Chapter 7 Atomic Structure and Periodicity

1. Chapter 7 Atomic Structure and Periodicity

2. 7.1 Electromagnetic Radiation • electromagnetic radiation: • form of energy that acts as a wave as it travels • includes: visible light, X rays, ultraviolet and infrared light, microwaves, and radio waves • travel at a speed of 2.9979 x 108 m/s in a vacuum • All forms are combined to form electromagnetic spectrum

3. Electromagnetic Spectrum

4. Low Frequency High Frequency Low Energy High Energy X-Rays Radiowaves Microwaves Ultra-violet GammaRays Infrared . Long Wavelength Short Wavelength Visible Light

5. - Page 139 “R O Y G B I V” Frequency Increases Wavelength Longer

6. Crest Wavelength Amplitude Trough Parts of a wave Origin

7. Electromagnetic radiation propagates through space as a wave moving at the speed of light. Equation: c = c = speed of light, a constant (2.998 x 108 m/s)  (lambda) = wavelength, in meters (nu) = frequency, in units of hertz (hz or sec-1)

8. Wavelength and Frequency • Are inversely related • As one goes up the other goes down. • Different frequencies of light are different colors of light. • There is a wide variety of frequencies • The whole range is called a spectrum

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

10. Electromagnic radiation propagates through space as a wave moving at the speed of light.

11. Wave nature of electromagnetic Radiation • wavelength: • λ= Greek letter lambda • distance between points on adjacent waves (consicutive peaks or troughs) • in nm (109nm = 1m) • frequency: • = Greek letter nu • number of wave cycles that passes a point in a second. 108 cycles/s= 108s-1 • =108 Hertz = 108 Hz • in 1/second (Hertz = Hz)

12. C=speed of light, a constant (3.00 x 108 m/s) =frequency, in units of hertz (hz, sec-1) = wavelength, in meters

13. Long Wavelength = Low Frequency = Low ENERGY Wavelength Table Short Wavelength = High Frequency = High ENERGY

14. Calculate the energy of red light vs. blue light. red light: 700 nm blue light: 400 nm red: blue: E = 2.85 x 10-19 J E = 4.96 x 10-19 J sunburn????? uv

15. 7.2 Nature of Matter • Before 1900, scientists thought that matter and energy were totally different

16. In 1900 • Matter and energy were seen as different from each other in fundamental ways. • Matter was particles. • Energy could come in waves, with any frequency. • Max Planck found that as the cooling of hot objects couldn’t be explained by viewing energy as a wave.

17. Explanation of atomic spectra • When we write electron configurations, we are writing the lowest energy. • The energy level, and where the electron starts from, is called it’s ground state - the lowest energy level.

18. Changing the energy • Let’s look at a hydrogen atom, with only one electron, and in the first energy level.

19. Changing the energy • Heat, electricity, or light can move the electron up to different energy levels. The electron is now said to be “excited”

20. Changing the energy • As the electron falls back to the ground state, it gives the energy back as light

21. Changing the energy • They may fall down in specific steps • Each step has a different energy

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23. The further they fall, more energy is released and the higher the frequency. • This is a simplified explanation! • The orbitals also have different energies inside energy levels • All the electrons can move around. Ultraviolet Visible Infrared

24. What is light? • Light is a particle - it comes in chunks. • Light is a wave - we can measure its wavelength and it behaves as a wave • If we combine E=mc2 , c=, E = 1/2 mv2 and E = h, then we can get:  = h/mv (from Louis de Broglie) • called de Broglie’s equation • Calculates the wavelength of a particle.

25. Wave-Particle Duality J.J. 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 an energy wave! The electron is a particle!

26. Confused? You’ve Got Company! “No familiar conceptions can be woven around the electron; something unknown is doing we don’t know what.” Physicist Sir Arthur Eddington The Nature of the Physical World 1934

27. The physics of the very small • Quantum mechanics explains how very small particles behave • Quantum mechanics is an explanation for subatomic particles and atoms as waves • Classical mechanics describes the motions of bodies much larger than atoms

28. Heisenberg Uncertainty Principle • It is impossible to know exactly the location and velocity of a particle. • The better we know one, the less we know the other. • Measuring changes the properties. • True in quantum mechanics, but not classical mechanics

29. Heisenberg Uncertainty Principle “One cannot simultaneously determine both the position and momentum of an electron.” You can find out where the electron is, but not where it is going. OR… You can find out where the electron is going, but not where it is! Werner Heisenberg

30. It is more obvious with the very small objects • To measure where a electron is, we use light. • But the light energy moves the electron • And hitting the electron changes the frequency of the light.

31. After Before Photon wavelengthchanges Photon Moving Electron Electron velocity changes Fig. 5.16, p. 145

32. Light is a Particle? • Energy is quantized. • Light is a form of energy. • Therefore, light must be quantized • These smallest pieces of light are called photons. • Photoelectric effect? Albert Einstein • Energy & frequency: directly related.

33. The energy (E ) of electromagnetic radiation is directly proportional to the frequency () of the radiation. Equation:E = h E = Energy, in units of Joules (kg·m2/s2) (Joule is the metric unit of energy) h = Planck’s constant (6.626 x 10-34 J·s) = frequency, in units of hertz (hz, sec-1)

34. Nature of Matter • Max Planck: a German physicist suggested that an object emits energy in the form of small packets of energy called quanta • Quantum- the minimum amount of energy that can be gained or lost by an atom Planck’s constant (h): 6.626 x 10-34 J*s

35. Nature of Matter • Einstein proposed that radiation itself is really a stream of particles called photons • Energy of each photon is : • also showed that energy has mass

36. Nature of Matter shows that anything with both mass and velocity has a corresponding wavelength

37. Nature of Matter • In 1924, Louis de Broglie (French scientist) • suggested that matter has both particle-like and wave-like characteristics

38. Main Ideas: • matter and energy are not distinct • energy is a form of matter • larger objects are mostly particle-like • smaller objects are mostly wave-like

39. 7.3 The Atomic Spectrum of Hydrogen Spectroscopic analysis of the visible spectrum… White light …produces all of the colors in a continuous spectrum

40. Atomic Spectra • White light is made up of all the colors of the visible spectrum. • Passing it through a prism separates it.

41. These are called the atomic emission spectrum • Unique to each element, like fingerprints! • Very useful for identifying elements

42. Continuous Spectra White light passed through a prism produces a spectrum – colors incontinuousform.

43. The Continuous Spectrum l ~ 650 nm • ~ 575 nm l ~ 500 nm l ~ 480 nm l ~ 450 nm The different colors of light correspond to different wavelengths and frequencies

44. Continuous Emission Spectrum • line-emission spectrum- series of wavelengths of light created when visible portion of light from excited atoms is shined through a prism • scientists using classical theory expected atoms to be excited by whatever energy they absorbed • continuous spectrum- • emission of continuous range of frequencies of EM radiation • contains all wavelengths of visible light

45. Spectroscopic analysis of the hydrogen spectrum… H receives a high energy spark H-H bonds Are broken and H atoms are excited …produces a “bright line” spectrum

46. Line Spectra Light passed through a prism from an element produces a discontinuousspectrum of specific colors

47. Hydrogen only four lines are observed

48. Line Spectra The pattern of lines emitted by excited atoms of an element is unique = atomic emission spectrum

49. These are called the atomic emission spectrum • Unique to each element, like fingerprints! • Very useful for identifying elements

50. H Line-Emission Spectrum • light is emitted by excited H atoms when bond is broken in the diatomic molecule • ground state- lowest energy state of an atom • excited state- when an atom has higher potential energy than it has at ground state