Week 1 Flashcards
Atomic # or Z
number of protons (retrieve from periodic table)
Mass #
N + Z (neutrons + protons)
isotopes
same # of protons but different # neutrons
alkali metals
group 1 - basic, react with H2O - explosive, exothermic and more reactive as you move down the group
alkaline earth metals
group 2 - basic, but not as reactive as alkaline metals
noble gases
group 8 - full octet, not very reactive
halogens
group 7 - single electron acceptor
transition metals
form bright colored compounds, i.e. copper = green/blue
exceptions to noble gas configurations
Cr, Mo, Cu, Ag, Au (because more stable if shells are half full or full)
s,p,d,f orbitals
s = 2e, p= 6e, d = 10e, f = 14e
paramagnetic
attracted to magnetic field (unpaired electrons)
diamagnetic
slight deflection of magnetic field (all electrons are paired)
*remember electrons like to fill each shell across before pairing up
Quantum #: n
principal - shell - [1….infinity]
Quantum #: l
azimuthal - subshell (type of orbital) - [0…(n-1)]
s = 0, p =1, d=2, f=3
Quantum #ml
magnetic - specific orbital (orientation in space), [-l…+l]
i.e. if l =2; [-2,-1,0,1,2] (typical d orbital has 5 orbitals)
Quantum #: ms
up or down - spin - (-1/2) or (+1/2)
Quantum # format
[n, l, ml, ms]
What forces hold the nucleus together?
strong nuclear force or nuclear binding energy
nuclear symbols
alpha (same as He), p (proton), n (neutron), beta (electron or positron), gamma (gamma ray light)
Stable Nucleus features
- even # protons and/or neutrons
- N/Z ratio ~ 1 for Z= 20 (up to Ca)
- daughter (product) of parent nuclear rxn is more stable due to release in energy
Alpha decay
reduces mass # of large nuclei (Z>83) - product = 4/2 alpha (like helium but not electrons)
Beta (-) decay/emmission
converts neutrons to protons (occurs when N/Z ratio is too high) product = electron or 0/-1 beta
beta (+) decay (positron emmission)
converts proton to neutrons (occurs when N/Z ratio is too low) product = 0/+1 beta
electron capture
converts protons to neutrons (occurs when N/Z ratio is too low) (reactant side = 0/-1 electron)
gamma decay
product is a photon (emitted when coming from excited state to ground state)
Most stable nucleus
iron-56 (56-Fe) —atoms closer to this is more likely to be more stable
Periodic Trends
- Atomic radius - down + left of the table; radius gets bigger because z(eff) is smaller
z(eff) = # of protons - shielding/core electrons - Ionic Radius - cations are usually smaller because they lose electrons, anions gain electrons and protons can’t hold as close
- non metals form anions
- metals form cations
- 1st ionization energy = energy req’d to remove 1 e-
(trend is opposite of atomic radius – up and right harder to remove e-)
*big jumps in ionization energy indicates all valence electrons have been removed and moving into core electrons
**EXCEPTIONS: Be + B and N + O
These are slightly lower because filled/half-filled (making them more stable (Be over B and N over O) - Electron affinity - energy in gaining an electron (exothermic) - increasing trend from left to right (no vertical trend, but 3rd row is more negative than others - Cl)
* **EXCEPTIONS: groups in Be, N, noble gases (endothermic due to filled shell/half-filled shell - Electronegativity - associated with sharing an electron (shows how close it can pull a shared e- in a covalent bond)
Trend: up and to the right (fluorine)
Q: which is most polar? look for most electronegative atom
Covalent bond
non metal with non metal (covalent network solids - diamond) or molecular compounds (low mp/bp)
Ionic bond
metal with non metal (crystalline, high mp/bp)
Metallic
metal with metal
- conductive (electric + thermal)
- ductile (draw them into wires)
- malleable (pound them to sheets)
- Luster (shiney)
Lewis base
ligand, anything with lone pair of e-
Lewis acid
usually a metal ion
Octet Rule
Exceptions:
- 2H, 4Be, 6B, 6Al
- 3rd row and lower can go over octet rule
- odd # e- overall (i.e. NO)
Formal Charge
Normal # of valence electrons - (each line + dots)
lowest formal charge = preferred species
Electron Domain = 2
sp, 180 degrees, linear, 0 nonbonding e, linear
Electron Domain = 3
sp2, 120 degrees, trigonal planar, 1 nonbonding e, trigonal planar
1 non-bonding e- = bent
Electron Domain = 4
sp3, 109.5 degrees, tetrahedral, 0 nonbonding e- pair, tetrahedral
1 non-bonding e- = trigonal bipyramidal
2 non-bonding e- = bent
Electron Domain = 5
sp3d, 90/120 degrees, trigonal bipyramidal, 0 nonbonding e- pair, trigonal bipyramidal
1 nonbonding e- pair = see-saw
2 nonbonding e- pair = T-shaped
3 nonbonding e- pair = linear
Electron Domain = 6
sp3d2, 90 degrees, octahedral, 0 nonbonding e- pair, octahedral
1 nonbonding e- pair = square pyramidal
2 nonbonding e- pair = square planar
sigma
single bonds
pi bond
double bond (has 1 sigma and 1 pi bonds) triple bond (has 1 sigma and 2 pi bonds)
Intermolecular forces - Hydrogen bonds
strongest over dipole-dipole and london, needs F-H, O-H, or N-H “FON”
Intermolecular forces - London Dispersion Forces
Van der Waals - all molecules have them, larger ones have MORE, nonpolar molecules have the most
intermolecular forces - dipole-dipole
partial + and partial - weak attractive forces (i.e. HCl with HCl) between polar molecules
intermolecular forces - ion-dipole
stronger than dipole-dipole, similar to hydrogen bond, occurs between something that is completely charged to a partially charged atom
i.e. Na+ in water
Increase in intermolecular forces
increase MP, BP, viscosity, surface tension
decrease vapor pressure (the measure of a liquids vapor that is in equilibrium above it)
solid to liquid
fusion “melting”; delta H>0; delta S>0 (more random), endothermic
liquid to gas
vaporization “boiling”; delta H>0; delta S>0 (more random), endothermic
solid to gas
sublimation (i.e. dry ice); delta H>0; delta S>0 (more random), endothermic
gas to liquid
condensation; delta H
liquid to solid
crystallization; delta H
gas to solid
deposition; delta H
Calorimetry
q = mc(delta)T
m= mass in grams
c = specific heat of substance
delta T = change in temperature (kelvin or C)
**During phase changes, cant use mcdeltaT, must use ndeltaH (n = mols)
Triple Point
solid, liquid, gas phases are in equilibrium
critical point
when beyond this point, there is no phase change between liquid and gas (substance is in “super critical fluid” state)
boiling point
when vapor pressure equals the atmospheric pressure pushing down on it
normal atmospheric pressure
1 atm = 760 torr
phase diagram of water
different because solid H20 is less dense than liquid H20 at 1 atm
Phase diagram of CO2
At 1 atm, there is no liquid phase– must increase pressure to get liquid; CO2 will continue to sublime at 1 atm (Solid to gas)
Ideal Gas Assumptions
- gas molecules have no volume (good at low pressures)
- no intermolecular attractive forces (all collisions are elastic) good at high temperatures because gas molecules are moving so fast there is no time for feel other forces of H-bonds, dipole-dipole, or london
STP for ideal gas
1 atm, 273 K, 22.4L for 1 mol gas
one mol of any gas takes up the same volume
Ideal Gas Law
PV = nRT
P= pressure (atm) V = volume (L) n = mol R = given constant (0.08206 for gases) T = Kelvin
Boyle’s Law
V proportional to 1/P
Charles’ Law
V proportional to T
increase T = increase Volume = decrease in density (hence a balloon rises)
Avogadro’s Law
V proportional to mols
Combined Gas Law
P1V1/n1T1 = P2V2/n2T2
use this when given two parts
understand proportions of the Laws’s – will help deduce answer without calculator
gas density
d = mass/volume
two main equations for molar mass
d = m/v and molar mass = m/n (g/mol)
Dalton’s Law of Partial Pressures
Pa + Pb + Pc … = Ptotal
Pa = (Chi) x Ptotal
(Chi) = mol fraction
Graham’s Law of Effusion
when gas escapes through a narrow opening
rate2/rate1 = sqrt(molar mass1/molar mass2)
**simply at the same temp, the lighter gas will move at a higher velocity and escape faster
saturated
max amt of dissolved solute in solvent
unsaturated
not the max amt dissolved solute in solvent
supersaturated
more than max solute dissolved by heating solvent
Molarity
mol of solute/Liter of soln
Molality
mol of solute/Kg of solvent
mol fraction
Mol(a)/Mol(total)
Solubility Rules (MEMORIZE)
- ALL group 1 metal salts, NO3-, NH4+, ClO4-, C2H3O2- (acetate) salts are SOLUBLE in water
- Most Ag+, Pb2+, Hg2 (2+) salts are INSOLUBLE in water unless it is with any salts in RULE #1
Phase solubility
- solids are more soluble at higher temp
- gases are less soluble at higher temp
- gases are more soluble in a liquid at higher pressures
(i. e. if we compressed air into a beaker, it pushes the molecules down into the liquid that were originally in the air)
Colligative Properties: Freezing point depression
deltaT(f) = -i Kf m
i = van't Hoff # (how many ions can a molecule dissociate into) Kf = constant (degree C/molality) m = molality Tf = change in freezing temp
i.e. NaCl = Na+ and Cl- = 2 = i
Na2SO4 = 3 = i
Colligative Properties: Boiling point elevation
delta Tb = i Kb m
i = van't Hoff # (how many ions can a molecule dissociate into) Kb = constant (degree C/molality) m = molality Tb = change in boiling temp
i.e. NaCl = Na+ and Cl- = 2 = i
Na2SO4 = 3 = i
Colligative Properties: Vapor Pressure Depression
Rauolt’s Law - basically pure water vapor pressure is 22 mm Hg, the vapor pressure for (8 mol H2O and 2 mol CH3OH) is just 8/10 mol ratio times the pure water vapor pressure 22 mm Hg)
Colligative Properties: Osmosis
Osmostic pressure (pi) = iMRT
i = van't Hoff factor M = molarity mol/volume (L) R = constant T = temp (K)
