Chapter 5 “Electrons in Atoms”

1. Chapter 5“Electrons in Atoms” Chemistry I Manchester High School

2. Light and The Electromagnetic Spectrum • OBJECTIVES: • Identify the inadequacies in the Rutherford atomic model. • Identify the new proposal in the Bohr model of the atom. • Learn the energy order of the Electromagnetic Spectrum. • Calculate the wavelength, frequency, and Energy using two key equations.

3. Ernest Rutherford’s Model • Discovered dense positive piece at the center of the atom- “nucleus” • Electrons would surround and move around it, like planets around the sun • Atom is mostly empty space • It did not explain the chemical properties of the elements – a better description of the electron behavior was needed

4. Niels Bohr’s Model • Why don’t the electrons fall into the nucleus? • Move like planets around the sun. • In specific circular paths, or orbits, at different levels. • An amount of fixed energy separates one level from another.

5. Light • The study of light led to the development of the quantum mechanical model. • Light is a kind of electromagnetic radiation. • Electromagnetic radiation includes many types: gamma rays, x-rays, radio waves… • Speed of light = 2.998 x 108 m/s, and is abbreviated “c” • All electromagnetic radiation travels at this same rate when measured in a vacuum

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

7. Wavelength and Frequency • Are inversely related • As one increases the other decreases. • Different frequencies of light are different colors of light. • There is a wide variety of frequencies • The whole range is called the electromagneticspectrum

8. 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)

9. Question 1. • The number of cycles per unit of time is the wave’s • Wavelength • Amplitude • Frequency • Crest

10. Question 2. • The SI unit for frequency • Meters • Hertz • Moles • Seconds

11. Question 3. • The symbol used for wavelength is • ώ • β • ν • λ

12. 4. Yes/No Wave A has the longest wavelength. A. B.

13. 5.Yes/No Wave A has the highest frequency. A. B.

14. - Page 140 Use Equation: c =

15. Practice Problems: • What is the wavelength of radiation with a frequency of 1.50 x 1013 Hz? Does this radiation have a longer or shorter wavelength than red light? 2.00 x 10-5 m; longer than red light Answer

16. Practice Problems: • What is the frequency of radiation with a wavelength of 5.00 x 10-8 m? In what region of the electromagnetic spectrum is this radiation? 6.00 x 1015 m; ultraviolet Answer

17. 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)

18. The Math in Chapter 5 • There are 2 equations: • c =  • E = h Know these!

19. Examples • What is the wavelength of blue light with a frequency of 8.3 x 1015 Hz?

20. 2. What is the frequency of red light with a wavelength of 4.2 x 10-5 m?

21. 3. What is the energy of a photon of each of the above?

22. The Bohr Model of the Atom I pictured the electrons orbiting the nucleus much like planets orbiting the sun. However, electrons are found in specific circular paths around the nucleus, and can jump from one level to another. Niels Bohr

23. Bohr’s Model • Bohr looked at the hydrogen atom, with only one electron, in the first energy level.

24. Bohr’s Model • Adding energy (heat, electricity, or light) can move the electron up to a higher energy level. The electron is now said to be “excited”

25. Bohr’s Model • As the electron falls back to the ground state, it emits the energy in the form of light.

26. The light emitted when passed through a prism creates an atomic emission spectrum or bright line spectrum • Unique to each element, like fingerprints! • Very useful for identifying elements

27. Bohr’s Model • Energy level of an electron • analogous to the rungs of a ladder • The electron cannot exist between energy levels, just like you can’t stand between rungs on a ladder • Aquantum of energy is the amount of energy required to move an electron from one energy level to another

28. The Quantum Mechanical Model • Energy is “quantized” - It comes in chunks. • A quantum is the amount of energy needed to move from one energy level to another. • Since the energy of an atom is never “in between” there must be a quantum leap in energy.

29. The Quantum Mechanical Model • Has energy levels for electrons. • Orbits are not circular. • It can only tell us the probability of finding an electron a certain distance from the nucleus.

30. The Quantum Mechanical Model • The atom is found inside a blurry “electron cloud” • An area where there is a chance of finding an electron.

31. Atomic Orbitals • Principal Quantum Number (n) = the energy level of the electron: 1, 2, 3…∞ • Within each energy level, the complex math of Schrodinger’s equation describes several shapes. • These are called atomic orbitals- regions where there is a high probability of finding an electron. • Sublevels- designated by the letters s, p, d, and f

32. Principal Quantum Number Generally symbolized by “n”, it denotes the energy level (shell) in which the electron is located. Each energy level is divided into sublevels. n=1 has 1 sublevel n=2 has 2 sublevels n=3 has 3 sublevels

34. Schrodinger’s Wave Equation Equation for the probability of a single electron being found along a single axis (x-axis) Erwin Schrodinger Erwin Schrodinger

35. Bohr’s Model The fixed energies of an electron are called: Photons Quantums Energy levels Orbitals

36. Bohr’s Model The amount of energy required to move an electron from a lower energy level to a higher one. Photon Quantum Energy level Orbital

37. The Quantum Mechanical Model Energy levels are quantized The atom is found inside an electron cloud The orbits are not circular All of the above

38. Which of the following is the symbol for the energy level? e n el c

39. What is the number of electrons for an atom of Na-23? 11 23 12 10

40. Section 5.2Electron Arrangement in Atoms • OBJECTIVES: • Describe how to write the electron configuration for an atom.

41. Section 5.2Electron Arrangement in Atoms • OBJECTIVES: • Explain why the actual electron configurations for some elements differ from those predicted by the aufbau principle.

42. Electron Configurations… • …are the way electrons are arranged in various orbitals around the nuclei of atoms. Three rules tell us how: • Aufbau principle- electrons enter the lowest energy first. • This causes difficulties because of the overlap of orbitals of different energies – follow the diagram! • Pauli Exclusion Principle- at most 2 electrons per orbital – opposite spins

43. Electron Configurations • Hund’s Rule- When electrons occupy orbitals of equal energy, they don’t pair up until they have to. • Let’s write the electron configuration for Phosphorus • We need to account for all 15 electrons in phosphorus

44. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s aufbau diagram - page 133 Aufbau is German for “building up”

45. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The first two electrons go into the 1s orbital Notice the opposite direction of the spins • only 13 more to go...

46. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The next electrons go into the 2s orbital • only 11 more...

47. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The next electrons go into the 2p orbital • only 5 more...

48. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The next electrons go into the 3s orbital • only 3 more...

49. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The last three electrons go into the 3p orbitals. • They each go into separate shapes (Hund’s) • 3 unpaired electrons • = 1s22s22p63s23p3 Orbital notation

50. Orbitals fill in an order • Lowest energy to higher energy. • Adding electrons can change the energy of the orbital. Full orbitals are the absolute best situation. • However,half filled orbitals have a lower energy, and are next best • Makes them more stable. • Changes the filling order