Ch 5 electrons in atoms notes


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Ch 5 electrons in atoms notes

  1. 1. CH 5 ELECTRONS IN ATOMS 5-1 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.
  2. 2. 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
  4. 4. 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 (λ, lambda) – The distance between identical points on successive waves. (crest to crest or trough to trough)  Units: meters (m) c = λ v c = speed of light, 3.00 x 108 m/s
  5. 5. 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. o Ex: Iron when hot appears red or blue, emits nrg that is quantized has a specific frequency. o Heating water – temp increases by molecules absorbing a specific amt or quanta.  Calculated as follows: Equantum= hv o E = Energy (J) o h = Planck’s constant 6.626 x 10-34 (J s) o v = frequency ( Hz or s-1 )
  6. 6. 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
  7. 7. 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.
  8. 8. 5-2 QUANTUM THEORY OF THE ATOM  Bohr Model of the Atom  Used to explain why AES was set of discontinuous lines of specific frequencies (color).  Proposed that Hydrogen atoms have only certain allowable energy states based on Planck’s and Einstein’s quantized energy.  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 nrg level, closest to nucleus.
  9. 9. 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 amts of nrg (quanta).  The smaller the e- orbit, the lower the energy state/level  The larger the e- orbit, the higher the energy state/level n =1 n =2 n =3 n =4 n =5 n =6 nucleus
  10. 10. BOHR MODEL OF THE ATOM  Hydrogen’s Line Spectrum (AES)  At n= 1 H atom is in ground state  When nrg 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 absorbed A photon is emitted with E= hυ
  11. 11. HYDROGEN’S LINE SPECTRUM  Lines which show up have specific energies which correspond to a frequency of a color of light. EnergyofHydrogenAtom 1 2 3 4 5 6 n A photon is emitted with E= hυ for each frequency E= 4.85 x 10-19 J E= 3.03 x 10-19 J
  12. 12. 5-2 QUANTUM THEORY AND THE ATOM  Quantum mechanical model is the modern atomic model and comes from A. Louis De Broglie: radiation (energy) behaves like particles and vice versa. 1. All particles w/ a mass have wave characteristics 2. E- move around nucleus in a wave-like manner B. Heisenberg uncertainty principle- impossible to know both the velocity and position of an e- at the same time. C. Shrodinger: e-’s energy are limited to certain values (quantum) but does not predict path 1. Treated e-’s as waves 2. Created wave function = predicts probability of finding e- in a volume of space (location) Moving Electron Photon Before Electron velocity changes Photon wavelength changes After
  13. 13.  Shrodinger’s wave eqn predicts atomic orbitals  Atomic orbital - 3D regions around the nucleus that describes the e-’s probable location. a. atomic orbital = fuzzy cloud b. Do not have a defined size c. Shape = volume that contains 90% of the probable location of e-’s inside that region. HYDROGEN’S ATOMIC ORBITALS
  14. 14. QUANTUM MECHANICAL MODEL  Like Bohr, electrons occupy space surrounding 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 nrg levels consist of energy sublevels with different nrg values.  Energy sublevels – shape of the atoms’ orbitals s = spherical p = dumbbell d, f= different shapes
  15. 15. QUANTUM MECHANICAL MODEL  Principal energy levels have specific allowed sublevels - shapes.  s sublevel is lower in energy and f has higher energy 1 2 3 4 n =s s p s p d s p d f
  16. 16. 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
  17. 17. 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 
  18. 18. 5-3 ELECTRON CONFIGURATIONS  Electron configuration – arrangement of e- in atoms; lower nrg arrangements  Arrangements defined by: 1. Aufbau principle – e- occupy lowest nrg orbital available a. All orbitals in a sublevel are = in nrg (px py pz ) b. Sublevels within an energy level have different energies  Ex: 2s lower in nrg than 2p a. Order of energy = s, p, d, f b. Sublevels in one energy level can overlap with sublevels in another principal energy level. a. Ex: 4s lower in nrg than 3d
  20. 20. ELECTRON CONFIGURATIONS 2. Pauli exclusion principle – a max of 2 e- may occupy a single orbital only if they have opposite spins. 3. 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
  21. 21. 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. 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p
  22. 22.  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 ORBITAL DIAGRAM AND E- CONFIGURATIONS ][1s2 2s2 2p6 3s2 3p6 4s2 3d6 [Ar] 4s2 3d6
  23. 23. 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: a. Symbol - represents the atomic nucleus & inner-level electrons b. Surrounded by dots- represent the valence electrons. c. Ex: O = 1s2 2s2 2p4 or [He]2s2 2p4 ve- =6 in grp 6A O
  24. 24. PERIODIC TABLE SHORTCUT Periods=EnergyLevel Groups (A only) = Valence e- 1A 2A 3A 4A 5A 6A 7A 8A Energy level = n-1 for d sublevel Energy level = n-2 for f sublevel