Chapter 10- Part One Modern Atomic Theory

1. Chapter 10- Part OneModern Atomic Theory Objectives: Review history… (10.1) Describe electromagnetic radiation (10.2) Describe the Bohr atom (10.3) Explain energy levels of electrons and diagram atomic structures for elements (10.4 & 10.5)

2. Review… • Dalton • Thomson • Rutherford • Model doesn’t explain how the negative electron can stay in orbit and not be attracted to the positive proton

3. Electromagnetic Radiation • Light travels in waves, similar to waves caused by a moving boat or a pebble tossed in a pond • Light is a form of Electromagnetic Radiation • Form of energy that exhibits wavelike behavior as it travels through space

4. Electromagnetic Radiation • All waves can be described in 3 ways: • Amplitude – the height of the wave, results in the brightness or intensity of the light • Wavelength (l): distance between consecutive peaks in a wave • Frequency (n): number of waves that pass a given point in a second

5. Electromagnetic Radiation • Speed of light in air: Electromagnetic radiation moves through a vacuum at speed of 3.00 x 108 m/s • Since light moves at constant speed there is a relationship between wavelength and frequency: c = ln Wavelength and frequency are inversely proportional

6. Electromagnetic Spectrum

7. Quantum Theory • Wave theory does not explain • Heated iron gives off heat • 1stred glow yellow glow white glow • How elements such as barium and strontium give rise to green and red colors when heated

8. Quantum Theory • Max Planck (1858-1947) • Proposed that there is a fundamental restriction on the amounts of energy that an object emits or absorbs, and he called each of these pieces of energy a quantum. • Energy is released in Quanta

9. Quantum Theory • A quantum is a finite quantity of energy that can be gained or lost by an atom E = hn E = energy v = frequency h = 6.626 x 10-34 J/s • This constant, h, is the same for all electromagnetic radiation

10. Photoelectric Effect • The emission of electrons by certain metals when light shines on them • Albert Einstein (1905) used Planck’s equation to explain this phenomenon; • proposed that light consists of quanta of energy that behave like tiny particles of light • Photon = individual quantum light (also known as a particle of radiation)

11. Photoelectric Effect • He (Einstein) explained that the photoelectric effect would not occur if the frequency and therefore the energy of each photon is too low to dislodge an electron. • Analogy: • 70 cents placed in soda machine: no soda • 30 cents more and you will get your soda

12. Now… • Light can be described as both particles and waves • Dual Wave-Particle Nature of Light was accepted • What does this mean for the atom???

13. LineSpectrum • Elements in gaseous states give off colored light • High temperature or high voltage • Always the same • Each element is unique • http://home.achilles.net/~jtalbot/data/elements/

14. Line Spectrum • Ground state • Lowest energy level available • Excited state • State in which electron has a higher potential energy than in its ground state • Farther from nucleus • Higher potential energy

15. Line Spectrum • Electron falls from higher energy level to lower one…emits light at a specific frequency • Color of light emitted depends on difference between excited state and ground state • See figure 10.5 page 201

16. Line Spectrum • Each band of color is produced by light of a different wavelength • Each particular wavelength has a definite frequency and has definite energy • Each line must therefore be produced by emission of photons with certain energies

17. Line Spectrum

18. Line Spectrum • Whenever an excited electron drops from such a specific excited state to its ground state (or lower excited state) it emits a photon • The energy of this photon is equal to the difference in energy between the initial state and the final state.

19. Niels Henrik David Bohr • 1885-1962 • Physicist • Worked with Rutherford • 1912 • Studying line spectra • of hydrogen

20. Niels Henrik David Bohr • 1913 – proposed new atomic structure • Electrons exist in specific regions away from the nucleus • Electrons revolve around nucleus like planets around the sun

21. The Bohr Atom • Nucleus with protons and neutrons • Electrons move in “stationary states” which are stable (paths or orbits) • When an electron moves from one state to another the energy lost or gained is done is ONLY very specific amounts • Each line in a spectrum is produced when an electron moves from one stationary state to another

22. The Bohr Atom • Model didn’t seem to work with atoms with more than one electron • Did not explain chemical behavior of the atoms

23. Wave Matters… • Louis de Broglie (1924) • Proposed that electrons might have a wave-particle nature • Used observations of normal wave activity

24. Wave Matters… • Erwin Schrodinger (1926) • Used mathematical understanding of wave behavior – devised an equation that treated electrons moving around nuclei as waves • Quantum Theory

25. Quantum Theory • Describes mathematically the wave properties of electrons and other very small particles • Applies to all elements (not just H)

26. Energy Levels of Electrons • Principal energy levels • Designated by letter n • Each level divided into sublevels • 1st energy level has 1 sublevel • 2nd energy level has 2 sublevels • Etc.

27. Energy Levels of Electrons

28. Orbitals • Electrons don’t actually orbit like planets • Orbital: region in space where there is a high probability of finding a given electron • Each orbital sublevel can hold 2 electrons

29. Orbitals Each sublevel (orbital) has a specific shape http://daugerresearch.com/orbitals/

30. Orbitals • Pauli exclusion principle: an atomic orbital can hold a maximum of two electrons which must have opposite spins • Electrons can only spin in two directions • Shown with arrows

31. Rules for Orbital Filling • Pauli’s Exclusion Rule • No two electrons have the same set of quantum numbers • Hund’s Rule • Electrons will remain unpaired in a given orbital until all orbitals of the same sublevel have at least one electron 1s 2s 2p 3s 3p

32. Rules for Orbital Filling • Diagonal Rule • The order of filling once the d & f sublevels are being filled • Due to energy levels

33. Rules for Orbital Filling

34. Quantum Numbers • Numbers that specify the properties of atomic orbitals and their electrons • Principle Quantum Numbers: • Symbolized by n, indicates the main energy levels surrounding a nucleus, which indicates the distance from the nucleus (shells or levels)

35. Quantum Numbers • Orbital Quantum Number: • Indicates the shape of an orbital • (subshell or sublevels) • s, p, d, f Principal Quantum # Orbital Quantum # 1 1s 2 2s, 2p 3 3s, 3p, 3d 4 4s, 4p, 4d, 4f

36. Quantum Numbers • Magnetic Quantum Number: • Indicates the orientation of an orbital about the nucleus • Orbital position with respect to the 3-dimensional x, y, and z axes

37. Quantum Numbers • Spin Quantum Number: • Indicates two possible states of an electron in an orbital Type of Orbital Number of Orbitals s 1 () p 3 (x, y, z) (, , ,) d 5 (, , , , ) f 7 Each orbital holds a maximum of 2 electrons

38. Application of Quantum Numbers • Several ways of writing the address or location of an electron • Lowest energy levels are filled first • Electron Configuration: using the diagonal rule, the principal quantum number (n), and the sublevel write out the location of all electrons 12C: 32S: 1s22s22p2 1s22s22p63s23p4

39. Application of Quantum Numbers • Orbital filling electron diagram: using Hund’s rule and the diagonal rule write out the location of all electrons • See examples on whiteboard

40. Chapter 10 – Part TwoThe Periodic Table Objectives: Understand the arrangement of the Periodic Table (10.6) Identify connections between electron configuration and placement on the periodic table

41. The Periodic Table • 1869 – arrangement proposed by Dmitri Mendeleev • And Lothar Meyer (different layout) • Still similar today • Based on increasing atomic masses and other characteristics • Was able to predict properties of elements not yet discovered….and was correct!

42. The Periodic Table • Horizontal rows • Periods • Corresponds to outermost energy level • Vertical Columns • Groups or families • Similar properties; reactions

43. The Periodic Table • Several systems for naming groups • Left to right, 1-18 • Roman numerals and A and B • Used in this book • Group A: Representative Elements • Noble Gases • IA – Alkali Metals • IIA – Alkaline Earth Metals • VIIA - Halogens • Group B: Transition Elements

44. The Periodic Table • Chemical behavior and properties of elements in a particular family similar • Have the same outer shell electron configuration • Figure 10.15 page 211 • Noble gas configuration (shortcut) • Use previous noble gas in square brackets • Finish with valence electrons

45. The Periodic Table • Examples: • K is 1s22s22p63s23p64s1 or [Ar]4s1 • Ca is 1s22s22p63s23p64s2 or [Ar]4s2 • Write abbreviated configuration for the following elements: • Fr • Y

46. The Periodic Table • Arrangement of Periodic Table also means that elements filling similar orbitals are grouped • s block • p block • d block • f block • Know these blocks…

47. The Periodic Table - Highlights • The number of the period corresponds to the highest energy level occupied by electrons in that period • The group numbers for the representative elements are equal to the total number of valence electrons in that group

48. The Periodic Table - Highlights • The elements of a family have the same outermost electron configuration • (just different energy levels) • The elements within each of the s, p, d, and f blocks are filling the corresponding orbitals • There are some discrepancies with order of filling • (not covered in this book)