1. Unit 2: Atomic Theory

2. Video links overview of atomic history http://www.youtube.com/watch?feature =player_detailpage&v=k1RHY8QcN1s

3. I. Atomic History A. The Greeks Democritus [Philosopher] All matter is made of tiny, indivisible parts called ‘atoms’ Developed word ‘atomos’ meaning not divisible

4. John Dalton (1803-1808) Used experiments with gases to develop the “Atomic Theory” Determined atoms looked like ‘cannonballs’ or solid masses

5. Dalton’s Atomic Theory 1) All elements are made of atoms 2) Atoms of each element are all the same, or have the same masses 3) Atoms of different elements are different, or have different masses 4) Atoms cannot be created or destroyed 5) Atoms combine in small, whole number ratios

7. J.J. Thomson (1897) Developed ‘Cathode Ray’ experiment Said atoms consisted of particles smaller than an entire atom Discovered that the smaller particles within an atom had a negative charge Discovered 1 st subatomic particle: Electron Founded “Plum Pudding Model”: Electrons were embedded within a positively charged mass

9. Cathode Ray Tube Experiment Thomson manipulated cathode rays with a magnet to discover that subatomic particles existed and that they had negative charges

10. Ernest Rutherford (1898) Discovered alpha and beta radiation emitted from certain radioactive substances Developed and used Gold Foil Experiment First to separate the smaller parts of the atom Discovered the nucleus Placed electrons outside the nucleus Stated that atoms are composed of lots of empty space

12. Rutherford’s Gold Foil Experiment micro.magnet.fsu.edu/electromag/java/rutherford/

13. Niels Bohr (1922) Bohr analyzed work of others and studied atomic spectra, or light, given off by the elements Described the “Atomic Spectra” of elements Developed ‘Solar System’ model Moved electrons from single, giant pathway into discrete energy levels around the nucleus Each energy level contained 2, 8, 18, 32, etc. electrons total

14. Bohr Model of the Atom Stated that electrons moved around the nucleus in ‘orbits’ or energy levels As electrons gain energy, they jump up energy levels, then release this energy to generate spectra

15. Bohr’s model and the atomic spectrum http://jersey.uoregon.edu/vlab/elements/Elements.html The spectral lines in the visible region of the atomic emission spectrum of barium are shown below. Spectral lines exist in series in the different regions (infra-red, visible and ultra-violet) of the spectrum of electromagnetic radiation. The spectral lines in a series get closer together with increasing frequency. Each element has its own unique atomic emission spectrum.

16. Erwin Schrodinger (1930) Developed mathematical equations representing electrons Electrons had wave and particle behaviors Created “Wave-Mechanical” or “Modern” model Most scientists use this model today Placed electrons in orbitals

17. “Electron Cloud” Model Created paths for electrons within Bohr’s energy levels Only 2 electrons per path ORBITALS Electron paths, or ORBITALS, are mathematical equations describing probability densities for electrons Developed sublevels with discrete paths within each energy level

18. II. Subatomic Particles A. Particles Protons 1) Protons found in the nucleus of an atom charge of +1, mass of 1.0073a.m.u. Neutrons 2) Neutrons found in the nucleus of an atom no charge, mass of 1.0087a.m.u.

19. A. Subatomic Particles 3) Electrons  Found outside the nucleus in regions of probability [orbitals]  Charge of –1, mass of 5.46 x 10 -4 a.m.u., or 1/1836 a.m.u.  Have particle and wave properties

20. B. Atomic Number Atomic number Atomic number = the number of protons in the nucleus All atoms of the same element have the same atomic number Atoms arranged on PT by increasing atomic numbers In neutral atoms: Atomic number equals number of electrons

21. C. Isotopes Isotopes Isotopes = atoms of the same element that have differing numbers of neutrons in their nucleus, different mass number, but same atomic number Same number of protons!!! Changing number of neutrons affects properties [radioactivity…]

22. D. Atomic Mass Atomic Mass Number Atomic Mass Number = number of protons plus the number of neutrons in the nucleus Whole number!! Mass Number changes when using different isotopes Written in isotopic notations, just subtract the top from bottom values:

23. E. Ions Ions Ions = atoms of the same element that have lost or gained electrons Have overall (+) or (-) charge Same numbers of protons, number of neutrons irrelevant Positive ions: have LOST electrons Negative ions: have GAINED electrons

24. F. Atomic Mass (average) Atomic Mass Atomic Mass = weighted average of the natural isotopes times their percent abundance Decimal value on PT Accounts for the natural existence of various isotopes Ex] calculate the atomic mass of carbon given that 98.92% is carbon-12 and 1.108% is carbon-13

25. Virtual textbook http://www.chem1.com/acad/web text/intro/int-1.html#SEC1

26. III. Electronic Structure A. EMS [Electromagnetic Spectrum]A. EMS [Electromagnetic Spectrum]

27. A. EMS [Electromagnetic Spectrum] EMS EMS = continuous series of various types of energy, separated by their wavelengths and frequencies Visible light Visible light = small portion; only part we can see without instruments Continuous spectrum Continuous spectrum = picture of all colors of visible light as they pass through a prism

28. EMS continued Wavelength = distance between 2 peaks or troughs of 2 consecutive waves Symbol = λ [Greek letter “lambda”] Units are usually in ‘m’ or ‘nm’ Frequency = the number of peaks or troughs that pass a single point in one second Symbol = ʋ [Greek letter “nu”] Units are usually in ‘1/s’ or ‘s -1’ or ‘Hz’

29. Calculations using lambda and nu c = λν C = speed of light C = 3.0 x 10 +8 m/s E = h ν E = energy of photon h = Planck’s constant h = 6.63 x 10 -34 Js All electromagnetic radiation travels at the speed of light Can calculate the energy of the radiation/electron given the wavelength

30. Planck’s Constant Planck observed hot, glowing matter Concluded: different substances glow different colors at different temperatures Determined: matter releases energy in tiny, discrete packets called ‘quanta’ Developed constant to relate energy and temperature, Planck’s constant, “h” h = 6.63 x 10 -34 J*s

31. Light traveling as waves All colors of light energy travel at the same speed, just different wavelengths!

32. Particle vs. Wave Behavior of Light

33. Wave behavior of light

34. B. Photoelectric Effect Einstein used Planck’s idea of quanta and photons to describe the photoelectric effect Light of a certain wavelength shines on clean metal, causing the metal to eject electrons

35. C.Bohr’s Model [conclusions made] Bohr used the idea of ‘quanta’ to explain the bright-line emission spectra Stated that each element’s atomic spectrum is unique Electrons exist in ground state energy levels, as listed via the periodic table

36. Bohr Model of the Atom Stated that electrons moved around the nucleus in energy levels Electrons will gain and lose energy at will This generated the element’s atomic spectrum

37. Bohr’s model

38. Useful Websites and References //www.avogadro.co.uk/light/bohr/spectra.htm shows formation of spectral lines for hydrogen idea of ground vs. excited state //jersey.uoregon.edu/vlab/elements/Elements.html Periodic table showing the absorption and emission spectra for each element Also check out Wikipedia under Bohr atom and Atomic spectra!

39. Creation of an emission spectrum absorb If electrons absorb packets of energy, quanta, they temporarily move to into a higher energy level, called the excited state release The electrons then release this quanta of energy and fall back down to ground state The release of energy generates the bright-line emission spectrum

40. Examples of Bohr Diagrams

41. IV. Electron Configurations A. Energy Levels These are areas with a high possibility of finding electrons with similar potential energies 7 energy levels total

42. Bohr Diagrams and Energy Levels Bohr Diagrams show the numbers of protons and neutrons in the nucleus Shows electrons in their respective energy levels Energy levels hold: 1 st holds 2 electrons 2 nd holds 8 electrons 3 rd holds 18 electrons 4 th holds 32 electrons Etc…..

43. B. Sublevels Sublevels are divisions within each energy level Represent the shapes and orientation in 3D space not good! Too many electrons within the energy levels & they lose momentum and will crash into the nucleus--- not good! 1 st energy level has 1 sublevel: “s” 2 nd has 2 sublevels: “s” and “p” 3 rd has 3 sublevels: “s, p, and d” 4 th has 4 sublevels: “s, p, d, and f”

44. Sublevels and Shapes “s” is spherical and has a max of 2 electrons “p” is dumbbell shaped and has a max of 6 electrons “d” is cloverleaf shaped and holds up to 10 electrons “f” is a split cloverleaf with a max of 14 electrons http://micro.magnet.fsu.edu/electromag/java/at omicorbitals/index.html

45. Order of Sublevel Filling It does not go in order… 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 4f 14 5s 2 5p 6 5d 10 5f 14 6S 2 6P 6 6d 10 7s 2 7p 6

46. Orbitals within Sublevels Each sublevel consists of 1 to 7 orbitals [areas of probability for finding an electron] Each path or orbital only holds 2 electrons The 2 electrons within in each orbital each have a different spin This allows the electrons to exist in the same area without conflicting

47. C.Extended and Abbreviated Configurations Electron Configurations = way to describe how the electrons are distributed around an atom and within the energy levels and sublevels Ground state configurations are same order as electrons on PT Excited state configurations have one electron shifted to a higher energy level

48. Writing Electron Configurations Electrons add in the same order as the atomic numbers of the PT Aufbau Principle = adding electrons in the exact order of the PT

49. Writing Configurations Neutral, Ground State atoms Add in order of arrows for Neutral, Ground State atomsExamples:

50. Abbreviated Configurations Abbreviated configurations Abbreviated configurations show only the placement of electrons added after the last ‘noble gas’ Ex] Bracket the configuration of the last noble gas [group 18] and add remaining electrons Ex]

51. D.Orbital Notations and Rules Orbital notations Orbital notations are specialized versions of a full electron configuration showing the spin of each electron within an orbital Draw the orbitals present for each sublevel and fill with ‘spin-paired’ electrons

52. Rules for Configurations Hund’s Rule 1. Hund’s Rule = electrons in the p, d, and f sublevels must be added to each orbital first, before one flips to spin-pair and fill the orbital Pauli Exclusion Principle 2. Pauli Exclusion Principle = no 2 electrons may be in the same orbital and have the same spin; no 2 electrons will have the same 4 quantum numbers

53. Rules cont’ 3. Heisenberg’s Uncertainty Principle 3. Heisenberg’s Uncertainty Principle = states that the electron’s momentum and position cannot be accurately determined at the same time Example…

54. Excited State vs. Ground State Excited State Excited State configurations show one electron has moved into a higher energy level, leaving an unfilled space below Ground state configurations are written in order of the periodic table **Total # of electrons = Atomic # for for BOTH !!

55. E. Lewis Dot Structures Lewis Dot Structures are pictures showing the placement and number of valence electrons for an element Structure: s 1 s 2 p 6 p 1 p 3 p 4 p 5 p 2 Valence electrons are s and p outer shell electrons Maximum of 8! Ex]

56. F. Quantum Numbers Each electron in an atom is assigned a set of 4 quantum numbers These numbers tell the exact “address” of an electron, regardless of the element No 2 electrons have the same 4 quantum numbers!

57. Quantum Numbers 1] Principle Quantum Number (n) First number Represents the energy level of the electron Values range from 1 to 7 2] Azimuthal Spin Number (l) Second Number Represents the sublevel Describes the shape of the orbital Values from 0 to 3

58. Quantum Numbers Magnetic Spin Number (m l ) 3] Magnetic Spin Number (m l ) Tells the orientation of the orbital along x, y, z axes Values for: l = 0, m l = 0 l = 1, m l = +1, 0, -1 l = 2, m l = +2, +1, 0, -1, -2 l = 3, m l =+3, +2, +1, 0, -1, -2, -3 Spin Number (m s ) 4] Spin Number (m s ) Tells if the electrons spin clockwise, or counterclockwise Values: +1/2 [spin up] or –1/2 [spin down]