| Question | Answer |
| atomic radius | Decreases from left to right across a period. Increase effective nuclear charge, decreased shielding. Increases down a group. Increases in the orbital sizes in successive principal quantum levels |
| electron affinity | Energy change associated with the addition of an electron to a gaseous atom. Negative if reaction exothermic, positive for endothermic. Energy increases down and to the left. Electrons further from nucleus, weaker attractions. To the left, easier to achieve noble gas formation by losing e- |
| first ionization energy | Energy to remove the highest-energy electron of an atom. Increases from left to right across the first six periods. Decreases down the groups. Further from the nucleus, easier it is to take an electron off |
| ionization energy | Energy required to remove an electron from a gaseous atom or ion |
| core electrons | Everything inside valence electrons |
| Valence electrons | The electrons in the outermost principal quantum level of an atom |
| Hund's Rule | The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate orbitals |
| Aufbau Principle | As protons are added one by one to the nucleus to build up the elements, electrons are similarly added to these hydrogen-like orbitals |
| Why do atoms have a tendency towards the s orbital | Although spends most of its time a little father from nucleus than p, d, or f, spends some time very near the nucleus. Penetration effect: In s orbital, electron penetrates to the nucleus more than other orbitals. Allows for lower energy |
| electron correlation problem | Schrodinger equation does not account for electron repulsions. To solve, treat each electron as if it were moving in a field of charge that is the net result of the nuclear attraction and the average repulsions of all the other electrons |
| polyelectronic atoms | Atoms with more than one electron |
| Pauli exclusion principal | In a given atom, no two electrons can have the same set of four quantum numbers (n, l, ml, ms) |
| electron spin quantum number | ms, Can only have 2 values. -1/2, 1/2 |
| electron spin | 2 possible orientations of electrons in an orbital. Two oppositely directed magnetic movement |
| orbital energies | Energy of particular orbital determined by its value of n. All orbitals with the same value of n have the same energy (degenerate) |
| nodes/nodal surfaces | Areas of zero probability of electrons. Number of nodes increases as n increases |
| magnetic quantum number | ml, Integral values between l and -l, Each set of orbitals with given value of l (or subshell) |
| angular momentum quantum number | l, Integral values, 0 for n=1, 0, 1 for n=2, 0, 1, 2 for n=3, 0 = s, 1 = p, 3 = d, 4 = f, Relates to shape of atomic orbital |
| principal quantum number | n, Integral values 1, 2, 3…, Dictates size and energy of orbital |
| quantum numbers | Wave functions/orbitals, describe various properties of the orbital |
| radial probability distribution | Total probability of finding the electron in each spherical shell is plotted versus the distance from the nucleus. Probability greatest near the nucleus, but the volume of the shell increases with the distance from the nucleus. Total probability increases to a certain radius, and then decreases as the electron probability at each position becomes very small. Based completely on probability |
| probability distribution | The square of the function indicates the probability of finding an electron near a particular point in space. Intensity of color indicates probability of value near a given point. Electron Density Map |
| Heisenberg Uncertainty Principal | There is a fundamental limitation to just how precisely we can know both the position and momentum of a particle at a given time. The more accurately we know a particle's position, the less accurately we can know its momentum. Not necessary for large things, Vital for small, like electrons |
| Quantum (wave) mechanical model | Wave function corresponding to the lowest energy for hydrogen atom |
| orbital | Specific wave function. Not Bohr orbital. Do not know how electron is actually moving. No detailed path of electron |
| wave function | Ψ. Function of the coordinates x, y and z = Electrons position in space |
| Standing wave | Similar to electron bound to nucleus. Stationary, do not travel along. Limitations to standing waves. Always a node at each end, must be a whole number of half wavelengths in allowed motions. Only certain circular orbits have a circumference into which a whole number of wavelengths will 'fit' |
| Why is the Bohr Model incorrect | Only works for Hydrogen. Important historically, paved the way for later theories. Current theory of atomic structure has nothing to do with Bohr model. No circular orbits for electrons |
| Bohr Model | 1. Correctly fits the quantized energy levels of the Hydrogen atoms and postulates only certain allowed circular orbits for the electron. 2. As the electron becomes more tightly bound, its energy becomes more negative relative to the zero-energy reference state. Closer to the nucleus, energy released. |
| ground state | Lowest possible energy state |
| Quantum Model | Niels Bohr. Electron in a hydrogen atom moves around the nucleus only in certain allowed circular orbits. Radii calculated. Travels in a circle, made some assumptions. Created expression for energy levels available to hydrogen atoms |
| Hydrogen line spectrum | Only certain energies are allowed for the electron in the hydrogen atom. Energy Quantized |
| line spectrum | Only a few lines which correspond to discrete wavelength |
| Emission spectrum | Created by excited electrons |
| Diffraction Pattern | Created when light diffracts. Can only be explained in terms of waves |
| Diffraction | Light scattered from a regular array of points or lines. Wavelengths of visible light are not all scattered in the same colors, but 'separate.' Works the best when spacing between the scattering points is about the same as the wavelength of the wave being diffracted |
| De Broglie's equation | m = h/ λv |
| Dual nature of light | Certain characteristics of light that particulate matter, certain that particulate waves |
| Energy of a photon equation | hc / λ. Photons do exhibit apparent mass calculated. No rest mass. |
| E = mc^2 | Eistein's special theory of relativity. Energy has mass. |
| The photoelectric effect | Electrons are emitted from the surface of a metal when light strikes it. 1. Threshold frequency v0, electrons not emitted below it regardless of intensity. 2. Number of electrons increase with intensity of light, if frequency above threshold. 3. Kinetic energy of emitted electrons increases linearly with frequency of light. Minimum energy required to remove electron = E0 = hv0 |
| Quantum | 'Packets' of energy. Energy can only transfer in whole quanta |
| Change in energy formula | nhv |
| Planck's Constant | h = 6.626x10^-34 Js |
| Max Plank | Postulated energy can be gained of lost only in whole-number multiples of the quantity hv where h is a constant |
| hertz | Unit per second. 1/s. Frequency measurement. Hz |
| 3 characteristics of electromagnetic radiation | wavelength, frequency, speed |
| Electromagnetic radiation | Energy traveling through space. Travels at the speed of light in a vacuum |
48 cards - created nov 18, 7:06pm