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45 Cards in this Set
- Front
- Back
Four quantum numbers |
1. Principal quantum number 2.angular momentum quantum number 3.magnetic quantum number 4.spin quantum number |
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Principal quantum number |
Can have any integer value greater than zero |
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Magnetic quantum number |
Can be 0 as well as both negative and positive integers up to +or - L |
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Spin quantum number |
+1/2 (up) and -1/2 (down) |
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Atomic orbital |
The probability of finding an electron of a certain energy on a particular region of space -s: spherical -p: dumbbell-shaped |
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Valence electron |
-The electrons in the outermost shell -least tightly held electrons -group number = # of valence electrons |
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Electron configuration |
- a list of the atomic orbitals occupied by electrons in an atom |
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Simple rules to fill atomic orbitals with electrons |
1. Aufbau principle 2. Pauli exclusion principle 3. Hund's rule |
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Aufbau principle |
Aufbau: building up -states that the lowest-energy orbital is filled first |
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pauli exclusion principle |
-each orbital can have a maximum of 2 electrons that have opposite spin |
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Hund's Rule |
-with degenerate orbitals, 1 electron is placed in each degenerate orbital first, before electrons pair up |
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valence bond model |
-a covalent bond between 2 atoms is formed in terms of an in-phase overlap of a half-lifted orbital of one atom with a half-filled orbital of the other |
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sigma bond |
-covalent bond formed by head-on overlap of atomic orbitals |
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pi bond |
-covalent bond formed by sideways overlap of atomic orbitals ex: carbon-carbon double bonds contain a pie bond formed sideways overlap of 2 p orbitals |
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octet rule |
-the tendency of atoms in molecules to have 8 electrons in their valence shells (2 for hydrogen atoms) |
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Valence shell electron pair repulsion (VSEPR) theory |
-bonds + electrons are arranged about central atom so bonds are far apart as possible -used to predict bond angles |
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Hybrid atomic orbitals |
1. Sp3, tetrahedral, 109 2. Sp2, trigonal planar, 120 3. Sp, linear, 180 |
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Hybridization |
- hybrid atomic orbitals are used to make sigma bonds -hold lone pairs in resonance -used as empty orbitals -add up number of lone pairs not used in resonance + number of sigma bonds |
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bond line-structure |
1) dash 2) condensed 3) bond-line |
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rules for drawing bond-line structures |
-assume theres a C agtom at each intersection of 2 or more lines at end of line -assume enough H around C to give it 4 bonds -draw all heteroatoms +show H atoms directly attached to them |
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functional group |
-a characteristic group of atoms/bonds that possess a predictable chemical behavior |
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organic compounds containing oxygen |
-alcohol + classification -phenol -ether -aldehyde -ketone -carboxylic acid -ester |
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organic groups containing nitrogen |
-amine and classification -amide -nitrile |
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Hydrocarbons |
-alkane -alkene -alkyne |
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covalent bond |
-sharing of electrons, usually two non-metals |
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ionic bond |
-electrostatic attraction of cations and anions -usually a metal + non metal |
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electronegativity |
-tendency of an atom to attract electrons to itself in a covalent bond -increases across periodic table from left to right and bottom to top |
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polar covalent bond |
-a bond between atoms with significantly different electronegativities |
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Formal charge |
valence electrons of atom - valence electrons of atom in molecule -carbons 4 -oxygen 2 -halogens 1 -with abnormal # of bonds assign a formal charge |
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lewis structures |
1) count number of valence electrons -if ion, add or subtract electrons to give proper change 2) draw skeleton structure for species, join atoms by single bonds 3) determine number of valence electrons available for distribution -to this deduct 2 valence electrons for each single bond in step 2 4) determine number of valence electrons required to fill out an octet or each atom -if number of electrons available is less than the number required then must add bonds |
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2 types of electron counting |
1. count all lone pair electrons + all bonding electrons to know whether atom has complete octet 2. for formal charge- count all lone pair electrons + half the bonding elctrons |
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rules for resonance |
1) do not break a single bond 2) do not exceed octet for second period elements 3) electronegative atoms like O and N must have octets 4) for a charged system, do not generate additional charge -positions stay unchanged; only pie or lone pair electrons move |
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major vs. minor resonance contributor guidelines |
1) octet 2) charge compatibility 3) charge separation |
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arrow pushing |
-always push electrons -never push positive charges -push electrons away from the negative charge and toward the positive charge -electrons must be pushed away from center of high electron density and toward low electron density |
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how do you generate a resonance structure of a given Lewis structure? |
1) look for a lone pair to a positive charge -push electrons toward positive 2) lone pair next to pie bond -nonbonding electrons pushed away from center of negative charge 3) pie bond next to positive charge -pie electron pairs in multiple bonds toward positive charge 4) pie bond between 2 atoms where 1 atom is electronegative -pie electron pairs in multiple bonds toward the more electronegative atom 5) alternating double and single bonds looped around in a ring |
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acid |
-a substance that donates a proton |
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conjugate base |
-base formed when an acid donates a proton |
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base |
-a substance that accepts a proton |
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conjugate acid |
-an acid hat is formed when a base accepts a proton |
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the smaller the pKa |
the stronger the acid |
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lewis base |
-substance donates an electron lone pair to an acid |
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lewis acid |
-substance accepts an electron pair from a base |
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curved arrow formalism |
-bonds are made when electron-rich atom donates a pair of electrons to an electron-poor -bonds are broken when one atom leaves with both electrons from the former bond |
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predicting outcome of acid-base reactions |
-equilibrium lies to the side of products or reactants -acid-base reactions always favor the formation of the weaker acid - the stability of conjugate base is good guide to acidity |
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compare acidity of any 2 acids |
-always draw conjugate bases -determine conjugate base is more stable -the more stable the conjugate base, the more acidic the acid 1) electronegativity 2) atom size 3) resonance 4) inductive effects 5) hybridization |