Modern Atomic Theory

1. Modern Atomic Theory

2. Please select the appropriate Team. 1. Girls 2. Guys

3. A_O_ 1. W H 2. M N 3. M T 4. G Z 0% 0% 0% 0%

4. R_T_E_F_R_ 1. H R U O D 2. N I V W O 3. U V M E R 4. Q O R O H 0% 0% 0% 0%

5. The fact that carbon dioxide always contains 73 percent oxygen by mass is an illustration of 0% 1. the ideas of Democritus. 2. the law of conservation of matter. 3. the law of definite composition. 4. Dalton's atomic theory of matter. 0% 0% 0% 10

6. Which of the following is not one of the three fundamental particles that makeup atoms? 0% 1. electron 2. proton 3. neutron 4. alpha particle 0% 0% 0% 10

7. The atomic number of an atom is defined as its 1. mass in amu. 2. number of electrons. 3. number of neutrons. 4. number of protons. 0% 0% 0% 0% 10

8. An ion always contains an 1. unequal number of protons and electrons. 2. equal number of protons and electrons. 3. unequal number of protons and neutrons. 4. equal number of protons and neutrons. 0% 0% 0% 0% 10

9. Two atoms are isotopes if they contain different numbers of electrons. different numbers of protons and different numbers of neutrons. the same number of protons but different numbers of neutrons. the same number of neutrons but different numbers of protons. 1. 2. 3. 4. 0% 0% 0% 0% different numbers of ... different numbers of ... the same number of ... the same number of ... 10

10. Team Scores 212.17 166 Girls Guys

11. • In the last 200 years, vast amounts of data have been accumulated to support atomic theory. When atoms were originally suggested by the early Greeks, no physical evidence existed to support their ideas. Early chemists did a variety of experiments, which culminated in Dalton’s model of the atom. Because of the limitations of his model, modifications were proposed first by Thomson and then Rutherford, which eventually led to our modern concept of the nuclear atom.

12. • As with the mystery box you cannot directly see what is inside, we cannot easily see what is inside an atom. • Scientists have studied energy and light for centuries, and several models have been proposed to explain how energy is transferred from place to place. One is through the electromagnetic spectrum.

14. Three Characteristics of Waves •Wavelength: λ, distance between consecutive peaks or troughs in a wave •Frequency: f, how many waves pass a particular point per second. •Speed: v, how fast a wave moves through space (m/s)

15. We have evidence for the wavelike nature of light. We know that a beam of light behaves like a stream of tiny packets of energy called photons.Flinn Spectra Slides

16. Which of the following is not a form of electromagnetic radiation? 1. X-rays 2. gamma rays 3. sound waves 4. visible light 0% 0% 0% 0% 10

17. Which of the following has the longest wavelength? 1. ultraviolet radiation 2. infrared radiation 3. X-rays 4. gamma rays 0% 0% 0% 0% 10

18. What does the wavelength of the substance tell you? 1. Energy of the atom 2. Identity of the atom 3. Color of light given off from the atom 4. All of the above 5. None of the above 0% 0% 0% 0% 0%

19. Where does this connect to the atom? • At high temperatures or when subjected to high voltages, elements in the gaseous state give off colored light. Brightly colored neon signs illustrate this property well. • When the light is emitted by a gas is passed through a prism or diffraction grating, a set of brightly colored lines called a line spectrum results.

20. Continuous and Line Spectra

21. Fireworks

22. Composition of Fireworks • Gunpowder – Sulfur, charcoal, potassium nitrate (saltpeter) • Salts (to give color) – Red = lithium – Green = copper

23. • Neils Bohr proposed a model of the atom, where electrons are found in energy levels and they jump from one energy level to another by adding or losing a quantum of energy.

24. • Neils Bohr studied the line spectrum of hydrogen which led him to believe that electrons exist in specific regions at various distances from the nucleus. He visualized that electrons revolved in orbits around the nucleus like the planets revolve around the sun.

25. • Bohr applied Planck’s concept of energy quanta. • Quanta: A small discrete package of energy, energy is not emitted in a continuous stream.

26. Bohr’s Ideas • Electrons have several possible energies • These energies correspond to several possible orbits at different distances from the nucleus • Therefore an electron has to be in a specific energy level, not between levels. The electron is quantized • When a hydrogen atom absorbed one or more quanta of energy, it jumped to a higher energy level or orbital.

27. The color of the light emitted corresponds to one of the lines of the hydrogen spectrum. level a quantum of energy in the form of light is emitted by the atom. When an electron falls from a higher energy level to a lower energy 10.4

28. Principal Quantum Number (n): Principal Quantum Number (n): tells the shell of the electron a. Ground state (n=1) lowest possible energy level, closest to the nucleus b. Excited state (n=2, 3, 4, etc) when the electron absorbs the appropriate amount of energy, it jumps to a level of higher energy. Radiation is emitted when the electron falls back from a higher energy level to a lower one.

29. • In Bohr’ ’s model of hydrogen, the energy levels that the electrons occupy are similar to lanes on a freeway. • The lowest energy level, ground state, is like the slow lane on the freeway. This is where electrons are normally found. • When the electron receives more energy from an outside source, it has to change to the next highest energy level. Similarly, a car in the slow lane that increases its speed will move to a faster lane. • Both electrons and cars will return to the ground state when their energy or speed decreases…we can’ ’t drive between the lanes, therefore electrons cannot stay between levels. • This is basically known as Aufbau’ ’s principle.

30. • The maximum number of electrons in a energy level = 2n2 • Energy level 1 = 2(1)2= 2 electrons • Energy level 2 = 2(2)2= 8 electrons • Energy level 3 = 2(3)2= 18 electrons • Energy level 4 = 2(4)2= 32 electrons

31. Bohr Model: Bohr Model: theoretical model where electrons orbit the nucleus in defined energy levels. 1. Gather information – number of protons, neutrons, and electrons 2. Build the nucleus – protons and neutrons 3. Draw the energy levels 4. Place the electrons Practice: H, He, Cl, Cl+2 • •

32. Modern Theory (1925) • Bohr’s calculations succeeded very well in correlating the experimentally observed spectral lines with electron energy levels for the hydrogen atom. • His methods did not succeed for heavier atoms. • Bohr’s Model no longer explained all observations. New theory proposed that atoms did not travel in definite paths but behaved more like waves on a vibrating string.

33. Electrons as Waves • Louis de Broglie (1924) Louis de Broglie ~1924 – Applied wave-particle theory to electrons – electrons exhibit wave properties QUANTIZED WAVELENGTHS Standing Wave Fundamental mode Second Harmonic or First Overtone 200 200 200 150 150 150 100 100 100 50 50 50 0 0 0 - 50 - 50 - 50 -100 -100 -100 -150 -150 -150 -200 -200 -200 0 50 100 150 200 0 50 100 150 200 0 50 100 150 200 Adapted from work by Christy Johannesson www.nisd.net/communicationsarts/pages/chem

34. Electrons as Waves QUANTIZED WAVELENGTHS Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

35. Electrons as Waves Evidence: DIFFRACTION PATTERNS ELECTRONS VISIBLE LIGHT Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem Davis, Frey, Sarquis, Sarquis, Modern Chemistry 2006, page 105

36. Quantum Mechanics Erwin Schrodinger ~1926 • (1926) Erwin Schrödinger created a mathematical model of electrons as waves. • Schrödinger’ ’s work led to a new branch of physics called wave or quantum mechanics. • Using Schrödinger’s wave mechanics, the probability of finding an electron in a certain region around the atom can be determined. • The actual location of an electron within an atom cannot be determined. Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

37. Quantum Mechanics Erwin Schrodinger ~1926 • Schrödinger Wave Equation – finite # of solutions  quantized energy levels – defines probability of finding an electron   0 3/2 σ e  1 Z a Ψ  1s π Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

38. Orbitals • This makes it quite difficult to pinpoint an exact location, but a region can be predicted. • An orbital is a region in an atom where there is a high probability of finding electrons.

39. Quantum Mechanics • Orbital (“electron cloud”) – Region in space where there is 90% probability of finding an electron 90% probability of finding the electron Electron Probability vs. Distance 40 Electron Probability (%) 30 20 10 0 0 50 100 150 200 250 Distance from the Nucleus (pm) Orbital Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

40. The specific location of an electron can be determined. 1. True 2. False 37% 63% True False

41. Electrons cannot be located between energy levels or orbitals. 1. True 2. False 6% 94% True False

42. Heisenberg’s Uncertainty Principle • The more precisely the position is determined, the less precisely the momentum is known in this instant, and vice versa. --Heisenberg, uncertainty paper, 1927 • http://youtu.be/noZWLPpj3to ?list=PL07E6A22017705261

43. Quantum Mechanics • Heisenberg Uncertainty Principle Werner Heisenberg ~1926 – Impossible to know both the velocity and position of an electron at the same time g Microscope Electron

46. Electron Configuration Allows us to map the location of all electrons Specify the “address” of each electron’s region in an atom UPPER LEVEL Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

47. An Atom’s Address • An electron subshell is a region of space within an electron shell that contains electrons that have the same energy.

48. grape : raisin :: plum : 1. peach 2. fig 3. apricot 4. prune 0% 0% 0% 0% fig prune apricot peach

49. Clark Kent : Superman :: Bruce Wayne : 1. Hulk 2. Batman 3. Spiderman 4. Wolverine 0% 0% 0% 0% Wolverine Hulk Batman Spiderman

50. Electron Subshell : energy level :: . . : region of space 1. shell 2. orbital 3. level 4. path 0% 0% 0% 0% level orbital shell path