CH 101 Final Exam

Avogadro's number

6.022 x 10^23

Isotopes

Have the same atomic number (# of protons) but different mass (change in # of neutrons)

Cations

Positively charged

Anions

Negatively charged

Neutron

1
n
0

Proton

1 1
H or P
1 1

Alpha Particle

4
He
2

Beta Particle

0
e
-1

Positron

0
e (electron with positive charge)
1

Wavelength

Distance between 2 adjacent maxima

Frequency

# of waves/sec

Speed of Light (C)

3.00 x 10^8 ms-1

h

6.626 x 10^-34 JS

Wave-like behavior

C = lambda x v

Particle-like behavior

E = hv

Quantum Number n

n = 1,2,3,4,...; defines row of periodic table; higher the n, further nucleus, higher energy

Quantum Number l

l= 1,2,...,(n-1); l= 0(s), 1(p), 2(d), 3(f); designates subshell; defines shape of the orbital

Quantum number ml

-l =< ml =< +l; describes orientation of the orbital in space

Quantum number ms

ms = +1/2 or -1/2; spin up or spin down

Zeff

Actual nuclear charge experienced by specific electron; Zeff = Z - sigma

0.35 (slater's rules)

electrons in same (ns, np) group

0.85 (slater's rules)

electrons in the (n-1) shell

1.00 (slater's rules)

electrons (n-2) or lower groups

Atomic Radius

Increase down group and decrease across period

Ionization Energy (IE)

Decrease down group and increase across period

Electronegativity

Increase diagonally across table

Paramagnetic

Unpaired electrons; magnetic

Diamagnetic

No unpaired electrons; not magnetic

Oxidation State

Charge of atom assuming all bonds are ionic

Ammonium

NH4^1+

Hydroxide

OH1-

Perchlorate

ClO4^1-

Nitrate

NO3^1-

Carbonate

CO3^2-

Sulfate

SO4^2-

Phosphate

PO4^3-

SP

1/2(ER -VE)

Bond Order

(# of bonds)/(# of bonding areas)

Formal Charge

VE - Assigned e-

Extended Octet

Exception for lewis structures in row 3

Resonance

Lewis structures that differ only in the placement of electrons

Linear

180; linear

Trigonal Planar

120; trigonal planar or bent

Tetrahedral

109; tetrahedral, trigonal pyramidal, bent

Sigma Bond

End on overlap; first

Pi Bond

Side on overlap; second

sp

2 e- groups

sp2

3 e- groups

sp3

4 e- groups

Bond Order (MO Theory)

1/2(BE -ABE)

Ideal Gas Law

PV=nRT

STP

273 K, 1.00 atm, 1.00 mol

Ion-Dipole

Attraction between ion and polar molecules

London Dispersion

All molecules have dispersion forces (even noble gases)

Dipole-Dipole

Molecules with dipole are polar; based on geometry of molecule

Hydrogen Bonding

F-O-N

Enthalpy (DeltaH)

Heat of Reaction; Broken-Formed

Exothermic

Energy released; DeltaH < 0

Endothermic

Energy absorbed; DeltaH > 0

Entropy (DeltaS)

Molecular motion; Product - Reactant

Gibbs Free Energy (DeltaG)

Describes reaction spontaneity; DeltaH - TDeltaS

Keq

[Product]/[Reactant]

Bronsted Acid

Proton donor

Bronsted Base

Proton acceptor

Strong Acid

Ka > 1

Weak Acid

Ka < 1

Lower the pKa

The stronger the acid

Neutral

[H3O+] = [OH-]

Acidic

[H3O+] > [OH-]

Basic

[H3O+] < [OH-]

Lewis Acid

Electron acceptor; electrophile

Lewis Base

Electron donor; nucleophile

Redox Reactions

Transfer of electrons from one reactant to another

Reduction

Gain of electrons; oxidation state decr.

Oxidation

Loss of electrons; oxidation state incr.; anode

Reducing Agent

Reductant; reduces another substance; loses electrons and therefore is oxidized

Oxidizing Agent

Oxidant; oxidizes another substance; gains electrons and therefore is reduced

Ereduced - Eoxidized

Erxn

-nFE

DeltaG (redox reaction)
F- 96,458 c/mol

Insoluble

Ksp < 1; reactants favored

Soluble

Ksp > 1; products formed

Strong Electrolytes

Ionic compounds, strong acids/bases

Weak Electrolytes

Weak acids/bases

Nonelectrolytes

All other covalent compounds