Electronic Structure of Atoms

1. Electronic Structure of Atoms Chapter 6 Chemistry 100

2. What is light? Light is obviously real - it is part of our world. Darkness is the absence of light Light is NOT a solid, a liquid, or even a gas So what is it? It is a form of energy We call it a form of radiant energy because it carries energy through space

3. Electromagnetic radiation Visible light is a type of electromagnetic radiation Other types include: infra-red, ultra-violet, X-rays, gamma rays, radar waves, microwaves, radio and TV waves Electromagnetic radiation has wave like properties

4. Waves Wavelength  (lamda) Frequency  (nu) Speed c (see?)   = c c = 3.00  108m/s for all types of electromagnetic radiation. So how is IR different from UV, for example?

6. Electromagnetic Spectrum Different types of electromagnetic radiation have different wave lengths () and different frequencies () Frequency: number of cycles (vibrations) per second. Unit is second-1 or s-1. The Hertz is the SI unit for frequency. 82,000 s-1 is the same as 82 kHz (kiloHertz)

7. Units for wavelength Unit Symbol Length (m) Type of Radiation -10 10 Angstrom Å X-ray -9 10 Nanometre nm UV & visible -6 m m 10 Micrometre IR -3 10 Millimetre mm IR -2 10 Centimetre cm microwave Metre m 1 TV, Radio

8. Max Planck and his constant h Suggested that energy is quantized - comes in small chunks E = h where n = 1, 2, 3 Compare the potential energy of a brick on a staircase to one on a slope

9. Can this be true? • We do not find that energy is quantized in everyday life - h is very small. • Cannot see the difference between 200,000,000h and 200,000,001 h • Einstein used Planck’s idea to explain the photoelectric effect • For electromagnetic radiation, E = h where  is the frequency of the radiation. High frequency  more energy

10. What is light? When we look how light behaves in experiments with lens, mirrors, etc., we are led to believe that light has wave properties In the photoelectric effect, light appears to consist of particles - which we call photons Dual nature of electromagnetic radiation

11. Bohr’s Atom Bohr said: if energy is quantized then the energy of an electron in an atom is quantized So the radius of its orbit cannot be any arbitrary value but must obey the quantum theory. Only certain orbits are allowed

12. Allowed Orbitals in Bohr’s Atom The quantity n is a quantum number

13. Bohr’s Atom 1913 Electrons move in orbitals with specified radii Each orbital is associated with a specific energy This explains why atoms emit (or absorb) light of well-defined frequency. Examples: the yellow sodium street light and the neon tube.

14. Wave Behaviour Louis de Broglie (1892-1987): If light can have both wave and particle behaviour, why not wave behaviour for all particles?  = h/m He talked about matter waves

15. Matter waves Find  for baseball moving at 100 km/h Find  for electron moving at 5.97 106 m/s

16. Heisenberg Postulated that there is a limit to how precisely we can measure both position and momentum The measurement effects the object being measured Heisenberg’s Uncertainty Principle

17. Schrödinger’s wave equation The quantity 2 provides information about the electron's position when it has energy E! In 1926, Schrödinger put de Broglie’s and Heisenberg’s ideas together and came up with the wave equation

18. Quantum Numbers • Schrödinger's wave equation has three quantum numbers. • Principal quantum number n. Has integer values 1, 2, 3 • Azimuthal quantum number, l. Allowed values values of 0, 1... up to n - 1 • Magnetic quantum number, ml. Allowed values -l … 0 … +l • There is also the Spin quantum number, ms. It can have a value of -½ or +½

19. Atomic orbitals The first shell n = 1 The shell nearest the nucleus l = 0 We call this the s subshell (l = 0) m = 0 There is one orbital in the subshell s = -½ The orbital can hold two electrons s = + ½ one with spin “up”, one “down” No two electrons in an atom can have the same value for the four quantum numbers: Pauli’s Exclusion Principle

20. The second shell l = 0 The s subshell m = 0 One orbital in the subshell s = -½ or + ½ Subshell can hold two electrons l = 1 The p subshell m = -1, 0, +1 Three orbitals in the subshell s = -½ or + ½ Each orbital can hold 2 electrons. p subshell can hold 6 electrons The second shell can hold 8 electrons: 2 in s orbitals and 6 in p orbitals n = 2 l = 0 or 1 There are two subshells

21. If the principal quantum number is n, the shell can hold up to 2n2 electrons

22. s Orbitals are Spherical

23. p Orbitals are Dumbbell Shaped

24. d Orbitals are Complex

25. Aufbau Principle 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 6f

26. Let’s do Sodium, Z = 11 Electronic configuration of Na is 1s22s22p63s1 Aufbau Principle 1s 2s 2p 3s …. First 2 electrons 1s2 that’s 2 Next 2 electrons 2s2 that’s 4 Six this time 2p6 that’s 10 1 more to go 3s1 that’s all, folks

27. Hund’s Rule Shall we use 1s 2s 2p () () ()() Or, shall we use 1s 2s 2p () () ()()() The configuration with the maximum spin is more stable.

28. Shorthand configurations The configuration of Neon is: 1s22s22p6 Na is 1s22s22p63s1, or in short form: [Ne]3s1 The configuration of Argon:1s22s22p63s23p6 K is: 1s22s22p63s23p64s1, which in short form becomes [Ar]4s1 Note the similarity of the two elements from the same group in the periodic table.The incomplete orbitals are 3s1 and 4s1.

29. Same group, similar configuration Fluorine: [He]2s22p5 Chlorine: [Ne]3s23p5 Bromine: [Ar]3d104s24p5 Iodine: [Kr]4d105s25p5 The outer-shell configuration in each case is s2p5 We need not be concerned with the d electrons here because d10 is a filled subshell.

30. Electronic Configuration & Periodic Table

31. I’m in a spin!!! Nitrogen has Atomic Number 7 Electronic Configuration: 1s22s22p3 Let’s draw an orbital diagram: