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NAME __________________________________ UNIT 6 (1): QUANTUM MECHANICAL MODEL Big Idea 1: The chemical elements are fundamental building materials of matter, and, all matter can be understood in terms of arrangements of atoms. These atoms retain their identity in chemical reactions. Enduring understanding 1.A: All matter is Essential knowledge 1.A.1: Molecules are composed of made of atoms. There are a limited number specific combinations of atoms; different molecules are of types of atoms; these are the elements. composed of combinations of different elements and of combinations of the same elements in differing amounts and proportions. Essential knowledge 1.A.2: Chemical analysis provides a method for determining the relative number of atoms in a substance, which can be used to identify the substance or determine its purity. Essential knowledge 1.A.3: The mole is the fundamental unit for counting numbers of particles on the macroscopic level and allows quantitative connections to be drawn between laboratory experiments, which occur at the macroscopic level, and chemical processes, which occur at the atomic level. Enduring understanding 1.B: The atoms of each element have unique structures arising from interactions between electrons and nuclei. Essential knowledge 1.B.1: The atom is composed of negatively charged electrons, which can leave the atom, and a positively charged nucleus that is made of protons and neutrons. The attraction of the electrons to the nucleus is the basis of the structure of the atom. Coulomb’s law is qualitatively useful for understanding the structure of the atom. Essential knowledge 1.B.2: The electronic structure of the atom can be described using an electron configuration that reflects the concept of electrons in quantized energy levels or shells; the energetics of the electrons in the atom can be understood by consideration of Coulomb’s law. Enduring understanding 1.C: Elements display periodicity in their properties when the elements are organized according to increasing atomic number. This periodicity can be explained by the regular variations that occur in the electronic structures of atoms. Periodicity is a useful principle for understanding properties and predicting trends in properties. Its modern-day uses range from examining the composition of materials to generating ideas for designing new materials. Essential knowledge 1.C.1: Many properties of atoms exhibit periodic trends that are reflective of the periodicity of electronic structure. Essential knowledge 1.C.2: The currently accepted best model of the atom is based on the quantum mechanical model. 506 Enduring understanding 1.D: Atoms are so small that they are difficult to study directly; atomic models are constructed to explain experimental data on collections of atoms. Essential knowledge 1.D.1: As is the case with all scientific models, any model of the atom is subject to refinement and change in response to new experimental results. In that sense, an atomic model is not regarded as an exact description of the atom, but rather a theoretical construct that fits a set of experimental data. Essential knowledge 1.D.2: An early model of the atom stated that all atoms of an element are identical. Mass spectrometry data demonstrate evidence that contradicts this early model….e.g. ISOTOPES!!!! Essential knowledge 1.D.3: The interaction of electromagnetic waves or light with matter is a powerful means to probe the structure of atoms and molecules, and to measure their concentration. e.g. PES and Spectrophotometry Enduring understanding 1.E: Atoms are Essential knowledge 1.E.1: Physical and chemical processes can be conserved in physical and chemical processes. depicted symbolically; when this is done, the illustration must conserve all atoms of all types. Essential knowledge 1.E.2: Conservation of atoms makes it possible to compute the masses of substances involved in physical and chemical processes. Chemical processes result in the formation of new substances, and the amount of these depends on the number and the types and masses of elements in the reactants, as well as the efficiency of the transformation. Learning objective 1.1 The student can justify the observation that the ratio of the masses of the constituent elements in any pure sample of that compound is always identical on the basis of the atomic molecular theory. [See SP 6.1; Essential knowledge 1.A.1] Learning objective 1.2 The student is able to select and apply mathematical routines to mass data to identify or infer the composition of pure substances and/or mixtures. [See SP 2.2; Essential knowledge 1.A.2] Learning objective 1.3 The student is able to select and apply mathematical relationships to mass data in order to justify a claim regarding the identity and/or estimated purity of a substance. [See SP 2.2, 6.1; Essential knowledge 1.A.2] Learning objective 1.4 The student is able to connect the number of particles, moles, mass, and volume of substances to one another, both qualitatively and quantitatively. [See SP 7.1; Essential knowledge 1.A.3] Learning objective 1.5 The student is able to explain the distribution of electrons in an atom or ion based upon data. [See SP 1.5, 6.2; Essential knowledge 1.B.1] Learning objective 1.6 The student is able to analyze data relating to electron energies for patterns and relationships. [See SP 5.1; Essential knowledge 1.B.1] Learning objective 1.7 The student is able to describe the electronic structure of the atom, using PES data, ionization energy data, and/or Coulomb’s law to construct explanations of how the energies of electrons within shells in atoms vary. [SP 5.1, 6.2; Essential knowledge 1.B.2] Learning objective 1.8 The student is able to explain the distribution of electrons using Coulomb’s law to analyze measured energies. [See SP 6.2; Essential knowledge 1.B.2] Learning objective 1.9 The student is able to predict and/or justify trends in atomic properties based on location on the periodic table and/or the shell model. [See SP 6.4; Essential knowledge 1.C.1] Learning objective 1.10 Students can justify with evidence the arrangement of the periodic table and can apply periodic properties to chemical reactivity. [See SP 6.1; Essential knowledge 1.C.1] Learning objective 1.11 The student can analyze data, based on periodicity and the properties of binary compounds, to identify patterns and generate hypotheses related to the molecular design of compounds for which data are not supplied. [SP 3.1, 5.1; Essential knowledge 1.C.1] 507 Learning objective 1.12 The student is able to explain why a given set of data suggests, or does not suggest, the need to refine the atomic model from a classical shell model with the quantum mechanical model. [See SP 6.3; Essential knowledge 1.C.2] Learning objective 1.13 Given information about a particular model of the atom, the student is able to determine if the model is consistent with specified evidence. [See SP 5.3; Essential knowledge 1.D.1] Learning objective 1.14 The student is able to use data from mass spectrometry to identify the elements and the masses of individual atoms of a specific element. [See SP 1.4, 1.5; Essential knowledge 1.D.2] Learning objective 1.15 The student can justify the selection of a particular type of spectroscopy to measure properties associated with vibrational or electronic motions of molecules. [See SP 4.1, 6.4; Essential knowledge 1.D.3] Learning objective 1.16 The student can design and/or interpret the results of an experiment regarding the absorption of light to determine the concentration of an absorbing species in a solution. [See SP 4.2, 5.1; Essential knowledge 1.D.3] Learning objective 1.17 The student is able to express the law of conservation of mass quantitatively and qualitatively using symbolic representations and particulate drawings. [See SP 1.5; Essential knowledge 1.E.1] Learning objective 1.18 The student is able to apply conservation of atoms to the rearrangement of atoms in various processes. [See SP 1.4; Essential knowledge 1.E.2] Learning objective 1.19 The student can design, and/or interpret data from, an experiment that uses gravimetric analysis to determine the concentration of an analyte in a solution. [See SP 4.2, 5.1, 6.4; Essential knowledge 1.E.2] Learning objective 1.20 The student can design, and/or interpret data from, an experiment that uses titration to determine the concentration of an analyte in a solution. [See SP 4.2, 5.1, 6.4; Essential knowledge 1.E.2] Review and Diagnostic Work …. http://www.bozemanscience.com/ap-chem-007-quantum-mechanical-model Quantum Mechanics is essentially a theory developed to explain or to predict, as best as possible, the behavior of light and atoms. Throughout this review section our goals will be, in part, to understand the meaning and implications of a ground state configuration such as that for an atom of iron (Fe): 1s2 2s2 2p6 3s2 3p6 3d6 4s2 OR [Ar] 3d6 4s2 First, study the following to review what is meant by principal energy level, sublevel, and orbital: Superscripts indicate the number of e- occupying the orbitals of the specific sublevel 1s22s22p5 Large integers indicate the PEL (n) Recall that every orbital can hold a maximum of 2 e-, in opposite “spins”. •Every s sublevel has 1 orbital. •Every p sublevel has 3 (*degenerate) orbitals •Every d sublevel has 5 (degenerate) orbitals •Every f sublevel has 7 (degenerate) orbitals the small-case letters represent sublevels that are divided into orbitals, with e- The Aufbau Diagram (listed here up to 5f)., helps to place (configure) electrons, but does NOT help to interpret the “why” of the configuration 1s 2s 3s 4s 5s 2p 3p 3d 4p 4d 4f 5p 5d 5f *degenerate orbitals are orbitals equal in energy to each other 508 Here a few other ideas to keep in mind…. The following ideas are generalizations … The total number of electrons in any Principle Energy Level (n) equals 2n2 Heisenberg Uncertainty Principle: We cannot know with certainty the position and the momentum of an electron, simultaneously. The Heisenberg Uncertainty Principle tosses out the idea of Bohr’s “orbits”, and replaces it with orbitals Orbitals represent a shaped volume of space (thus, in essence, an energy) surrounding the nucleus. An orbital may hold 0, 1 , or a maximum 2 electrons. Aufbau Principle: Electrons are ordered (configured) from lowest energy levels to higher energy levels. The diagonal rule (prior page) helps to explain the order of filling – but it is limited and NOT absolute. There are exceptions to the Aufbau Principle, although these exceptions are not tested on the AP exam. Hund’s Rule: When a sublevel contains degenerate orbitals, electrons are configured into the orbitals, one at a time, and are paired only when energy concerns become dominant. Pauli Exclusion Principle: No two electrons may share the exact same quantum numbers – hence, when two electrons occupy a single orbital, their spins are in opposition to each other. Now, using the ground state electron configuration representing an atom of the element Fe0, with 26 electrons and thus 26 protons, answer the following questions a - n to the best of your abilities. This is NOT graded. Ten of fourteen is outstanding … 7 of fourteen is solid… 1s2 2s2 2p6 3s2 3p6 3d6 4s2 OR [Ar] 3d6 4s2 a. What is meant by “the ground state electron configuration” * The electrons are organized so that they represent that order, at their lowest energy levels. b. The element, will be found in period * 4 of the periodic table c. Based solely upon the configuration Fe, a transition metal (hint) is in group *8 periodic table of the d. Based upon the configuration, iron is a member of the family called the *transition metals you can infer this because of the configuration’s * incomplete d sublevel e. The atom appears to have * 4 (a number) incomplete or half-filled orbitals, in the ground state f. The first two electrons used in bonding are probably from the *4s oxidation state of 2+ sublevel, creating an 509 g. The oxidation state of 3+ is a possibility, due to the possibility that *one 3d electron can be used/lost h. The electron configuration of the 2+ oxidation state species is: * [Ar] 3d6 i. The electron configuration of the 3+ oxidation state species is: *[Ar] 3d5 j. The atom has electron shells completely filled up through the *2nd level. k. The configuration implies that while occupied, the * 3rd and 4th NOT filled. principal energy principal energy levels are l. The loss of the 4s2 electrons is best explained in that these electrons (select one) *ii) i) are the least shielded electrons from the nuclear attractive force, thus are held more tightly to the nucleus ii) experience the greatest shielding effect, relative to the nuclear attractive force, and are bound less tightly to the nucleus iii) are the only possible electrons to form bonds due to their excessively low level of energy, relative to the rest of the electrons of the atom iv) provides the resulting ion with a stable valence octet m. The arrangement of the 3 d electrons: 3d ↑↓ ↑ ↑ ↑ ↑ with multiple inner d electrons exhibiting similar spins suggests that iron i) ii) iii) iv) Keep Going…. *iv is more easily oxidized than atoms of sodium has a greater electronegativity than atoms of sulfur can be easily melted exhibits (ferro)magnetism n. As electrons are removed from the 4s and then from the 3d sublevel of an iron atom which of the following occurs? e.g.) as Fe4+ is formed from Fe3+ *iv i) the radius of the resulting ion becomes smaller ii) the ionization energy required for the removal of the next electron is greater iii) the tendency to attract electrons increases iv) all of the above ASSIGNMENT: Seriously check out Bozeman Science … a favorite http://www.youtube.com/watch?v=2AFPfg0Como It is 10 minutes well spent….it will introduce you to Coulomb’s Law and remind you of all you have learned from last year… It will begin to blend quantum and periodic trends. 510 I) Energy can travel through space by electromagnetic radiation, which are those forms of energy represented on the electromagnetic spectrum. All forms travel at the speed of light (in a vacuum) and exhibit a similar type of wavelike behavior; with wavelength, frequency and amplitude. All forms of electromagnetic radiation exhibit: wavelength (λ), frequency (ѵ measured in Hertz), and speed. Wavelength is just that …when measured from crest to crest or trough to trough Frequency is the number of cycles or waves per second that pass a given point in space. Small wavelength radiation has a high frequency; longer wavelength radiation has a lower frequency. This inverse relationship is expressed as: λѵ = c Where c is the speed of light in a vacuum. This equation is on your reference tables. You can compare the relationships by studying the following diagrams from Zumdahl (p 297) A) Keep in mind: shorter wavelengths = greater frequencies (a greater number of waves occur in 1 sec = greater frequency, but that must equate to a smaller wavelength) longer wavelengths = lower frequencies Recall: Converting from Hertz (Hz) to Megahertz (MHz), using Dimensional Analysis: 1) conversion factor: 1 megahertz (MHz) = 1 x 106 (or 1 million hertz) 2) Microwaves have a frequency of 2.5 x 109 Hz. Use dimensional analysis to calculate the equivalent of that frequency in MHz MHz = * 2.5 x 109 Hz | 1 MHz | = 2.5 x 103 MHz 1 x 106 511 3) Max Planck, the grandmaster of quantum, found, in the very late 19th century that: energy is quantized, and can only occur in whole number units equal to hѵ … See the work on standing waves and Schrodinger / De Broglie (about pages 515-16) where h is known as Planck’s constant of 6.626 x 10-34 J•s. Each small unit of exchanged energy was called a quantum. (plural = quanta) Energy must be transferred in whole “packets”… thus energy has some commonality to “particle” behavior as well as wave properties! (This is on your tables) B) Enter Einstein and the Photoelectric Effect …Einstein embraces Planck’s work: 1) Einstein assumed that the radiant energy striking a metal surface behaves like a stream of tiny packets of energy … He called these quanta of energy: photons. Using Planck’s work (and constant) he determined that a minimum frequency of light, different for different metals is required for to kick off (or emit) electrons. For instance, that frequency of light which will cause cesium metal to emit electrons must be equal to or greater than 4.60 x 1014s-1 Thus, electromagnetic radiation is itself quantized and can be seen as a stream of particles called photons. The energy of a photon is directly proportional to its frequency … 2) Ephoton = hc λ Where: E = the energy of the EM radiation c = 2.998 x 108 m/s h = 6.626 x 10-34 J•s (Planck’s constant) λ = the wavelength of EM radiation in meters e.g.) What is the energy of blue light with a wavelength of 450.0 nanometer? Where 1 nanometer = 1 x 10-9 meter a) Convert to meters: meters = *450.0 nm | 1 x 10-9 m | =4.50 x 10-7 meter 1 nm b) Then use the equation: E = hc λ -34 * E = (6.626 x 10 J•s )( 2.998 x 108 m/s) = 4.41 x 10-19J 4.50 x 10-7 m EM radiation can be reflected (bounced off something, like a silver mirror), refracted (as it passes through water …it can become bent) Humans can generally detect, visually the wavelengths in the visible spectrum (700 nm to 400 nm) 512 Humans and other organisms (e.g honey bees) can also detect UV radiation (Humans cannot see UV light, but bio-molecules (DNA) can react to it….therefore detect it, passively) Also certain drugs taken internally (e.g. tetracycline) or applied externally (salicylic acid) can absorb UV radiation, and then releases it into the skin, killing cells. This is why those using acne medications, must be careful about sun exposure Salicylic acid TRY THIS!!! Brown & Lemay 216 a.k.a 2-hydroxybenzoic acid: a beta-hydroxy acid essentially, helps to remove or the outer skin layer (the epidermis) which absorbs the vast amount of UV radiation) revealing the far more sensitive dermis layer to the power of UV. It additionally can absorb UV radiation, creating skin cell death. The colors of a butterfly are due in large part to specialized protein structures, trapping or bending light interacting with those protein structures. EMS radiation can also interact with the electrons of the atoms, and thus the bonds made between atoms, by the sharing or transfer of electrons. Think of a fairly simplified means of explaining the color of a blouse or shirt or the changing colors of leaves, in the autumn. The yellow light given off by a sodium vapor lamp used for public lighting, and those lamps found along the highway, has a wavelength of 589 nm. a) What is the frequency of this radiation? *convert to meters: meters = 589 nm| 1 x 10-9 m | 1 nm *use: λѵ = c or: ѵ = c/λ *ѵ = (2.998 x 108 m/s) = 5.09 x 1014 5.89 x 10-7 m Hint1: *Convert nm to meters. Hint 2: *You are asked for frequency, and given wavelength. You know of an equation using both variables … It is in your reference tables. ans: 5.09 x 10 b) What is the most correct unit for the answer? *s-1 Assignment: Read over, in Brown and LeMay p. 219-23 … Line Spectra and the Bohr Model For Fun: The Big Bang Theory: Schrodinger’s Cat: https://www.youtube.com/watch?v=pNTMYNj2Ulk It’s pretty right on track ….a nice “every person’s” explanation. 513 II) Enter: Werner Heisenberg and The Uncertainty Principle: When applied to the electrons in an atom, the Uncertainty principle states that it is impossible for us to know simultaneously both the exact momentum of an electron and its exact location in space. A) Your Brown and LeMay text runs you through a neat little calculation on page 225. It shows that the level of uncertainty in position for an electron in a H atom is 1 x 10-9 m Given that the diameter of a hydrogen atom is about 1 x 10-10 m, you will note that the uncertainty factor is greater than the diameter of the atom! Hence …we are clueless as to the location. 1) Between Louis De Broglie’s wave-particle duality concept (Nobel Prize in 1929), and Heisenberg’s Uncertainty Principle (Nobel Prize in 1932), Erwin Schrodinger’s wavefunction equation and Paul Dirac’s re-working of it to meet relativistic issues (resulting in a shared Nobel Prize in 1933) … the wave nature of the electron is embraced … & we begin to speak of the energy of the electron while describing its location, in terms of probabilities. Exit, Einstein. II) Atomic Orbitals: Due to the inferences made using the Uncertainty Principle, we can’t identify the exact location of an individual electron around the nucleus to which it is bound. We speak therefore of the probability of an electron being in a region of space around the nucleus …. A) (psi)2 or ψ2 = probability distribution (or electron density), represents the probability that an electron will be found in a region of space. B) The solution of the Schrodinger equation does not give us a line …as we might assume. Rather it gives us a set of wave functions which describe the e- bound to a specific nucleus ψ (psi) represents a wave function and ψ2 provides information about the e-‘s location when it is in an allowed energy state. Think of these diagrams as a stop-motion set of photos envisioning in 3D the travels of an eThe term, atomic orbital, is often envisioned in this manner. Brown & Lemay p 227 514 C) De Broglie and Schrodinger treated the electron in a hydrogen atom like a standing wave. (one that does not travel through space) The oft used analogy is a plucked guitar string. … The plucked guitar string produces a standing wave that has a fundamental frequency and higher overtones … Schrodinger envisioned something similar for the electron … which exhibits both a lowest-energy standing wave, and higher energy ones. There are also NODES … where the magnitude of the wave is zero … meaning that there is a 0 probability of finding the electron. Nodes and Standing Waves (Brown and LeMay p 226) I have used the example of a rope tied to a tree in the past: https://joulespersecond.wordpress.com/ Check out: https://www.allaboutcircuits.com/textbook/semiconductors/chpt-2/quantum-physics/ 515 Another means of visualizing the electron as “standing wave” String vibrating at resonant frequency between two fixed points forms standing wave. The atom according to de Broglie consisted of electrons existing as standing waves, a phenomenon well known to physicists in a variety of forms. As the plucked string of a musical instrument (Figure above) vibrating at a resonant frequency, with “nodes” and “antinodes” at stable positions along its length. De Broglie envisioned electrons around atoms standing as waves bent around a circle as in Figure below. “Orbiting” electron as standing wave around the nucleus, (a) two cycles per orbit, (b) three cycles per orbit. Electrons only could exist in certain, definite “orbits” around the nucleus because those were the only distances where the wave ends would match. In any other radius, the wave should destructively interfere with itself and thus cease to exist. 516 There are other means by which this is envisioned. We may “see” it as a radial probability distribution Most probable distance from the nucleus. Zumdahl 9th ed. p. 312 Check out: YouTube: Cassiopeia Project segment on radial probability: http://www.youtube.com/watch?v=Fw6dI7cguCg&feature=related More Cassiopeia Project https://www.youtube.com/watch?v=n4LnvLAjmcU At 12 minutes (Bohr Model), 16 min de Broglie 18 min Heisenberg, particle as wave: 20 min, Schrodinger 21.30, Heisenberg and radial probability 23 min, spin 29 min III) Quantum Numbers: n, ℓ, mℓ, ms A) The solution to the Schrodinger equation for the hydrogen atom yields a set of wave functions called ORBITALS (not orbits …as in the Bohr Model) *each orbital has a characteristic shape and energy 1) The quantum mechanical model uses 3 quantum numbers n, l, ml to describe an orbital. a 4th quantum number ms is used to describe the intrinsic property of the electron(s) of an orbital. This property is electron spin … we’ll see this when we re-visit the Pauli Exclusion Principle 2) The 3 quantum numbers used to describe the orbital are: a) n = principal quantum number i) n may equal positive integral values of 1, 2, 3, 4… as n increases the orbital becomes larger and the electron spends more time farther from the nucleus. as n increases, the electron has a higher energy and is therefore *less tightly bound to the nucleus. b) l = angular (azimuthal) momentum quantum number. i) l may have integral values from 0 to (n-l), for each value of n. l defines the *shape of the orbital. the value of l is generally designated by the letters s, p, d, f which correspond to l values of 0, 1, 2 and 3 … or rather 0 to n-1 517 Hence: value of l 0 1 2 3 4 5 symbol s p d f g h c) mℓ = magnetic quantum number with integral values between – l to l, including 0 This quantum number indicates the orbital’s orientation in space 3) The collection of orbitals with the same value of n is called an electron shell. The set of orbitals that have the same n and l values is called a subshell … designated by an integer for n and a letter (s, p, d, f) for l SUMMARY OF QUANTUM NUMBERS: In essence for quantum number work: n = 1 -7 l = 0 to (n-1) I have memorized: where s = 0, p = 1, d = 2, f = 3 mℓ = - l to l ms = ½ or -½ Quantum Numbers Name Designation Property of the Orbital Related to the size and energy of the orbital Principal quantum number n May be any positive integer 1, 2, 3, 4, …. Related to the shape of the orbital (s, p, d, f) Angular momentum quantum number l Magnetic quantum number Electron spin quantum number ml ms Related to the position of the orbital in space in relation to other orbitals. For instance, this deals with px, py, pz or those positions related with the d sublevel and/or f sublevel. Related to the spin of the electron, which can be only one of two values +1/2 or -1/2 Based essentially in the Pauli exclusion principle. 518 SUMMARY OF ALLOWED QUANTUM NUMBERS n l ml 1 0 0 2 0 0 1 -1, 0, +1 3 0 0 1 -1, 0, +1 2 -2, -1, 0, +1, +2 4 0 0 1 -1, 0, +1 2 -2, -1, 0, +1, +2 3 -3, -2, -1, 0, +1, +2, +3 Number of orbitals Orbital Name Number of electrons 1 1 3 1 3 5 1 3 5 7 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 2 2 6 2 6 10 2 6 10 14 http://www.angelo.edu/faculty/kboudrea/general/quantum_numbers/Quantum_Numbers.htm For a very nice summary (with practice problems) on quantum numbers try: I also really like: https://www.youtube.com/watch?v=kS-U0R3Kx8E https://www.allaboutcircuits.com/textbook/semiconductors/chpt-2/quantum-physics/ Reading: Take a look at the table and readings on page 229 of your Brown and LeMay text, for more. 519 IV) Representations of Orbitals: A) You’ll recall from Honors, that every s subshell is spherical … differing in volume, as we move from 1s to 2s to 3s …. http://www.dlt.ncssm.edu/tiger/diagrams/structure/s-orbitals_3-up.jpg B) No one will forget the “dumbbell” shaped p orbitals traveling oriented along the x, y, z axes. surface boundary representation 3 possible "p" orbitals and orientation in space Probability Distributions http://wps.prenhall.com/wps/media/objects/724/741576/Instructor_Resources/Chapter_01/Text_Images/FG01_0139UN.JPG https://www2.bwdsb.on.ca/~f_schlenker/4U/4U%20quantum%20chemistry/university%20website/imgres_files/a.htm http://upload.wikimedia.org/wikipedia/commons/1/1b/Eixos.jpg http://upload.wikimedia.org/wikipedia/commons/thumb/1/1c/Px_py_pz_orbitals.png/525px-Px_py_pz_orbitals.png 520 Shapes of three of the five d orbitals Checkout!! YouTube: http://www.youtube.com/watch?v=45KGS1Ro-sc Synthesis of Bohr/Heisenberg Particles with Schrodinger Waves YouTube: https://www.youtube.com/watch?v=F-xLQ1WBIlQ Shape of orbitals (not bad! …narration is a bit tougher, but good) YouTube: http://www.youtube.com/watch?v=VfBcfYR1VQo Shape of s, p YouTube: http://www.youtube.com/watch?v=sMt5Dcex0kg Scandium configuration Assignment: Take about 20 minutes to go through the 4 video clips from YouTube … just to re-familiarize yourself with the general concepts. Assignment: Get to page 231 of the text and read over pages 230 (bottom of) – 231. Try the “Go Figure”. It is a nice little exercise in analysis … nothing grand … the answer is on page A-39. C) Remember that the electrons of any “p” sublevel are equal in energy to each other… That is they are degenerate. 1) Orbitals which are degenerate are those of a common subshell … having the same energy. Hence the orbitals of 2p are degenerate. The orbitals of 3d are degenerate. 2) This holds true for the electrons of a specific “d” or “f” sublevel as well. 3) If we get to the Actinides and Lanthanides … we’ll see that configuration into the d and f sublevels play a significant role…. Mind you … it is one of the reasons as to why those elements have been pulled out of the main table. * e.g. Lanthanide series: Period 6: 6s2 stabilizes while 4f is filled (or rarely 5d..Ce, Gd, Lu) 521 D) Up to now your text has dealt exclusively with hydrogen atoms … but many-electron atoms have a few twists and turns worth reviewing. There are NO SURPRISES here … It may be worded a bit more formally than that to which we are familiar … but it is all the same tune…. In a many-electron atom, for a given value of n, the energy of an orbital increase with increasing value of l… e.g. 3s < 3p < 3d. The CONVERSE of this idea may be used to help explain the ionization (or oxidation) of species … in that we must recognize that electrons that are closer to the nucleus experience greater coulombic forces of attraction and therefore require more energy to be removed. In a hydrogen atom the energy of an orbital depends only on n … but in many-electron atoms, there is electron – electron repulsion (leading to shielding effect) … and this repulsion leads to the various sublevels being different in terms of energy. E) Pauli Exclusion Principle (ms) Recall that a moving charge creates a magnetic field … and e- have “spin” Essentially, no 2 e- may have all 4 quantum numbers, exactly the same) AND * have the same same directional spin. So the 2 e- of any single orbital must have opposite spins. (Translation: When 2 e- occupy the same orbital they must have opposite “spins” because this lessens the effects of the magnetic fields developed by moving charges 1) However, it is important to note that this pairing is not as stable a lone e-, because paired (correlated) electrons have more energy, than unpaired eAccording to Pauli, electrons of “like spin” tend to get as far from each other as is possible. This concept has been the most influential factor in determining the shapes and properties of molecules. Organic Chemistry Morrison & Boyd 3 ed p 8 (see Lewis Structures Unit 5) rd let : 1 e- be symbolized by or so a filled orbital = or a filled orbital = 2) But, the Pauli Exclusion Principle is the basis for our belief that an orbital may only hold a maximum of 2 electrons. a) For a given orbital, the values of n, l ml are fixed. If we want to put more than one electron in an orbital and satisfy the Pauli Exclusion Principle, we must assign a different ms value (-1/2 or +1/2). There IS NOT THIRD OPTION! Hence an orbital can only hold a maximum of two electrons. 522 3) The Pauli Exclusion Principle answers one of the more basic questions of physics, chemistry & biology: Why does matter have volume? First, there are 2 types of fundamental particles: Fermions and Bosons spin: odd half integral spin 1/2, 3/2 spin: whole integral spin 0,1,2 Characteristics: may be elementary or composite particles but, only 1 can exist per quantum state ... Obey Dirac-Fermi Equations and Pauli Exclusion Principle Characteristics: may be elementary or composite particles but, many may occur in the same quantum states (which helps explains, lasers, for instance) ... Obey Bose-Einstein equations and do NOT obey Pauli Exclusion Principle e.g.) quarks, electrons, protons, neutrons, neutrinos, and nuclei with an odd number of fermions like a nucleus of 7Li) e.g.) all force carriers, like photons & gluons citations: Look up these for more information.... http://www.particleadventure.org/fermibos.html http://en.wikipedia.org/wiki/Boson http://www.pa.msu.edu/courses/1997spring/phy232/lectures/atomic/bosons.html http://sciencepark.etacude.com/particle/forces2.php eg.) mesons, any nucleus with an even number of fermions like 4He or 12C. Thus: Matter has volume (but energy does not) because... ... Photons are bosons, not fermions, and as such do not obey the Pauli Exclusion Principle. They can and do travel "through" other photons. Fermions are particles such as Protons, Neutrons, and Electrons. The exclusion principle describes the observation that no two of those particles can occupy the same spot in space AND have the same Quantum Numbers. (Which describe the properties of each particle, such as energy level and spin) It is a direct result of that principle that matter has volume. Most of the atoms that make up all matter is simply empty space. It only has volume because the electrons cannot occupy the same spots and cannot get close to the nucleus on average. http://www.physicsforums.com/showthread.php?t=513639 4) (I believe) ... by 1924 - 1926 Pauli had used such results to develop what is now called the Exclusion Principle. The Exclusion Principle directly underpins the construction of the periodic table. It is that important to chemistry... 5) Around 1930 Pauli proposed the existence of the neutrino. The neutrino as a fundamental particle was confirmed in 1956, by Reines and Cowan. Professor Charles P Enz was Pauli's last research assistant and his biographer. According to Enz's book, No Time To Be Brief when Pauli was written with the news confirming the existence of the neutrino, Pauli cabled back: "Thanks for message. Everything comes to him who knows how to wait. Pauli." 6) Pauli was awarded the Nobel Prize in Physics for his contributions to the field in 1945 523 F) Hund's Rule: When accounting for the e- in sublevels with multiple orbitals the electrons fill each orbital one at a time and then pair up using Pauli Exclusion Principle (Bus Rule) p _____ ______ _______ not p ______ _______ _______ 1) In sublevels with multiple orbitals (such as : p,d,f) , the orbitals of the sublevel are ... degenerate (equal in energy) (e.g. px energy = py energy = pz energy ) thus, making the electrons of those orbitals of the specific sublevel EQUAL in energy to each other. a) e.g. the electrons of 2px 2py 2pz . are equal in energy to each other 2) Translation: My friend Elaine Battaglino explained this using the Battaglino Bus Rule: When boarding a bus, all the seats are equal, and passengers fill seats of equal energy one at a time... and double up only when each seat has one passenger already (but the second passenger/seat must sit upside down to obey the Pauli Exclusion Principle :-) Or, you could look at it like a movie theater... All the seats (orbitals) of a specific row (sublevel) in a movie theater give you the same view (have equal energy). 3) According to Hund: A symmetric spin state (multiple electrons in separate orbitals) forces an anti-symmetric spatial state (e.g. the different spatial directions of x,y,z) where the electrons are on average further apart and provide less shielding for each other, yielding a lower energy. http://hyperphysics.phy-astr.gsu.edu/hbase/hframe.html a) What Hund was stating is very important...The electrons entering a degenerate sublevel (a sublevel containing more than one orbital (like p,d,f sublevels)), will have THE MOST STABLE arrangement when the electrons occupy the orbitals singly, rather than in pairs. 4) The lessening of this shielding effect becomes important in our next unit ... But essentially, a lesser (lower) shielding effect, helps to stabilize an atom a) Shielding effect the interference for the nucleus’s pull on a valence electron, due to the "inner (or core) electrons". The lessening of the shielding effect, (that is, allowing the nucleus to attract the electron as maximally as possible) creates a more stable electron configuration and greater chemical stability, because the electron(s) is/are at lower energy 524 H) Condensed Quantum Electron Configurations Recall that there are the inner (core) electrons and the outer valence electrons. We are beginning to see that the valence electrons are NOT the only electrons that can take part in a chemical reactions … Rather, with the transition metals especially, it is not uncommon to see the inner d electrons play some sort of role in bonding. What then, are the valence electrons? Frankly and simply, valence electrons are the outermost electrons. In the case of transition metals, the valence electrons and the inner d electrons can be very close in energy to each other. The inner core electrons comprise the levels of electron configuration equal to the noble gas prior to the element. For instance, when dealing with iron, as we did in our diagnostic: 1s2 2s2 2p6 3s2 3p6 3d6 4s2 OR [Ar] 3d6 4s2 The core configuration is that of argon … the noble gas capping Period 3, and built upon for iron, found in period 4. Hence for ruthenium, in period 5 and just below iron, it is not surprising to see the configuration abbreviated as: [Kr] 4d7 5s1 Ru atoms have the same e- configuration up to an atom of Krypton, but extends beyond the stable octet of krypton, to the 4d and 5s sublevel. Now, because configurations can become so long ... we often abbreviate a configuration, by using the symbol of the noble gas of the element's prior period ... and then finish up the configuration with inner d (or f electrons), and of course, the valence electrons in an "s" or "s and p" set of orbitals. TRY THIS!!!: Use only your periodic table and your grasp of the concepts. You may use the condensed configuration (or not) … 1) Write the ground state electron configuration for an atom of titanium (atomic # 22) *1s2 2s22p6 3s23p63d2 4s2 or [Ar] 3d2 4s2 2) Write the electron configuration for an atom of nickel (atomic # 28) *1s2 2s22p6 3s23p63d8 4s2 or [Ar] 3d8 4s2 3) Write the electron configuration for an atom of zinc (atomic #30) *1s2 2s22p6 3s23p63d10 4s2 or [Ar] 3d10 4s2 a) Zinc will react with Cl2 to produce ZnCl2. From which sublevel are the electrons of zinc most likely lost? *4s 525 TRY THIS!!! Condensed Configurations for the Alkaline-Earth Metals Alkaline- Earth Symbol Atomic Electron Configuration Number Using the Condensed Method Beryllium Be 4 [He] 2s2 Be has the inner core of He plus 2s2 Magnesium Mg 12 [Ne] 3s2 Mg has the inner core of Ne, plus 3s2 Calcium Ca 20 [Ar] 4s2 Strontium Sr 38 [Kr] 5s2 Barium Ba 56 [Xe] 6s2 Radium Ra 88 [Rn] 7s2 1) When given the ground state configuration: [Rn] 5f 36d1 7s2 the atom is of the element: *uranium 2) When given the ground state configuration: [Ar] 3d10 4s2 4p5 the atom is of the element *bromine 3) When given the ground state configuration: [He] 2s2 2p5 the atom is of the element *fluorine 4) When given the ground state configuration: [Ar] 3d10 4s1 the atom is of the element *copper 5) When given the ground state configuration: [Ne] 3s2 3p4 the atom is of the element *sulfur answers: uranium, bromine, fluorine, copper, sulfur H) Quantum Odds & Ends: Before we begin to actually interpret configurations I want to re-state that the Aufbau Principle is great and handy for writing configurations, but it really does not help explain some of the seeming inconsistencies. (Remind me, to discuss Cr vs. W. There’s a mind-bender!!...or you can let Dr E Scerri do it, when you tackle one of the references) For anyone interested in going further I would like to suggest a slow, studied read of the following, and of a few of the provided references. Quantum theory shows that each atom's electronic structure is a unique compromise between several different effects. Electronic configurations of the fourth period elements can be appreciated by considering these effects. These issues surround what I have already included in your notes. Inconsistencies can be explained by connecting (yet recognizing them as competing) ideas such as: lessening of shielding effect (or increasing ENC) the relative restrictive pathways of some sublevels (like d) compared to the more “open” s, the natural closeness in energies of some sublevels (such as 3d and 4s), along with the fact that the sublevels shift positions (and thus change in their energies), as the effects of added electrons are exhibited. Recall … subshells start repelling each other, generating a shielding effect and this affects the energies of the subshells!!! Check out: http://employees.oneonta.edu/viningwj/sims/atomic_electron_configurations_s1.html 526 Or, according to Dr. Fred Senese: http://antoine.frostburg.edu/chem/senese/101/electrons/faq/4s-3d.shtml Raising n raises orbital energy. Electrons are attracted to the nucleus. To pull an electron farther away from the nucleus, you have to work against that attraction. That means an electron farther from the nucleus has more energy than electron closer in; energy is required to move the electron out, and energy can be released when the electron moves in. So we expect outer shells to have higher energies than inner shells, because increasing n increases the average distance between the nucleus and the electron. For atoms heavier than copper this effect dominates, and 4s electrons have higher energy than 3d electrons. Raising l raises orbital energy. Higher l values result in orbitals with more nuclear nodes (a node being a place where the probability of finding the electron is zero). We say high-l orbitals are "less penetrating" because their electrons have a lower probability of being found at or near the nucleus. That gives high-l orbitals (like d orbitals) more energy than low- l orbitals (like s orbitals) within the same shell. This effect causes 4s orbitals to have lower energy than 3d orbitals for elements lighter than copper. (Although for hydrogen, the unoccupied 4s and 3d orbitals have nearly identical energies). Within a subshell, more unpaired spins means a lower overall energy. Quantum mechanics predicts that the motions of electrons with paired spins are "correlated". Paired electrons move together, while electrons with unpaired spins can stay farther away from each other on average. Since electrons repel each other, paired electrons have more energy than unpaired electrons, all other things being equal. This spin correlation effect explains why Cr has a [Ar] 4s13d5 configuration rather than a [Ar]4s23d4 configuration- the former has more unpaired electron interactions than the latter. Other references are: 1) http://www.chemguide.co.uk/atoms/properties/3d4sproblem.html 2) http://ericscerri.blogspot.com/2012/06/trouble-with-using-aufbau-to-find.html 3) http://crescentok.com/staff/jaskew/ISR/chemistry/class4.htm 4) W. H. Eugen Schwarz: The Full Story of the Electron Configurations of the Transition Elements: Journal of Chemical Education, Vol. 87 No. 4 April 2010 (copies available) 5) R. N. Keller: Textbook errors Energy Level Diagrams and Extranuclear Building of the Elements: J. Chem. Educ., 1962, 39 (6), p 289, published June 1962) (copies available) 6) Quantum numbers and Orbital diagrams: https://www.youtube.com/watch?v=nexoMIZK7cE For a peak at the de Broglie-Bohm model, check out: http://www.spaceandmotion.com/Physics-DavidBohm-Holographic-Universe.htm and look at the section called Bohmian Mechanics. 527 V) Writing Ground State Configurations (Note, in AP Chemistry e- configurations are NOT on your tables) DIRECTIONS: I have asked a number of questions which will require you to interpret and re-write a number of orbital notations. Diagonal Rule (Aufbau) 7s 6s 5s 4s 3s 2s 1s 7p 6p 5p 4p 3p 2p 7d 6d 5d 4d 3d 7f 6f 5f 4f For questions 1 - 5 use the orbital configuration: 1s22s22p63s23p4 You may re-write it like this ... (re-write the configuration, using the "arrows" first) or like this vertical design 1s _____ 2s _____ 2p _____ _____ _____ 3s _____ 3p _____ _____ _____ low energy high energy _____ 1. What is the total number of electrons found in this atom? _____ 2. Which principal energy level is considered the valence principal energy level ? _____ 3. How many electrons are in the valence principal energy level? _____ 4. How many orbitals are only half - filled? _____ 5. How many sublevels are filled completely? http://www.nlcs.k12.in.us/bnlhs/staff/coy/notes/c1notes/l32.pdf For questions 6 - 12 use the orbital configuration: 1s22s22p63s23p63d34s2 1s ___ 2s ___ 2p ___ ___ ___ 3s ___ 3p ___ ___ ___ 3d ___ ___ ___ ___ ___ 4s ___ low energy high energy _____ 6. What is the total number of electrons found in this atom? _____ 7. Which principal energy level is considered the valence principal energy level ? _____ 8. How many electrons are in the valence principal energy level? _____ 9. How many orbitals are half - filled? _____ 10. How many sublevels are filled completely? _____ 11. What is the group number in which this element is found? _____ 12. In what period of the periodic table is this element found? 528 For questions 13-20 use the orbital configuration: 1s22s22p63s23p63d104s24p3 1s ___ 2s ___ 2p ___ ___ ___ 3s ___ 3p ___ ___ ___ 3d ___ ___ ___ ___ ___ 4s ___ 4p ___ ___ ___ low energy high energy _____ 13. What is the total number of electrons found in this atom? _____ 14. Which principal energy level is considered the valence principal energy level ? _____ 15. How many electrons are in the valence principal energy level? _____ 16. How many orbitals are only half - filled? _____ 17. How many sublevels are filled completely? _____ 18. The above atom is representative of which element ? (Write the symbol) _____ 19. What is the group number in which this element is found? _____ 20. In what period of the periodic table is this element found? 529 VI) Periodic Table and Wave-Mechanic Model A) Using quantum, the elements are organized, beyond groups and periods, into "blocks". The block of an element is determined by the last orbital to be used to configure an electron. That is, the block corresponds to the angular momentum quantum number (l) 1) Now, don't confuse the "block"-designation with the valence level. Recall Aufbau... For instance, it is accurate to say that the s-block elements do have a Valence Principal Energy Level with an s sublevel e.g. Lithium: 1s 2s The p block elements do indeed have s and p orbitals in the Valence Principal Energy Level .... eg.) Fluorine: 1s 2s 2p However, while the transition metals are in the d block, they have a valence level with an s sublevel. . . BUT, the last configured electron(s) go into an orbital of a d sublevel e.g.) Iron: [Ar] 3d 4s these are the last to be configured BUT!! These are the valence e- s block p block d block f block http://www.mikeblaber.org/oldwine/chm1045/notes/Struct/EPeriod/Struct09.htm 2) As stated by Dr. Fred Senese at http://antoine.frostburg.edu/chem/senese/101/periodic/ Each block contains a number of columns equal to the number of electrons that can occupy that sublevel The s-block has 2 columns, because a maximum of 2 electrons can occupy the single orbital The p-block has 6 columns, because a maximum of 6 electrons can occupy the three p orbitals The d-block has 10 columns, because a maximum of 10 electrons can occupy the five d orbitals. The f-block has 14 columns, because a maximum of 14 electrons of the seven orbitals in a f-sublevel. 530 VII) Quantum helps to answer questions regarding a number of topics: from magnetism to chemical reactivity to the natural production of (or lack thereof) certain ions. A) For example: How is iron seen as being ferromagnetic ... but copper, is not? Why is sulfur more chemically stable as an anion, of S2-, than as an atom, S0? How do we explain carbon atoms making 4 equal covalent bonds to atoms such as H, and not just two? What seems to be the explanation for multiple oxidation states of some transition metals, such as Fe3+ ions and Fe2+ but only on stable oxidation state of sodium ion, Na1+? We have evidence of Na1+ & Mg2+ cations, but why is there so little evidence of Mg1+ ions? B) Quantum can help explain phenomena that are related to far more difficult questions such as: Why is mercury, in the liquid phase, at STP? How can we account for the non-white (non-silver) color of solid gold and the (lesser, but still evident) golden-hue of metallic cesium? How can we account for the smaller-than-predicted atomic radii of francium and radium? Why is gold so resistant to oxidation? Why is lead metal (group 14), not found in tetrahedral crystals, like carbon (diamonds)? First, recognize that the chemical and physical properties of any atom is due to the energies and spatial distributions of the electrons, of the atom. Secondly, as atoms become larger in nuclear charge, Einstein's Theory of Special Relativity can be employed to explain certain issues. To this end, we need to assume two issues: Special Relativity states that the mass of any moving object changes as its velocity changes! Under most "normal" circumstances, this velocity is so small that any relativistic issues are minimal (ridiculously so...), thus elements of Periods 1 -4 are really not an issue. (Thayer J. Chem. Educ., 2005, 82 (11) As the mass of the atoms increases (as with the atoms of the elements in periods 6 and 7), the velocity of the electrons increases to far more important levels, approaching the speed of light. (To help simplify the issue, recall that Bohr showed that all electrons exist around a nucleus in the same amount of time, thus supporting the idea that valence electrons have greater kinetic (and potential energies. Perhaps a more sophisticated and correct view is that the standing waves of the electrons are exhibited around the nucleus simultaneously). Anyway, with the increase in mass relativistic causations / explanations of main group chemistry are valuable. The relativistic effects are not explained by solutions to the Schrodinger Equation(s). It was Paul Dirac that added or adjusted the equations to deal with these relativistic issues of the very heavy atoms (with, perhaps the initiating force of Wolfgang Pauli...but that is debated). The following page addresses only a few specific issues - but it hopefully offers you enough intriguing thoughts to see why some men and women just love this field. As a bit of advice... keep this mind: When considering the relativistic effects, I think there is a domino effect essentially ... I am sometimes lost in a "which came first chicken/egg” issue. This is evidenced in the following ideas that the velocity of electrons increase, causing shrinkage of orbitals in terms of the s and p, and size increases at the d and f ...etc. To start putting these ideas together, read on… 531 The relativistic phenomena exist primarily due to a large nuclear charge, causing an increase in electron velocity. Thus, a nucleus with a large charge (lots of protons) will cause an electron to have a high velocity. A higher electron velocity means an increased electron relativistic mass, as a result the electrons will be near the nucleus more of the time and thereby contract the radius for the smaller principal energy levels (e.g... n =1,2). http://en.wikipedia.org/wiki/Relativistic_quantum_chemistry When we begin to look at the electrons of these Period 6 species, the velocity of the "s" electrons are approaching 58% of the speed of light. At such terrific velocities a few things occur: s-orbitals decrease substantially in energy as do the p-orbitals (but to a lesser extent) & this decrease in energy is due to an increase of the relativistic mass of the e- and a contraction of the s and p orbitals & a complimentary increase in electron cloud density of these energy levels. (Thayer J. Chem. Educ., 2005, 82 (11), p 1721) These factors increase shielding effect, causing an increase in the energy of the d and f orbitals. These changes in orbital energies (due to orbital contraction of the s and p orbitals), affect bond lengths and inter-atomic separations (Thayer J. Chem. Educ., 2005, 82 (11), p 1721) This phenomenon is seen in the color of gold. The 5d electrons, shift towards those wavelengths required for their excitation, due to the orbital contraction of the 6s. The liquidity of mercury at STP results from the lowered energy of the 6s electrons, weakening the Hg-Hg interatomic bonds (Norrby, LJ JCE Volume 68 Number 2 February 1991), thus ostensibly, increasing the potential energy between the bonded atoms shifting it towards values that are exhibited by other metals in the liquid phase. Other Terrific Articles and Websites: 1) I happen to have a copy of Lars J Norrby's: Why is Mercury a Liquid? (JCE Volume 68 Number 2 February 1991) If you wish for a copy …just you ask .. and you got it 2) I also have a copy of Dr. JS Thayer's chapter titled: Relativistic Effects and the Chemistry of the Heavier Main Group Elements. (JCE Volume 82, Number 11) 2005, p 1721) Or, p.66 of the Chapter 2 piece, of the same title, and author Ask me to get you a copy, if you are interested. 3) As a final visualization, you may find this link helpful.... http://tinyurl.com/668g8 or, http://www.colorado.edu/physics/2000/applets/a2.html You may wish to re-visit this visualization after some practice with orbital notation 4) As a helpful alternative voice for explaining the issues, you may find http://tinyurl.com/249bk6r or http://www.mhhe.com/physsci/chemistry/chang7/ssg/chap07_8sg.html , meaningful. TRY THIS!!! ___1 Which of the following notations best represent the valence principal energy level of a silicon atom, in the ground state ? a) b) c) d) 3s 3s 3s 3s 3p 3p 3p 3p 532 ___2 How many orbitals are completely filled in an atom of lithium in the ground state ? a) 1 b) 2 c) 3 d) 4 ___3 Which of the following notations best represents the valence principal energy level of an oxygen atom, in the ground state ? a) b) c) d) 2s 2s 2s 2s 2p 2p 2p 2p For 4 – 7 use choices a) alkali metal b) alkaline-earth metal . __ . . __ c) halogen d) noble gas ___4 An atom with the electron configuration of 1s2 2s2 2p6 3s23p6 4s2 is best classified as a(n) ___5 An atom with the electron configuration of 1s2 ___6 An atom with the electron configuration of 1s2 2s2 2p5 is best classified as a(n) is best classified as a(n) ___7 Which is the electron configuration of a fluorine atom in the ground state? a) 1s22s22p4 b) 1s22s22p5 c) 1s22s22p63s1 d) 1s22s22p63s2 ___8 The formation of an aluminum ion occurs when 3 electrons are removed. Which electron would require the most energy to remove? a) b) c) d) 3s because it is filled before 3p 3p because it is farther from the nucleus 3p because it only has 1 electron in the orbital 3s because it experiences higher attractive forces ___9 Why does chlorine gain an electron to become a chloride ion? a) It has only 7 valence electrons b) Less shielding means there is room for another electron c) Adding an electron completely fills the s and p orbitals of energy level 3 d) An electron is attracted to the large effective nuclear charge of the atom. ______10 How many electrons can occupy the 5d sublevel? ______11 How many electrons may occupy any one of the orbitals in 3p? Answers: 1) c Si is in group 14 with 4 valence e- 2) a 3) b 4) b 5) d 6) c 7) b 8) d 9) d 10) 10 (5 orbitals, with a max. number of 2 e-) 11) 2 … no orbital may hold more than 2 e- Assignment: Scan pages 240 -245 of your text for questions or issues which you would like to have addressed. 533 VIII) Electron Configuration and Magnetic Properties of Ions A) Magnetism: The attractive power of magnets. It's caused by electron spin in materials containing unpaired electrons. http://www.engineering-timelines.com/how/electricity/definitions.asp Most of us should recall that an unpaired electron (a charged particle) generates a magnetic field. http://www.fao.org/docrep/003/t0355e/t0355e02.htm This phenomenon links us to the electromagnetic spectrum and radio waves. Thus, moving electrons create radio waves & astronomers "listen" to areas of space with radio telescopes (as with the Very Large Array [VLA]), analyzing the waves for coherent, organized, patterned activity trying to verify observational data, supernovae activity, gravitational lenses (a large celestial body, capable of bending light...) and even intelligent life(???) Back to Magnetism: "Although known about since antiquity as with, for example, the lodestone (a natural magnet), explanations of magnetism tended to be descriptive. Even in the 19th century, magnetism was ascribed as a property of its lines of force. Michael Faraday showed that all matter, including gases, possesses magnetic properties." http://www.engineeringtimelines.com/how/electricity/definitions.asp http://jpkc.whut.edu.cn/web18/main/wangluo/webelements/webelements/elements/text/bi/key.html All atoms have inherent sources of magnetism because electron spin contributes a magnetic field and electron orbits act as current loops which produce a magnetic field. http://hyperphysics.phy-astr.gsu.edu/hbase/solids/magpr.html#c1 As an important aside to this is that the presence, or absence, of unpaired electrons is detected experimentally by the behavior of a element when placed in a magnetic field… hence the two issues, paired/unpaired e- and magnetism are intertwined! 534 B) 3 forms of magnetism: Diamagnetism Paramagnetism weakest Ferromagnetism strongest (overwhelms the other forms) When materials fail to react perceptively to an applied magnetic field, we tend to rate the materials as "non-magnetic" ... when really, the materials are probably diamagnetic or paramagnetic, but very weakly so.... On this issue, let’s just Cut To The Chase Personally, diamagnetic and paramagnetic substances drive me batty… (okay, battier) There is a good deal of gunk out there on the web… but the lowdown is this: Essentially, diamagnetism is associated with * paired electrons. and paramagnetism is associated with * unpaired electrons. Any substance which has both unpaired and paired electrons will skew towards a net paramagnetism, since the effect is much stronger than diamagnetism. 1) Diamagnetism: All of the electrons in an atom are paired. …Thus diamagnetism is caused by only the orbital motion of electrons creates tiny atomic current loops, which produce magnetic fields. Diamagnetic substances have only paired electrons. When an external magnetic field is applied to a material, these current loops will tend to align in such a way as to oppose the applied field. Translation: Diamagnetism is the property of repulsion of a material by a magnetic field. http://www.mhhe.com/physsci/chemistry/chang7/ssg/chap07_8sg.html Hence, any superconductor (such as the materials we hope to use in maglev trains) would be a superior diamagnetic material (because the goal is to oppose any magnetic disturbance ... and thus "push away and float") ....(note: this is highly simplified …. issues re: Lenz's Law and the Meissner Effect and the London Equation are each involved) See: http://hyperphysics.phy-astr.gsu.edu/hbase/solids/magpr.html#c1 or http://tinyurl.com/27d9fna) http://msnbcmedia4.msn.com/j/msnbc/Components/Photo_StoryLevel/080619/urban%20maglev.hlarge.jpg http://www.fzu.cz/~jirsa/htm/principles.htm 535 a) Some diamagnetic substances are: dihydrogen (H2(g)), ammonia (NH3(g)), mercury (Hg(l)), water (H2O), Ne(g), He(g), Zn0 AND Zn2+ TRY THIS!!! Given: Zn0 has a ground state electron configuration of [Ar] 4s2 3d10 Use electron configurations to prove the diamagnetic nature of Zn0 and Zn2+. Use the definition of diamagnetism to support your answer. * remove the box for an answer A species which is diamagnetic has only paired electrons. Both species of zinc have only paired electrons. For instance: Zinc atom: [Ar] 4s ↑↓ 3d ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ Because the 4s2 electrons will be lost first from the atom, upon ionization, yielding a Zn2+ ion, it too is diamagnetic … for Zn2+ ion: [Ar] 4s___ 3d ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ All electrons are still paired. 2) Paramagnetism: Some materials exhibit a magnetization which is proportional to the applied magnetic field in which the material is placed... In most materials the magnetic moments of the electrons cancel, but in materials which are classified as paramagnetic, the cancelation is incomplete. Translation: Paramagnetism is the property of attraction to a magnetic field, in proportion to the field. The atoms of a paramagnetic substance contain UNpaired e-. The property is not permanent, in that the material cannot be “magnetized”, and disappears once the magnetic field is removed. a) Some paramagnetic substances are: Uranium, Platinum, Tungsten, and Aluminum FeO (iron (II) oxide), and dioxygen gas (O2) O O & this little fact, creates a real problem, when considering the double bond between oxygen of a dioxygen molecule But this reasoning is found in molecular orbital theory http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch8/mo.html b) When weighed in a magnetic field, paramagnetic substances appear to weigh more they do, when not in a magnetic field (a diamagnetic material appears to weigh less) 536 TRY THIS!!! Given: Al3+, S2- and Fe3+. Use electron configurations for each ion to determine whether each is paramagnetic or diamagnetic. Hint:* Draw orbital diagrams (using arrows) to illustrate the configuration of the ions. Apply the definitions. ans: * Al3+ and S2- are diamagnetic … they end up being isoelectronic to a noble gas …no unpaired e-! Fe3+ is paramagnetic. http://www.bozemanscience.com/ap-chem-007-quantum-mechanical-model/ Assignment: Please prep for the work on periodic trends, by going to the Trivedi flash and messing about with 7.22 re: size, electron affinity and ionization energy. 537