1. 1 Atomic Structure and Periodicity AP Chemistry Chapter 7

2. 2 The ELECTRON: Wave – Particle Duality Graphic: www.lab-initio.com

3. 3 The Dilemma of the Atom Electrons outside the nucleus are attracted to the protons in the nucleus Charged particles moving in curved paths lose energy What keeps the atom from collapsing? Electrons outside the nucleus are attracted to the protons in the nucleus Charged particles moving in curved paths lose energy What keeps the atom from collapsing?

4. 4 Wave-Particle Duality JJ Thomson won the Nobel prize for describing the electron as a particle. His son, George Thomson won the Nobel prize for describing the wave-like nature of the electron. The electron is a particle! The electron is an energy wave!

5. 5 The Wave-like Electron Louis deBroglie The electron propagates through space as an energy wave. To understand the atom, one must understand the behavior of electromagnetic waves.

6. 6 c = c = c = speed of light, a constant (3.00 x 10 8 m/s) = frequency, in units of hertz (hz, sec -1 ) = wavelength, in meters Electromagnetic radiation propagates through space as a wave moving at the speed of light.

7. 7 Electromagnetic Spectrum Long wavelength  small frequency Short wavelength  high frequency increasing frequency increasing wavelength

8. 8 Electromagnetic Spectrum

9. 9 Scientific belief around the 1900’s was that there was NO relationship between matter and light Light given off by objects that were heated to high temperatures could not be explained. Problems with Wave Theory of Light

10. 10 Black Body Radiation http://www.youtube.com/watch?v=1pfqIcSydgE &list=PLEF4658F09DF08831&index=30

11. 11 7.2 Nature of Matter Objects radiate energy in small packets of energy called quanta. In case you didn’t know, a quantum is a specific amount of energy that can be gained or lost by an atom. Max Planck

12. 12 E = h E = h E = Energy, in units of Joules (kg·m 2 /s 2 ) h = Planck’s constant (6.626 x 10-34 J·s) = frequency, in units of hertz (hz, sec -1 ) = frequency, in units of hertz (hz, sec -1 ) The energy (E ) of electromagnetic radiation is directly proportional to the frequency ( ) of the radiation.

13. 13 Photoelectric Effect Remember me? I was also the first person to observe the photoelectric effect which is the emission of electrons from a metal surface when exposed to light of a specific energy. JJ Thomson

14. 14 EMR can be viewed as a stream of particles called “photons”. The energy of these photons could be calculated by Planck’s equation. Photons strike the electrons therefore ejecting them from metal. Albert Einstein

15. 15 The success of Einstein’s work in explaining the photoelectric effect was largely responsible for the acceptance of the particle behavior of light E photon = h E = mc 2 Dual Wave-Particle Behavior Of Light

16. 16 Using Einstein’s and Planck’s equations, de Broglie derived: The momentum, mv, is a particle property, whereas is a wave property. In one equation de Broglie summarized the concepts of waves and particles as they apply to low mass, high speed objects. Can matter act as a wave?

17. 17 Compare the wavelength for an electron (mass = 9.11 x 10 -31 kg) traveling at a speed of 1.0 x 10 7 m/s with that for a ball (mass = 0.10 kg) traveling at 30 m/s. Sample Problem

18. 18 Energy is a form of matter. All matter exhibits both particle and wave properties. Large pieces of matter (i.e. baseball) exhibits mostly particle properties. Tiny pieces of matter (i.e. photons) exhibits mostly wave properties. Pieces of matter somewhere in the middle (i.e. electrons) clearly show both types of properties! Dual Wave-Particle Behavior Of Matter

19. 19 Observed that light was given off when they heated different chemicals in their designed burner They passed the light through a prism and saw separate lines instead of a continuous spectrum. Kirchoff and Robert Bunsen (1854)

20. 20 Emission spectra- the colors produced by an object when burned or heated. Absorption spectra- the colors that are not shown, rather absorbed in the spectrum http://chemistry.beloit.edu/BlueLight/moviepages/ab_em_el.htm Absorption and Emission Spectra

21. 21 Only four lines are emitted: –Red, green, blue, violet Only certain energies are allowed. 7.3 Hydrogen Spectrum

22. 22 To understand emission spectrum, we must understand these two terms: Ground state: the lowest energy state for the electron Excited state: state where electron has higher energy than ground state. Why do elements produce these lines?

23. 23 Atoms are heated, which adds energy. The electron become excited (thus unstable). They want to return to their normal, or ground state. To do so, they give off energy in the form of EMR. Why do elements produce these lines?

24. 24 Balmer: developed a numerical relationship between the wavelength of the lines in the spectrum and the amount of energy Lyman: discovered lines produced in the UV range. Paschen: discovered lines produced in the IR range Scientists associated with the H spectrum

25. 25 Electron transitions involve jumps of definite amounts of energy. This produces bands of light with definite wavelengths.

26. 26 7.4 The Bohr Model The electron orbits the nucleus in fixed paths called orbits (like the planets orbit the sun). The electron can jump to higher orbitals when energy is added. The electron’s energy can be calculated in the different orbits. People sometimes call my model of the atom the Planetary Model! The electron orbits the nucleus in fixed paths called orbits (like the planets orbit the sun). The electron can jump to higher orbitals when energy is added. The electron’s energy can be calculated in the different orbits. People sometimes call my model of the atom the Planetary Model!

27. 27 Bohr Model of the Atom How does this relate to the Hydrogen spectrum? Bohr calculated the energy that the electron would lose as it fell from higher orbits to lower orbits. Bohr’s calculations agreed exactly with Lyman, Balmer and Paschen’s observations.

28. 28 Bohr Model of the Atom Bohr’s model worked very well for the Hydrogen atom. Through Bohr’s work, as well as the other scientists mentioned, a very good understanding of the electron within the atom was now in place. Bohr Model of the Atom

29. 29 Downfalls to Bohr’s Model 1.Bohr’s model of the atom worked very well for the hydrogen atom and the He +, but failed when applied to multielectron atoms. 2.Bohr’s model could not explain why the electron could not exist between orbits. Downfalls to Bohr’s Model of the Atom

30. 30 We need a new approach to the atom! Big Three: de Broglie, Heisenberg & Schrodinger Developed wave mechanics AKA quantum mechanics (7.5) Now What?

31. 31 Heisenberg Uncertainty Principle The more certain you are about where the electron is, the less certain you can be about where it is going. The more certain you are about where the electron is going, the less certain you can be about where it is. “One cannot simultaneously determine both the position and momentum of an electron.” Werner Heisenberg

32. 32 Schrödinger proposed an equation that contains both wave and particle terms. Solving the equation leads to wave functions. The wave function gives the probability distribution of an electron. orbitalsWe call wave functions orbitals. Schrodinger’s Wave Equation

33. 33 7.6 Quantum Numbers When we solve the Schrödinger equation for the hydrogen atom, we find many wave functions (orbitals) that satisfy it. Each orbital is characterized by a series of numbers called quantum numbers that describe various properties of the orbital.

34. 34 Related to the size and energy of the orbital – think energy level n has integer values: 1,2,3… As n becomes larger, the atom becomes larger and the electron is further from the nucleus. A larger n value also corresponds to higher energy because the electron is less tightly bound to the nucleus. Principal Quantum Number, n

35. 35 Electron Energy Level (Shell) Principal Quantum number Generally symbolized by n, it denotes the probable distance of the electron from the nucleus. “n” is also known as the Principal Quantum number Number of electrons that can fit in a shell: 2n 2

36. 36 Orbital shapes are defined as the surface that contains 90% of the total electron probability. An orbital is a region within an energy level where there is a probability of finding an electron. Electron Orbitals The angular momentum quantum number, generally symbolized by l, denotes the orbital (subshell) in which the electron is located.

37. 37 The s orbital ( l = 0) has a spherical shape centered around the origin of the three axes in space. s Orbital shape

38. 38 There are three dumbbell-shaped p orbitals ( l = 1) in each energy level above n = 1, each assigned to its own axis (x, y and z) in space. p orbital shape

39. 39 Things get a bit more complicated with the five d orbitals ( l = 2) that are found in the d sublevels beginning with n = 3. To remember the shapes, think of “double dumbells ” …and a “dumbell with a donut”! d orbital shapes

40. 40 Shape of f ( l = 3) orbitals

41. 41 Energy Level (n) Sublevels in main energy level (n sublevels) Number of orbitals per sublevel Number of Electrons per sublevel Number of electrons per main energy level (2n 2 ) 1s122 2spsp 1313 2626 8 3spdspd 135135 2 6 10 18 4spdfspdf 13571357 2 6 10 14 32 Energy Levels, Sublevels, Electrons

42. 42 Magnetic Quantum Number magnetic quantum number The magnetic quantum number, generally symbolized by m l, denotes the orientation of the electron’s orbital with respect to the three axes in space.

43. 43 For hydrogen, energy is determined by value of n All orbitals with the same value of n have the same energy – they are degenerate. Energies of orbitals in Hydrogen

44. 44 Orbital filling table

45. 45 Electron Spin The Spin Quantum Number The Spin Quantum Number describes the behavior (direction of spin) of an electron within a magnetic field. Possibilities for electron spin:

46. 46 Pauli Exclusion Principle Two electrons occupying the same orbital must have opposite spins Wolfgang Pauli

47. 47 Hund’s Rule Friedrich Hund All orbitals in the same sublevel will each receive one electron before pairing will occur.

48. 48 ElementConfiguration notation Orbital notationNoble gas notation Lithium1s 2 2s 1 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 1 Beryllium1s 2 2s 2 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 Boron1s 2 2s 2 p 1 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 p 1 Carbon1s 2 2s 2 p 2 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 p 2 Nitrogen1s 2 2s 2 p 3 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 p 3 Oxygen1s 2 2s 2 p 4 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 p 4 Fluorine1s 2 2s 2 p 5 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 p 5 Neon1s 2 2s 2 p 6 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 p 6

49. 49 Many properties of atoms depend on electron configurations and how strongly valence electrons are attracted to the nucleus. Coulomb’s Law – strength of the interaction between 2 electrical charges depends on the size of the charges and the distance between them. Zeff = Z – S where Z is # protons in nucleus and S is number of core electrons Explains differences in sublevel energies but also describes periodic trends. Effective Nuclear Charge

50. 50 The effective nuclear charge increases as we move across any row (period) of the periodic table (Z gets larger while S stays the same) The effective nuclear charge also increases as we go down a column (group) of the periodic table, but the effect is far less than going across a row. Effective nuclear charge

51. 51 Atomic Radius Simple diatomic molecule The distance between the two nuclei is called the bond distance. If the two atoms which make up the molecule are the same, then half the bond distance is called the covalent radius of the atom.

52. 52 Atomic size varies consistently through the periodic table. As we move down a group, the atoms become larger. As we move across a period, atoms become smaller. There are two factors at work: –principal quantum number, n (down a group) –the effective nuclear charge, Z eff (across a period) Atomic Radius

53. 53 Atomic Radius

54. 54 Ionization energy – minimum amount of energy required to remove an electron from the ground state of an isolated gas atom or ion. Na (g)  Na + (g) + e - First ionization energy Na + (g)  Na 2+ (g) + e - Second ionization energy The greater ionization energy, the more difficult it is to remove the electron. Ionization Energy

55. 55 Ionization energy increases for each additional electron removed from an atom. There is a sharp increase in ionization energy when a core (non-valence) electron is removed. Ionization Energy

56. 56 Same factors influence ionization energy – effective nuclear charge & distance of electron from nucleus. Increasing effective charge or decreasing distance from nucleus increases attraction between electron & nucleus – more difficult to remove an electron so ionization energy increases. (Both happen when move across row) As we move down group, the atomic radius increases (due to larger n) while effective nuclear charge only increases slightly. Attraction between nucleus & electron decreases, so ionization energy decreases. Ionization Energy Trend