ACS study Flashcards
1
Q
mega
A
10^6
2
Q
kilo
A
10^3
3
Q
centi
A
10^-2
4
Q
milli
A
10^-3
5
Q
micro
A
10^-6
6
Q
nano
A
10^-9
7
Q
pico
A
10^-12
8
Q
isotopes
A
same # of protons and differing # of neutrons
9
Q
equation for average atomic mass
A
[(mass isotope 1) * % isotope 1/100%] + [(mass isotope 2) * % isotope 2/100%]
10
Q
7 diatomic molecular elements
A
Br2, O2, Cl2, I2, F2, H2, N2
11
Q
Energy equations
A
E = hv, v = c/λ, E = hc/λ
12
Q
main group elements
A
1-2, 13-18
13
Q
group 1
A
alkali metals
14
Q
group 2
A
alkaline earth metals
15
Q
group 3
A
rare earth elements, 89 is first actinoid
16
Q
actinoids
A
89-102
17
Q
groups 4-12
A
transition metals
18
Q
groups 13-16
A
metals/non-metals (zintl line)
19
Q
group 17
A
halogen
20
Q
group 18
A
noble gases
21
Q
strong electrolytes
A
high conc. of ions in solution (strong acids/bases, soluble salts)
22
Q
weak electrolytes
A
low con. of ions in solution (weak acids/bases)
23
Q
non-electrolytes
A
don’t form ions (insoluble salts, all non-metals without a charge)
24
Q
7 common strong acids
A
HBr, HCl, HI, HClO3, HClO4, HNO3, H2SO4
25
strong bases (all alkaline metals)
LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)2, Sr(OH)2, Ba(OH)2
26
weak bases
NH3 (ammonia)
27
q
thermal energy transferred
28
w
work involved
29
equation for energy
ΔE = q + w
30
specific heat
resistance to temp change. Higher changes temp more slowly, will absorb more energy to do so.
31
ΔE
change in internal energy of a system
32
thermal energy transferred equation
q = cmΔT
33
c
specific heat (J/g°C) = amount of heath necessary to raise one gram of substance by one degree
34
m
mass
35
solubility rule 1
all alkali compounds and NH4 are soluble
36
solubility rule 2
all nitrates (NO3-1), perchlorates (ClO4-1), chlorates (ClO3-1), and acetates (C2H3O2-1) are soluble
37
solubility rule 3
bromides (Br-1), chlorides (Cl-), and iodides (I-) are soluble except Ag+, Hg2+, and Pb2+
38
solubility rule 4
fluorides (F-) are soluble except for Mg2+, Ca2+, Sr2+, Ba2+, and Pb2+
39
solubility rule 5
sulfates (SO42-) are soluble except Sr2+, Ba2+, Pb2+, Cu2+, and Ag+
40
solubility rule 6
carbonates (CO32-), phosphates (PO43-), oxalates (C2O42-), and chromates (CrO42-) are insoluble except rule #1
41
solubility rule 7
sulfides (S2-) are insoluble except rule #1 and alkaline earth metals
42
solubility rule 8
hydroxides (OH-) are insoluble except rule #1 and Ca2+, Sr2+, and Ba2+
43
precipitation rxn
two aqueous reactants form at least one solid product (anions and cations trade partners)
44
neutralization reaction
when acids and bases react, the net ionic equation is H+(aq) + OH-(aq) --> H2O(l)
45
acid-base rxn
acid + base --> salt + water
46
oxidation states rule 1
where two rules contradict, follow the highest rule
47
oxidation states rule 2
ox. state of an atom in the pure state (uncombined w/ anything else) is zero
48
oxidation states rule 3
total ox. states of all atoms is zero or the charge of the molecule
49
oxidation states rule 4
alkali metals = +1, alkali earth metals = +2
50
oxidation states rule 5
H = +1, F = -1
51
oxidation states rule 6
O = -2
52
oxidation states rule 7
in binary compounds, halogens = -1, oxygen family = -2, nitrogen family = -3
53
oxidation states shorthand
I = +1, II = +2, H = +1, F = -1, O = -2
54
redox reaction:
something changing oxidation state
55
acidic rules 1
break into half RXNs
56
acidic rules 2
balance non H's and O's
57
acidic rules 3
balance O's by adding H2O
58
acidic rules 4
balance H's by adding H+
59
acidic rules 5
balance charge by adding e's
60
acidic rules 6
recombine half reactions if the e's cancel
61
basic rules
use acidic rules, then add equal amount of OH-s as H+s to each side, forming H2O.
62
solution dilution equation
M1V1 = M2V2
63
pH equations
pH +pOH = 14, pOH = -log[OH-], conc H+ = -log[H3O+], [H3O+] = 10^-pH, [OH-] = 10^-pOH
64
pH < 7
basic/alkaline
65
pH > 7
acidic
66
rxns of metal oxides + water (ox state does not change)
makes a base. e.g. MgO(s) + H2O --> Mg(OH)2(aq)
67
rxns of non-metal oxides + water (ox state does not change)
makes an acid. e.g. CO2(g) + H2O(l) --> H2CO3(aq)
68
combination rxn
more than one reactant and only one product
69
decomposition rxn
a single reactant and more than one product
70
displacement rxn
one element displaces another (e.g. Zn(s) + CuSO4(aq) --> Cu(s) + ZnSO4(aq)
71
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)
72
zeroth law of thermodynamics
if A⇌B and B⇌C, then A⇌C
73
first law of thermodynamics
Δu = uf-ui = q+w
74
u
internal energy
75
Δ
state function - independent of path
76
thermo equations
q = mcΔT (heat will be in joules), w = -PoppΔv
77
C (capital)
heat capacity - amount of heat necessary to raise the temp of a whole container by one degree (J/°C)Δ
78
Δv
vf - vi
79
H (enthalpy)
heat absorbed or released in an open atmosphere. Cannot be found directly, you can only find ΔH.
80
H equation
H = U + PV
81
ΔH
ΔH = qp (q at constant pressure)
82
Δu
Δu = qv (q at a constant volume)
83
if ΔH is positive,
heat is absorbed (endothermic)
84
if ΔH is negative,
heat is released (exothermic)
85
standard enthalpy of formation
amount of heat absorbed or released when 1 mol of product is formed from its elements
86
standard states
1 atm, 25°C, 1M
87
elements in their most stable form have a ΔH = 0
88
combustion rxn
add O2 to reactants, products are CO2(g) + H2O(l)
89
stable form of P, S, Cu
S8, P4, Cu(s)
90
Hess' Law
A ͢. B, B ͢. C, C ͢. D - net is A ͢. D.
91
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.
92
A gas behaves most ideally at
low pressures and high temperatures
93
density equations
PV=nRT, n=PV/RT
94
combined gas law
(P1V1)/T1 = (P2V2)/T2 for absolute temperature!! kelvin
95
25 C = 298 K
96
for temp graph, highest molar mass is highest peak of curve
97
as intermolecular forces increase, boiling point increases
