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- 1. Topic 3 Electron Configurations
- 2. Line Emission for Hydrogen <ul><li>Bohr said that electrons could only travel in fixed orbits around the nucleus…..only works for Hydrogen </li></ul><ul><ul><li>Gaseous elements emitted electromagnetic radiation when heated </li></ul></ul><ul><ul><li>Quanta – small packets of light </li></ul></ul><ul><ul><li>Ground State (release) Excited State (absorbed) </li></ul></ul><ul><ul><li>Line spectrum – distinct bands (gaseous elements) </li></ul></ul><ul><ul><li>Continuous spectrum – one continuous band (fluorescent and white light) </li></ul></ul><ul><ul><li>Wavelength of line created by the electron returning to ground state is also fixed </li></ul></ul><ul><ul><li>Gives a specific pattern for each element </li></ul></ul>
- 3. What Bohr Proposed <ul><li>1. The electron on the hydrogen atom can exist only in certain spherical orbits. </li></ul><ul><li>2. As the distance from the nucleus increases, the energy of an electron in that orbit increases. </li></ul><ul><li>3. The closest orbit (energy level) is called the ground state. Higher energy levels are called excited states. </li></ul><ul><li>4. When an electron falls from a higher energy level to a lower energy level, it emits a definite amount of energy that is equal to the difference in the energy of the two levels. </li></ul>
- 4. Bohr’s Model <ul><li> E photon = </li></ul><ul><li>energy of level n final -energy of level n initial </li></ul>
- 5. <ul><li>Scientists figured that since only specific frequencies of light were emitted then the energy differences between the atoms’ energy were fixed. </li></ul><ul><li>This is what lead Bohr to believe that a hydrogen atom exists only in very specific energy states </li></ul>
- 6. <ul><li>These are additional lines that were discovered in the ultraviolet and infrared regions of hydrogen’s line spectrum </li></ul>
- 7. Wave Nature of electrons <ul><li>Wavelength – distance between 2 repeating points on a wave ( λ ) •units are nm or m (m must be used in calculations) </li></ul><ul><li>Frequency - # of waves that pass in a second ( ν ) • units are s -1 or Hz </li></ul><ul><li>Electromagnetic radiation are all waves that can be defined with: </li></ul><ul><ul><li>E = h ν (h = Planck’s constant = 6.626 x 10 -34 J•s and ν = frequency) </li></ul></ul><ul><ul><li>C = λν (c = speed of light = 3.00 x 10 8 m s -1 ) </li></ul></ul>
- 8. <ul><li>E = hc/ λ (E has the unit of J) </li></ul><ul><li>mc 2 = hc/ λ (m = mass in kg) </li></ul>High frequency Long wavelength Any radiation will fall somewhere in this electromagnetic spectrum Cosmic X rays ultraviolet visible infrared microwaves Radio waves
- 9. <ul><li>DeBroglie suggested that the electron had some wavelike characteristics </li></ul><ul><li>Bohr suggested that the electron was a discrete particle </li></ul><ul><li>Schrodinger developed the idea that the electron had wave equations that involved the electron with a particle nature • dual wave-particle nature </li></ul><ul><ul><li>Heisenberg Uncertainty Principle: impossible to determine simultaneously both the position and velocity of an electron or any other particle </li></ul></ul><ul><ul><li>Orbitals give the probability of finding an electron at a given place around the nucleus </li></ul></ul><ul><ul><li>Quantum Theory : describes mathematically the wave properties of electrons and other very small particles </li></ul></ul>
- 10. Quantum Numbers <ul><li>Principle Quantum # (n) </li></ul><ul><ul><li>1 st shell has quantum # of 1, 2 nd shell has 2 </li></ul></ul><ul><ul><li>Maximum # of electrons is given by 2 n (where n = the energy level) </li></ul></ul><ul><li>Angular Momentum Quantum # (ℓ) aka azimuthal quantum # </li></ul><ul><ul><li>Subshells: s, p, d, f correspond to 0, 1, 2, 3 (0 to n-1) </li></ul></ul><ul><ul><li>Letters refer to 3-D shape </li></ul></ul><ul><ul><ul><li>s orbital = spherical shaped </li></ul></ul></ul><ul><ul><ul><li>p orbital = dumbbell shaped and align on x, y, z axes </li></ul></ul></ul><ul><ul><ul><li>d orbitals are 4-leaf clover shaped </li></ul></ul></ul><ul><ul><ul><li>f orbitals have more complicated shapes </li></ul></ul></ul>
- 11. <ul><li>Magnetic Quantum # (m ℓ ) </li></ul><ul><ul><li>refers to orientations </li></ul></ul><ul><ul><li>Each sub shell is in orbitals </li></ul></ul><ul><ul><li>The # of orbitals that are possible is = to twice the azimuthal quantum # plus 1 (2ℓ + 1) </li></ul></ul><ul><ul><li>Possible values are -ℓ to +ℓ including 0 </li></ul></ul><ul><li>Spin Quantum # (m s ) </li></ul><ul><ul><li>Each orbital holds a max of 2 electrons </li></ul></ul><ul><ul><li>Pauli exclusion principle says that no one electron can have the same set of quantum numbers, so since each orbital can hold a max of 2 electrons, they must be distinguished </li></ul></ul><ul><ul><li>Can have values of + ½ or – ½ </li></ul></ul>
- 12. <ul><li>Choice of Quantum Numbers </li></ul><ul><ul><li>When there is a choice of magnetic quantum numbers, the lowest values are chosen 1 st and + ½ is chosen before – ½ </li></ul></ul>
- 13. Rules for Filling Orbitals <ul><li>Aufbau Principle - an e - occupies the lowest orbital that can receive it. </li></ul><ul><ul><li>Find how many electrons are present </li></ul></ul><ul><ul><li>Lowest energy orbitals are filled 1 st – 1s, 2s, etc… </li></ul></ul><ul><ul><li>4s has lower energy than 3d orbitals, as is the 5s and the 4d </li></ul></ul><ul><li>Hund’s Rule - orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron, and they must have parallel (same) spins </li></ul><ul><ul><li>All 2p orbitals have same energy (all 3d, all 4f) </li></ul></ul>
- 14. Electron Configurations <ul><li>Determining Electron Configuration </li></ul><ul><ul><li>Period # shows shell </li></ul></ul><ul><ul><li>Block shows type of orbital </li></ul></ul><ul><ul><li>Add 1 electron until the orbital is full </li></ul></ul><ul><li>Anomolies to electron configurations </li></ul><ul><ul><li>Cr and Cu have configurations of 4s 1 3d 5 and 4s 1 3d 10 because a half or completely filled d shell is considered to have extra stability </li></ul></ul><ul><ul><li>The s and d orbitals are very close in energy so it is fairly easy for an electron to shift between them </li></ul></ul>
- 15. <ul><li>Noble Gas Configuration </li></ul><ul><ul><li>Phosphorus becomes [Ne]3s 2 3p 3 </li></ul></ul><ul><ul><ul><li>Write previous noble gas in square brackets and then fill in orbitals as before </li></ul></ul></ul><ul><ul><ul><li>You will always start with the s that has the row number of the element you are working with </li></ul></ul></ul><ul><li>Orbital diagrams </li></ul><ul><ul><li>Arrows represent the electron (and its spin) and boxes/lines represent orbitals </li></ul></ul>
- 16. <ul><li>Paramagnetic and diamagnetic </li></ul><ul><ul><li>Paramagnetic species are those that are attracted by a magnet (created by unpaired electrons present in the atom) </li></ul></ul><ul><ul><li>Diamagnetic species are slightly repelled by magnets and occur when all electrons are paired </li></ul></ul>
- 17. Rydeberg Equation <ul><li>Used to calculate the E changes when electrons are promoted to higher energy levels and subsequently fall back to the lower energy levels </li></ul><ul><li>E = the energy associated with a particular quantum # </li></ul><ul><ul><li>E = -2.178 x 10 -18 / n 2 n is the diff. between 2 levels </li></ul></ul><ul><li>Can also calculate energy released by: </li></ul><ul><ul><li>E = (-2.178 x 10 -18 / n 2 ) – (-2.178 x 10 -18 / n 2 ) </li></ul></ul><ul><ul><ul><li>Where the first n 2 is the higher energy level and the second n 2 is the lower energy level </li></ul></ul></ul><ul><ul><ul><li>By calculating the energies for 2 quantum levels and finding the difference, one can calculate the E required to promote an electron from one to another </li></ul></ul></ul><ul><ul><ul><li>Released energy = + value </li></ul></ul></ul>
- 18. <ul><li>Energy changes during transitions are proportional to the (atomic #) 2 </li></ul><ul><li>This means that if an electron is promoted from level 1 to level 5 in a species that has less p + in the nucleus, then the same transition for a species with more protons would be more difficult </li></ul><ul><li>This is because the protons in the nucleus are attracting the electron to the lower E level and more E is required to promote them </li></ul><ul><li>Consequently, a greater amount of E is released from the one with the larger amount of protons </li></ul>
- 19. <ul><li>D block metal ions </li></ul><ul><ul><li>When forming metal ions, d block elements lose their outer s electrons before any d electrons </li></ul></ul><ul><li>Isoelectronic – have the same electronic configuration, as a result they must be distinguished by some other means, for example the # of protons present. </li></ul>

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