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| where is carbon located in the periodic table? |
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| atoms to the left of carbon tend to |
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| atoms to the right of carbon tend to |
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| what prevents the positively charged nucleus from drawing in the negative electrons? |
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Definition
| kinetic energy of the electrons |
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Term
| where does the mass of an atoms mostly come from? the volume? |
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Definition
mass-protons and neutrons=nucleus
volume-electrons |
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Definition
| number of protons in a nucleus OR number of electrons on a neutral atom |
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| number of protons and neutrons |
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| have the same atomic numbers BUT different mass numbers because they have a different number of neutrons |
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| time it takes for 1/2 of the nuclei to decay |
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| average weighted mass of an element's atoms |
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| sum of atomic weights of all the atoms in a molecule |
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Definition
| describes the behavior of an electron |
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Definition
is the solution to a wave equation and tells
1. energy of electron
2.where it is most likely to be found |
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Definition
| characterizes movement of an electron around the nucleus like the wave motion of a guitar string |
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Definition
| 3-D region of volume where an e- is most likely to be found |
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Term
| the closer an atomic orbital is to the nucleus, the _______ it's energy |
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Definition
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Definition
| orbitals with the same energy |
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| what are the 3 principles that are considered in ground state electron configurations? |
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Definition
1. aufbau
2. pauli-exclusion
3. hund's rule |
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Term
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Definition
| an electron would rather occupy an available atomic orbital of the least energy (closer to the nucleus) |
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Term
| pauli-exclusion principle |
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Definition
1. max of 2 e- in each atomic orbital
2. they must be of opposite spin |
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Term
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Definition
| an e- would rather occupy an empty atomic orbital before one that already has an e- to minimize electron repulsion |
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Term
electropositive
where are they found? |
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Definition
| elements that readily lose an e- and become + charged; 1st column: the alkali metals |
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Term
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Definition
| elements that will readily gain an electron and become positive; column 17 |
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Definition
| bond due to only electrostatic attractions (attraction b/t opposite charges) and where electrons are transferred and NOT shared |
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Term
| if a hydrogen atom loses its only e- then it becomes |
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Definition
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Term
| if a hydrogen atom gains an e- to have 2 electrons in its outer shell it is a |
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Definition
| negatively charged ion: a hydride ion |
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Definition
| electrons are shared equally |
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Definition
| the atoms of the bond have different electronegativities,therefore the e- are NOT shared equally |
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Term
| the greater the difference of the electronegativity in a polar bond the more ___________ the bond is |
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Definition
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Term
| what type of bond has a dipole? |
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Definition
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Definition
| a polar bond that has a - and + end |
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Term
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Definition
| the size of the dipole in a polar bond |
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Term
| what is the formula used to determine the dipole moment? |
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Definition
μ=e x d
μ= dipole moment
e= magnitude of charge
d= distance between the charges |
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Term
| what is the charge on an electron in electrostatic units (esu)? |
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Definition
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Term
| a bond length of 1.39Å would be what value in cm? |
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Definition
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Term
| 6.97 esu cm or 6.97 x 10-18 esu cm would be what value in D (debye) units? |
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Definition
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Term
| electrostatic potential maps |
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Definition
| show the distribution of charge in a molecule |
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Term
| in an electrostatic map what does blue and red represent? what is most attracted to them? |
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Definition
red- most negative electrostatic potential: attracts + charge
blue- most positive electrostatic potential
attracts negative charge |
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Term
| why in the electrostatic potential maps for LiH and HF was the H in LiH larger? |
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Definition
| Because a potential map marks the edges of a molecule's electron cloud and the electron cloud around the H in LiH is largest because it has more electrons around it than the other H in HF. |
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Definition
| uses dots to represent valence electrons |
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Term
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Definition
| FC=valence electrons - (lone pair electrons + 1/2 bonding electrons) |
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Term
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Definition
| any species with a single lone pair electron |
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Term
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Definition
| lone pair e- are not shown except to make a point |
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Term
in a neutral atom how many lone pairs should these atoms ALWAYS have?
nitrogen-
oxygen-
halogen (such as chlorine)- |
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Definition
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Definition
| don't show covalent bonds and list atoms bonded together |
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Term
| Heisenberg uncertainty principle |
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Definition
| cannot determine location and momentum of an electron simultaneously |
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Term
| is the electron density more or less in 2s than in 1s? |
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Definition
| the e- density in 2s is less because it is a larger sphere that is farther away from the nucleus |
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Definition
| node found in spherical atomic orbital (s), probability of finding a e- is 0 |
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Definition
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Term
| molecular orbital theory (MO) |
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Definition
| covalent bonds form when AO combine to form MO |
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Term
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Definition
| formed when two "s" orbitals overlap or two "p" orbitals overlap end on end |
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Definition
| formed when two "p" atomic orbitals overlap side to side |
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Term
| which bonds are cylindrically symmetrical? which aren't? |
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Definition
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Definition
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Definition
| when maximum stability is achieved (which means minimal energy) |
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Definition
| energy needed to break a bond or energy that is released when the bond forms |
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Term
| when a covalent bond is formed, what happens in terms of energy? |
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Definition
| energy is released when a covalent bond is formed |
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Definition
| is constructive; e- are most likely in between nuclei, attracting them and increasing e- density which binds atoms |
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Term
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Definition
| destructive, there is a node so no e- will be there. the two + nuclei will repel each other |
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Term
| how does the amount of overlap affect the strength of a bond? |
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Definition
| the greater the overlap=stronger covalent bond |
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Term
| the strongest covalent bonds are formed by e- that occupy MO with the ______ energy. Why? |
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Definition
| least; because the lower the energy, the more stable |
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Term
how many nodes do these have (for s and p)?
σ
Π
σ*
Π* |
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Definition
σ- s:0 p:2
Π- p:1
σ*- s:1 p:3
Π*- p:2 |
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Term
| in terms of overlap, which is greater: p orbitals overlapping side to side or end to end? |
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Definition
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Term
which is more stronger?
σ
Π
σ*
Π*
place them in order |
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Definition
a σ is stronger than Π because it is more stable (less energy).
σ>Π>Π*>σ* |
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Term
| When 2 p AO form 2 MO, the AO of the more EN will contribute more to ______ while the least EN will for the _______ |
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Definition
more EN: bonding
lesser: antibonding |
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Term
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Definition
| valence shell electron pair repulsion: minimization of e- repulsion by the positioning of lone pairs as far as possible |
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Term
| what are the bond angles in methane, or any tetrahedron? |
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Definition
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Term
| in organic chemistry, all single bonds are ____ bonds |
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Definition
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Term
trigonal planar bond angle
what compound uses this? |
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Definition
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Term
| what are in the different bonds in terms of σ and Π? |
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Definition
single: 1 σ
double: 1 σ and 1 Π
triple: 1 σ and 2 Π |
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Term
| what is the strongest and shortest bond? why? |
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Definition
| triple; b/c its held together by 6 e- |
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Term
| what is the angle for a triple bond? |
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Definition
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