Chem 6A (UCSD-) FINAL

Reactivity

Reactivity increases as you go down the periodic table.

Metallic

Metallic increases down a group and to the left.

Ionization energy

The energy to remove an electron from a molecule.

Increases to the right and up a group.

frequency and wavelength relationship

Energy increases with frequency.

Long wave=small frequency

short wavelength= large frequency

Quanta

energy is released in small packets.

incident frequency

Need minimum frequency to emit an e-.

-unqiue to element

-velocity of ejected e- increases with frequency

mass vs. wavelength

as mass increases the wavelength decreases

particle in box

increase atom size = increase energy level

-zero point energy = n=1 (lowest ground state)

atomic radius

increases down a group and to the left

electron affinity

the amount of energy released or spent when an electron is added to a neutral atom

Ea=(energy of anion)-(energy of anion + e-)

-increases across period and up a group

Uncertainty Principle

-you can not know the position and momentum at the same time

-larger mass= smaller uncertanty

deltaxdeltav=(h/4piem)

N

-principal quantum number

-energy and size

-is # of the period or row

-number of nodes = n-1

-nxn = possible orbitals

ex) 1,2,3

l

-different possible values for a given value of n

-specifics the shape of orbital

code: 0 1 2 3 4 5

s p f d g h

- from 0 to n-1

ml

-orientation

ml=+l to -l

ms

-spin #

+1/2 to -1/2

orbital energys

-s

electron configuration

-3d= higher energy than 4s-> once filled= lowers in energy

-once F orbital is filled=lower energy

diamagnetic

has no unpaired e-

-unattracted to field

paramagnetic

has unpaired e-

-attracted to field

effective nuclear charge

-nuclear charge experienced by outermost electrons

-rough calculation= valence e- and subtract the e- from the previous closed shells

-valence e- shielded by the e-. Feel a positive charge = size smaller as you go from left to right.

Atomic radii

-e- on outside feel + charge

-transiton elements feel around the same charge bc d orbital is bad @ shielding

-lose e-=size decrease bc e-/proton attration increases

-gain e-=size increases b/c e-/proton attration decreases

why bond forms?

bonds occur b/c it lowers the energy of the system

Ionic

-completely transfer of e-

-metal/ nonmetal

-strongest bonds/ higher boiling and melting

covalent

-share of electrons

-non metal/non metal

-intermolecular forces = not a real bond

-lower boiling and melting

coulombs law

e=k(q1q2)/d

electronegativity

-(highest)FONClBrTSCH(lowest)

-tendency of an atom to attract a bonding pair of electrons

Lewis structure

1.count valence electrons

2.symbol + single bond

3.complete the octect for periphal atom

4.left over e- go on central atom

5. if not enough e- try more bonds

resonance

-only occurs btw structures with the same arrangment of atoms

-resonance stabilizes molecule by lowering it's total energy

-one that shares better = more stable

formal charge

-measures how much e- atoms own

Fc=v-(l+(s/2))

v=#e-

l=# of lone pair e-

s=# of shared e-

exception to the octet rule

1. hydrogen (duet rule)

2. radicals (odd # of e-/keep making radicials)

3.expanded valence e- shell

-must be period two and down b/c they have expanded D orbital to expand

4.electron deficient

-representation of lewis base/acid

ex)B and Al

5. transition metals and actinides

acid/base

-acid=electron pair acceptor

-base= electron pair donor

electronegativity in bonding

-pulling power of atom

-greater pulling power= gets e-

-High e- affinity and ionization= high electronegativity

formula: 1/2(Ie+Ea)

polar covalent

- partially negative or positive

(+) sign over positively charged molecule and arrrow toward negative molecule

determining Bonding sharing

delta x< 0.4 covalent

0.22.0. ionic

Polarizing power

-cations

-increase with decrease in cation size and increase with charge

Polarizability

-anions

-mainly dependent on size

-bigger = more polarizable

-more e-= cloud shifts

Bond strength

decrease in bond length = increase in energy

VESPR

-(replusion) = lp-lp>lp-bp>bp-bp

-e- reduce ideal bond length b/c more lp= more replusion = shorter bond length

-most sense to put e- on equitorial to lower repulsion

electron geometry

-include e-

-can only be the 5 basic shapes

-if you have lone pair= molecular never matches electronic

molecular polarity

- like dissolves like

-dipole molecule= orientation of polar molecules in eletric field

-molecules will be polar if:

A) bonds are polar

B) if mole is not symetrical

shape can contribute to polarity

-non polar= shape without e- pairs

-5 basic shapes will never be polar

Dipole

-the greater the Dipole = the more polar the molecule

polarity

-more polar = more chemical /physical properties

Isomers

-same number and type of atom but different in how atoms are arranged -> different properties

-two types : structural and stereoisomers

Structural (constitutional)

-same chemical formula but different connectivity (physcially located at different part of the molecule)

-react differently

-if you have to pick up your pen when tracing then it is a different isomer

stereoisomers

-same physical attachments but different in arrangement

-Types: geometric and optic

style:

1) draw the reflection of it

2) keep static but switch two adjacent

Geometric

-cis = same side

-trans=opposite side

-square planar and octahedral

Optic isomers

-same formula and same connectivity

-nonsuperimpossibale (mirror image of that molecule is a different molecule)

aka enantiomers of one another - chiral

ex) can't have for square planar

square planar

-turns on 2 axis so sit can change to be superimposable

-always have cis and trans

tetrahedral

-must have 4 different things attached to be chiral

(only have optical)

-doesnt have cis and trans

valence e- theory

-e- are localized

-set of overlapping orbitals have 2 e- at most

-half filled orbitals overlap to make bonds-> more overlap= stronger

-hybridization

sigma bond

-can only make one sigma between two atoms

-can rotate around

pie bond

-single bond and a pie bond

-can not rotate pie bonds

-double bond = leave 1 orbital

-triple= leave 2 orbitals

determining hybridization

1.lewis dot

2.count # of things attatched

3.remember order of atomic orbitals

SPPPDDDD

4.starting with s , count off an orbital for every thing

-hybrid orbitals can ONLY form sigma bonds

molecular orbital theory

-valence e- are delocalized

-valence are in orbitals called molecular orbitals spread over entire molecule

bonding order

-def: the number of pairs of e- the molecules are sharing

=1/2(bonding e- - anti bonding e-)

wave particle duality

-electrons exhibit properties of both wave and particles

-around nucleus= wave

-away/moving freely=particle

S orbital

-sphereical shape

-l=0

P orbital

-plannar node through nucleus

-l=1

D orbital

-2 planar nodes

-l=2

-daisy shape

Ion Configurations

-lose electrons in order : np,ns and (n-1)d

-exceptions e- taken out of S orbital to fill D orbital

isoelectronic

-same number of electrons

metallic

-electrons share a pool of e-

-usually between metal and metal

Electrolytes

Def:conduct electricity by having e- present

-strong= many ions->soluble in water ( dissociates with greater charges= stronger)

-weak =only few ions->weak acid and weak bases

-nonelectrolytes= no ions

MO theory diagram

1.write electron configuration

2.draw basic diagram

3.fill in where electrons are

( can use valence or all orbitals)

4. fill from lower orbital to higher orbital

MO sigma

-heteronuclear diatomic a does not equal b

-homonuclear diatomic a=b

-the more electron negative gets the larger constant