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16 Cards in this Set
- Front
- Back
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Bohr's Explanation of Electrons and Orbitals |
- Electrons can only occupy certain orbits because they have certain energies - Electrons in permitted orbits have specific energies that will not be radiated from the atom - Energy is absorbed or emitted which allows the electron to move from one energy state to another |
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Electrons and its relationship with spectrum |
Closest to nucleus is n=1, then n=2, then n=3
Higher n #, the more its likely to emit a red color (high wavelength)
Lower n#, more likely to emit a violet color |
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Heisenberg Uncertainty Principle |
It is impossible to know the position of a particle and the speed and its direction at the same time
The uncertainty of an electron's position can be greater than the size of the atom |
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Quantum Mechanics: Def and who developed it? |
Def: Mathematical treatment into which both th ewave and particle nature of matter is incorporated
He figured out where electrons are most likely to be using orbitals and quantum clouds |
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Quantum Numbers: N, L, M |
Describes an orbital. Each orbital describes a spatial distribution of electron density
N: Describes energy level that the orbital resides, equl to the number of the row on the periodic table
L: Describes the shape of the orbital. 0(s), 1(p), 2(d), 3(f)
M: Describes the three-dimensional orientation of the orbital. On any given level, there can be up to 1s, 3p, 5d, 7f. Each with the ability to hold 2 electrons each |
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Nodes |
Regions in orbitals where there is zero probability of finding an electron |
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Zeff (Effective Nuclear Charge): Trends across a periodic table? |
The positive charge of nuclear protons acting on valence electrons
- Electrons closer to the nucleus sheild electrons further out from feeling the full positive charge of the nucleus
- Increases left to right (period) - Decreases top to bottom (group) |
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Periodic Trends: Size of atoms, and Zeff |
Left to right: Size decreases, Zeff increases (electrons are pulled closer to the nucleus)
Top to bottom: Size increases as n increases |
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Electron Configuration of Ions |
Remove electrons from the orbitals with highest n value
Ex: Li (1s2,2s1) : Li+ (1s2)
For Anions, add eletrons to unfilled orbits with the lowest n value first
Ex: F (1s2. 2s2, 2p5) : F (1s2, 2s2, 2p6) |
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Ionization Energy: 1st & 2nd |
Amount of energy required to remove an electron from the ground state
1st: Energy required to remove the first electron 2nd: Energy required to remove the second e
- All ionization energies are endothermic. Energy needs to be put into the system
- Filled shells are most stable and hardest to ionize
- When all valence electrons are removed, the next ionization energy becomes much larger |
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Electron Affinity |
Change in energy when an electron is added to a gaseous atom to form a negative ion.
- This reaction gives off heat energy (exothermic, and a negative change in Enthalpy) |
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Periodic Table Trends: Ionization and Electron AFfinity |
IE: Increases Left to Right (because it gets harder to remove an electron with increasing Zeff) Decreases Top to Bottom (because electrons are farther away from nucleus)
EA: Increases Left to Right (because it becomes more exothermic) Slightly decreases Top to Bottom (atom is bigger so less attraction)
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Pauli Exclusion Principle |
- No 2 electrons in an atom can have the same 4 quantum numbers
- Two electrons' spins must be paired. No more than 2 electrons can occupy the same orbital |
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Hund's Rule |
- Fills Orbital from lowest energy to highest energy
- Add electrons to degenerate orbitals with parallel spins before pairing the electrons |
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Aufbau Building Up Principle |
The combination of the Pauli Exclusion Principle and Hund's Rule |
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Bonding Atomic Radius vs. Non-Bonding Atomic Radius |
Bonding: based on the distance between nuclei in molecular compounds
Non-Bonding: the distance from the nucleus to the edge of the electron cloud of the valence electrons |
