1. 7 - 1 Atomic Structure The Electron The Nucleus Traveling Waves Electromagnetic Radiation Bohr Model of the Hydrogen Atom Wave Theory of the Electron Heisenberg Uncertainty Principle Quantum Model of the Atom

2. 7 - 2 Discovery of the electron 1807 1807Davy suggested that electrical forces held compounds together. 1833 1833Faraday related atomic mass and the electricity needed to free an element during electrolysis experiments. 1891 1891Stoney proposed that electricity exists in units he called electrons. 1897 1897Thomson first quantitatively measured the properties of electrons.

3. 7 - 3 Cathode rays Thomson’s ‘discovery’ of electrons was based on studies of cathode rays. These are produced when a gas is ionized. Ionized gas cathode anode Small dot

4. 7 - 4 Cathode rays Thomson observed that the position of the dot was altered when either electrical or magnetic fields were applied.

5. 7 - 5 Cathode rays Thomson was not able to measure the mass or charge of an electron. He could only determine the mass to charge ratio -- 6 x 10 -12 kg/C Columb (C) - the SI unit of charge. It is defined as the amount of charge that flows past a fixed point in a wire per second when the current is one ampere.

6. 7 - 6 Charge of the electron Millikan experimented with electrically charged oil drips. By varying the electrical field, he observed that the drops had charges that were whole number multiples of 1.5924 x 10 -19 C This represented the charge of an electron. The modern value is now known to be -1.602 1773 x 10 -19 C

7. 7 - 7 Mass of the electron Once both the mass to charge ratio and the actual charge of an electron were known, finding the mass was pretty easy. m e =( mass to charge ratio ) ( charge ) = ( 6 x 10 -12 kg C -1 ) ( 1.5924 x 10 -19 C) = 1 x 10 -30 kg The modern value for the mass of an electron is:9.109 390 x 10 -31 kg 5.485 799 x 10 -4 u

8. 7 - 8 Discovery of the nucleus 1909 1909Rutherford bombarded thin metal foils with alpha particles (helium ions). He felt that the particles would pass through the foils. When tested, he observed that about 1 alpha particle in 8000 was deflected by the foil. This deflection indicated the existence of a small, dense, positively charged nucleus.

9. 7 - 9 Discovery of the nucleus

10. 7 - 10 Determination of nuclear charge Rutherford estimated that the charge of the nucleus of an atom was about one half of the atomic mass. Moseley, while working for Rutherford, developed a more accurate measurement. While working with cathode rays, he measured the wavelength of the X-rays produced. He found that a direct relationship exists between the atomic number and the square root of the frequency.

11. 7 - 11 Determination of nuclear charge Moseley concluded that the charge of the nucleus was an integer. Further, it was the same as the number of electrical units (electrons) but of opposite charge. Moseley concluded that the charge of the nucleus was an integer. Further, it was the same as the number of electrical units (electrons) but of opposite charge. Atomic number X-Ray Frequency 1/2

12. 7 - 12 Discovery of protons and neutrons Measurements of the mass to charge ratio of the nucleus were made in a manner similar to earlier work with electrons. The ratio was found to be dependent on the gas that was used for the experiment. Hydrogen was found to produce particles with the lowest mass. These particles were assumed to be common to all atoms and were called protons.

13. 7 - 13 Discovery of protons and neutrons 1932 1932Chadwick observed that when beryllium-9 was exposed to alpha particles, particles with the same mass as protons but no charge were given off. These were called neutrons and are present in all atoms except hydrogen-1. They contribute to the force that holds the nucleus together and reduce the repulsive force between positively charged protons.

14. 7 - 14 Traveling waves Much of what has been learned about atomic structure has come from observing the interaction of visible light and matter. An understanding of waves and electromagnetic radiation would be helpful at this point. Let’s start with some basic definitions.

15. 7 - 15 Waves Some definitions Wavelength, Wavelength, The distance for a wave to go through a complete cycle.Amplitude Half of the vertical distance from the top to the bottom of a wave. Frequency, Frequency, The number of cycles that pass a point each second.

16. 7 - 16 Waves wavelength node + - ++ ---

17. 7 - 17 Electromagnetic radiation A form of energy that consists of perpendicular electrical and magnetic fields that change, at the same time and in phase, with time. The SI unit of frequency ( ) is the hertz, Hz 1 Hz = 1 s -1 Wavelength and frequency are related = c = c c is the speed of light, 2.998 x10 8 m/s

18. 7 - 18 Electromagnetic radiation 10 20 10 15 10 10 10 5 10 0 Gamma rays X- raysUltravioletVisible Infrared Microwave Television Radio 10 -10 10 -5 10 0 10 5 10 10 Wavelength ( ), m Frequency ( ), s -1

19. 7 - 19 Separation of light ‘White’ light is actually a blend of all visible wavelengths. They can separated using a prism.

20. 7 - 20 Electromagnetic radiation Electromagnetic radiation (EM) and matter Transmission Transmission - EM will pass through matter - - no interaction. Absorption Absorption - EM is absorbed by an atom, ion or molecule, taking it to a higher energy state. Emission Emission - the release of energy by an atom, ion or molecule as light, taking it to a lower energy state.

21. 7 - 21 Particle properties Although EM has definite wave properties, it also exhibits particle properties. Photoelectric effect. First observed by Hertz and then later explained by Einstein. When light falls on Group IA metals, electrons are emitted (photoelectrons). As the light gets brighter, more electrons are emitted. The energy of the emitted electrons depends on the frequency of the light.

22. 7 - 22 Photoelectric effect The cathode has a photoemissive surface. When light hits the cathode electrons are ejected. They are collected at the anode and can be measured. cathode anode

23. 7 - 23 Photoelectric effect photons Studies of this effect led to the discovery that light existed as small particles of electromagnetic radiation called photons. The energy of a photon is proportional to the frequency. Photon energy = h Photon energy = h The energy is inversely proportional to the wavelength. Photon energy = h c -1 h - Plank’s constant, 6.626 x 10 -34 J. s

24. 7 - 24 Photon energy example Determine the energy, in kJ/mol ofa photon of blue-green light with a wavelength of 486 nm. energy of a photon = = = 4.09 x 10 -19 J / photon h c (6.626 x 10 -34 J. s)(2.998 x 10 8 m. s -1 ) (4.86 x 10 -7 m)

25. 7 - 25 Photon energy example We now need to determine the energy for a mole of photons (6.022 x 10 23 ) Energy for a mole of photons. = (4.09 x 10 -19 J / photon) (6.022 x 10 23 photons/mol) = 246 000 J/mol Finally, convert to kJ = ( 244 000 J/mol ) = 244 kJ / mol 1 kJ 10 3 J

26. 7 - 26 Bohr model of the atom Bohr studied the the spectra produced when atoms were excited in a gas discharge tube. He observed that each element produced its own set of characteristic lines.

27. 7 - 27 Bohr model of the atom Balmer later determined an empirical relationship that described the spectral lines for hydrogen. 1 = 1.097 x 10 7 m -1 ( ) 1221222 1n21n2 - n = 2, 3, 5,... Spectra of many other atoms can be described by similar relationships.

28. 7 - 28 Bohr model of the atom Bohr proposed a model of how electrons moved around the nucleus. He wanted to explain why electrons did not fall in to the nucleus. He also wanted to account for spectral lines being observed. He proposed that the energy of the electron was quantized - only occurred as specific energy levels.

29. 7 - 29 Bohr model of the atom In the Bohr model, electrons can only exist at specific energy levels (orbit). Each energy level was assigned a principal quantum number, n. Energy

30. 7 - 30 Bohr model of the atom The Bohr model is a ‘planetary’ type model. Each principal quantum represents a new ‘orbit’ or layer. The nucleus is at the center of the model.

31. 7 - 31 Bohr model of the atom Bohr was able to use his model hydrogen to: Account for the observed spectral lines. Calculate the radius for hydrogen atoms. His model did not account for: Atoms other than hydrogen. Why energy was quantized. His concept of electrons moving in fixed orbits was later abandoned.

32. 7 - 32 Wave theory of the electron 1924 1924De Broglie suggested that electrons have wave properties to account for why their energy was quantized. He reasoned that the electron in the hydrogen atom was fixed in the space around the nucleus. He felt that the electron would best be represented as a standing wave. As a standing wave, each electron’s path must equal a whole number times the wavelength.

33. 7 - 33 De Broglie proposed that all particles have a wavelength as related by: =wavelength, meters h=Plank’s constant m=mass, kg v=frequency, m/s De Broglie waves = hmvhmv

34. 7 - 34 De Broglie waves Using De Broglie’s equation, we can calculate the wavelength of an electron. = 6.6 x 10 -34 kg m 2 s -1 (9.1 x 10 -31 kg)(2.2 x 10 6 m s -1 ) The speed of an electron had already been reported by Bohr as 2.2 x 10 6 m s -1. = 3.3 x 10 -10 m

35. 7 - 35 Heisenberg uncertainty principle In order to observe an electron, one would need to hit it with photons having a very short wavelength. Short wavelength photons would have a high frequency and a great deal of energy. If one were to hit an electron, it would cause the motion and the speed of the electron to change. Lower energy photons would have a smaller effect but would not give precise information.

36. 7 - 36 Heisenberg uncertainty principle According to Heisenberg, it is impossible to know both the position and the speed of an object precisely. He developed the following relationship:  x  v As the mass of an object gets smaller, the product of the uncertainty of its position and speed increase. h 4  m >

37. 7 - 37 Quantum model of the atom Schrödinger developed an equation to describe the behavior and energies of electrons in atoms. His equation is similar to one used to describe electromagnetic waves. While the equation is too complicated to write here, we can still use the results. Each electron can be described in terms of its quantum numbers.

38. 7 - 38 Quantum numbers Principal quantum number, n Tells the size of an orbital and largely determines its energy. n = 1, 2, 3, …… Angular momentum, l The number of subshells that a principal level contains. It tells the shape of the orbitals. l = 0 to n - 1

39. 7 - 39 Quantum numbers Magnetic quantum number, m l Describes the direction that the orbital projects in space. m l = l to + l (all integers, including zero) For example, if l = 2, then m l would have values of -2, -1, 0, 1 and 2. Knowing all three numbers provide us with a picture of all of the orbitals.

40. 7 - 40 Quantum numbers subshell # of nlm l label orbitals 1001s1 2002s1 1 -1, 0, 12p3 3003s1 1 -1, 0, 13p3 2-2, -1, 0, 1, 23d5 4004s1 1 -1, 0, 14p3 2-2, -1, 0, 1, 24d5 3 -3, -2, -1, 0, 1, 2, 34f7

41. 7 - 41 The s orbital The s orbital is a sphere. Every level has one s orbital.

42. 7 - 42 p orbitals There are three p orbitals: p x, p y and p z

43. 7 - 43 Representative d orbitals

44. 7 - 44 Representative f orbitals

45. 7 - 45 Combined orbitals - n=2

46. 7 - 46 Combined orbitals - n=3

47. 7 - 47 Electron spin Pauli added one additional quantum number that would allow only two electrons to be in an orbital. Spin quantum number, m s. It can have values of +1/2 and -1/2 Pauli also proposed that no two electrons in an atom can have the same set of four quantum numbers -- Pauli exclusion principle Pauli exclusion principle.