1. Big Idea #1 Structure of the Atom

2. John Dalton (1766 – 1844) Proposed the first scientifically supported atomic theory.

3. Dalton’s Model of the Atom Dalton's model was that the atoms were tiny, indivisible, indestructible particles and that each one had a certain mass, size, and chemical behavior that was determined by what kind of element they were. What made us change Dalton’s Atomic Model?

4. J. J. Thomson (1856 – 1940) Credited for the discovery of the electron. Invented the mass spectrometer which led to his discovery of isotopes.

5. Scientific Inquiry Scientists use experimental results to test scientific models, such as the model of the atom. When experimental results are not consistent with the predictions of a scientific model, the model must be revised or replaced.

6. Lord Ernest Rutherford (1871 – 1937) Discovered the nucleus of the atom.

7. Rutherford’s Gold Foil Experiment

9. Rutherford’s Atomic Model The “Planetary Model” of the Atom The nucleus is very small, dense, and positively charged. Electrons surround the nucleus. Most of the atom is empty space

10. Subatomic Particles PARTICLESYMBOLCHARGEMASS (amu) LOCATION electrone-e- 00 orbit nucleus protonp+p+ +1 11 inside nucleus neutronn0n0 0 11 inside nucleus

11. Niels Bohr In 1913 Bohr published a theory about the structure of the atom based on an earlier theory of Rutherford's. Bohr expanded upon this theory by proposing that electrons travel only in certain successively larger orbits.

12. Bohr Model of the Atom Electrons orbit the nucleus in orbits that represent specific quantities of energy. The energies of the electrons in the atom are quantized. Only certain electron orbits (energy levels) are allowed. The Bohr Atom

14. Electromagnetic Waves Properties of waves include speed, frequency, wavelength and energy All electromagnetic waves including light travel at a speed of 3 x 10 8 m/s. However the frequency, wavelength and energy of the waves vary.

15. ) Wavelength ( ) Measured in units of length: m, nm, A º

16. ) Frequency ( ) Measured in cycles/second = hertz (Hz)

18. Visible Light

19. For all waves = cFor all waves = c c = the speed of light = 3.00 x 10 8 m/s Electromagnetic Radiation

20. A photon of red light has a wavelength of 665 nm. What is the frequency of this light? An x-ray has a frequency of 7.25 x 10 20 Hz. What is the wavelength?

21. For all waves: E = hFor all waves: E = h h = 6.63 x 10 -34 J s This is known as Planck’s constant. This is known as Planck’s constant. Energy of Electromagnetic Radiation

22. A photon of red light has a wavelength of 665 nm. What is the energy of this light? An x-ray has a frequency of 7.25 x 10 20 Hz. What is it’s energy?

23. Wavelength, frequency and energy Wavelength and frequency have an indirect relationship. Energy and frequency have a direct relationship. Electromagnetic radiation of short wavelength will have high frequency and high energy. Electromagnetic radiation of long wavelength will have low frequency and low energy.

24. Niels Bohr Bohr also described the way atoms emit radiation by suggesting that when an electron jumps from an outer orbit to an inner one, that it emits light.

25. Bohr Model Electrons move around the nucleus in orbits of definite energies. The energy of the orbit is related to its distance from the nucleus. The lowest energy is found in the orbit closest to the nucleus. Radiation is absorbed or emitted when an electron moves from one orbit to another.

26. Ground State The lowest energy state of an atom.

27. Excited State Any energy state of an atom that is of higher in energy than the ground state.

28. Energy Absorbed

29. Absorption (Dark – Line) Spectra

30. Energy Emitted Electron jumps to a lower orbit

31. What if we wanted to calculate the energy of electron transitions? E = - 2.178 x 10 -18 J (n 2 ) Where n is the principle energy level. So  E = E final – E initial Calculate the energy required to move an electron from n=1 to n=2. Calculate the wavelength of light that must be absorbed to make this transition.

32. Emission (Bright – Line) Spectra

33. Emission Spectra

34. The lines present in an emission spectrum are the lines missing in an absorption spectrum.

35. The Heisenberg Uncertainty Principle

36. The Uncertainty principle Heisenberg determined that it was impossible to experimentally determine both the position and the speed of the electron at the same time. This became known as the Heisenberg Uncertainty Principle. It simply means that the electron is so small and moving so fast, that the simple act of trying to measure its speed or position would change either quantity.

37. Problems with the Bohr Model It violates the Heisenberg Uncertainty Principle because it considers electrons to have known orbits. It makes poor predictions regarding the spectra of atoms larger than hydrogen.

38. Schrodinger The Austrian scientist, Erwin Schrödinger, pursued the idea of the electron having wave properties and it seemed to him that the electron might be like a standing wave around the nucleus.

39. Schrodinger’s Model This idea agreed very well with Bohr's idea of quantized energy levels: only certain energies and therefore, wavelengths would be allowed in the atom. This explained why only certain colors (wavelengths) were seen in the spectrum of the hydrogen atom.

40. Schrodinger’s Model Schrodinger set out to make a mathematical model that assumed the electron was a standing wave around the nucleus. His solutions to that model agreed not only with the experimental evidence for hydrogen (as Bohr’s did too), but gave excellent results for all atoms when compared to their actual spectrum.

41. Schrodinger’s Model This is our modern model of the atom and is known as the Quantum Mechanical Model. Calculations involving the QM model of the atom can be approximated using computers. The solutions to these calculations is the basis for modern software that calculates the structure and reactivity of molecules.

42. The Quantum Mechanical Model The quantum mechanical model is based on quantum theory, which says matter also has properties associated with waves.

43. The Quantum Mechanical Model According to quantum theory, it’s impossible to know the exact position and momentum of an electron at the same time. This is known as the Uncertainty Principle.

44. The Quantum Mechanical Model The quantum mechanical model of the atom uses complex shapes of orbitals (sometimes called electron clouds), volumes of space where an electron is likely to be found.

45. The Quantum Mechanical Model Therefore the quantum mechanical model is based on probability rather than certainty.

46. Absorption and Emission of Energy by Molecules A photon is a particle representing a specific amount of electromagnetic energy. When a photon is absorbed by a molecule, the energy of the molecule is increased. When a photon is released by a molecule, the energy of the molecule is decreased.

47. Absorption and Emission of Energy by Molecules Different types of molecular motion lead to absorption or emission of photons in different spectral regions.

48. Principle Energy Levels – Sublevels - Orbitals Electrons in an atom are within atomic orbitals which are within sublevels which are within energy levels. Principle Energy Level (VA) – 1-7 Sublevels (Manassas) – s,p,d,f Orbitals (8909 Euclid Ave)

50. © 1998 by Harcourt Brace & Company s p d (period-1) f (period-2) 6767 Periodic Table and Electron Configuration 12345671234567 6767

51. A. General Rules Pauli Exclusion Principle –Each orbital can hold TWO electrons with opposite spins.

52. Correct A. General Rules Hund’s Rule –Within a sublevel, place one e - per orbital before pairing them. –All electrons in singly filled orbitals have the same direction of spin. Incorrect

53. A. General Rules Aufbau Principle –Electrons always enter the lowest energy location first –This is why every EC from He on starts 1s 2

54. Magnetism Magnetism can be a complicated concept. We will only deal with magnetism as it is predicted by electron spin which is an oversimplified way of dealing with this subject.

55. Electron Spin & Magnetism Opposite spins produce opposite magnetic fields.

56. Electron Spin & Magnetism Most materials with one or more unpaired electrons are at least slightly magnetic. These substances are said to be paramagnetic. The overall magnetic field strength of atoms with all paired electrons is zero. These substances are said to be diamagnetic.

57. Photoelectric Effect The photoelectric effect is the emission of electrons from substances that are exposed to light. Ionization energy is the amount of energy it takes to remove an electron from an atom.

58. Photoelectric Effect More energy is required to remove successive electrons from atoms. This is due to Coulomb’s Law.

59. Coulomb’s Law Coulomb’s Law quantifies the general rule of electrostatics that opposites charges attract and like charges repel. The electrostatic force between two charged bodies is proportional to the product of the amount of charge on the bodies divided by the square of the distance between them.

60. Coulomb’s Law F= force of attraction or repulsion k = constant q 1 and q 2 = the charges of the two bodies r = radius between the charges

61. Great Video integrating mulitple topics www.bozemannscience.com/ap-chem- 004-coulombs.www.bozemannscience.com/ap-chem- 004-coulombs The link will be on the wiki page as well.

62. Photoelectric Effect E = hLight consists of photons of a certain energy (E = h ). In the photoelectric effect, light ejects electrons from a material. This requires the photon to have sufficient energy to eject the electron. This is known as Photoelectron Spectroscopy (PES) PES data give support to the electron configurations that we use to place electrons within atoms.

63. Photoelectron Spectroscopy (PES) Photoelectron Spectroscopy (PES) determines the energy needed to eject electrons from a substance. Photoelectron Spectroscopy (PES) supports the shell, subshell model of the atom. PES data for multi-electron atoms show that certain electrons have ionization energies that are relatively close to one another. We therefore group these electrons into shells and subshells as a result.

64. Photoelectron Spectroscopy (PES) Photoelectron Spectroscopy (PES) data for atoms also helps us to confirm the number of electrons within particular shells and subshells. The intensity (shown as the height on a graph) of the photoemission signal at a given energy is a measure of the number of electrons in that shell or subshell. In other words the height of the PES signal is proportional to the number of electrons in that sublevel.

70. http://www.chem.arizona.edu/ch emt/Flash/photoelectron.html Go to the website above and go through the first 10 elements: Show different sublevels and different numbers of electrons in each sublevel based on peak intensity (height). If you choose dual mode you can compare the PES spectra of two elements.