1. Electronic Structure of Atoms Electronic Structure of Atoms Resources Our TB: Ch. 6 of Chemistry: The Central Science AP version (10 th edition) Powerpoint * (from pearson) and in-class work POGIL activities: (1) Analysis of Spectral Lines and (2) Interaction of Radiation and Matter Online resources for our TB (in particular online quiz) Chem tours from ch. 7 of the W.W. Norton online book by Gilbert:Chem tours from ch. 7 of the W.W. Norton online book by Gilbert: http://www.wwnorton.com/college/chemistry/gi lbert2/contents/ch07/studyplan.asp Animations from Glencoe site: http://glencoe.mcgraw- hill.com/sites/0023654666/student_view0/chapter7/ Extra quizzes from Glencoe http://glencoe.mcgraw- hill.com/sites/0023654666/student_view0/chapter7/ Video lectures from chem guy http://www.kentchemistry.com/moviesfiles/chemguy/AP/C hemguyAtomicTheory.htm Handouts and practice problems from M. Brophy’s web site

2. Electronic Structure of Atoms Chapter 6 Electronic Structure of Atoms Chemistry, The Central Science, 10th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten John D. Bookstaver St. Charles Community College St. Peters, MO  2006, Prentice Hall, Inc.

3. Electronic Structure of Atoms Waves To understand the electronic structure of atoms, one must understand the nature of electromagnetic radiation. The distance between corresponding points on adjacent waves is the wavelength ( ).

4. Electronic Structure of Atoms Waves The number of waves passing a given point per unit of time is the frequency ( ). For waves traveling at the same velocity, the longer the wavelength, the smaller the frequency.

5. Electronic Structure of Atoms Electromagnetic Radiation All electromagnetic radiation travels at the same velocity: the speed of light (c), 3.00  10 8 m/s. Therefore, c =

6. Electronic Structure of Atoms The Nature of Energy The wave nature of light does not explain how an object can glow when its temperature increases. Max Planck explained it by assuming that energy comes in packets called quanta.

7. Electronic Structure of Atoms The Nature of Energy Einstein used this assumption to explain the photoelectric effect. He concluded that energy is proportional to frequency: E = h where h is Planck’s constant, 6.63  10 −34 J-s (i.e. units for h are Js)

8. Electronic Structure of Atoms The Nature of Energy Therefore, if one knows the wavelength of light, one can calculate the energy in one photon, or packet, of that light: c = E = h

9. Electronic Structure of Atoms For electromagnetic radiation animation and problems see: http://www.wwnorton.com/coll ege/chemistry/gilbert2/tutorial s/interface.asp?chapter=chap ter_07&folder=frequency_wa velength For All Chem tours for the electrons in atoms and periodic properties topic see: http://www.wwnorton.co m/college/chemistry/gilb ert2/contents/ch07/study plan.asp Recommeded chem tours animations: Electromagnetic radiation Light Emission and Absorbtion Bohr Model of the Atom De Broglie Wavelngth Quantum numbers Electron configuration

10. Electronic Structure of Atoms The Nature of Energy Another mystery involved the emission spectra observed from energy emitted by atoms and molecules.

11. Electronic Structure of Atoms The Nature of Energy One does not observe a continuous spectrum, as one gets from a white light source. Only a line spectrum of discrete wavelengths is observed.

12. Electronic Structure of Atoms Go To Glencoe Animation http://glencoe.com/sites/common_a ssets/advanced_placement/chemist ry_chang9e/animations/chang_7e_ esp/pem1s3_1.swf POGIL activity on Spectral Lines (To Complete)

13. Electronic Structure of Atoms The Nature of Energy Niels Bohr adopted Planck’s assumption and explained these phenomena in this way: 1.Electrons in an atom can only occupy certain orbits (corresponding to certain energies).

14. Electronic Structure of Atoms The Nature of Energy Niels Bohr adopted Planck’s assumption and explained these phenomena in this way: 2.Electrons in permitted orbits have specific, “allowed” energies; these energies will not be radiated from the atom.

15. Electronic Structure of Atoms The Nature of Energy Niels Bohr adopted Planck’s assumption and explained these phenomena in this way: 3.Energy is only absorbed or emitted in such a way as to move an electron from one “allowed” energy state to another; the energy is defined by E = h

16. Electronic Structure of Atoms The Nature of Energy The energy absorbed or emitted from the process of electron promotion or demotion can be calculated by the equation:  E = −R H ( ) 1nf21nf2 1ni21ni2 - where R H is the Rydberg constant, 2.18  10 −18 J, and n i and n f are the initial and final energy levels of the electron.

17. Electronic Structure of Atoms Go To Glencoe and Norton Animations http://glencoe.com/sites/common_a ssets/advanced_placement/chemist ry_chang9e/animations/chang_7e_ esp/pem1s3_1.swf POGIL activity on Interaction of Radiation and Matter (To Complete) Go to Chem tour for Bohr Model of atom (and Rydberg equation) http://www.wwnorton.com/college/c hemistry/gilbert2/tutorials/interface. asp?chapter=chapter_07&folder=hy drogen_energies

18. Electronic Structure of Atoms The Wave Nature of Matter Louis de Broglie posited that if light can have material properties, matter should exhibit wave properties. He demonstrated that the relationship between mass and wavelength was = h mv

19. Electronic Structure of Atoms The Uncertainty Principle Heisenberg showed that the more precisely the momentum of a particle is known, the less precisely is its position known: In many cases, our uncertainty of the whereabouts of an electron is greater than the size of the atom itself! (  x) (  mv)  h4h4

20. Electronic Structure of Atoms Quantum Mechanics Erwin Schrödinger developed a mathematical treatment into which both the wave and particle nature of matter could be incorporated. It is known as quantum mechanics.

21. Electronic Structure of Atoms The Quantum Mechanical Model Energy is quantized - It comes in chunks. A quantum is the amount of energy needed to move from one energy level to another. Since the energy of an atom is never “in between” there must be a quantum leap in energy. In 1926, Erwin Schrodinger derived an equation that described the energy and position of the electrons in an atom (this slide from: J. Hushen’s presentation on Atomic Structure at http://teachers.greenville.k12.sc.us/sites/jhushen/Pages/AP%20Chemistry.aspx)

22. Electronic Structure of Atoms Schrodinger’s Wave Equation probability Equation for the probability of a single electron being found along a single axis (x-axis) Erwin Schrodinger (this slide from: J. Hushen’s presentation on Atomic Structure at http://teachers.greenville.k12.sc.us/sites/jhushen/Pages/AP%20Chemistry.aspx)

23. Electronic Structure of Atoms Quantum Mechanics The wave equation is designated with a lower case Greek psi (  ). The square of the wave equation,  2, gives a probability density map of where an electron has a certain statistical likelihood of being at any given instant in time.

24. Electronic Structure of Atoms Quantum Numbers Solving the wave equation gives a set of wave functions, or orbitals, and their corresponding energies. Each orbital describes a spatial distribution of electron density. An orbital is described by a set of three quantum numbers.

25. Electronic Structure of Atoms Principal Quantum Number, n The principal quantum number, n, describes the energy level on which the orbital resides. The values of n are integers ≥ 0.

26. Electronic Structure of Atoms Azimuthal Quantum Number, l This quantum number defines the shape of the orbital. Allowed values of l are integers ranging from 0 to n − 1. We use letter designations to communicate the different values of l and, therefore, the shapes and types of orbitals.

27. Electronic Structure of Atoms Azimuthal Quantum Number, l Value of l0123 Type of orbitalspdf

28. Electronic Structure of Atoms Magnetic Quantum Number, m l Describes the three-dimensional orientation of the orbital. Values are integers ranging from -l to l: −l ≤ m l ≤ l. Therefore, on any given energy level, there can be up to 1 s orbital, 3 p orbitals, 5 d orbitals, 7 f orbitals, etc.

29. Electronic Structure of Atoms Magnetic Quantum Number, m l Orbitals with the same value of n form a shell. Different orbital types within a shell are subshells.

30. Electronic Structure of Atoms Level n 1 2 3 Sublevel l Orbital m l Spin m s 00 00 1 0 0 1 0 21 0 -2 2101 = +1/2 = -1/2 Allowed Sets of Quantum Numbers for Electrons in Atoms

31. Electronic Structure of Atoms s Orbitals Value of l = 0. Spherical in shape. Radius of sphere increases with increasing value of n.

32. Electronic Structure of Atoms s Orbitals Observing a graph of probabilities of finding an electron versus distance from the nucleus, we see that s orbitals possess n−1 nodes, or regions where there is 0 probability of finding an electron.

33. Electronic Structure of Atoms p Orbitals Value of l = 1. Have two lobes with a node between them.

34. Electronic Structure of Atoms d Orbitals Value of l is 2. Four of the five orbitals have 4 lobes; the other resembles a p orbital with a doughnut around the center.

35. Electronic Structure of Atoms Energies of Orbitals For a one-electron hydrogen atom, orbitals on the same energy level have the same energy. That is, they are degenerate.

36. Electronic Structure of Atoms Energies of Orbitals As the number of electrons increases, though, so does the repulsion between them. Therefore, in many- electron atoms, orbitals on the same energy level are no longer degenerate.

37. Electronic Structure of Atoms Spin Quantum Number, m s In the 1920s, it was discovered that two electrons in the same orbital do not have exactly the same energy. The “spin” of an electron describes its magnetic field, which affects its energy.

38. Electronic Structure of Atoms Spin Quantum Number, m s This led to a fourth quantum number, the spin quantum number, m s. The spin quantum number has only 2 allowed values: +1/2 and −1/2.

39. Electronic Structure of Atoms Pauli Exclusion Principle No two electrons in the same atom can have exactly the same energy. For example, no two electrons in the same atom can have identical sets of quantum numbers.

40. Electronic Structure of Atoms Go To www.ptable.com www.ptable.com IMPORTANT  Use periodic Table to help you write electron configurations of atoms (and ions) Dynamic Periodic Table and Investigate (play with) the Orbitals option (on Top Tabs) for quantum numbers, orbitals and electron configurations of various elements Go To Glencoe site for animations on electron configuration http://glencoe.mcgraw- hill.com/sites/0023654666/student_ view0/chapter7/animations_center. html#

41. Electronic Structure of Atoms Electron Configurations Distribution of all electrons in an atom Consist of  Number denoting the energy level

42. Electronic Structure of Atoms Electron Configurations Distribution of all electrons in an atom Consist of  Number denoting the energy level  Letter denoting the type of orbital

43. Electronic Structure of Atoms Electron Configurations Distribution of all electrons in an atom. Consist of  Number denoting the energy level.  Letter denoting the type of orbital.  Superscript denoting the number of electrons in those orbitals.

44. Electronic Structure of Atoms Orbital Diagrams Each box represents one orbital. Half-arrows represent the electrons. The direction of the arrow represents the spin of the electron.

45. Electronic Structure of Atoms Hund’s Rule “For degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized.”

46. Electronic Structure of Atoms Periodic Table We fill orbitals in increasing order of energy. Different blocks on the periodic table, then correspond to different types of orbitals.

47. Electronic Structure of Atoms Some Anomalies Some irregularities occur when there are enough electrons to half- fill s and d orbitals on a given row.

48. Electronic Structure of Atoms Some Anomalies For instance, the electron configuration for chromium is [Ar] 4s 1 3d 5 rather than the expected [Ar] 4s 2 3d 4.

49. Electronic Structure of Atoms Some Anomalies This occurs because the 4s and 3d orbitals are very close in energy. These anomalies occur in f-block atoms, as well.

50. Electronic Structure of Atoms ELECTRON SPIN 1920--chemists realized that since electrons interact with a magnetic field, there must be one more concept to explain the behavior of electrons in atoms. m s --the 4th quantum number; accounts for the reaction of electrons in a magnetic field MAGNETISM magnetite--Fe 3 O 4, natural magnetic oxide of iron 1600--William Gilbert concluded the earth is also a large spherical magnet with magnetic south at the north pole (Santa's habitat). NEVER FORGET: opposites attract & likes repel PARAMAGNETISM AND UNPAIRED ELECTRONS diamagnetic--not magnetic [magnetism dies]; in fact they are slightly repelled. All electrons are PAIRED. paramagnetic--attracted to a magnetic field; lose their magnetism when removed from the magnetic field; HAS ONE OR MORE UNPAIRED ELECTRONS ferromagnetic--retain magnetism upon introduction to, then removal from a magnetic field All of these are explained by electron spins Each electron has a magnetic field with N & S poles electron spin is quantized such that, in an external magnetic field, only two orientations of the electron magnet and its spin are possible +/- 1/2 H is paramagnetic; He is diamagnetic, WHY? H has one unpaired electron He has NO unpaired electrons; all spins offset and cancel each other out (Taken from summary notes posted on M. Brophy’s website)

51. Electronic Structure of Atoms What about ferromagnetic? clusters of atoms have their unpaired electrons aligned within a cluster, clusters are more or less aligned and substance acts as a magnet. Don't drop it!! When all of the domains, represented by these arrows are aligned, it behaves as a magnet. This is what happens if you drop it! The domains go indifferent directions and it no longer operates as a magnet. (Taken from summary notes posted on M. Brophy’s website)

52. Electronic Structure of Atoms Activities and Problem set __ TB ch. 6 – all sections required for SAT II and AP exams and most are required for regents exam View and take notes on the recommended animations POGIL activities on (1) Analysis of Spectral Lines and (2) Interaction of Radiation and Matter Online practice quiz and work due by ______ Ch 6 Problems: write out questions (or photocopy them) ; write out answers & show work First carefully study the sample exercises in chapter 6 (you don’t have to copy them out) and then DO all in- chapter practice exercises according to the directions above. Do all GIST, and Visualizing concepts, problems end of chapter 6 exercises: To be assigned in class