Part 1- Inorganic Chemistry Flashcards
Hydroxide
OH 1-
Carbonate
CO3 2-
Perchlorate
ClO4 1-
Cyanide
CN 1-
Bicarbonate
HCO3 1-
Chlorate
ClO3 1-
Sulfate
SO4 2-
Peroxide
O2 2-
Chlorite
ClO2 1-
Phosphate
PO4 3-
Superoxide
O2 1-
Hypochlorite
ClO 1-
Nitrate
NO3 1-
Permanganate
MnO4 1-
Ammonium
NH4 1+
IA IIA
Na 1+ K 1+ Mg 2+ Ca 2+ Ba 2+
Important B group element
Fe 2+ or Fe 3+ Co 2+ or Co 3+ Ni 2+ Zn 2+ Ag 1+
IIIA VIA
Al 3+ O 2- S 2-
VIIA
F 1- Cl 1- Br 1- I 1-
Electron shell
1s, 2s, 3s, 4s, 5s, 6s, 7s
2p, 3p, 4p, 5p, 6p, 7p
3d, 4d, 5d, 6d, 7d
4f, 5f, 6f, 7f
Pv=nRT
P, Pressure if in mmHg or torr divide it by 760
R, gas constant .082 L atm/mole
K n, number of moles T, temperature kelvin if in C ad 273
V, is volume is in L
Dalton’s law of partial pressures
The total pressure of a gas is a simple sum of the partial pressures of the individual gases that are present in the gas mixture
London forces (Van Der Waal’s forces)
Are the weakest intermolecular forces Are by which the means by which neutral, non-polar molecules attract each other.
Dipole dipole
Polar molecules attract each other strong
Hydrogen bonding
Very strong F, O, or N Polar molecules
Diatomic atoms
Not polar example H2 O2 N2 F2 Cl2 Br2 I2
Water
H2O Bent polar Hydrogen bonding
Carbon dioxide
CO2 Linear non-polar
Sulfur dioxide
Bent and polar SO2
Sulfur trioxide
Trigonal planar and non polar SO3
Ammonia
NH3 Pyramidal and polar and hydrogen bonding
Methane
Tetrahedral and non polar
Sulfur hexafluoride
SF6 Octahedral and non polar
Ionic bonds
1) metal/nonmetal NaCl 2) common polyatomic ions NH4NO3, NH4Cl
Strength of intermolecular forces
Ion-ion > h-bonding > dipole-dipole > London
Stronger IMF mean
Higher boiling and meting
Size on IMF
Neutral, nonpolar atoms- larger molecules the stronger attraction Polar molecules- smaller molecules the stronger the attraction
homogenous
completely uniform (sugar dissolved in water)
solvent solute
solvent= the which does the dissolving solute= being dissolved solution always takes on the physical state of the solvent
concentration
solute/solvent
Molarity
moles of solute/ V of solution in L Moles x V(L)= # of moles Moles x M(ml) = # of mmoles
colligative properties raoult’s law
solution value is based on number of solutes molecules not the nature of the solute molecules example: adding non-volatile to volatile solvent lowest he vapor pressure. P= Xsolvent x P degree
- Xsolvent= mole fraction of the solvent, (moles of solvent is moles of solvent/total moles of solution)
- P is vapor pressure
- P degree is vapor pressure of solvent
boiling point elevetation freezing point
delta Tb= Kb x m delta Tb= how much boiling point elevated Kb= constant characteristic of solvent how much the boiling point is elevated per unit of molarity m= molarity of solute add this amount to the regular BP of water or solution….. this will give you the new BP delta Tf= Kf x m subtract this form regular freezing temp for new temp
greatest colligative effect
the greatest colligative effect is the solution with the largest solute concentration this includes lowest vapor pressure or highest BP, or the lowest FP or the largest osmotic pressure
osmotic pressure
the flow of solvent through a semi-permeable membrane, in an effort to equalize concentration proportional to solute concentration…. so more concentraiton, more ion the better. like KCl vs. K2SO4
solute concentration
ionic solutes (electrolytes) - will have a greater concentration than non-ionic solutes (non-electrolytes)
common strong electrolyte solute
water soluble strong acids- HCl, HNO3, H2SO4 water soluble strong bases- NaOH, KOH Water soluble salts- NaCl, KNO3, K2SO4
common weak electrolyte solutes
water soluble molecular weak acids- CH3COOH water soluble molecular weak bases- NH3
Common non-electrolyte solutes
glucose C6H12O6 or sucrose sugar alcohols such as methanol CH3OH
common strong acid
hydrochloric acid HCL nitric acid HNO3 Sulfuric acid H2SO4 Perchloricacid HClO4
common weak acid
acetic acid HC2H3O2 or CH3COOH carbonic acid H2CO3 hydrofluoric acid HF phosphoric acid H3PO4
common strong base
IA and IIAA hydroxide salts, KOH NaOH etc.
common weak bases
NH3
calorie
defined as the energy needed to raise the temperature of one gram of liquid water by one degree Celcius
joule
is the SI unit of energy and is defined as one cal= 4.18 J
SH specific heat
is defined as the heat needed to raise the temperature of one gram of the substance by one degree Celsius SH of h2o = 1 cal/g degrec celsius= 4.18 J/g degree celsius
heat calculations involving no change in physical state
q (heat)= mass x SH x change in temp
heat of fusion
q (FUS) amount of heat needed to convert one gram of solid to liquid at the substance’s melting point water q (fus)= 80 cal/g= 334 J/q
heat of vaporation
amount of heat needed to convert one gram of liquid to vapor at the liquid’s boiling point q vap= 540 cal/g= 2260 J/g
entropy
delta s positive- disorder increases
negative- disorder decrease
change in delta is= product- reactant
enthalpy
delta H, heat change in the reaction exothermic- negative endothermic- positive
gibbs free energy delta G
delta G= delta H - T delta S
what is not in the equilibrium
solids, pure liquids and water
LeChatelier’s principle
- changes in concentration 2. changes in pressure 3. changes in temperature
decrease in pressure
equalibritim will shift to the side with more moles of gase this change means there is an increase in volume
anhydride
formed by the removal of one of H2O from the molecuels H2SO4 turning into SO3 plus H20
what phsyical state does helium fail to exhibit
solid
sublimation deposition
sublimation= solid to gas deposition= gas to solid
K(sp)
smaller K(sp) the less soluble the slat will be
one salt disolving in pure water
x=solubilty= sq(K(sp))
common ion effect
if salt is initially present in solution, the solubility of the salt in the solution will be significantly less than its solubility in pure water
oxidation reduction ocurres to waht
oxidation occurs to anod reduction occurs to cathode
