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. 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

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

6. 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)

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

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

9. 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

10. 7.2 Nature of Matter • Before 1900, scientists thought that matter and energy were totally different

11. 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.

12. Photoelectric effect: e- Metal surface Light of the right frequency (energy) can strike a metal and cause an electron to be ejected (n = infinity).

13. 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

14. 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

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

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

17. 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

18. 7.3 The Atomic Spectrum of Hydrogen Spectroscopic analysis of the visible spectrum… White light …produces all of the colors in a continuous spectrum

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

20. 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

21. 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

22. 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

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

24. Hydrogen only four lines are observed

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

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

27. 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

28. H Line-Emission Spectrum • when an excited electron falls back to ground state, it emits photon of radiation • the photon is equal to the difference in energy of the original and final states of electron • since only certain frequencies are emitted, only certain energies are allowed for electrons in H atom

29. Electron transitionsinvolve jumps of definite amounts ofenergy. This produces bands of light with definite wavelengths.

30. 7.4 The Bohr Model • Niels Bohr (Danish physicist) in 1913 • Developed a quantum model for H atom that explained the emission line spectrum • Electron moves around the nucleus only in certain allowed circular orbits, in which it has a certain amount of energy

31. The Bohr model • Energy level of an electron analogous to the steps of a ladder • The electron cannot exist between energy levels, just like you can’t stand between steps on a ladder • Aquantumof energy is the amount of energy required to move an electron from one energy level to another

32. Niels Bohr • Developed the quantum model of the hydrogen atom. • He said the atom was like a solar system. • The electrons were attracted to the nucleus because of opposite charges. • Didn’t fall in to the nucleus because it was moving around.

33. The Bohr Atom • He didn’t know why but only certain energies were allowed. • He called these allowed energies energy levels. • Putting Energy into the atom moved the electron away from the nucleus. • From ground state to excited state. • When it returns to ground state it gives off light of a certain energy.

34. The Model: Summary • Space around nucleus is divided into spherical (circualr) paths (orbits) each has a number called “Principal Quantum number” • The electron can exist only in one of these orbitals but not in between • Orbits possess fixed size and energy, therefore electron has a definite energy characteristic of its orbit

35. Orbits allowed for electron are those in which electron has an angular momentum= • An electron can pass only from one bit to another. Absorption or emission will occur • Energy of the outermost orbit is zero

36. The Bohr Atom n = 4 n = 3 n = 2 n = 1

37. Bohr Model • To create an accurate model, he had to use quantum theory instead of classical • created an equation used to calculate the energy levels available to electrons in a certain atom: where n= integer and Z=atomic number negative sign makes the energy more negative the closer it is to the nucleus

38. Bohr Model • Can gain energy by moving to a higher energy level • Can lose energy by moving to lower energy level

39. Bohr Model a photon is released that has an energy equal to the difference between the initial and final energy orbits

40. Bohr Model • equation can be used twice to find the ∆E when an electron moves energy levels

41. Bohr Model • can wavelength of photon released by using • E=0 is set at an distance of ∞ away from the nucleus and becomes more negative as the electron comes closer to the nucleus

42. Example 1 Calculate the energy required to move the hydrogen electron from n=1 to n=2. Find the wavelength of radiation that had to be absorbed by the electron.

43. Example 2 Calculate the energy required to remove the electron from the hydrogen atom in its ground state. Energy was absorbed by the electron so the value of ∆E value is positive.

44. The Bohr Model • Doesn’t work. • Only works for hydrogen atoms. • Electrons don’t move in circles. • The quantization of energy is right, but not because they are circling like planets.

45. Bohr Model • problems: • did not work for other atoms • did not explain chemical behavior of atoms

46. Heisenberg’s Uncertainty Principle • According to de Broglie: Electron behaves like a wave • It is possible to specify the position of a wave at a particular instant? • Energy, wavelength and amplitude can be determined • But exact position is impossible to be determined • The electron cannot be imagined as : • moving particle • In a path of the same radius (well defined orbits) • Thus, location, direction and speed of motion of a particle cannot be determined • Then Bohr Model had to be “Abandoned

47. Heissenberg Uncertainty Principle “It is impossible to determine both the position and momentum of a subatomic particle (such as the electron) with arbitrarily high accuracy” • The effect of this principle is to convert the laws of physics into statements about relative, instead of absolute, certainties.

48. Heisenberg Uncertainty Principle • we cannot know the exact position and momentum (motion) of the electron • as more is known about position, less is known about momentum • uncertainties are inversely proportional • where • ∆x: uncertainty in position • ∆m : uncertainty in mometum • minimum uncertainty is h/4

49. 7.5 The Quantum Mechanical Model • Exactt position of electron can not be defined • Exact bath of electron about nucleus can not be defined • Werner Heisenberg, Louis de Broglie and Erwin Schrodinger made the approach called “Quantum Mechanics” • They assumed that the electron is a standing wave

50. The Quantum Mechanical Model • Waves are associated with electrons • Information about energies of electrons and their positions are obtained from studying the associated waves • Description of electron is based upon “Probability of finding a particle within a given region of space”“but not on the exact position”