1. Electrons in Atoms Ch. 13

2. Models of the Atom 13-1

3. The evolution of Atomic Models Dalton (1766-1844): atom indivisible J.J. Thomson (1856-1940): –“Plum-pudding” model – negative electrons stuck in positively charged material Rutherford (1871-1937): –Electrons surround dense nucleus, rest of atom is empty space Bohr (1885-1962): –“Planetary model,” electrons fixed in energy levels around nucleus

5. The Quantum Mechanical Model Quantum Mechanical Model: –Estimates the probability of finding an electron in a certain place using the Schrodinger equation. –“Fuzzy cloud” model; where the cloud is more dense the probability of finding the electron is high, where the cloud is less dense the probability is low.

6. Energy Levels Electrons move around the nucleus in energy levels. Quantum of energy = amount of E required to move to a higher level. When they move towards the nucleus (down a level) they release energy When they move away from the nucleus (up a level) they require energy. The farther from the nucleus the energy level is, the more energy is required to move up a level (away from nucleus).

7. Energy Levels Quantum number (n) refers to an energy level –n = 1, 2, 3, 4, …7, values increase going away from nucleus. –Each energy level fits a certain amount of electrons: Level 1 = 2 electrons Level 2 = 8 electrons Level 3 = 18 electrons Level 4 = 32 electrons

8. Sublevels + Orbitals Within each energy level there are sublevels; the number of sublevels is equal to the quantum number. –Ex: Energy level 4 has 4 sublevels within it. A sublevel is made up of atomic orbitals: s, p, d, f –Orbital s fits 2 electrons total –Orbital p fits 6 electrons total –Orbital d fits 10 electrons total –Orbital f fits 14 electrons total –s fills up first, then p, then d, then f

9. Energy Level (n) Sublevel/Orbi tal Electrons in each Sublevel/Orbita l Total # of electrons in Level 11s22 22s 2p 26268 33s 3p 3d 2 6 1018 44s 4p 4d 4f 2 6 10 1432

10. Atomic Orbitals OrbitalShape# of Electrons sSpherical2 pDumbbell6 dClover-leaf10 fComplex14

11. DRAW! s - orbital p - orbital d - orbital f - orbital

12. Electron Arrangement in Atoms 13-2

13. d-1 f-2 5d 1 6d 1 Label + Color on white, blank, large periodic table!

14. Electron Arrangement in Atoms Period 1 - 1s 2 Period 2 - 2s 2 2p 6 Period 3 - 3s 2 3p 6 Period 4 - 4s 2 3d 10 4p 6 Period 5 - 5s 2 4d 10 5p 6 Period 6 - 6s 2 4f 14 5d 10 6p 6 Period 7 - 7s 2 5f 14 6d 10 7p 6 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 7s 2 5f 14 6d 10 7p 6

15. Electron Configuration Notation Notation used to represent electron configurations: –H: 1s 1 –He: 1s 2 –Li: 1s 2 2s 1 –Be: 1s 2 2s 2 –B: 1s 2 2s 2 2p 1 Energy level Sub level/orbital # of electrons in sublevel/orbital

16. Color the 4 sublevels + make a key d-1 f-2

17. …Then write the configuration of each element!

18. You Try! C: –(6): 1s 2 2s 2 2p 2 F: –(9): 1s 2 2s 2 2p 5 Ne: –(10): 1s 2 2s 2 2p 6 Na: –(11)1s 2 2s 2 2p 6 3s 1 P: –(15)1s 2 2s 2 2p 6 3s 2 3p 3 Ca: –(20)1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 Ir: –(77)1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 7 Cm: –(96)1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 7s 2 5f 7 Write the electron configurations for the following:

19. Abbreviated form: shows preceding noble gas and the configuration of only the last energy level! –Mg: 1s 2 2s 2 2p 6 3s 2 or [Ne] 3s 2 –B: 1s 2 2s 2 2p 1 or [He] 2s 2 2p 1 –Si: 1s 2 2s 2 2p 6 3s 2 3p 2 or [Ne] 3s 2 3p 2 –Al: 1s 2 2s 2 2p 6 3s 2 3p 1 [Ne] 3s 2 3p 1 –Xe: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 [Kr] 5s 2 4d 10 5p 6

20. What happens in the fourth period? –After 4s 2, comes 3d 10, then 4p 6 –Scandium (#21): 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 1 or [Ar] 4s 2 3d 1 –Copper is [Ar] 4s 2 3d 9 –Bromine is [Ar] 4s 2 3d 10 4p 5 What happens in the sixth period? –After 6s 2, comes 4f 14, then 5d 10, then 6p 6 –Tungsten (W) is [Xe] 6s 2 4f 14 5d 4

21. Why is 3d on the 4 th row after 4s? ______ energy than 4s and ______energy than 5p. Why is 4f on the 6 th row after 6s? ______ energy than 6s and ______energy than 5d Aufbau principle: lowest energy orbitals are filled first! Sublevel order = 1s,2s,2p,3s,3p,4s,4p,5s,4d,5p,6s,4f,5d,6p

22. Orbital Notation Rules 1)Aufbau principle: electrons enter orbitals of lowest energy first. 2)Pauli exclusion principle: an atomic orbital may describe at most 2 electrons. 3)Hund’s rule: one electron enters each orbital until ALL orbitals contain 1 electron with parallel spins.

24. Light and Atomic Spectra 13-3 (only pg. 372-375) PLAY: electromagnetic spectrum song:

25. Electromagnetic Spectrum Energy in the form of electromagnetic radiation (radiant energy) travels in waves Waves transfer the energy from one place to another Ex: radio waves, TV, microwave, visible light, x-rays, gamma rays, infrared, UV All forms of radiant energy are part of the electromagnetic spectrum

26. Low energyHigh Energy

27. Wavelength + Frequency Two main properties of electromagnetic waves: 1) Frequency 2) Wavelength Wavelength is the distance between two corresponding peaks or troughs.

28. Frequency is the number of wave cycles per second. Wavelength is inversely proportional to frequency wavelength frequency Wave length Frequency Wave length Frequency

29. Higher frequency waves (short wavelength) have high energy –Ex: gamma rays, x-rays, ultraviolet rays –Ex: Violet light in visible spectrum Low frequency waves (long wavelength) have low energy –Ex: radio waves, microwaves, infrared (heat) waves –Ex: Red light in visible spectrum

30. Light and Atomic Spectra Electrons absorb energy and move to higher energy states/levels Electrons give off that energy in the form of light when they fall back down to lower energy states, or ground state. ALL electromagnetic waves travel at the speed of light in a vacuum – 300 million meters per second or (3.0 x 10 8 m/s)

31. When atoms are energized by an electric current they emit light. When this light is passed through a prism they produce an emission spectrum. Each element has its own unique atomic emission spectrum fingerprint

34. http://phys.educ.ksu.edu/vqm/html/emission.html