ACS study Flashcards
mega
10^6
kilo
10^3
centi
10^-2
milli
10^-3
micro
10^-6
nano
10^-9
pico
10^-12
isotopes
same # of protons and differing # of neutrons
equation for average atomic mass
[(mass isotope 1) * % isotope 1/100%] + [(mass isotope 2) * % isotope 2/100%]
7 diatomic molecular elements
Br2, O2, Cl2, I2, F2, H2, N2
Energy equations
E = hv, v = c/λ, E = hc/λ
main group elements
1-2, 13-18
group 1
alkali metals
group 2
alkaline earth metals
group 3
rare earth elements, 89 is first actinoid
actinoids
89-102
groups 4-12
transition metals
groups 13-16
metals/non-metals (zintl line)
group 17
halogen
group 18
noble gases
strong electrolytes
high conc. of ions in solution (strong acids/bases, soluble salts)
weak electrolytes
low con. of ions in solution (weak acids/bases)
non-electrolytes
don’t form ions (insoluble salts, all non-metals without a charge)
7 common strong acids
HBr, HCl, HI, HClO3, HClO4, HNO3, H2SO4
strong bases (all alkaline metals)
LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)2, Sr(OH)2, Ba(OH)2
weak bases
NH3 (ammonia)
q
thermal energy transferred
w
work involved
equation for energy
ΔE = q + w
specific heat
resistance to temp change. Higher changes temp more slowly, will absorb more energy to do so.
ΔE
change in internal energy of a system
thermal energy transferred equation
q = cmΔT
c
specific heat (J/g°C) = amount of heath necessary to raise one gram of substance by one degree
m
mass
solubility rule 1
all alkali compounds and NH4 are soluble
solubility rule 2
all nitrates (NO3-1), perchlorates (ClO4-1), chlorates (ClO3-1), and acetates (C2H3O2-1) are soluble
solubility rule 3
bromides (Br-1), chlorides (Cl-), and iodides (I-) are soluble except Ag+, Hg2+, and Pb2+
solubility rule 4
fluorides (F-) are soluble except for Mg2+, Ca2+, Sr2+, Ba2+, and Pb2+
solubility rule 5
sulfates (SO42-) are soluble except Sr2+, Ba2+, Pb2+, Cu2+, and Ag+
solubility rule 6
carbonates (CO32-), phosphates (PO43-), oxalates (C2O42-), and chromates (CrO42-) are insoluble except rule #1
solubility rule 7
sulfides (S2-) are insoluble except rule #1 and alkaline earth metals
solubility rule 8
hydroxides (OH-) are insoluble except rule #1 and Ca2+, Sr2+, and Ba2+
precipitation rxn
two aqueous reactants form at least one solid product (anions and cations trade partners)
neutralization reaction
when acids and bases react, the net ionic equation is H+(aq) + OH-(aq) –> H2O(l)
acid-base rxn
acid + base –> salt + water
oxidation states rule 1
where two rules contradict, follow the highest rule
oxidation states rule 2
ox. state of an atom in the pure state (uncombined w/ anything else) is zero
oxidation states rule 3
total ox. states of all atoms is zero or the charge of the molecule
oxidation states rule 4
alkali metals = +1, alkali earth metals = +2
oxidation states rule 5
H = +1, F = -1
oxidation states rule 6
O = -2
oxidation states rule 7
in binary compounds, halogens = -1, oxygen family = -2, nitrogen family = -3
oxidation states shorthand
I = +1, II = +2, H = +1, F = -1, O = -2
redox reaction:
something changing oxidation state
acidic rules 1
break into half RXNs
acidic rules 2
balance non H’s and O’s
acidic rules 3
balance O’s by adding H2O
acidic rules 4
balance H’s by adding H+
acidic rules 5
balance charge by adding e’s
acidic rules 6
recombine half reactions if the e’s cancel
basic rules
use acidic rules, then add equal amount of OH-s as H+s to each side, forming H2O.
solution dilution equation
M1V1 = M2V2
pH equations
pH +pOH = 14, pOH = -log[OH-], conc H+ = -log[H3O+], [H3O+] = 10^-pH, [OH-] = 10^-pOH
pH < 7
basic/alkaline
pH > 7
acidic
rxns of metal oxides + water (ox state does not change)
makes a base. e.g. MgO(s) + H2O –> Mg(OH)2(aq)
rxns of non-metal oxides + water (ox state does not change)
makes an acid. e.g. CO2(g) + H2O(l) –> H2CO3(aq)
combination rxn
more than one reactant and only one product
decomposition rxn
a single reactant and more than one product
displacement rxn
one element displaces another (e.g. Zn(s) + CuSO4(aq) –> Cu(s) + ZnSO4(aq)
metathesis rxn
cations/anions change partners without changing oxidation states. a) acid base RXNs, b) precipitation RXNs, c) gas formation (e.g. 2HCL(aq) + CaCO3(s) –> CO2(g) + H2O(l) + CaCl2(aq)
zeroth law of thermodynamics
if A⇌B and B⇌C, then A⇌C
first law of thermodynamics
Δu = uf-ui = q+w
u
internal energy
Δ
state function - independent of path
thermo equations
q = mcΔT (heat will be in joules), w = -PoppΔv
C (capital)
heat capacity - amount of heat necessary to raise the temp of a whole container by one degree (J/°C)Δ
Δv
vf - vi
H (enthalpy)
heat absorbed or released in an open atmosphere. Cannot be found directly, you can only find ΔH.
H equation
H = U + PV
ΔH
ΔH = qp (q at constant pressure)
Δu
Δu = qv (q at a constant volume)
if ΔH is positive,
heat is absorbed (endothermic)
if ΔH is negative,
heat is released (exothermic)
standard enthalpy of formation
amount of heat absorbed or released when 1 mol of product is formed from its elements
standard states
1 atm, 25°C, 1M
elements in their most stable form have a ΔH = 0
combustion rxn
add O2 to reactants, products are CO2(g) + H2O(l)
stable form of P, S, Cu
S8, P4, Cu(s)
Hess’ Law
A ͢. B, B ͢. C, C ͢. D - net is A ͢. D.
lattice energy
the amount of energy needed to completely separate 1 mole of a solid ionic compound into its gaseous ions. increases as the charges on the ions increase and decreases as the size of ions increases.
A gas behaves most ideally at
low pressures and high temperatures
density equations
PV=nRT, n=PV/RT
combined gas law
(P1V1)/T1 = (P2V2)/T2 for absolute temperature!! kelvin
25 C = 298 K
for temp graph, highest molar mass is highest peak of curve
as intermolecular forces increase, boiling point increases
