1. Rutherford’s Atom Electromagnetic Radiation Emission of Energy by Atoms Energy Levels of Hydrogen Atomic Models Hydrogen Orbitals Electron Arrangements Electron Configuration Atomic Properties and the Periodic Table pages 278-315 1 Unit 10 – Atomic Theory

2. Atomic Theory The concept of atoms explains many important observations, such as why compounds always have the same composition (a specific compound always contains the same types and numbers of atoms) and how chemical reactions occur (they involve a rearrangement of atoms). We learned to picture the atom as a positively charged nucleus composed of protons and neutrons at its center and electrons moving around the nucleus in a space very large compared to the nucleus. In this unit, we develop a picture of the electron arrangements in atoms. 2

3. Rutherford’s Atom Replay Video 3 - Ernest Rutherford and the nucleus 3

4. Rutherford’s Atom What are the electrons doing? How are they arranged and how do they move? Rutherford suggested the electrons might revolve around the nucleus like the planets revolve around the sun in our solar system. He couldn’t explain why the negatively charged electrons aren’t attracted into the positive nucleus, causing the atom to collapse. More observations of the properties of atoms would be needed to understand the structure of the atom. To help understand these observations, we need to discuss the nature of light and how it transmits energy. 4

5. Electromagnetic Radiation Energy is transmitted by light, more properly called electromagnetic radiation. Think of a light bulb, the flame of a Bunsen burner or the sun. Many different forms of this energy exist, including x- rays, microwaves and radio waves as well as light. A wave is characterized by three properties: wavelength, frequency and speed. Wavelength (symbolized by the Greek letter lambda, λ) is the distance between consecutive wave peaks. Frequency (symbolized by the Greek letter nu, ν) indicates how many waves pass a certain point in a given period of time. The speed of a wave indicates how fast a peak travels. 5

6. Electromagnetic Radiation 6 Examples using ocean waves.

7. Electromagnetic Radiation Light (electromagnetic radiation) also travels as waves. The various types of electromagnetic radiation (x-rays, microwaves) differ in their wavelengths. X-rays have very short wavelength, while radio waves have very long wavelengths. 7

8. Electromagnetic Radiation 8 Radiation provides an important means of energy transfer. Energy from the sun reaches the earth mainly as visible and ultraviolet radiation. We visualize radiation (light) as a wave that carries energy through space. Sometimes, though, light does not act as a wave. Electromagnetic radiation can sometimes have properties that are characteristic of particles. Another way to think of light is as a stream of packets of energy called photons.

9. Electromagnetic Radiation 9 Is light a wave or a stream of particles of energy? It appears to both. This is called the wave-particle nature of light.

10. Electromagnetic Radiation 10

11. Emission of Energy by Atoms 11 When compounds are burned, they emit a color characteristic of the cation. Li +, for example, emits a red flame when burned. Na + burns with a yellow flame, Cu 2+ with a green flame. The colors of the flames result from atoms releasing energy in the form of visible light of specific wavelengths, or colors. The heat from the flame causes the atom to absorb energy. The atom becomes excited. Some of the excess energy is released as light. The atom moves to a lower energy state as it emits a photon of light.

12. Emission of Energy by Atoms 12

13. Emission of Energy by Atoms 13 When atoms receive energy, they become excited. They can release the energy by emitting light. The emitted energy is carried away by a photon. The energy of the photon corresponds exactly to the energy change of the emitting atom. High energy photons correspond to short wavelength light. Low energy photons correspond to long wavelength light. The photons of red light have less energy than the photons of blue light because red light has a longer wavelength than blue light.

14. Energy Levels of Hydrogen 14 An atom with excess energy is said to be in an excited state. An excited atom can release some or all of its excess energy by emitting a photon and thus move to a lower energy state. The lowest possible energy state of an atom is called its ground state. Different wavelengths of light carry different amounts of energy per photon. When a photon is emitted, the energy contained in the photon correspond to the change in energy that the atom experiences in going from the excited state to the lower state.

15. Energy Levels of Hydrogen 15

16. Energy Levels of Hydrogen 16

17. Energy Levels of Hydrogen 17 When we study the photons of visible light emitted, we see only certain colors. Only certain types of photons are produced. Because only certain photons are emitted, only certain energy changes are occurring. So, hydrogen atoms must have certain discrete energy levels. We say the energy levels of hydrogen are quantized, that is, only certain values are allowed. Energy levels of all atoms are quantized.

18. Energy Levels of Hydrogen 18

19. Energy Levels of Hydrogen 19

20. Energy Levels of Hydrogen 20

21. The Bohr Model of the Atom 21 This model proposed by Niels Bohr worked well to explain the hydrogen atom, but the model did not explain other atoms.

22. Wave Mechanical Model 22 Louis de Broglie and Erwin Schrodinger developed the wave mechanical model. The model gives no information about when the electron occupies a certain point or how the electron moves.

23. The Hydrogen Orbitals 23 The probability map is called an orbital. The orbital shown in Figure 10.20 is called the 1s orbital and describes the ground (lowest) state of energy for hydrogen.

24. The Hydrogen Orbitals 24 Hydrogen has discrete energy levels. They are called principal energy levels and labelled with an integer. Each principal energy level has sublevels.

25. The Hydrogen Orbitals 25 Principal level 2 has 2 sublevels. They are called 2s and 2p. Principal level 3 has 3 sublevels called 3s, 3p and 3d. Principal level 4 has 4 sublevels called 4s, 4p, 4d and 4f.

26. The Hydrogen Orbitals 26 The principal levels describe size and shape. The s orbital is spherical. Level 1 is smaller than level 2, which is smaller than level 3.

27. The Hydrogen Orbitals 27 The three 2p orbitals are lobed, not spherical. They are oriented along the x, y or z axis.

28. The Hydrogen Orbitals 28 The shapes of the five 3d orbitals are shown below.

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30. Electron Arrangements 30 An atom has as many electrons as it does protons, so all atoms beyond hydrogen have more than one electron. Each electron appears to spin like a top on its axis. It can only spin in one direction. We represent spin with an arrow, ↑ or ↓. Electrons in the same orbital must have opposite spins. This leads to the Pauli exclusion principle: an atomic orbital can hold a maximum of two electrons and those two electrons must have opposite spins.

31. Electron Arrangements 31 Hydrogen has an atomic number of 1 (Z =1) and therefore a single electron to have a net charge of zero. To show its electron configuration, we write the principal energy level followed by the sublevel, 1s. The number of electron in the orbital is placed as a superscript, 1s 1. The electron configuration can also be shown using an orbital diagram, or box diagram, as below.

32. Electron Arrangements 32 Hydrogen (Z=1)1s 1 Helium (Z=2)1s 2 Lithium (Z=3) 1s 2 2s 1 Berylium (Z=4) 1s 2 2s 2 Boron (Z=5) 1s 2 2s 2 2p 1 Carbon (Z=6) 1s 2 2s 2 2p 2 Nitrogen (Z=7) 1s 2 2s 2 2p 3 Oxygen (Z=8) 1s 2 2s 2 2p 4 Fluorine (Z=9) 1s 2 2s 2 2p 5 Neon (Z=10) 1s 2 2s 2 2p 6 The orbital diagram for nitrogen is below.

33. Electron Arrangements 33 Sodium (Z=11) 1s 2 2s 2 2p 6 3s 1 or [Ne] 3s 1 Magnesium (Z=12) [Ne] 3s 2 Aluminum (Z=13)[Ne] 3s 2 3p 1 Silicon (Z=14)[Ne] 3s 2 3p 2 Phosphorous (Z=15)[Ne] 3s 2 3p 3 Sulfur (Z=16)[Ne] 3s 2 3p 4 Chlorine (Z=17)[Ne] 3s 2 3p 5 Argon (Z=18)[Ne] 3s 2 3p 6

34. Electron Arrangements 34 Valence electrons are the electrons in the outermost (highest) principal energy level of an atom. These are the electrons involved in bonding of atoms to each other. Also note that the atoms of elements in the same group have the same number of electrons in a given type of orbital, except that the orbitals are in different principal energy levels. Elements with the same valence electron arrangement show very similar chemical behavior.

35. Electron Arrangements 35 The order of filling orbitals changes for Z=19. Experiments show that the chemical properties of potassium are very similar to lithium and sodium. We predict that the 4s orbital will fill before the 3d orbital. This means that principal energy level 4 begins to fill before level 3 is full. Potassium (Z=19)[Ar] 3s 2 3p 6 3d 1 Potassium (Z=19)[Ar] 3s 2 3p 6 4s 1 Calcium (Z=20)[Ar] 3s 2 3p 6 4s 2 Scandium (Z=21)[Ar] 3s 2 3p 6 4s 2 3d 1

36. Electron Arrangements 36

37. Electron Arrangements 37

38. Periodic Trends 38 As you go down a group, metals are more likely to lose an electron.

39. Periodic Trends 39 Ionization energy is the energy required to remove an individual atom in the gas phase. Metals have relatively low ionization energies.

40. Periodic Trends 40

41. Electron Configuration and Periodic Trends 41 Homework Pages 311-312; problems 50, 54, 60, 80 and 82