Chapter 4: Arrangement of Electrons in Atoms

1. Chapter 4:Arrangement of Electrons in Atoms Chemistry

2. Development of a New Atomic Model • There were some problems with the Rutherford model…It did not answer: • Where the e- were located in the space outside the nucleus • Why the e- did not crash into the nucleus • Why atoms produce spectra at specific wavelengths

3. Properties of Light • Wave-Particle Nature of Light – early 1900’s • A Duel Nature • It was discovered that light and e- both have wave-like and particle-like properties

4. Wave Nature of Light • Electromagnetic radiation – form of energy that exhibits wave-like behavior as it travels through space • Electromagnetic spectrum • All the forms of electromagnetic radiation • Speed of light in a vacuum • 3.0 x 108 m/s

5. Wave Nature of Light • Wavelength • Distance between two corresponding points on adjacent waves • λ • nm • Frequency • Number of waves that pass a given point in a specified time • ν • Hz - Hertz

6. Wave Nature of Light • Figure 4-1, page 92 • Equation • c=λν • Indirectly related! • Spectroscope • Device that separates light into a spectrum that can be seen • Diffraction Grating – the part of the spectroscope the separates the light

7. Particle Nature of Light • Quantum • Minimum quantity of energy that can be lost or gained by an atom • Equation • E=hν • Direct relationship between quanta and frequency • Planck’s Constant (h) • h=6.626 x 10-34 Js

8. Particle Nature of Light • Photon • Individual quantum of light; “packet” • The Hydrogen Atom • Line emission spectrum (Figure 4-5, page 95) • Ground State • Lowest energy state (closest to the nucleus) • Excited State • State of higher energy **When electron drops from its excited state to its ground state, a photon is emitted! This produces a bright-line spectrum. Each element has a characteristic bright-line spectrum – much like a fingerprint!**

9. http://jersey.uoregon.edu/vlab/elements/Elements.html

10. Particle Nature of Light • Why does an emission spectrum occur? • Atoms get extra energy – voltage • The e- jumps from ground state to excited state • Atoms return to original energy, e- drops back down to ground state • Continuous spectrum • Emission of continuous range of frequencies

11. Particle Nature of Light • Bohr Model of the H atom • 1913 – Danish physicist – Niels Bohr • Single e- circled around nucleus in allowed paths or orbits • e- has fixed E when in this orbit (lowest E closest to nucleus) • Lot of empty spacebetween nucleus and e- in which e- cannot be in • E increases as e- moves to farther orbits • http://chemmovies.unl.edu/ChemAnime/BOHRQD/BOHRQD.html

12. Particle Nature of Light • Bohr Model (cont) • ONLY explained atoms with one e- • Therefore – only worked with hydrogen!!

13. Particle Nature of Light • Spectroscopy • Study of light emitted by excited atoms • Bright line spectrum

14. The Quantum Model of the Atom • e- act as both waves and particles!! • De Broglie • 1924 – French physicist • e- may have a wave-particle nature • Would explain why e- only had certain orbits • Diffraction • Bending of wave as it passes by edge of object • Interference • Occurs when waves overlap

15. The Quantum Model of the Atom • Heisenberg Uncertainty Principle • 1927 – German physicist • It is impossible to determine simultaneously both the position and velocity of an e- 12:28-14:28

17. The Quantum Model of the Atom • Schrodinger Wave Equation • 1926 – Austrian physicist • Applies to all atoms, treats e- as waves • Nucleus is surrounded by orbitals • Laid foundation for modern quantum theory • Orbital – main energy level; 3D region around nucleus in which an e- can be found • Cannot pinpoint e- location!!

18. Quantum Numbers • Quantum Numbers • Solutions to Schrodinger’s wave eqn • Probability of finding an e- • “address” of e- • Four Quantum Numbers • Principle • Anglular Momentum • Magnetic • Spin

19. Principle Quantum Number • Which main energy level? (“orbital” “shell”) • Symbol- n • n is normally 1-7 (corresponds to period on periodic table) • Higher the n, the greater the distance from the nucleus

20. Angular Momentum Quantum Number • What is the shapeof the orbital? • F shape • Symbol – l • l = s,p,d,f • When n = 1, l = s n = 2, l = s,p n = 3, l = s,p,d n = 4, l = s,p,d,f • http://www.chemeng.uiuc.edu/~alkgrp/mo/gk12/quantum/

21. Magnetic Quantum Number • Orientation of orbital around nucleus • Symbol – m • s – 1 p – 3 d – 5 f – 7 • Every orientation can hold 2 e-!! • Figures 4-13, 4-14, 4-15 on page 102-103

22. Spin Quantum Number • Each e- in one orbital must have opposite spins • Symbol – s • + ½ , - ½ • Two “allowed” values and corresponds to direction of spin

23. Electron Configuration • Electron configurations – arrangements of e- in atoms • Rules: • Aufbau Principle – an e- occupies the lowest energy first • Hund’s Rule – each orbital is filled with 1e- first and then the 2nd e- will fill • Pauli Exclusion Principle – no 2 e- in the same atom can have the same set of QN 14:30-18:25

24. Electron Configuration • Representing electron configurations • Use the periodic table to write! • Know the s,p,d,f block and then let your fingers do the walking!

25. Electron Configuration

26. Representing Electron Configurations • Three Notations • Orbital Notation • Electron Configuration Notation • Electron Dot Notation

27. Orbital Notation • Uses a series of lines and arrows to represent electrons • Examples

28. Orbital Notation • More examples

29. Electron Configuration Notation • Eliminates lines and arrows; adds superscripts to sublevels to represent electrons • Long form examples

30. Electron Configuration Notation • Short form examples – “noble gas configuration”

31. Electron Dot Notation • Outer shell e- • Inner shell e- • Highest occupied energy level / highest principle quantum number • Valence electrons – outermost e- • Examples

32. Electron Dot Notation • More examples

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