 1s22s22p63s23p3 Notice that three of the electrons are unpaired. What valence electron configuration is common to all halogens? Halogens are the elements found in column VIIA. Their valence (or outermost) electron shell has 7 ve-­‐ (column number). Let’s try them:
F #9
1s22s22p5 Cl #17 1s22s22p63s23p5 Br #35 1s22s22p63s23p64s23d104p5 I #53
1s22s22p63s23p64s23d104p65s24d105p5 So the common valence configuration is p5. Soon, we’ll see that all halogen ions have a charge of -­‐1 to fill that p orbital. p.235-­‐ 22
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Electron configuration & the periodic table p.242 Lecture 6: Sub-­‐atomic & quantum structure Lecture 6 Topics: Brown, chapter 6 & 7 1. Atomic properties from e-­‐ configuration 2. The true nature of the atom? •  Light (and electrons) behave as waves & particles 3. Developing a new physics for atoms
•  A quick tour of quantum mechanics 4. Bohr’s quantum planetary model
5. Applying quantum mechanics to the atom •  Electrons inhabit orbitals 7.1 6. Orbitals filling and electron configuration •  AuQau & orbital diagrams 6.8 – 6.9 6.1 6.2 – 6.4 6.2 6.6 – 6.7 23
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demonstrates wave nature demonstrates particle nature Maxwell’s electromagnetic theory also subject t0 intensity as function of λ spectra thermal continuous spectra reveal temperature particles when mass is zero atomic or more mass photoelectric effect subatomic classical leads to mass (e-­‐) (Newtonian) Einstein’s mechanics theory of quantum mechanics photons behavior emit/absorb in governed by refracted in prism Interference effects EM spectrum light atomic line spectra reveal allowed quantum energy states of atoms uncertainty principle whose consequence in atoms is that e-­‐ can be located only as probability functions (orbitals) Stephen Lower http://www.chem1.com/acad/webtext/atoms/atpt-­‐2.html Concept map for quantum nature of electrons Lecture 6: Terms to Know •  Periodic table – atomic properties •  Column •  Row •  Electron configuration •  Bohr’s model of the electron •  Quantum •  Wave vs. particle •  Principle •  Azimuthal •  Magnetic •  Shell •  Subshell •  Valence shell •  Orbital •  Electron density plot •  Orbital shapes: s, p, d •  Orbital energy levels •  Electron spin •  Pauli’s exclusion principle •  Orbital diagrams •  Hund’s rule •  Order of filling •  Electron configuration •  Abbreviated configuration 24