Organic Chem

Front 1

Four quantum numbers

Back 1

1. Principal quantum number

2.angular momentum quantum number

3.magnetic quantum number

4.spin quantum number

Front 2

Principal quantum number

Back 2

Can have any integer value greater than zero

Front 3

Magnetic quantum number

Back 3

Can be 0 as well as both negative and positive integers up to +or - L

Front 4

Spin quantum number

Back 4

+1/2 (up) and -1/2 (down)

Front 5

Atomic orbital

Back 5

The probability of finding an electron of a certain energy on a particular region of space

-s: spherical

-p: dumbbell-shaped

Front 6

Valence electron

Back 6

-The electrons in the outermost shell

-least tightly held electrons

-group number = # of valence electrons

Front 7

Electron configuration

Back 7

- a list of the atomic orbitals occupied by electrons in an atom

Front 8

Simple rules to fill atomic orbitals with electrons

Back 8

1. Aufbau principle

2. Pauli exclusion principle

3. Hund's rule

Front 9

Aufbau principle

Back 9

Aufbau: building up

-states that the lowest-energy orbital is filled first

Front 10

pauli exclusion principle

Back 10

-each orbital can have a maximum of 2 electrons that have opposite spin

Front 11

Hund's Rule

Back 11

-with degenerate orbitals, 1 electron is placed in each degenerate orbital first, before electrons pair up

Front 12

valence bond model

Back 12

-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

Front 13

sigma bond

Back 13

-covalent bond formed by head-on overlap of atomic orbitals

Front 14

pi bond

Back 14

-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

Front 15

octet rule

Back 15

-the tendency of atoms in molecules to have 8 electrons in their valence shells (2 for hydrogen atoms)

Front 16

Valence shell electron pair repulsion (VSEPR) theory

Back 16

-bonds + electrons are arranged about central atom so bonds are far apart as possible

-used to predict bond angles

Front 17

Hybrid atomic orbitals

Back 17

1. Sp3, tetrahedral, 109

2. Sp2, trigonal planar, 120

3. Sp, linear, 180

Front 18

Hybridization

Back 18

- 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

Front 19

bond line-structure

Back 19

1) dash

2) condensed

3) bond-line

Front 20

rules for drawing bond-line structures

Back 20

-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

Front 21

functional group

Back 21

-a characteristic group of atoms/bonds that possess a predictable chemical behavior

Front 22

organic compounds containing oxygen

Back 22

-alcohol + classification

-phenol

-ether

-aldehyde

-ketone

-carboxylic acid

-ester

Front 23

organic groups containing nitrogen

Back 23

-amine and classification

-amide

-nitrile

Front 24

Hydrocarbons

Back 24

-alkane

-alkene

-alkyne

Front 25

covalent bond

Back 25

-sharing of electrons, usually two non-metals

Front 26

ionic bond

Back 26

-electrostatic attraction of cations and anions

-usually a metal + non metal

Front 27

electronegativity

Back 27

-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

Front 28

polar covalent bond

Back 28

-a bond between atoms with significantly different electronegativities

Front 29

Formal charge

Back 29

valence electrons of atom - valence electrons of atom in molecule

-carbons 4

-oxygen 2

-halogens 1

-with abnormal # of bonds assign a formal charge

Front 30

lewis structures

Back 30

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

Front 31

2 types of electron counting

Back 31

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

Front 32

rules for resonance

Back 32

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

Front 33

major vs. minor resonance contributor guidelines

Back 33

1) octet

2) charge compatibility

3) charge separation

Front 34

arrow pushing

Back 34

-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

Front 35

how do you generate a resonance structure of a given Lewis structure?

Back 35

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

Front 36

acid

Back 36

-a substance that donates a proton

Front 37

conjugate base

Back 37

-base formed when an acid donates a proton

Front 38

base

Back 38

-a substance that accepts a proton

Front 39

conjugate acid

Back 39

-an acid hat is formed when a base accepts a proton

Front 40

the smaller the pKa

Back 40

the stronger the acid

Front 41

lewis base

Back 41

-substance donates an electron lone pair to an acid

Front 42

lewis acid

Back 42

-substance accepts an electron pair from a base

Front 43

curved arrow formalism

Back 43

-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

Front 44

predicting outcome of acid-base reactions

Back 44

-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

Front 45

compare acidity of any 2 acids

Back 45

-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