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Chapter 2

Chapter 2 Atomic Theory.pdf

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University of Florida
CHM 2045

Chapter 2 Atomic Theory Tuesday, July 02, 2013 12:07 AM 1. Subatomic Particles a. Protons,neutrons and electrons b. c. An atom usually carries a neutral charge unless it has loss or gained an electron 2. Average Atomic Mass a. 3. ThomsonExperiment a. Demonstratedthe existence of opposite charges in an atom and that charge is a fixed quantity b. Thomsonobserved when he applied an electric field perpendicular to the electron beam, he could deflect tit by the exact amount each time i. The magnitude of deflection depend on the strength of the field and the mass of the electron c. d. Because the arc fo the deflection was constant, Thomsonconcluded that electrons have a fixed charge to mass ratio e. Because the direction of the deflection changed when the orientationof the field changed, Thomsonconcluded that there must be two types of charge that oppose one another 4. Mass spectrometry a. Produces a spectra of the masses of the moleculescomprising a sample of material i. Measure the charge to mass ratio of a charged particle b. A charged particle is sent into a perpendicular magnetic field and observing the degree to which it curves i. The degree of arching for a particle can vary with mas, vi, magnitude of charge and the magnetic field ii. As momentum(mv) increases, the particle deflects less, so the radius of curvature increases (bigger circle) iii. As the charge magnitude increases, the force that interacts with the magnetic field increases, causing the radius of curvature to decrease and the particle to deflect more c. By comparing the curvature of an atomic molecularion to a known standard, the mass of the unknown ion can be determined i. d. Can be used to determine isotopic abundance i. The MS sends high energy electronsto collide with the molecule, knocking away one of the molecular electrons of the molecular electrons ii. This leaves behind a molecular ion, a radical cation iii. The molecular ion may fragment further into neutral pieces iv. The mass spectrum will display as a vertical bar graph, in which each bar represents an ion having a specific mass-to-charge ratio (m/z) and the length of the bar represents the abundance 1) Most of the ions form in mass spec have a single charge, so the m/z value is equal to the mass itself v. vi. The molecular ion peak is the peak representing the molecular ion, excluding any heavier isotopes 1) In the case of Bromine, the molecular ion peak is 160 e. 5. Millikan Oil Drop Experiment a. Electrons are added to oil droplets b. Since oil droplets are not charged this is hard i. Falling oil drops pass through a beam of electronswhere some oil drops are penetrated at random by electrons ii. Eventually, an oil drop retains the charge and falls through a hole in a capacitor plate c. There is another plate of the capacitor below the one with the hole d. If the falling droplet is suspended in air, than the force applied must equal the gravitational force e. f. Experiment allows Millikan to solve for the charge of an electron 6. The three previous experiments demonstratedWHAT subatomic particles are 7. The rutherford experiment a. Demonstratedwhere subatomic particles are i. Dense nuclei b. An incident beam is focused and aimed at a thin slice of gold metal i. Gold is used because it has a large nucleus and its atoms pack in such a manner where light can pass through the lattice more easily than other metals c. If α particles are used, a luminsecent screen is placed around the gold foil to detec where the particle pass through the foil and strike the screen, which glow when struck the particle pass through the foil and strike the screen, which glow when struck i. Some particles are deflected by the gold atoms, resulting in parts of the luminescing screen never illuminating d. Disproved the diffuse particle model AKA the plum pudding model e. 8. The location of the electrons IS NOT DETERMINED BY RUTHERFORD Experiment!! 9. Heisenberg Uncertainty Principle a. It's not possible to determine the exact location and the exact speed of an electron at the same time b. 10. Atomic Model (Bohr Model) a. Electrons occupy specific circular orbits about the nucleus i. Thus, electrons have specific energy levels and can only exist in specified orbits (shells) ii. b. According to classical physics, a charged particle moving in a circular path should continuously lose energy and then spiral into the positively charged nucleus i. Bohr assumed the prevailing laws of physics were inadequate to describe all aspects of the atom 1) "an electronin a permitted orbit is in an "allowed" energy state, which does not radiate energy and thereforedoes not spiral into the nucleus c. Energy levels are spaced according to the energetics of transition between the levels i. More energy is required to carry out transitions when the electron is nearest the nucleus d. When energy is absorbed by an electron,it can be elevated to a higher energy level, only if the exact amount of energy is absorbed e. The more negative the energy is, the more stable the atom is f. Energy is proportional to Z /n , Z = nuclear charge, n = electronicenergy level g. Light energy is often given off when an electron drops back down the lower level (orbit) i. The smaller the gap between energy levels, the less energy that is give off, the longer the wavelength of light that is emitted the wavelength of light that is emitted h. 11. Hydrogen Energy Levels a. For hydrogen, Energy E= b. The transition from one energy level to the next is therefor i. By calculating the difference in energy, we can calculate the wavelength of light emitted ii. 12. Electronic Structure a. Electrons orbit in distinct shells. b. Not all shells can hold the same number of electrons! c. Shells farther from the nucleus have a greater radius, and thus a greater capacity to hold electrons 2 d. Number of electrons in a shell= 2(n) e. Electrons are held in their orbit by an attractive force to the nucleus i. They are also repelled by other electrons ii. The net charge responsible for holding the valence electrons in place is the effective nuclear charge 1) Accounts for attraction to the nucleus, repulsion from core electrons, and minimal repulsion by other valance electrons 2) To calculate, the nuclear charge (number of protons) is added to the core electron charge (a negative number) 3) f. Z effreases while moving left to right while moving across periodic table i. The effect of the extra valence electronis not as significant as the effect of the additional proton g. Each electron travels a unique pathway i. They are hard to track, so we study the magnetic field that they produce h. Electron spin pairing i. Electrons fill orbitals in a specific sequence, filling evenly into orbitals with equal energy with a like spin (up) before following with the down electrons ii. Electrons can either spin clockwise or counterclockwiseabout their axis 1) 2) These spinning electrons generate magnetic momentsin opposite directions 3) A spin up produces a magnetic field pointing north while a spin down produces a magnetic field pointing south i. The shell represents the energy levels an electron can occupy i. The orbitals, which make up the shells, represent the region in which the electron is likely to be found j. Spin Pairing and magnetism i. Paramagneticspecies is an atom or molecule that contains at least one unpaired electron (radicals!) 1) Because the electron is unpaired, it is susceptible to magnetic fields 1) Because the electron is unpaired, it is susceptible to magnetic fields 2) If an external magnetic field is applied to a paramagnetic species, the electron spins align with the field 3) This induces a magnetic momentinto the compound, thus making it magnetic (i.e. paramagnetic species can have magnetism induced into them) ii. Diamagnetic species is an atom or moleculethat contains no unpaired electrons 1) All electrons are spin paired 2) Because they are all spin paired, diamagnetic compounds are not susceptible to magnetic fields a) If a magnetic field is applied to the species, half of the electron spins align with the field and the other half align against the field k. Electron density and orbitals i. Atomic orbitals are 3d pictorial representations of the region where an electron is likely to be found NOT THE PATH AN ELECTRON TAKES! 1) We look at where the electron usually is and draw a probability map 2) ii. The collection of orbitals with the same value of n is called the electron shell 1) All the orbitals with n=3 are said to be in the third shell 2) iii. S orbitals result from the spherical distribution of the electronsabout the nucleus 1) 2) For each value of n, there is only 1 s orbital! iv. P orbitals result from barbell-like distribution of the electrons about the nucleus 1) Density is concentratedin two regions on either side of the nucleus, separated by a node at the nucleus 2) Each shell has three p orbitals 3) v. D orbitals result from double barbell like distribution of the electrons about the v. D orbitals result from double barbell like distribution of the electrons about the nucleus 1) Have two nodal planes 2) l. Pauli's exclusion principle i. No two electronscan have the same set of quantum numbers m. Hund's rule i. Electrons completely fill lower energy levels before starting to fill higher energy levels ii. "degenerate", have the same energy n. Aufbau principle i. Electrons are added one by one, starting with the lowest energy level ii. Commonexceptions! 1) Electrons will jump up to the higher energy level in order to create a more even D orbital 2) o. The 4s level is filled before the 3d level i. This is because the energy level is based on an average position of the electron, not the extremeposition 1) Since 4s have a closer average to the nucleus, it has a lower energy p. OUTER SHELL ELECTRONS ARE ALWAYS REMOVED FIRST WHEN FORMING CATIONS i. AN ELECTRON IS REMOVEDFROM 4S before 3D q. Quantum Numbers i. Address for an electron ii. Consists of the shell, the orbital, the orientation of the orbital, and the alignment of the magnetic field r. Principle Quantum number, n i. Describes the shell+ ii. Can be any integer greater than zero s. Angular momentum,l (subshell) i. Describes the orbital shape ii. Must be less than the value of n (n-1) iii. Can be positive or zero iv. For n
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