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UNIT VI: Atomic and Molecular Structure

UNIT VI: Atomic and Molecular Structure. By Jake Grodsky and Sarine Hagopian. Atomic and Electronic Structure/ Quantum Mechanics. Electromagnetic Radiation and Spectrum. Image From: http://www.eoearth.org/files/115601_115700/115629/350px-Spectrum.jpg. Atomic Spectra.

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UNIT VI: Atomic and Molecular Structure

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  1. UNIT VI: Atomic and Molecular Structure By Jake Grodsky and SarineHagopian

  2. Atomic and Electronic Structure/ Quantum Mechanics

  3. Electromagnetic Radiation and Spectrum Image From: http://www.eoearth.org/files/115601_115700/115629/350px-Spectrum.jpg

  4. Atomic Spectra • An electron in the atom absorbs energy from heat, electricity, radiation, etc. • That electron moves to an orbital at a higher energy level • Later, the excited electron returns to a lower energy level • Excess energy lost by electron is released as light or other electromagnetic radiation  Since each element has its orbitals at slightly different energies, each spectrum has a unique finger print.

  5. Atomic Spectra http://www.youtube.com/watch?v=QI50GBUJ48s

  6. Energy Quantization • According to Niels Bohr’s theory: electrons can only exist in certain possible energy levels. • Energy of an electron is proportional to its distance from the nucleus Image From: http://hyperphysics.phy-astr.gsu.edu/hbase/imgmod/bohr1.gif

  7. Bohr Model Image From: http://reich-chemistry.wikispaces.com/file/view/rutherford_bohr_model.gif/103784023/rutherford_bohr_model.gif

  8. Photoelectric Effect • When light is shone on metal, electrons are emitted from the metal • The effect can be used to switch a light signal into an electric current Bright light KEejectede-=Ephoton-Ethreshold current Dim light 0 fthreshold

  9. Wave-Particle Duality • For a long time it was believed that light was solely a wave • Both light and electrons have a dual nature • They exhibit characteristics of both waves and particles • The photoelectric effect proves that light has a particle nature as well • The wave properties of electrons are shown through the DeBroglie Hypothesis

  10. DeBroglie Wavelength Constant: h = 6.626 × 10-34 J•s λ = velocity wavelength of a particle Constant: me = 9.11 × 10-31 kg

  11. Wave-Mechanical Model • The wave mechanical model is the most recent model of the atom • Improvements were made on Bohr’s model, specifically dealing with electrons • Electrons are treated as waves instead of particles- electron has more in common with light, tv, radio waves, microwaves, and x-rays than it does with protons and neutrons • Orbitals are the regions in atoms which are most likely to have electrons in them • The model is more statistical than visual • This model includes energy levels which are numbered 1-7(closest to farthest) which indicates how far a given electron is from the nucleus • The energy level can be viewed in the same way as Bohr’s model viewed the shell

  12. Electron Configurations of Atoms and Ions • Lower energy levels are always filled first • Ions have less electrons than the neutral parent atom • this means that electron configurations of ions look like those of other neutral elements (even from a different atom) • Image From: http://www.mikeblaber.org/oldwine/chm1045/notes/Struct/EPeriod/IMG00011.GIF

  13. Example: Electron Configuration of Zinc: 1s22s22p3s23p64s2 3d10 Abbreviated Electron Configuration of Zinc: [Ar] 4s2 3d10 • Configuration can be abbreviated • The last Noble Gas element symbol is put in brackets and remainder of the electron configuration is written out Image From: http://www.mpcfaculty.net/mark_bishop/abbreviated_electron_configuration_help.htm

  14. Quantum Numbers Key: n= principal quantum number l= angular momentum number [0- (n-1] ml= magnetic quantum number [-l – l] • The final quantum number is the ms number. This is ± ½, depending on the spin of the electron Image From: http://library.thinkquest.org/19662/images/eng/pages/improved-bohr-2.jpg

  15. Orbital Diagrams • Electron configurations can be expressed as orbital diagrams as pictured below by visualizing each individual electron and its corresponding spin, as well as the orbital and energy level that it is a part of. • To form an orbital diagram: • Determine the electron configuration of the atom and the total amount of electrons. Following Hund’s Rule, begin to fill orbitals from lowest energy level to highest, remembering the Pauli Exclusion Principal and having an upward and downward arrow in each orbital, representing the positive and negative electron spin (responsible for the +½ and -½ values of ms. Oxygen 2s 2px 2py 2pz 1s Electron Configuration: 1s22s22p4 What would the quantum number be for thise-? Total number of e-: 8 2, 1, -1, -½

  16. PeriodicTrends

  17. Atomic Size Atomic Radius: Size of an atom which is influenced by the volume of the e- orbitals (clouds) Decreases • Why does atomic radius increase as you go down a group? • more energy levels so the new levels are “blocked” and therefore not as tightly pulled to the center Increases • Why does atomic radius decrease as you go across a period? • Only one p+ and one e- are added • Increasing nuclear charge pulls outermost e-s closer and closer to the nucleus  reduces atom size • All additional e- go into same principle energy level so shielding is not an issue  nucleus just gets stronger and squeezes everything closer

  18. Ionic Size • Cations are smaller than their neutral parent atoms • Cations have more protons than electrons (hence their positive charge) • Protons more tightly pull the electrons towards the nucleus therefore reducing the size of the atom • Anions are larger than their neutral parent atoms • Anions have more electrons than neutrons (hence their negative charge) • Electrons are not as attracted to the nucleus therefore increasing atomic radius

  19. Ionization Energy and Electronegativity Ionization Energy: amount of energy needed to remove an e- from an atom or ion Electronegativity: a measure of the ability of an atom in a chemical compound to attract/gain e-s Increase Decrease • As you go down a group, more orbitals are added  valence e- are farther from nucleus so pull of p+s on the e-s is reduced • As you go across a period, more protons are added to the nucleus  valence electrons are held more tightly

  20. Molecular Structure

  21. Covalent Bonding • Formed when e- pairs are shared amongst atoms • Generally a metal and a nonmetal pair • Key Terms • Lewis structure: representation of a covalently bonded molecule and its valence electrons • Octet Rule: In a covalent molecule, each atom has eight electrons around it • Lone Pair: pair of e-s not involved in bonding • Bond pair: pair of e-s shared between two atoms • Double bond: two pairs of e-s shared between atoms • Triple bond: three pairs of e-s shared between atoms } • As the number of shared e- pairs goes up, bond length goes down

  22. Lewis Structures • Symmetrical arrangements are more likely than asymmetrical ones • The less electronegative atom tends to be in the middle • Subtract valence e-s from total electrons needed to complete octet/duet and divide by two  this is the number of bonds that will need to be made • If e- needed > e- remaining, add bonds • If e- remaining > e- needed, add lone pairs to central atom Image from: https://vinstan.wikispaces.com/file/view/lewis_structure.gif/46694367/lewis_structure.gif

  23. Polarity Are any bonds polar? No Yes Are polar bonds arranged symmetrically? Non-Polar Molecule Yes No Polar Molecule • Electronegativity difference of: • 0.5 or less bond is nonpolar • greater than .5 bond is polar • greater than 1.7bond is ionic

  24. Hybridization, Shape, and Angle Trigonal Pyramidal Image from: http://www.chem.hbnu.edu.cn/jysweb/whjys/wangwd/jghxywkj/AX3E1.gif Seesaw Image from: http://www.chem.hbnu.edu.cn/jysweb/whjys/wangwd/jghxywkj/AX4E1.gif Square Planar Image from: http://www.chem.hbnu.edu.cn/jysweb/whjys/wangwd/jghxywkj/AX4E2.gif

  25. Formal Charge • Helps decide which Lewis structure is most reasonable • It is the charge the atom would have if all the atoms in the molecule had the same electronegativity • To calculate: 1. Count all nonbonding electrons per atom. 2. Count half of any bond. 3. Subtract valence electrons by the number assigned to each atom.

  26. Bonding Theory

  27. Valence Bond Theory (Hybrid Orbitals) Valence bond theory says that electrons in a covalent bond can be found in a section that is the overlap of the individual atomic orbitals that are bonding Image From: http://courses.chem.psu.edu/chem210/quantum/pictures/sigms.gif

  28. Molecular Orbital Theory

  29. Sigma and Pi Bonds Sigma bonds = formed by the overlap of two s orbitals, an s and a p orbital, or two p orbitals Pi bonds = formed by the overlap between two p orbitals oriented perpendicularly to the internuclear axis Ex) O=O has 1 sigma bond and 1 pi bond N≡N has 1 sigma bond and 2 pi bonds • Where do Pi bonds come from? • Only period 2 elements form pi bonds because they are small in size and therefore form short bonds. • Due to these short bonds when 2 of these atoms form a sigma bond, they are so close together, that an additional energy level(orbital) overlap and a pi bond is formed

  30. Bond Order ½ [ (#bonding e-) - (#anti-bonding e-)]

  31. Equations and Constants • E = hf • KE = ½mv2 • c = fλ • λ = • c = 2.9979 × 108 m/s • h = 6.626 × 10-34 J•s • me = 9.11 × 10-31 kg • NA = 6.022 × 1023

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