5-1 Quantum Theory of the atom

# 5-1 Quantum Theory of the atom

## 5-1 Quantum Theory of the atom

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##### Presentation Transcript

1. 5-1 Quantum Theory of the atom • Bohr Model of the Atom • Proposed that electrons orbited the nucleus in circular paths. • Ground state- lowest allowable energy states of an atom. • Excited state- atom gains energy; H atoms can have many different excited states although it contains 1 e-. • Electrons move around a H atom in circular orbit • Orbits equal to a principal quantum number n, where n=1 is lowest energy level, closest to nucleus.

2. Bohr model of the atom • Orbits/ levels are like rungs in step ladder • Cannot stand b/w rungs, e- can’t exist b/w levels (orbits). • E- move from 1 orbit to the next emitting or absorbing certain amounts of energy (quanta). • The smaller the e- orbit, the lower the energy state/level • The larger the e- orbit, the higher the energy state/level n =6 n =5 n =4 n =3 n =2 n =1 nucleus

3. 5-1 Quantum Theory and the atom • Quantum mechanical model is the modern atomic model and comes from • Louis De Broglie: radiation (energy) behaves like particles and vice versa. • All particles w/ a mass have wave characteristics • E- move around nucleus in a wave-like manner • Heisenberg uncertainty principle- impossible to know both the velocity and position of an e- at the same time. • Shrodinger: e-’s energy are limited to certain values (quantum) but does not predict path • Treated e-’s as waves • Created wave function = predicts probability of finding e- in a volume of space (location)

4. Hydrogen’s Atomic Orbitals • Shrodinger’s wave eqn predicts atomic orbitals • Atomic orbital - 3D regions around the nucleus that describes the e-’s probable location. • atomic orbital = fuzzy cloud • Do not have a defined size • Shape = volume that contains 90% of the probable location of e-’s inside that region.

5. Quantum mechanical model • Like Bohr, electrons occupy space surrounding the nucleus and exist in several principal energy levels = principal quantum number (n) • Relative size and energies of atomic orbital • n = 1,2, 3, etc. = period • Principal energy levels consist of energy sublevels with different energy values. • Energy sublevels – shape of the atoms’ orbitals s = spherical p = dumbbell d, f= different shapes

6. Quantum Mechanical model • Principal energy levels have specific allowed sublevels - shapes. • s sublevel is lower in energy and f has higher energy s p d f s p d s p s 4 3 2 n = 1

7. Quantum Mechanical Model • Sublevels consist of orbitals of different orientation. • Orbitals in same sublevel are = in energy (no matter orientation) • Orbitals only hold 2e- maximum with opposite spins (+ or – spins). Sublevel Orientations/ Orbitals Max # e- s 1 2 p 3 6 d 5 10 f 7 14

8. Orientations/ orbitals per sublevel • s- spherical only 1 orbital orientation • p- dumbbell has 3 orbital orientations • d- 2dumbbells with 5 orbital orientations • f- 3dumbbells with 7 orbital orientations • http://winter.group.shef.ac.uk/orbitron/AOs/1s/index.html

9. Bohr model of the atom • Hydrogen’s Line Spectrum (AES) • At n= 1 H atom is in ground state • When energy is added, e- moves to higher energy level, n=2 (excited state). • e- drop back to lower energy level n=1 and emitts a photon equal to the difference b/w levels. A photon is emitted with E= hυ A photon is absorbed

10. Hydrogen’s line spectrum • Lines which show up have specific energies which correspond to a frequency of a color of light. A photon is emitted with E= hυ for each frequency E= 4.85 x 10-19 J E= 3.03 x 10-19 J n 6 5 4 3 2 Energy of Hydrogen Atom 1

11. 5-2 Electron configurations • Electron configuration – arrangement of e- in atoms; lower nrg arrangements • Arrangements defined by: • Aufbau principle – e- occupy lowest nrg orbital available • All orbitals in a sublevel are = in nrg (px py pz ) • Sublevels within an energy level have different energies • Ex: 2s lower in nrg than 2p • Order of energy = s, p, d, f • Sublevels in one energy level can overlap with sublevels in another principal energy level. • Ex: 4s lower in nrg than 3d

12. Aufbau diagram

13. Electron Configurations • Pauli exclusion principle – a max of 2 e- may occupy a single orbital only if they have opposite spins. • Hund’s rule – energy charged e- repel each other. • All same nrg orbitals are filled first with e- containing same spin before extra e- can occupy the same orbital with opposite spins. • Ex: 3 orbitals of 2p 2px 2py 2pz

14. Filling sublevels with electrons • Energy sublevels are filled from lower energy to higher energy following the diagram. • ALWAYS start at the beginning of each level and follow it until all e- in an element have been placed. Increasing Energy 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p

15. Orbital diagram and E- Configurations • Orbital diagram for Fe: • Iron has how many e- ? • 26 e- 1s 2s 2p 3s 3p 4s 3d • Electron configuration for Fe: • Iron has 26 e- • 1s2 2s2 2p6 3s2 3p6 4s2 3d6 • Shortcut to the E- config. for Fe is Noble gas notation • Group 18 or 8A are the Nobel Gases • Argon has 18 e- • Iron has 26 e- • Noble gas notation: [ ] 1s2 2s2 2p6 3s2 3p6 1s2 2s2 2p6 3s2 3p6 4s2 3d6 [Ar] 4s2 3d6

16. Valence electrons and Electron dot structures • Valence electrons – outer energy level/orbital electrons which are involved in bonding. • Valence electrons = groups 1A to 8A • B groups do not count • E- dot structures- consists of the element’s: • Symbol - represents the atomic nucleus & inner-level electrons • Surrounded by dots- represent the valence electrons. • Ex: O = 1s2 2s2 2p4 or [He]2s2 2p4ve- =6 in grp 6A O

17. Periodic table shortcut 1A 8A Periods = Energy Level Groups (A only) = Valence e- 2A 3A 4A 5A 6A 7A Energy level = n-1 for d sublevel Energy level = n-2 for f sublevel

18. 5-3 Light and Quantized Energy Some elements emit visible light when heated with a flame. This chemical behavior is due to the arrangement of e- in atoms.

19. Electromagnetic radiation • Form of energy that exhibits wave-like behavior as it travels through space. • There are many types of electromagnetic radiation and all are represented in the electromagnetic spectrum

20. Electromagnetic spectrum

21. Parts of a wave • Frequency (v, nu) –The number of complete wavelengths that pass a given point each second. • Units: wave/second = 1/s = s-1 = Hertz (Hz) • Wavelength (l, lambda) – The distance between identical points on successive waves. (crest to crest or trough to trough) • Units: meters (m) c = l v c = speed of light, 3.00 x 108 m/s

22. Wave nature of Light • Max Planck theorized that all matter can gain/ lose energy in small “chunks” of light (quanta). • Quantum- minimum amt of energy that can be gained or lost by an atom. • Ex: Iron when hot appears red or blue, emits energy that is quantized has a specific frequency. • Heating water – temp increases by molecules absorbing a specific amt or quanta. • Calculated as follows: Equantum= hv • E = Energy (J) • h = Planck’s constant 6.626 x 10-34 (J s) • v = frequency ( Hz or s-1)

23. Particle Nature of Light • Photoelectric effect – electrons are emitted from a metal’s surface when light of a specific frequency shines on the surface. • Albert Einstein (1905) assumed that light travelled as a stream of tiny particles or packets of energy called photons. • Photons- EM radiation w/ no mass that carries a quantum of energy. • EM radiation has both wave- like and particle- like nature. • Ephoton= hv • Photon = quantum of energy

24. Atomic emission spectra • Set of frequencies of light waves emitted by an atom of an element. • Line spectrum – consists of several individual lines of color from light energy emitted by excited unstable atoms • Only certain colors (frequencies) appear in an element’s AES & it can be used to identify the element.