Chapter 4 Electron Structure of the Atom

1. Chapter 4Electron Structure of the Atom

2. Review of Atomic Structure • The center of the atom is called the nucleus. • In it are the particles with mass: the protons and neutrons • Protons determine the identity of an atom • Electrons determine the properties of an atom • Where are the electrons?

3. Light • What is light, and what does it have to do with the electrons in an atom? • Light is electromagnetic radiation.

4. Electromagnetic Radiation (EMR) • Form of energy with wavelike behavior as it travels through space at the speed of light • Seven major types: • Gamma () Rays • X rays • Ultraviolet (UV) • Visible Light (ROYGBIV) • Infrared (IR) • Microwaves • Radio waves

5. Electromagnetic Spectrum(p. 98)

6. Wavelength (λ) • The distance between two corresponding points on a wave • Units are same as length - m, or commonly nm (10-9 m) Figure 7.5

7. Know the wavelengths for the visible spectrum R O Y G B IV 700 650 600 550 500 450 400 nm

8. Frequency • number of wave cycles that move through a point in space in 1 s • hertz (Hz) • same as inverse seconds (1/s) or (s-1)

9. Figure 7.7 • Which has the greater wavelength the red light or the green light? • Which has a greater frequency?

10. Frequency & Wavelength • inversely proportional • i.e. as one increases the other decreases c = λν c = speed of light (3.0 x 108 m/s) λ = wavelength (in meters) ν = frequency (in Hz)

11. Practice • What is the frequency and wave type of a wave with a wavelength of 5.2 x 10-7 m (5.2 x 102 nm)? wave type: green visible light

12. Practice • What is the frequency and wave type of a wave with a wavelength of 5.2 x 10-7 m (5.2 x 102 nm)? wave type: green visible light

13. What is the approximate frequency of blue light? • Wavelength = 475nm R O Y G B IV 400 700

14. Convert nm to m R O Y G B IV 400 700

15. What is the approximate frequency of blue light? R O Y G B IV 400 700

16. Homework • Calculate the frequency of light with a wavelength of 3.41 x 103 cm? • Calculate the wavelength of light with a frequency of 3.21 x 1016 Hz? What type of EMR is it? • What is the frequency of orange light?

17. Continuous spectrum all the wavelength of in the visible spectrum Produced by white light Line Spectrum distinct colored lines each a single wavelength of light Visible when an element has been heated “atomic fingerprint” Line Spectra

18. Energy is Quantized! Quantized = quantity = specific measured amount • Max Planck • energy produced by atoms can only have certain values • only distinct lines are seen in element line spectra • can only exist at certain wavelengths.

19. Max Planck • Proposed that objects emit energy in small packets called “quanta” • Quantum: min. quantity of energy lost or gained by an atom • This quantized energy is related to the frequency of the energy

20. Photons – waves and particles Energy of a photon • directly proportional to the frequency Ephoton = hν • inversely proportional to the wavelength Ephoton = hc/l Ephoton = (in Joules) h = Planck’s constant (6.626 x 10-34 Js) ν= frequency (in Hz) l = wavelength in meters

21. Low High Low E High Low High E

22. Practice • What is the energy of a photon with a frequency of 1x1017Hz?

23. Practice • What is the energy of a photon with a wavelength of 8.1m?

24. Atomic Spectra • When visible light passes through a prism, its components separate into a spectrum. • White light, such as sun light or light from a regular light bulb, gives a continuous spectrum:

25. Photoelectric Effect • Emission of electrons from a metal when light shines on it

26. This was explained by Einstein • Based on Planck's work, Einstein proposed that light also delivers its energy in chunks; light would then consist of little particles, or quanta, called photons, each with an energy of Planck's constant times its frequency • Electrons are only emitted if the photons have a high enough energy (high enough frequency)

27. Atomic Spectra • Colored light gives only specific colors in a line spectrum:

28. What does this have to do with electrons? • Hydrogen atom: 1 proton, 1 electron • Passing electricity through a tube containing hydrogen gas gives off a pink light • Light is given off as electrons fall from an excited state to a ground state • The light can be separated into four frequencies

29. Spectra of other elements(spectral signatures applet)

30. Energy Level Transition • Excited State: higher energy state • Ground State: lowest energy state Ephoton Ephoton= Eexcited- Eground=h

31. Bohr Model links electron and emissions • Niels Bohr (Danish physicist) • planetary model • electron are in orbits • orbit 1 closest to the nucleus • increasing numbers as the orbits get further away from the nucleus.

32. Bohr Model • Orbits have a fixed radius. • Electrons cannot exist between orbits • lowest energy is closest to the nucleus • increases as the orbits get further away • Electrons absorb or emits energy when they change orbitals ∆E = Ef – Ei

33. Deficiencies of Bohr Model • Did not work for other atoms • Doesn’t explain chemical behavior of atoms • More complex model needed • However, the Bohr model did give insight into the quantized behavior of the atom that was better understood in later days

34. Modern Models of the Atom • Bohr • deBroglie • Schrödinger (Wave Mechanical Model) • To view at home, click here. Click on Run Now!

35. Modern Model of the Atom(Quantum Mechanical Model) • electrons exist in orbitals. • Orbitals are 3-dimensional regions in space where an electron is likely to be found • not a circular pathway • The electron is thought of having wave properties

36. The exact location of the electron cannot be known. (Heisenberg Uncertainty Principle) • The location is described as a probability • This is a Probability Map for lowest-energy state of the electron in an H atom

37. A typical representation of an s-orbital

38. Principal Energy Levels • Orbitals of similar size exist in the same principal energy level (n=1, 2, 3…) • The principal energy levels correspond to Bohr’s energy levels and represent a distance from the nucleus

39. Energy levels contain orbitals • Lower energy orbitals are smaller. • Higher energy orbitals are larger; further away from the nucleus. • An orbital can hold at most 2 electrons.

40. Sublevels • Orbitals are arranged in sublevels, which have specific shapes • s, p, d, and f are sublevels s p d f

41. s Orbitals Figure 7.14

42. p Orbitals

43. d Orbitals

44. f Orbitals

45. Memorize terms! • Principal Energy level: distance from nucleus • Sublevel: series of orbitals having equal energy • Orbital: particular region where an electron exists • Orbitals make up sublevels, and sublevels make up energy levels

46. What each energy level holds(Table 2, p. 110) • The first energy level (n=1) • A single s orbital • The second energy level (n=2) • 2s orbital and 2p sublevel (three 2p orbitals) • The third energy level (n=3) • 3s orbital, 3p sublevel (three 3p orbitals) and 3d sublevel (five 3d orbitals) • The fourth energy level (n=4) • 4s orbital, 4p sublevel, 4d sublevel and 4f sublevel (seven 4f orbitals)

49. Hydrogen Orbital Diagram • Aufbau Orbital diagrams • Show the sublevels and orbitals that can exist at each principal energy level • Each box represents an orbital • Groups of boxes represent sublevels • In the hydrogen atom only, the sublevels within a principal energy level all have the same energy.

50. Multielectron Orbital Diagram • In the multielectron atoms, the sublevels within a principal energy level have different energy levels.