general chem Flashcards
1
Q
charge of proton
A
1.6 x 10^-19 C
2
Q
atomic number
A
of protons
3
Q
mass number
A
sum of protons + neutrons
4
Q
13
C, which number is which
6
A
13 is mass number
6 is the atomic number
5
Q
atomic mass = ________ number
A
mass
6
Q
which is reported on periodic table? atomic mass or atomic weight
A
weight
7
Q
1 proton = ___ amu = _________
A
one
amu = 1.66 x 10^-24 g
8
Q
elements with varying mass numbers
A
isotopoes
9
Q
weighted average of all isotopes
A
atomic weight
10
Q
Avogadros number
A
6.06 x 10^23
11
Q
Planck relation
A
E = hf
E = energy of quantum h = 6.626 x 10^-34 Js f = frequency
12
Q
angular momentum of orbiting hydrogen equation
A
L = nh / 2π
n = principal quantum # h = 6.626 x 10^-34 Js
13
Q
Bohr’s angular momentum equations says electric momentum changes
A
only in discrete amounts, with respect of principal quantum #
14
Q
energy of electron equation
A
E = - 2.18 x 10^-18 (J/e-) / n^2
15
Q
energy of an electron ______ (becomes _____ negative) as it moves away from nucleus
A
increases, less negative
16
Q
EM energy of photons
A
E = hc / λ
c = 3.00 x 10^8 m/s h = 6.626 x 10^-34 Js
17
Q
hydrogen emission lines if you go from any n down to 1
A
Lyman
18
Q
hydrogen emission lines if you go from n= 3 or more down to 2
A
Balmer
19
Q
hydrogen emission lines if you go from n=4 or more down to 3
A
Paschen
20
Q
larger energy transitions = ______ wavelength
A
shorter
21
Q
Bohr did not take into account
A
atoms with more than one electron
22
Q
heisenberg uncertainty principle
A
it is impossible to determine the momentum and position of an electron simultaneously
if you wanna know position, it has to stop.
23
Q
pauli exclusion principle
A
no two electrons in an atom can have the same set of quantum numbers (cant be in the same position/energy)
24
Q
principal quantum number (n)
A
any positive #, gives the shell number
25
max # of electrons within a shell
= 2n^2
26
azimuthal quantum number (l)
```
0 to (n-1)
shape and # of sub shells within a shell
```
27
spectroscopic notation of azimuthal quantum numbers
```
Sam Please Dont Fail
l = 0 = s
l = 1 = p
l = 2 = d
l = 3 = f
```
28
max # of electrons within a subshell
4l + 2
29
magnetic quantum number (ml)
-l to +l
| particular orbital within a subshell the electron may be found
30
shapes of orbitals
s = spherical
p = dumbbell (x, y, z)
d and f don't need to know
31
max # of electrons
2n^2
32
max # of electrons in orbital
2
33
spin quantum number (ms)
+1/2 or -1/2
34
2p^4, what does this mean
```
n = 2
l = 1
# of electrons in subshell = 4
```
35
n + l rule
when asked which will fill first, do n + l
```
5d = 5 + 2 = 7
6s = 6 + 0 = 6
```
36
when doing electron configuration for cation, what do you do?
start with neutral atom
remove from highest n value
if tied, remove from highest l value
37
Hund's rule
orbitals are filled up before adding a second electron
38
two common exceptions to electron configuration
chromium and copper
39
paramagnetic
have unpaired electrons, will orient spins to be weakly attracted to magnetic field
40
diamagnetic
have paired electrons, repel magnetic field
41
effective nuclear charge
pull between valence e- and nucleus
42
effective charge trend
increases to right
| constant in group
43
more inner shells causes effective charge to
become constant, even though there are more protons, the separation is canceled by inner shells
44
inner shells =
principal quantum number
45
atomic radius trend
decrease to right
| atomic radius increases down
46
alkali _____ dense than other metals
less
47
incomplete octet
hydrogen (2), helium (2), lithium (2)
| beryllium (4) and boron (6)
48
expanded octet
period 3 or more
| ex: phosphorous, sulfur, chlorine
49
bond between opposite charges
ionic bond
50
what shape do ionic bonds form
lattice structures
51
covalent bonding
shared e- between two similar charged ions (usually nonmetals)
52
non polar bond
electron in covalent bond shared equally
53
polar bond
electron in covalent bond shared uneqally
54
coordinate covalent
only one ion giving both electrons in covalent bond
55
bond structure of covalent bonds
individually bonded molecules
56
properties of ionic bonds (4)
1. high melting and boiling point
2. dissolve in polar solvents
3. good conductors when molten
4. form crystalline lattice structures to minimize repulsions
57
properties of covalent bonds (2)
1. low melting and boiling points
| 2. poor conductors of electricity
58
EN difference < 0.5 =
non polar bond
59
to be polar covalent, what is difference in EN
0.5 - 1.7
60
dipole moment equation
p = qd
q = charge, d = displacement
61
formal charge
Valence electrons - dots - sticks
62
formal change _______ EN
underestimates, you assume electrons are shared
63
oxidation number ________ EN
overestimates, you give all e- to more EN atom
64
electronic geometry
all bonds
65
molecular geometry
only bonding (no LP)
66
parallel orbits =
pi bond
67
head to head orbitals =
sigma bond
68
three types of intermolecular forces
1. LDF
2. dipole dipole
3. hydrogen
69
dispersion forces caused by
dipoles changing and shifting fast
70
when are DD interactions absent
gases, not close enough
71
what causes DD interactions
polar molecules aligning to opposite charge of a nearby molecule
72
when is H able to form hydrogen bonds
when bound to N, O, F
73
what does H attract in hydrogen bonds
negative charge on N O F
74
equivalent weight
a mass that provides one mole of whatever you need
= molar mass / n (n= # particles of interest)
75
if they give you an amount produced, what equation do you use to find the equivalents
mass of compound / gram equivalent weight (mass collected)
76
molarity =
normality / n (n=# of thing produced)
77
Normality units
equivalents / L
78
law of constant composition
any compound will contain the same ratio of elements
ex: H2O will always have 2x H than O
79
empirical formula
simplest whole # ratio of elements
80
molecular formula
exact # of atoms for each element
81
percent composition
% of a compound made up by a given element
82
percent composition equation
mass of element / mass of compound x 100
83
A + B -->
compound reaction
84
C --> A + B
decomposition reaction
85
CH4 + 2O2 -> CO2 + 2H201
combustion reaction
86
what do you need for a combustion reaction (2)
1. a fuel (hydrocarbon)
| 2. an oxidant (oxygen)
87
single-displacement reaction
atom or ion is replaced by an atom or ion of another element
88
double-displacement reaction
elements from two different compounds swap places
89
acid + base -> salt + H20
neutralization reaction
90
theoretical yield
amount of product from the balanced equation
91
actual yield
amount of product you actually get
92
percent yield
actual / theoretical x 100
93
endings for oxyanions
```
ite = less
ate = more
```
94
electrolytes
solutes that enable solutions to carry currents
95
collision theory
rate of reaction = # of CORRENT collisions per second
96
according to collision theory, rate =
Z x f
```
Z = collisions
f = fraction of effective collisions
```
97
Arrhenius equation
k = Ae ^ (-Ea / RT)
```
A = frequency factor (s-1), amount of collisions
Ea = activation energy
```
98
the exponent in the Arrhenius equation
if it becomes a smaller number, it becomes more positive
| so, rate INCREASES
99
transition state
higher energy than reactants and products, cannot be isolated/theoretical
100
energy required to meet transition state =
activation energy
101
for most forward, irreversible equations, rate =
k{A}^x[B]^y
k = rate constant
x / y = orders of reaction
102
overall order of a reaction =
x + y
103
orders of reaction must be determined
experimentally, NOT STOICHIOMETRIC
104
zero-order reaction
formation of C does not depend on A or B
only way to change rate is by increasing K by changing temperature or adding catalyst
105
zero-order reaction graph
linear, slope opposite of k
106
first-order reaction
rate depends only on changing one reactant
107
first-order graph
nonlinear, depends on reactant concentration
on ln[A] vs. time, it is linear with a slope opposite of K
108
second-order reaction
rate depends on squaring one reactant or on both reactants
109
second-order on 1 /[A] vs time graph
linear, slope equals k
110
mixed-order reactions
rate orders that vary over time
111
broken-order
rate order is a fraction
112
equlibrium will have max _____ and min _____
max entropy
| min Gibbs free energy
113
in equilibrium, exponents are =
to K
114
Q < K
equation not at equilibrium yet, more reactants
115
Q =. K
at equilibrium
116
Q > K
forward reaction has exceeded equlibirum, go back
117
If Keq is a negative, small value
[A] = 1
118
if you decrease pressure,
you move towards the side with less moles of gas
119
if you increase pressure
you move towards side where there is less moles of gas
120
if you increase volume,
you move towards the greatest number of moles of gas
121
kinetic pathway requires ____ energy to reach transition state but has ______ energy product
```
less energy
higher energy (unstable)
```
122
thermodynamic pathway has a _____ energy transition state but produces a ______ energy product
```
higher
lower energy (more stable)
```
123
fast products
kinetic, form faster
124
closed vs open system
closed cant exchange matter, open can
125
first law of thermodynamics
ΔU = Q - W
```
Q = heat added
W = work done by system
ΔU = change in internal energy of system
```
126
isothermal process
temperature = ΔU = constant = 0
Q = W
area under PV curve = Q and W
127
adiabatic process
no heat is exchanged
Q = O
Δ U = -W
work done ON system
128
isobaric process
pressure is common
PV curve has a flat line
129
isovolumetric process
volume constant, no work occurs
```
ΔU = Q
W = 0
```
130
STP
```
T = 273 K
P = 1 atm
```
131
standard conditions
```
T = 298 K
P = 1 atm
M = 1 M
```
132
sublimation
solid to gas
133
deposition
gas to solid
134
critical point
temperature and pressure when there is no distinction between phases, point to the far right (gas and liquid)
135
heat of vaporization past the critical point =
0
136
zeroth law of equilibrium
objects in thermal equilibrium when their temperatures are equal
137
state functions
TV HUGS
temp, volume
enthalpy, internal energy, Gibbs free, entropy
independent of path taken
138
process functions
Q and W
139
enthalpy = Q under
constant pressure
140
q =
mcΔT
141
heat capacity
mass x specific heat
142
when can we not use q = mcΔT
during a phase change
143
latent heat
q = mL
L = enthalpy of isothermal process
144
entropy equation
ΔS = Qrev / T
145
density
mass / V
146
at STP, 1 mole of gas = ____ L
22.4
147
Avogadros principle
as n increases, V increases
148
Boyle's law
as P increases, V decreases
149
Charle's Law
as V increases, T increases
150
Gay-Lussacs law
as P increases, T increases
151
Boyle
PV
152
Charles
TV
153
Gay Lussacs
TP
154
Dalton's Law
pressure of container = sum of all partial pressures
155
mole fraaction
moles of gas A / total moles of gas
156
partial pressure equation
PA = (Xa)(PT)
157
vapor pressure
pressure exerted by evaporated participles above a liquis
158
solubility of gas increases with
partial pressure
159
[A] = Kh x PA
concentration dissolved = Henry's constant x partial pressure
160
a in real gases will be
small for small/less polarizable gases
large for big/polarizable gases
it corrects for attraction
161
b in real gases will be
smaller in general, corrects for volume
162
molar solubility > ___ will dissolve
0.1 M
163
two main solubility rules
1. all groups I metals are soluble
| 2. all nitrate salts are soluble (NO3-)
164
chelation
central cation bound to multiple places on the ligand (electron donor)
165
solubility constant ______ with temperature for gas
decreases
166
same molecular formula, different arrangement of bonds
structural isomer
167
same molecular formula, same connectivity of bonds
stereoisomers, differ in space
168
configurational isomers
can only change by breaking bonds
169
conformational isomers
differ in rotation around a single bond
170
which isomers are most similar
conformational
171
which isomers are most different
structural
172
staggered anti
two bigger groups are antiperiplanar (same plane, opposite sides)
173
gauge staggered
opposite of each other, but 60 degrees (not 180) apart
174
highest-energy state of Newman straight chain
total eclipse
175
totally eclipsed
two methyls overlap eachother
176
eclipsed
two methyls are 120 apart and overlap with hydrogens
