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Principles of Chemistry 1 Test 3 Review

Principles of Chemistry 1 Test 3 Review. By Bo Marshall. The Fine Print….

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Principles of Chemistry 1 Test 3 Review

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  1. Principles of Chemistry 1 Test 3 Review By Bo Marshall

  2. The Fine Print… • This is NOT intended to be your only source of information for the test - this is simply an outline for me to follow to help you review…As usual, I recommend you read your text, attend and pay attention in class, re-do the homework, do suggested homework problems, and take the practice test. • Also…It is a pain to put all the solutions to the problems into power point, so we will do practice problems after we review the concepts. -Bo-

  3. Gases => Lots of available space for motion, therefore solids and liquids are referred to as “condensed states.” Pressure = force/area, therefore smaller area gives larger pressure Use atm =>know how to convert from kPa, Pa, psi, mmHg, torr (on eqn sheet)

  4. Gas Law Dudes • Boyle: P1V1=P2V2 : So P and V are inversely proportional, if one is large, the other is small • Charles: (V1/T1)=(V2/T2) : V and T directly proportional, so double one gives double the other • Amonton: (P1/T1)=(P2/T2) • Avagadro: (V1/n1)=(V2/n2) • So be familiar with these, but the ones you really need to know how to use are…

  5. Ideal gas law: PV=nRT • One sample of gas- use to find the missing variable - may have to find something else from this (ex. Find moles, use it to find mass) • Combined Gas Law: (P1V1/T1)=(P2V2/T2) • Use when changing pressure, temp or moles • Hot air balloon question in class (do we want to do that one again?)

  6. Gas density from ideal gas law… • D = PM/RT => Use to compare densities of different gases, or find P or M when given density

  7. So, at 40ºC and 1 atm, 43.8 g of a gas occupies 1L. What is the molar mass of the gas?

  8. STP • Standard Temp and Pressure • Don’t confuse with • thermo “standard states” in thermo chem • Stone Temple Pilots • Fuel Additive • So what are STP for gas laws?

  9. STP • P = 1atm • T = 273.15 K • So when a problem says to “use standard temp and pressure” you know what to do… • Gas laws can also be used for stoichiometry…example done in class and we’ll get to another one later.

  10. Effusion / Diffusion • Effusion - Escape of gas through a tiny hole into an evacuated space • Diffusion - spread of one substance throughout space or throughout a 2nd substance • Eqn for rate of effusion • Note the A’s and B’s flipped

  11. Effusion • 3.0 L of helium effused through a membrane in 24 hrs. How many hours would it take a 3.0L sample of oxygen to effuse through the membrane? • Molar Mass He = 4.003 g/mol • Molar Mass Oxygen = 16, but exists as O2, so need to use 32 g/mol

  12. Effusion • So He is about 2.8 times as fast as Oxygen. If He took 24 hrs… • 24hrs(2.8) = 67.2 hrs for Oxygen to effuse • Another way to look at it if anyone is confused…let me know

  13. Partial Pressures • Sum of the partial pressures of all the gases involved gives the total pressure • Ptotal=P1+P2+…Pn (Dalton’s law of partial pressures)

  14. Kinetic Molecular Theory • 1. Particle volume very small compared with container - so small they have no volume • 2. Particle motion is constant and random in a straight line until they collide with another particle or wall • 3. At a given temp - all gases have the same average kinetic energy (more to it, this was the main summary…) • Showed graph of gas speed and number of particles at that speed

  15. Lighter gases faster, so He has more particles at the faster speed…Book actually has a good explanation of this…

  16. Real Gases • Particles DO have volume • So need to adjust volume downward when using ideal gas law b/c gases can occupy volume • They DO attract each other (so motion is not totally random) • Must adjust the pressure upwards • Van der Waals eqn accounts for this

  17. Van der Waals eqn • a and b are constants that vary for each gas (see table 10.3) • So a gas molecule is about to hit the wall - this molecule is attracted by neighboring molecules, so the impact of the collision is less, so the pressure is less

  18. Electromagnetic Radiation • Wavelength = = distance / wave cycle • Frequency =  =number of waves per unit time • Main idea:   = c = 3.00 x 10^8 m/s • So given , you can solve for  and vice versa…

  19. Should probably know the  and  of the visible spectrum… •  measured in Hz (s-1 or 1/s) => can prefix with any of the metric stuff (KHz = 1000Hz, etc) • When doing conversions, don’t forget to convert  to meters and  in Hz

  20. Warning • I am NOT an expert in the wave/particle duality of light. In the following slides, I will try to explain it in the most basic way that I can, but if you are super interested in this, I suggest reading your Chemistry book or asking a Chemistry prof. • I CAN work the problems regarding this material and I WILL give you as many examples as I can of them…

  21. The wavelength stuff is cool, but some aspects of light can’t be explained using it… • Blackbody Radiation: emission of light from hot objects • Planck showed that atoms emit energy in packets - a packet is called a quantum • Eqn for smallest possible energy change, i.e., the smallest amount of energy an object can absorb or emit: ∆E=h  • Notice how  is used in the eqn. What if they have you the  instead… • h = 6.626 x 10^-34 J*s => Really small number, but atoms are really small too…

  22. Photoelectric effect: emission of electrons from metal surfaces on which light shines • Some aspect of light dependent on the  => a minimum frequency was needed to cause the metal to emit e-’s => Didn’t need to be on there for extended length of time => the e-’s were emitted immediately • Some light never caused emission of any e-’s • Radiant energy traveled as photons, one photon’s energy: E=h 

  23. Line spectra: Emission of a certain color (wavelength) from electronically excited atoms • Suggested that e-’s can only be present in the atom in certain energy levels and that the light emitted corresponds to the energy levels • Bohr model: e-’s orbit nucleus, energy is not emitted as long as the e-’s stay in the same energy level - When an e- moves from one level to another, a photon is either absorbed or emitted

  24. Bohr model cont. • Remember that zero energy is when e- is “infinitely removed” from the nucleus and gets more negative as it gets closer to the nucleus • The closest energy level to the nucleus will vary depending on the nucleus • The lowest energy state is the ground state. When an electron is at a higher energy level, it is said to be in an excited state. • Some eqns on pg.4 of the notes, we’ll do an example with them later…

  25. To summarize… • When an e-’ move from high to low energy, the atom emits a photon of light and that photon has a  and  that corresponds to the energy emitted • When an e- moves from lower to higher energy level, it absorbs a photon • Colors of the lines in line spectra correspond to energy emitted, everywhere else black • Absorption spectra would have black where these colored lines are and colors everywhere else, b/c energy being absorbed not emitted

  26. So light has wave and particle properties • de Broglie thought that matter may have wave properties then too…it does, but the wavelength is very compared to macroscopic level • Ex in class with baseball vs. electron •  = h/mu

  27. Heisernberg Uncertainty Principle • Not applicable to macroscopic objects • Bohr model - great for 1 e- systems, but need another model to predict energy of systems w/ >1 e- • Quantum Mechanical Model (QM) • Wave functions used to predict motion of e- • The e-’s are located in orbitals! (Diff from orbits of the Bohr Model)

  28. Orbitals - area where electrons are most likely found • 3 quantum #’s needed to describe an orbital • Principle Quantum #, n, defines size of orbital • Angular momentum quantum #, l, defines shape of orbital • Can be a number from 0 to n-1 • If l=0, that is an s orbital • 1 = p orbital, 2 = d orbital, 3 = f orbital • “SO, pi, deuce, free” • Magnetic Quantum #, ml, defines the orientation of the orbital • Can go from -l….0…+l

  29. Spin quantum #, ms, tells the spin of the e- • +1/2, or -1/2 • This quantum # deals specifically with the e- • Need all 4 to describe an e- • Pauli exclusion principle (joke to remember) • Should probably be familiar with the shapes of the orbitals too…(anyone look at the practice test yet?)

  30. Electrostatic effects • Things that effect the energy of an e- • 1) Z effect: more + charge in Nucleus makes e- harder to remove - therefore energy needed to remove is more, so e- is more stable • 2) Electron Shielding (Same orbital): if you add electrons to the same orbital, the nucleus cannot attract both of them as strong as it can just one - so the e- is easier to remove • 3) Electron Shielding (diff orbitals): electrons in inner orbitals “take away” from the nuclear attraction on an outer electron

  31. Order of Sublevel energies • For a given n value, the lower the l value means a lower sublevel energy • s < p < d < f • Orbital filling - lets go to the periodic table, but first..

  32. Remember exceptions to orbital filling • Valence Electrons: Electrons in the highest energy levels (largest n value) • Inner (Core) electrons: those in prev. noble gas config AND those in filled sets of d and f orbitals • Also - don’t forget about Hund’s Rule • Lets look at this sweet periodic table again to practice this…http://www.dayah.com/periodic/

  33. Practice • What is the correct set of quantum #’s for an e- in a 3d orbital • A) n=3, l=0, ml= -1 • B) n=3, l=1, ml= +2 • C) n=3, l=2, ml= -2

  34. Practice • We know n = 3 • So l can be 0, 1 or 2 • The d tells us that l= 2 • If l = 2, then ml can be: -2, -1, 0, 1, 2 • So C is the only possible answer

  35. More Practice • See supplemental practice problems sheet…

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