1. The Electron Dr. M Hazlett Mandeville High School

2. A Brief Review J.J. Thompson and others at first believed that the electrons floated in a sphere filled with a positive substance Later, Neils Bohr came up with the Planetary Model where the electrons orbited the nucleus in 7 layers Erwin Schrodinger showed that since the electrons are orbiting at the speed of light, it creates an electron cloud around the nucleus

3. Werner Heisenberg concluded that since electrons are affected by photons (light), no one can be certain as to their actual location Louis de Broglie explained that electrons exhibit both particle and wave behavior and determined its wave function (ψ) The electron is considered a lepton, part of non- baryonic matter Electrons are created through neutron decay Finally, Millikan determined the electron’s mass (9.109 x 10 -28 g) and charge (1.602 x 10 -19 C).

4. Electrons and Ions Atoms losing or gaining electrons become ions – If the atom has lost an electron, its has more p + than e -, thus a positive charge  cation – If the atom gains an electron, its has more e - than p + and a negative charge  anion This loss and/or gaining of electrons determines chemical bonds and reactions

5. The Standard Model and Electrons In the standard (quantum) model, the e - in the atoms have discrete energies – There are seven energy levels (n) – Electrons start in their ground state – their lowest energy level If the electron gains energy, it moves up one or more levels The e - is now in its excited state When the e - releases its energy in the form of a photon (ϒ), it returns to its ground state

6. The wavelength of light (λ) is related to the energies of the electron’s ground and excited states – E High - E Low = h c λ where h c = 1.196 x 10 5 KJ●nm/mol

7. The Bohr Model and Lightwaves Using spectroscopy, can determine the λ series of electrons and photons of atoms Bohr, using a CRT filled with H 2 noted that certain light wavelengths appeared when 5000 v of electricity connected to the tube Conclusion: The light represented the energy of the e - as it jumped from its ground to excited state and back

8. Different wavelengths and frequencies, different colors and thus – different elements

9. Other scientists used Bohr’s idea and created ranges or series of wavelength (λ) and frequency (ϒ) corresponding to the e - ground state, its excited state, and its final energy level Each series is based on light wavelength and frequency Series can measure absorbed and emitted energy in the form of photons

10. Series: 1.Lyman Series Final energy level (n 1 ) = 1 Light will be in UV range e - goes from n ≥ 2 to n 1 = 1 e - Start Level λ (nm) 2122 3103 497.2 594.9 693.7 +91.1

11. 2.Balmer Series n 1 = 2 Light is partially visible – R, B, V e - goes from n ≥ 3 to n 1 = 2 e - Start Level λ (nm) 3656 4486 5434 6410 7397 +365

12. 3.Paschen Series n 1 = 3 Infrared light e - Start Level λ (nm) 41870 51280 61090 71000 8954 +820

13. Other series are used: – Brackett Seriesn 1 = 4 – Pfund Seriesn 1 = 5 – Hymphreys Seriesn 1 = 6 Equations: c = λϒ (Wavelength x Frequency) 1 / λ = R ( 1 / (n 1 ) 2 - 1 / n 2 ) Wavelength of emitted/absorbed photon R = Rydberg Constant (1.097373 x 10 7 m) E photon = h ϒ = h c / λ

15. – If the e - goes to its ground state (n = 1)  UV light and use the Lyman Series – If the e - goes to its excited state (n = 2)  visible light and the Balmer Series – If the electron moves three levels (n = 3)  infrared light and the Paschen Series is used

16. Electron Wavelengths

19. Hydrogen’s Emission Spectrum

20. In the Standard Model...... There are four basic quantum numbers: 1. n = period = e - energy level, 1 through 7 Maximum electron capacity per level = 2n 2 2. l = orbital type or sublevel = s, p, d, f s - sharp or spherical ○ p – principle ∞ d – diffuse, four-leaf clover shaped f – fundamental, 6 leaf clover shaped

21. 3. m L = orientation of orbitals 4. m s = spin, either +1/2 or -1/2  Pauli Exclusion Principle – No two e - in the same atom can have the same quantum numbers In each suborbital, there will be two e -, and each will have opposite spins  Aufbau Principle – Atoms will fill up their orbitals starting with the lowest energies to highest

22.  Hund’s Rule – Each suborbital will have a maximum of 2 e - ; with opposite spins. Each suborbital in an energy level must have one e - before any receive a second one

23. Valency Valence Electrons – the electrons in the outermost shell or energy level – Octet Rule – atoms seek to fill to completion the outermost energy level by gaining or giving away electrons – Valence electrons will determine bonds and reactions – Valency provides grouping of elements on Periodic Table (Groups I through VIII)

24. Noting Electron Configuration We need to be able to show how an atom’s electrons are configured (distributed) in order to explain chemical bonds and reactions There are several ways to do this, but first: – You must know and understand the element’s valence energy level (the period) – You must know the # valence e - (Group I to VIII or if a Transitional Metal) – And you must know the orbital grouping: s, p, d, or f

25. Aufbau Notation This gives the valence energy level, the valence orbital shape and the # of valence electrons in that level Example: 2 p 5 Energy Level Orbital # e -

26. Electron Diagram – Aufbau Principle 7 p 7 s 6 d 6 p 5 d 4 f 6 s 5 p 4 d 5 s 4 p 3 d 4 s 3 p 3 s 2 p 2 s 1 s Start

27. Aufbau Notation: Long Format: – Na w/ 11 electrons would be: 1s 2 2s 2 2p 6 3s 1 – Short Format only gives the LAST Energy Level and the number of electrons in it Na  3s 1 – Without this e -, Na would isoelectric w/ Ne (have the same e - configuration) Cl -17 would be 1s 2 2s 2 2p 6 3s 2 3p 5 in Long Format Cl – 17 would be 3s 2 3p 5 in Short Format and its Lewis Dot Structure would have 7 dots around the elemental symbol

28. Often, you will see another shorthand way of writing an element’s electron configuration – This is called the Noble Gas Configuration – Noble Gases are in Column VIII – For example, using Na, Na’s configuration would be: [Ne]3s 1 – The [ ] means isoelectric (the starting point) – Li would be [He]2s 1 - La  [Xe]6s 2 5d 1 – Mn would be [Ar]4s 2 3d 5 - Ce  [Xe]6s 2 5d 1 4f 1 – Zn would be [Ar]4s 2 3d 10 - Pr  [Xe]6s 2 4f 3

29. Lewis Dot Diagrams Dot diagrams demonstrate the type of covalent bonds an element may make under certain conditions The element’s symbol is surrounded by up to 8 dots, each dot representing a valence e - Maximum of 2 dots per side (2 dots x 4 sides = 8 total, an octet valence level) It can be used for single atoms or to show molecules and their bonds and shapes

30. Determine dots from either doing the e - mapping OR learn to read the periodic table! Often, for transitional metals – need to do mapping (problem if ion) Remember Hund’s Rule Watch for + and – ions and either add or subtract the correct number of electrons Each side of elemental symbol represents one of the last 4 suborbitals (□) of the SAME Energy Level! Do not have to have 8 e - dots, can have openings – This is where bonds to other atoms can occur

31. Lewis Dot Diagrams

35. Orbitals

36. Next Stop..... Once we have the diagramming conquered, then it is onto the types of bonds – ionic and covalent From there – we can diagram bonded atoms (molecules and polyatomic ions) Then, we move onto reactions