3. 1. The earliest models http://www.youtube.com/watch?v=BhWgv0STLZs&feature=player_embedded 2. Smaller than the Smallest http://www.youtube.com/watch?v=WmmglVNl9OQ&feature=player_embedded 3. The Rutherford Model http://www.youtube.com/watch?v=FfY4R5mkMY8&feature=player_embedded 4. The Bohr Model http://www.youtube.com/watch?v=hpKhjKrBn9s&feature=player_embedded 5. Spectra http://www.youtube.com/watch?v=5z2ZfYVzefs&feature=player_embedded 6. Wave Mechanics http://www.youtube.com/watch?v=1bpG1lEjJfY&feature=player_embedded Take notes over each video. Check calendar for due date.

5.  atomic weight - weighted average of the masses of its constituent isotopes  Ex 2) Naturally occurring chromium consists of four isotopes. It is 4.31% 24 50 Cr, mass = 49.946 amu, 83.76% 24 52 Cr, mass = 51.941 amu, 9.55% 24 53 Cr, mass = 52.941 amu, and 2.38% 24 54 Cr, mass = 53.939 amu. Calculate the atomic weight of chromium.

6.  Naturally occurring Lithium exists as two isotopes, 6 Li (mass = 6.015 amu) and 7 Li (mass = 7.016 amu). The atomic weight is 6.941 amu. Which isotope is more abundant? Why?  Ex 3) The atomic number of boron is 10.811 amu. The masses of the two naturally occurring isotopes 5 10 B and 5 11 B, are 10.013 and 11.009 amu, respectively. Which isotope is most common? Calculate the fraction and percentage of each isotope. › requires a little algebra › remember X + (1-X) = 1

8.  Mass spectrometry is an analytical technique that measures the mass-to-charge ratio of charged particles. It is most generally used to find the composition of a sample by generating a mass spectrum representing the masses of sample components. The mass spectrum is measured by a mass spectrometer. http://www.youtube.com/watch?v=J- wao0O0_qM

9.  An analytical technique for identification of chemical structures, determination of mixtures, and quantitative elemental analysis, based on application of the mass spectrometer.  Francis Aston - devised first mass spectrograph

10. › In a mass spectograph ions pass down an evacuated path inside a magnet Four factors which determine particle’s path in mass spectrometer 1. accelerating voltage 2. magnetic field strength 3. masses of particles 4. charge on particles

11.  Mass spectrometers consist of three basic parts: an ion source, a mass analyzer, and a detector system. The stages within the mass spectrometer are: › Production of ions from the sample › Separation of ions with different masses › Detection of the number of ions of each mass produced › Collection of data to generate the mass spectrum

12. This technique is applicable in:  identifying unknown compounds by the mass of the compound molecules or their fragments  determining the isotopic composition of elements in a compound  determining the structure of a compound by observing its fragmentation  quantifying the amount of a compound in a sample using carefully designed methods  studying the fundamentals of gas phase ion chemistry (the chemistry of ions and neutrals in vacuum)  determining other important physical, chemical, or even biological properties of compounds with a variety of other approaches

13. Since different chemicals have different masses, this info is used to determine what chemicals are present in a sample. For example, table salt:  NaCl, may be vaporized, turned into gas and ionized into ions, electrically charged particles (Na + and Cl - ), in the first phase of the mass spectrometry.

14.  The sodium ions are monoisotopic, with mass 23 u. Chloride ions have two isotopes of mass 35 u (~75%) and mass 37 u (~25%). Since they have a charge, the speed and direction may be changed with an electric or magnetic field.  An electric field accelerates the ions to a high speed. Next they are directed into a magnetic field which applies a force to each ion perpendicular to the plane of the particles' direction of travel and the magnetic field lines.

15.  This force deflects the ions (makes them curve instead of traveling in a straight line) to varying degrees depending on their mass-to-charge ratio. Lighter ions get deflected more than the heavier ions. This is due to Newton’s 2 nd law of Motion. The acceleration of a particle is inversely proportional to its mass.  Therefore, the magnetic field deflects the lighter ions more than it does the heavier ions.

16.  The detector measures the deflection of each resulting ion beam.  From this measurement, the mass-to- charge ratios of all the ions produced in the source can be determined. So the chemical composition of the original sample (both sodium and chlorine are present in the sample) and the isotopic compositions of its constituents (the ratio of 35 Cl to 37 Cl atoms) can be determined.

17. relationship for electromagnetic radiation c =  c = velocity of light - 3.00 x 10 8 m/s

18. Ex. 5) What is the frequency of green light of wavelength 5200 Å ?  Max Planck - 1900 › energy is quantized › light has particle character  Planck’s equation Ex. 6) What is energy of a photon of green light with wavelength 5200 Å ?

20. Albert Einstein explained the photoelectric effect in 1905 › 1921 Nobel Prize in Physics  electrons are particle like b/c energy from a photon transfers to e - during collisions. If you increase energy, more electrons get kicked off. Each individual photon makes a spark, 1 e - per photon. The more intense the light, the more photons. Light strikes the surface of various metals and causes electrons to be ejected.

21. emission spectrum electric current passing through a gas in a vacuum tube (at very low pressure) causes the gas to emit light emission or bright line spectrum

22. absorption spectrum shining a beam of white light through a sample of gas gives an absorption spectrum shows the wavelengths of light that have been absorbed

23. use spectra to identify elements can even identify elements in stars

24. Light of a characteristic wavelength (& frequency) is emitted when electron falls from higher E (orbit) to lower E (orbit) origin of emission spectrum light of a characteristic wavelength (& frequency) is absorbed when electron jumps from lower E (orbit) to higher E (orbit) origin of absorption spectrum

25.  how atoms talk to us” › we have to interpret their language  Bohr, Schrodinger, and Heisenberg were some of the first scientists to translate the language of atoms Ex. 7) An orange line of wavelength 5890 Å is observed in the emission spectrum of sodium. What is the energy of one photon of this orange light?

26.  Werner Heisenberg - 1927 › Uncertainty Principle  It is impossible to determine simultaneously both the position & momentum of an electron. Why? The act of measuring a very small particle changes its position, so it is impossible to precisely determine both the position and momentum of that object. › electron microscopes use this phenomenon › devices for detecting motion of electron disturbs its position

28.  There are four quantum numbers which describe the relative position and energy of an electron in an atom. 1. Principal quantum number 2. Angular momentum quantum number 3. Magnetic quantum number 4. Spin quantum number

29.  Symbol “n” – refers to energy level n = 1, 2, 3, …   The principal quantum number describes the relative distance from the nucleus. It is often referred to as Energy Level or Shell.  electron’s energy depends principally on n

30.  angular momentum, tells shape of the atomic orbitals  Atomic Orbitals - regions of space where the probability of finding an electron around an atom is highest ~ volume that the electrons occupy 90-95% of the time  symbol ℓ to find ℓ plug into n-1 n=1 n=2 n=3 …. n=8 ℓ =0 ℓ = 0, 1 ℓ = 0, 1, 2 ℓ = 0, 1, 2, 3, 4, 5, 6, 7

31.  represents the sublevels within an energy level: s, p, d, f (code letters for the shapes of orbitals) è s=0 è p=1 è d=2 è f=3  Quantum number - ℓ ℓ = 0, 1, 2, 3, 4, 5,.......(n-1) ℓ = s, p, d, f, g, h,.......(n-1)

32. u s orbitals are spherical in shape. u Every energy level has an s orbital  s orbitals have angular momentum quantum number ( l ) equal to 0.

33.  p orbitals are shaped like dumbbells or peanuts.  They are oriented along the x, y, and z coordinates.  They have an angular momentum quantum number ( l ) equal to 1. 3 per n level, p x, p y, p z l = 1

34. p orbitals

35.  start with n = 3  4 clover leaf shaped and 1 peanut shaped with a doughnut around it on Cartesian axes and rotated 45 o  5 per n level l = 2 m l = -2,-1,0,+1,+2 5 values of m l m l = Magnetic quantum number (info in a few slides)

37.  start with n = 4  most complex shaped orbitals  7 per n level, complicated names l = 3  m l = -3,-2,-1,0,+1,+2, +3 7 values of m l  important effects in lanthanides & actinides

39.  3 rd quantum number  symbol m l  Helps tell orientation of orbitals m l = - l to + l

40. l =0 m l =0 › only 1 value s orbital l =1 m l = -1, 0, 1 › 3 values p orbitals l =2 m l = -2, -1, 0, 1, 2 › 5 values d orbitals l =3 m l = -3, -2, -1, 0, 1, 2, 3 › 7 values f orbitals

41.  4 th quantum number, symbol = m s › m s = +1/2 or -1/2  tells us the spin and orientation of the magnetic field of the electrons

42.  spin effects › every orbital can hold up to two electrons › one spin up  one spin down   spin describes the direction of their magnetic field  e- have charges two allowable magnetic states

43.  paired electrons have spins unaligned  › no net magnetic field  diamagnetic - repelled by a magnetic field, all electrons are paired  unpaired electrons have their spins aligned  or  › enhanced magnetic field  paramagnetic - attracted to a magnetic field, has unpaired electrons

44. 1. Always fill orbitals in lowest energy level first. (Aufbau Principle) Aufbau Principle - The electron that distinguishes an element from the previous element enters the lowest energy atomic orbital available. (in other words fill one energy sublevel before moving up) 2. No two e- can have same 4 quantum numbers in an atom. (Pauli Exclusion Principle) Wolfgang Pauli - 1925 › No two electrons in an atom may have identical sets of 4 quantum numbers.

45. 3. Spread e- out on a sublevel if possible. (Hund’s Rule) Electrons will spread themselves out among the orbitals individually and give unpaired, parallel spins. The pairing of electrons is an unfavorable process; energy must be expended in order to make it occur. Exception: If you can achieve full or half-full orbitals by moving one e- between s ~ d or s ~ f orbitals, do so. It’s lower in energy because there is an increased stability due to the decrease in the screening of electron/nuclear attractions

48. 1 st row Orbital Order 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p 8s

49. 2 nd row

50. 3 rd row

51. 4 th row

52. 4 th row continued

54.  Let’s find the complete set of quantum numbers for the electrons in Na and Fe (look at electron configurations/orbital diagrams 1 st )  11 Na › must have one set of 4 quantum numbers for each of the 11 electrons in Na  26 Fe › should have one set of 4 quantum numbers for each of the 26 electrons in Fe

55.  Ground State – lowest energy/most stable state of an atom, molecule, or ion. Fills one sublevel before moving up, follows Aufbau Principle.  Excited State – orbitals skipped, does not follow Aufbau Principle  Forbidden State – very wrong or not possible

56.  Isoelectronic – different elements that have the same electron configuration b/c of gaining or losing electrons.  Shielding effect - negatively charged electrons prevents higher orbital electrons from experiencing the full nuclear charge by the repelling effect of inner-layer electrons  Effective nuclear charge - the net positive charge experienced by an electron in a multi-electron atom

57.  Representative elements: group A (the last electron fills s or p levels)  IA – alkali metals  IIA – alkaline earth metals  IIIA – boron famly  IVA – carbon family  VA – pnictogens or nitrogen family  VIA – chalcogens or oxygen family  VIIA – halogens  VIIIA – noble gases

58.  Transition metals (filling d level)  3 - the scandium family  4 - the titanium family  5 - the vanadium family  6 - the chromium family  7 - the manganese family  8 - the iron family  9 - the cobalt family  10 - the nickel family  11 - the coinage or copper family  12 - the zinc family  Lanthanide series (filling 4f level)  Actinide series (filling 5f level)

59. Extra credit 1) In a universe far far away, the laws of quantum mechanics are the same as ours with one small change. Electrons in this universe have three spin states, -1, 0, and +1, rather than the two, +1/2 and -1/2, that we have. What two elements in this universe would be the first and second noble gases? (Assume that the elements in this different universe have the same symbols as in ours.)

60. 2) A) What is the atomic number of the element that should theoretically be below Ra? B) Its chemical behavior would be most similar to which elements? C) How many valence electrons would it have? D) Its electron configuration would be? E) An acceptable set of 4 quantum numbers for the last electron in this element would be? F) What would you name it?