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CHM 201 > ACS study > Flashcards

ACS study Flashcards

1
Q

mega

A

10^6

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2
Q

kilo

A

10^3

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3
Q

centi

A

10^-2

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4
Q

milli

A

10^-3

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5
Q

micro

A

10^-6

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6
Q

nano

A

10^-9

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7
Q

pico

A

10^-12

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8
Q

isotopes

A

same # of protons and differing # of neutrons

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9
Q

equation for average atomic mass

A

[(mass isotope 1) * % isotope 1/100%] + [(mass isotope 2) * % isotope 2/100%]

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10
Q

7 diatomic molecular elements

A

Br2, O2, Cl2, I2, F2, H2, N2

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11
Q

Energy equations

A

E = hv, v = c/λ, E = hc/λ

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12
Q

main group elements

A

1-2, 13-18

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13
Q

group 1

A

alkali metals

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14
Q

group 2

A

alkaline earth metals

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15
Q

group 3

A

rare earth elements, 89 is first actinoid

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16
Q

actinoids

A

89-102

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17
Q

groups 4-12

A

transition metals

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18
Q

groups 13-16

A

metals/non-metals (zintl line)

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19
Q

group 17

A

halogen

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20
Q

group 18

A

noble gases

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21
Q

strong electrolytes

A

high conc. of ions in solution (strong acids/bases, soluble salts)

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22
Q

weak electrolytes

A

low con. of ions in solution (weak acids/bases)

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23
Q

non-electrolytes

A

don’t form ions (insoluble salts, all non-metals without a charge)

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24
Q

7 common strong acids

A

HBr, HCl, HI, HClO3, HClO4, HNO3, H2SO4

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25
Q

strong bases (all alkaline metals)

A

LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)2, Sr(OH)2, Ba(OH)2

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26
Q

weak bases

A

NH3 (ammonia)

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27
Q

q

A

thermal energy transferred

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28
Q

w

A

work involved

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29
Q

equation for energy

A

ΔE = q + w

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30
Q

specific heat

A

resistance to temp change. Higher changes temp more slowly, will absorb more energy to do so.

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31
Q

ΔE

A

change in internal energy of a system

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32
Q

thermal energy transferred equation

A

q = cmΔT

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33
Q

c

A

specific heat (J/g°C) = amount of heath necessary to raise one gram of substance by one degree

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34
Q

m

A

mass

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35
Q

solubility rule 1

A

all alkali compounds and NH4 are soluble

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36
Q

solubility rule 2

A

all nitrates (NO3-1), perchlorates (ClO4-1), chlorates (ClO3-1), and acetates (C2H3O2-1) are soluble

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37
Q

solubility rule 3

A

bromides (Br-1), chlorides (Cl-), and iodides (I-) are soluble except Ag+, Hg2+, and Pb2+

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38
Q

solubility rule 4

A

fluorides (F-) are soluble except for Mg2+, Ca2+, Sr2+, Ba2+, and Pb2+

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39
Q

solubility rule 5

A

sulfates (SO42-) are soluble except Sr2+, Ba2+, Pb2+, Cu2+, and Ag+

40
Q

solubility rule 6

A

carbonates (CO32-), phosphates (PO43-), oxalates (C2O42-), and chromates (CrO42-) are insoluble except rule #1

41
Q

solubility rule 7

A

sulfides (S2-) are insoluble except rule #1 and alkaline earth metals

42
Q

solubility rule 8

A

hydroxides (OH-) are insoluble except rule #1 and Ca2+, Sr2+, and Ba2+

43
Q

precipitation rxn

A

two aqueous reactants form at least one solid product (anions and cations trade partners)

44
Q

neutralization reaction

A

when acids and bases react, the net ionic equation is H+(aq) + OH-(aq) –> H2O(l)

45
Q

acid-base rxn

A

acid + base –> salt + water

46
Q

oxidation states rule 1

A

where two rules contradict, follow the highest rule

47
Q

oxidation states rule 2

A

ox. state of an atom in the pure state (uncombined w/ anything else) is zero

48
Q

oxidation states rule 3

A

total ox. states of all atoms is zero or the charge of the molecule

49
Q

oxidation states rule 4

A

alkali metals = +1, alkali earth metals = +2

50
Q

oxidation states rule 5

A

H = +1, F = -1

51
Q

oxidation states rule 6

A

O = -2

52
Q

oxidation states rule 7

A

in binary compounds, halogens = -1, oxygen family = -2, nitrogen family = -3

53
Q

oxidation states shorthand

A

I = +1, II = +2, H = +1, F = -1, O = -2

54
Q

redox reaction:

A

something changing oxidation state

55
Q

acidic rules 1

A

break into half RXNs

56
Q

acidic rules 2

A

balance non H’s and O’s

57
Q

acidic rules 3

A

balance O’s by adding H2O

58
Q

acidic rules 4

A

balance H’s by adding H+

59
Q

acidic rules 5

A

balance charge by adding e’s

60
Q

acidic rules 6

A

recombine half reactions if the e’s cancel

61
Q

basic rules

A

use acidic rules, then add equal amount of OH-s as H+s to each side, forming H2O.

62
Q

solution dilution equation

A

M1V1 = M2V2

63
Q

pH equations

A

pH +pOH = 14, pOH = -log[OH-], conc H+ = -log[H3O+], [H3O+] = 10^-pH, [OH-] = 10^-pOH

64
Q

pH < 7

A

basic/alkaline

65
Q

pH > 7

A

acidic

66
Q

rxns of metal oxides + water (ox state does not change)

A

makes a base. e.g. MgO(s) + H2O –> Mg(OH)2(aq)

67
Q

rxns of non-metal oxides + water (ox state does not change)

A

makes an acid. e.g. CO2(g) + H2O(l) –> H2CO3(aq)

68
Q

combination rxn

A

more than one reactant and only one product

69
Q

decomposition rxn

A

a single reactant and more than one product

70
Q

displacement rxn

A

one element displaces another (e.g. Zn(s) + CuSO4(aq) –> Cu(s) + ZnSO4(aq)

71
Q

metathesis rxn

A

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
Q

zeroth law of thermodynamics

A

if A⇌B and B⇌C, then A⇌C

73
Q

first law of thermodynamics

A

Δu = uf-ui = q+w

74
Q

u

A

internal energy

75
Q

Δ

A

state function - independent of path

76
Q

thermo equations

A

q = mcΔT (heat will be in joules), w = -PoppΔv

77
Q

C (capital)

A

heat capacity - amount of heat necessary to raise the temp of a whole container by one degree (J/°C)Δ

78
Q

Δv

A

vf - vi

79
Q

H (enthalpy)

A

heat absorbed or released in an open atmosphere. Cannot be found directly, you can only find ΔH.

80
Q

H equation

A

H = U + PV

81
Q

ΔH

A

ΔH = qp (q at constant pressure)

82
Q

Δu

A

Δu = qv (q at a constant volume)

83
Q

if ΔH is positive,

A

heat is absorbed (endothermic)

84
Q

if ΔH is negative,

A

heat is released (exothermic)

85
Q

standard enthalpy of formation

A

amount of heat absorbed or released when 1 mol of product is formed from its elements

86
Q

standard states

A

1 atm, 25°C, 1M

87
Q

elements in their most stable form have a ΔH = 0

A
88
Q

combustion rxn

A

add O2 to reactants, products are CO2(g) + H2O(l)

89
Q

stable form of P, S, Cu

A

S8, P4, Cu(s)

90
Q

Hess’ Law

A

A ͢. B, B ͢. C, C ͢. D - net is A ͢. D.

91
Q

lattice energy

A

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
Q

A gas behaves most ideally at

A

low pressures and high temperatures

93
Q

density equations

A

PV=nRT, n=PV/RT

94
Q

combined gas law

A

(P1V1)/T1 = (P2V2)/T2 for absolute temperature!! kelvin

95
Q

25 C = 298 K

A
96
Q

for temp graph, highest molar mass is highest peak of curve

A
97
Q

as intermolecular forces increase, boiling point increases

A

CHM 201 (5 decks)

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