Chem I: 1-6 Flashcards
(459 cards)
1
Q
atomic number
A
(Z) number of protons
2
Q
mass number
A
- (A) total number of protons and neutrons in nucleus
- Z+N
- aka atomic mass
3
Q
isotopes
A
atoms that share an atomic number but have different mass numbers (same number of protons, diff number of neurons)
4
Q
valence electrons
A
- electrons that are farthest from nucleus
- electrons that have the greatest amount of potential energy (bc they are held less tightly by the nucleus)
5
Q
cation
A
positively charged aton
6
Q
anion
A
negatively charged atom
7
Q
isotopes are referred to…
A
by the name of the element followed by the mass number
8
Q
atomic weight
A
- weighted average of the naturally occurring isotopes
- mass of one mole of the element in grams
9
Q
binding energy
A
- energy required to disassemble an atom into its constituents
- allows nucleons to bind together in nucleus
10
Q
mass defect
A
- the difference between the mass of an atom and the sum of the masses of its protons, neutrons, and electrons
- result of matter that has been converted to energy
11
Q
number of protons and electrons in neutral atom
A
of protons = # of electrons
12
Q
Einstein’s equation that relates energy and mass (+ units)
A
E = mc^2 converts mass and energy
E in J
m in kg > mass defect
c in m/s
13
Q
E in E = mc^2
A
energy released when the nucleus of an atom is formed (nuclear binding energy) and energy needed to break nucleus apart
14
Q
nuclear binding energy
A
- the energy released when a nucleus is formed from nucleons
- energy required to disassemble an atom into its constituents
15
Q
nucleons
A
protons and neutrons in the nucleus
16
Q
electrostatic force
A
force between electrically charged objects (like charges repel and opposite charges attract each other).
17
Q
nuclear strong force
A
- force that overcomes repulsion of the protons in nucleus to hold the nucleus together
- only acts over extremely short distances
18
Q
N/Z ratio when Z < 20
A
N/Z = 1 –> stable nucleus
19
Q
N/Z ratio when Z > 20
A
N/Z = 1.5 –> stable nucleus
20
Q
N/Z ratio when Z > 83
A
p much all nuclei are unstable
21
Q
alpha decay
A
ejection of an alpha particle (He) from the nucleus of an atom
22
Q
beta decay
A
- ejection of a beta particle (electron) from the nucleus of an atom
- neutron converted to proton and electron
23
Q
gamma decay
A
- release of gamma rays (no charge and no mass) from a nucleus
- energy is released in form of radiation
24
Q
positron emission aka positive beta decay
A
a proton decays and “creates” a neutron
25
what increases the rate of radioactive decay?
1. more protons in nucleus = more unstable = more decay
2. for smaller atoms, the N/Z ratio determines stability (1:1)
26
half life
time required for a certain amount of pure substance to fall to half its original amount
27
what makes atom more likely to undergo alpha decay?
large number of protons in nucleus
28
electron capture
atom captures and electron and combines it with a proton to form a neutron
29
half life equation
t1/2 = 0.693/k
30
ionizing radiation
the release of energy that allows an unstable nucleus to attain a more stable form
31
how does the mass defect relate to the binding energy?
* there is a transformation of nuclear matter to energy with a resultant loss of matter
* mass defect - energy released
* binding energy - energy needed
* they are related by E = mc^2
32
fusion
when small nuclei combine to form a larger nucleus
33
fission
* large nucleus splits into smaller nuclei
* can occur through absorption of low energy neutron --\> making an excited state
34
T/F: Nuclear fission and nuclear fusion both release energy
TRUE
35
Which type of nuclear decay could be detected in an atomic absorption spectrum?
because gamma radiation produces electromagnetic radiation (rather than nuclear fragments), it can be detected on an atomic absorption spectrum.
36
isotope notation
37
avogadro's number
multiplying
6.02 x 10^23 atoms/mol
number of atoms/molecules in a mole
multiplying Avogadro's number by the moles of an element gives the number of atoms of that element.
38
equation relating energy and frequency
E = hf
39
equation
angular momentum of an electrno orbiting a hydrogen nucleus
L = nh / 2π
n = principal quantum number
40
how will protium behave in an electric field as compared with deuterium?
a) protium will accelerate slower than deuterium
b) neither protium nor deuterium will accelerate
c) protium will accelerate faster than deuterium
d) protium and deuterium will both accelerate at the same rate
b) neither protium nor deuterium will accelerate
only charged particles experience force due to an electric field
since they are both neutral, neither will be influenced by electric field
41
Phosphorus has 2 known isotopes that weigh 30 amu (90% abundance) and 31 amu (10% abundance). What is the most likely atomic mass of a single Phosphorus atom?
a) 30.1 amu
b) 31 amu
c) 30 amu
d) 30.9 amu
Atomic mass of a single atom is a discrete whole number based on the number of nucleons found in a single atom. The most abundant isotope is made of 30 nucleons, thus the most likely atomic mass of an individual atom is 30 amu.
42
Nuclei (Z \< 20) with N/Z ratio greater than one can undergo \_\_\_\_\_\_\_\_\_\_.
beta decay
43
Nuclei (Z \< 20) with N/Z ratio less than one can undergo \_\_\_\_\_\_\_\_\_.
positron emission
44
Phosphorus-32 is commonly used to label DNA and RNA for
use in Southern and Northern blots, respectively. A
researcher has a 70 g sample of Phosphorus-32 stored for
future use. If the researcher needs at least 10 g of
Phosphorus-32 to do his experiment, what's the greatest
number of days that he could store this sample? (Phosphorus-
32 has a half-life of 14.3 days)
a) 25 days
b) 45 days
c) 30 days
d) 40 days
d) 40 days
45
A radioactive sample is giving off gamma rays. What is occurring on the atomic scale?
a) a nucleus in an excited state releases a photon, becoming a new element in its ground state
b) a nucleus in an excited state releases a photon, going to its ground state
c) a nucleus in its ground state releases a photon, becoming a new element in its ground state
d) a nucleus in its ground state releases a photon and remains in its ground state
b) a nucleus in an excited state releases a photon, going to its ground state
In gamma decay, nucleus is in unstable high energy state. It releases this energy by emitting a photon, going to more stable ground state.
46
A researcher adds 5g of NaCl to a beaker containing 120 g of water. How many water molecules are present?
a) 4 x 10^24
b) 7.2 x 10^21
c) 6 x 10^23
d) 1.3 x 10^24
a) 4 x 10^24
47
a) emission spectrum with dark lines where hydrogen is emitting light
b) absorption spectrum with dark lines where hydrogen is absorbing light
c) an emission spectrum with bright lines where hydrogen is emitting light
d) an absorption spectrum with bright lines where hydrogen is absorbing light
48
mass to charge ratio
The mass of an ion divided by its charge
symbol m/z
49
purpose of mass spectrometer
determine the mass of a particle using electric and magnetic fields
50
how the mass spectrometer works
* the movement of excited (ionized) particles will be affected as they pass through a magnetic field
* the degree to which these particles are deflected from their original path will depend on their mass/charge
51
large vs small particles in mass spectrometer
* large particles travel with a larger radius and slower acceleration -\> travels longer distance than smaller particle
* smaller particle -\> greater centripetal acceleration --\> smaller radius
52
mass spectra
plot of relative abundance (y plane) vs mass to charge ratio (x plane)
53
Mass Spectrometry: if the charge on the ion is +1, the mass to charge ratio is.....
if the charge on the ion is +1, the mass to charge ratio is equal to the mass of that ion
54
Mass Spectrometry: relative abundance
the intensity of the ion as it collides with ion detector
55
Mass Spectrometry: p+1 peak
the molecule that contains heavier isotopes
56
Mass Spectrometry: fragmentation pattern
lower weight molecules that are fragments the og molecule
57
Mass Spectrometry: base peak
* the highest peak, which may or may not correspond to the parent ion
* the base peak is always made equal to 100% relative abundance
58
photoelectric effect
emission of electrons from a metal when light shines on the metal
59
Photoelectric effect equation
E = hv= hc/lambda
60
frequency related to wavelength equation
c = lambda \* v
lambda = wavelength
61
photoelectric effect: if shining light on an electron does not excite that electron to a new energy state, what will increasing the intensity of light do?
* intensity is the same as amount
* it will have no effect on the electron
62
photoelectric effect: if shining light on an electron does not excite that electron to a new energy state, what will increasing the frequency of light do?
* the light particle has more energy when the frequency increases (deltaE = hf)
* the electron has the possibility of being excited to a higher state
63
Coulomb's Law
* (bohr model) electric force between charged objects depends on the distance between the objects and the magnitude of the charges
* (electrons get neg sign)
64
force equation
F=ma (force = mass x acceleration)
65
the bohr model predicts...
the wavelength of light emitted by the hydrogen atom
66
binding energy per nucleon peaks at the element \_\_\_, which implies that...
iron
iron contains the most stable nucleus
67
weak nuclear force
* also contributes to stability of nucleus
* much weaker than strong nuclear force
68
4 fundamental forces of nature
strong nuclear force, weak nuclear force, gravitation, electrostatic forces
69
electrons are emitted by the nucleus when...
a nutron decays into a proton (a beta particle)
70
exponential decay (half life)
71
exponential decay equation
72
bohr model rules
1. electrons move about nucleus in circular orbit and each orbit corresponds to a discrete quantity of energy
2. electrons emit energy only when an electron moves from a higher orbit to a lower orbit
73
to calculate how much energy an electron loses in bohr model
Ehigher - Elower = Ephoton = hf
74
what happens when an electron moves from higher stationary state to a lower state?
releases a single photon of light
75
ionization energy
* the energy required to remove an electron from a gaseous species
* cation now has more energy than neutral atom -\> less stable
* the less likely an atom gives up an electron, the more energy required to take that electron away
* energy is put in to remove electron
* positive energy
* endothermic process
* F = Kq1q2 / r2
76
absorption
electron absorbs energy and jumps up to higher energy level
77
emission
when an electron falls from a higher energy level to a lower energy level, and a photon is emitted
78
the energy of the electron changes in discrete amounts with respect to...
energy is directly proportional to principal quantum number
79
bohr model: electron in any of its quantized states in the atom will have an attractive force towards....
the proton
80
as the energy of an electron increases (becomes \_\_\_\_)...
(becomes less negative)... the farther out from the nucleus that it is located (increasing n in eq)
81
as electrons go from a lower energy level to a higher energy level, they get...
AHED
absorb light, higher potential, excited, distant (from nucleus)
82
current
net charge flow per unit time
83
threshold frequency
* minimum frequency of light that causes ejection of electrons
* depends on type of metal being exposed to the radiation
84
if the frequency of the incident photon is less than the threshold frequency....
then no electron will be ejected because the photons do not have sufficient energy to dislodge the electron from the atom
85
waves with higher frequency have... (wavelengths, energy, light spectrum)
shorter wavelengths and higher energy (toward blue and UV end of spectrum)
86
waves with lower frequency have... (wavelengths, energy, light spectrum)
longer wavelengths and lower energy (toward red and infrared end of spectrum)
87
if the frequency of an incident photon is above the threshold frequency of a metal, the photon...
will have more than enough energy to eject a single electron, and the excess energy will be converted to kinetic energy in the ejected electron
88
maximum kinetic energy formula
Kmax = hf - W
89
work function
minimum energy required to eject electron (similar to activation energy)
90
work function related to frquency of metal formula
W = hfthreshold
91
what does the threshold frequency depend upon?
the chemical composition of the material
92
what electrical phenomenon results from the application of the photoelectric effect?
the accumulation of moving electrons creates a current during the photoelectric effect
93
ground state of an atom
state of lowest energy, in which all electrons are in the lowest possible orbitals
94
excited state of an atom
when at least one electron has moved to a subshell of higher than normal energy
95
when electrons return to their ground states...
each will emit a photon with a wavelength characteristic of the specific energy transition it undergoes
96
line spectrum
each line on the emission spectrum corresponds to a specific electron transition
97
lyman series
group of hydrogen emission lines corresponding to transitions from energy levels n ≥ 2 to n=1
98
balmer series
corresponds to transitions from energy levels n ≥ 3 to n=2
99
lyman vs balmer series
* lyman series includes larger energy trnasitions
* shorter photon wavelengths in UV region
100
paschen series
corresponds to transitions from n≥4 to n=3
101
equation that says: the energy of the emitted photon corresponds to the difference in the energy between the higher energy initial state and the lower energy final state
102
photoelectric effect: the higher the intensity of the light beam...
the greater the number of photons per unit time that fall on an electrode, producing a greater number of electrons per unit time liberated from the metal
103
photoelectric effect: when the light's frequency is above the threshold frequency...
the magnitude of the resulting current is direction proportional to the intensity and amplitude of the light beam
104
if the frequency of a photon of light incident on a metal is at the threshold frequency for the metal...
the electron barely escapes from the metal
105
Kmax is only achieved when...
all possible energy from the photon is transferred to the ejected electron
106
how does the work function relate to the energy necessary to emit an electron from a metal?
* work function describes the minimum amount of energy necessary to emit an electron
* any additional energy from a photon will be converted to excess kinetic energy during photoelectric effect
107
electric potential energy
bohr model
q1 = charge on electron
q2 = charge on nucleus
108
potential energy of electron in ground state
negative potential energy
109
What is the equation for magnetic force in terms of magnetic field strength and velocity of an ion?
F = qvB
```
F = Magnetic Force
q = Charge of ion
V = Velocity of ion
B = Magnetic Field Strength
```
110
In a mass spectrometer, the ion moves in a circle. Why?
The ion moves in a circle because the magnetic force is perpendicular to the velocity of the ion, pointing toward the center of the circle. The magnetic force essentially acts as a centripetal force.
111
Which equation can be used in tandem with F = qvB in order to calculate the mass of an ion in a mass spectrometer based on the radius of the circle?
Fc = m(v^2/r)
112
What is the resulting equation when everything is set equal to r?
r = mv/qB
113
The energy of the photon is greater than, equal to, or less than the energy of the emitted electron? Why?
The energy of the photon is greater than the energy of the emitted electron. This is because some of the energy of the photon was required (and used up) in order to free the electron from the metallic surface.
114
What is the equation for the energy of a photon in terms of the energy of the emitted electron?
KEp = KEe + WF
```
KEp = Kinetic energy of the photon
KEe = Kinetic energy of the electron
WF = Work function
```
115
A metal has a work function equal to 3.42 ⋅ 10^-19 J and the energy of the electron is equal to 3.47 ⋅ 10^-19 J. What is the frequency of the photon? (remember Plank's constant = 6.626 ⋅ 10^-34).
(A) 3.04 ⋅ 10 ^ -19 Hz
(B) 6.02 ⋅ 10 ^ -13 Hz
(C) 8.06 ⋅ 10 ^ 12 Hz
(D) 1.15 ⋅ 10 ^ 15 Hz
D) 1.15 ⋅ 10 ^ 15 Hz
KEp = KEe + WF
KEp = 3.47 ⋅ 10^-19 J + 3.42 ⋅ 10^-19 J
KEp = 6.89 ⋅ 10^-19 J
hv = KEp
(6.626 ⋅ 10^-34)v = 6.89 ⋅ 10^-19 J
v = (7 ⋅ 10^-19)/(6 ⋅ 10^-34)
v = about 1 ⋅ 10^15 hz (Actual 1.15 ⋅ 10^15 hz)
116
Which physics equation relates the energy of a photon to its frequency?
E = hf
```
E = Energy of a photon or quantum
h = Planck's Constant (6.626 ⋅ 10^-34 J⋅s)
f = Frequency of wave/radiation
```
117
What two equations can be used in concert with Fc = m(v^2/r) to determine the radius of a Bohr model electron?
Fe = k((q1q2)/r^2) and L = rp
118
An electron's total energy is composed of what two types of energy?
(A) Thermal and kinetic energy
(B) Kinetic and electric energy
(C) Electric and potential energy
(D) Thermal and electric energy
(B) Kinetic and electric energy
An electron's total energy is composed of kinetic and electric energy.
119
What equation can be used to determine the energy of any electron in terms of its principal quantum number?
En = E1/(n^2)
```
En = Energy of electron at n
n = Principal quantum number
E1 = Energy of electron at n=1 (-2.17 ⋅ 10^-18 J)
```
120
The energy at the first orbital is equal to -2.17 ⋅ 10^-18 J. What is the energy in terms of electron-volts (eV)?
(A) 19.2 eV
(B) 3.7 eV
(C) -1.4 eV
(D) -13.6 eV
(D) -13.6 eV
(-2.17 ⋅ 10^-18 J) / (1.6 ⋅ 10^-19) = approximately -10 eV (actual: -13.6 eV)
121
The energy at the first orbital is equal to -13.6 eV. What is the energy in terms of electron-volts (eV) of an electron in the third orbital?
(A) -9.6 eV
(B) -4.8 eV
(C) -3.4 eV
(D) -1.5 eV
(D) -1.5 eV
```
En = E1/(n^2)
E3 = -13.6 eV / 3^2
E3 = approx. -1.5 eV
```
122
An electron drops from the second orbital to the first orbital (-13.6 eV). A photon with how much energy is given off?
(A) -9.6 eV
(B) 6.3 eV
(C) 10.2 eV
(D) 13.6 eV
(C) 10.2 eV
```
En = E1/(n^2)
E2 = -13.6 eV / 2^2
E2 = approx. -3 eV
```
E1 - E2 = -13.6 eV - (-3eV) = -10.6 eV from the electron, so the photon has an energy of approx. 10.6 eV (actual: 10.2 eV)
123
What does it mean in terms of energy levels of electrons to say that the ionization energy of an atom is -13.6 eV?
It would require 13.6 eV to move an electron from n=1 to n=infinity, making the Hydrogen atom an ion.
124
Ernest Rutherford conducted an experiment in which he fired alpha particles at gold foil. He noticed that the majority of alpha particles were not deflected by the gold foil, and all alpha particles would exit the foil. What did these results reveal about an atom's structure?
The Rutherford experiments concluded that an atom has a dense, positively charged nucleus taking up a small fraction of an atom's volume.
125
Based on the previous description of the Bohr model, which of the following atoms does not represent the Bohr model?
(A) H
(B) He+
(C) Li+
(D) Be 3+
The Bohr model relies upon there being only a single electron. Li+ has two electrons, whereas all other options have only one electron.
126
True or false? The Rydberg unit of energy is equal to 2.18 ⋅ 10 ^-18 J/electron, and is the experimentally determined energy of an electron at the smallest possible orbital.
True. The Rydberg unit of energy is equal to 2.18 ⋅ 10 ^-18 J/electron, and is the experimentally determined energy of an electron at the smallest possible orbital.
127
a) 1:3:5
b) 1:1:1
c) 1:4:9
d) 1:2:4
c) 1:4:9
128
what happens during mass spectroscopy?
sample is bombarded with electrons in order to ionize the sample
ions are then subjected to both electric and magnetic fields
129
mass spectrometer separates ions based on...
mass to charge ratio
130
amount of deflection in mass spectrometer is ______ to mass to charge ratio
inversely propotional
131
how to tell which ions would be least deflected in mass spectrometer?
highest mass to charge ratio --\> least deflected
132
Which of the following is NOT true of an electron returning to its ground state?
a) it emits photons
b) distance between nucleus and electron decreases as the electron returns to its ground state
c) releases energy
d) gains potential energy as it falls to the ground level
d) gains potential energy as it falls to the ground level
133
according to the bohr model, which of the following transitions of an electron produce a photon with the lowest energy?
a) n=3 to n=4
b) n=4 to n=3
c) n=4 to n=2
d) n=2 to n=4
b) n=4 to n=3
134
a) 0
b) 360
c) 120
d) 60
a) 0
135
all of the following statements are false about photoelectric effect EXCEPT?
a) kinetic energy of the emitted electron depends on the frequency of the photon and work function of the metal
b) energy of the incident photon is inversely proportional to its freq
c) number of electrons emitted depends only on the frequency of the incident photons
d) energy of the incident photon is directly proportion to its wavelength
a) kinetic energy of the emitted electron depends on the frequency of the photon and work function of the metal
136
a) inc the wavelength of photons
b) inc the number of photons
c) inc the freq of photons
d) dec the freq of photons
b) inc the number of photons
137
photoelectric affect
what affects the velocity of electrons?
* inc freq, inc energy, inc velocity
* inc wavelength, dec energy, dec velocity
138
a) red and violet
b) green and blue
c) violet and red
d) blue and green
a) red and violet
139
a) Be3+
b) He+
c) H
d) Li+
d) Li+
140
a) Each series specifies the wavelength emitted when an electron drops down a specific number of orbitals, decreasing by exactly one, two, or three orbitals, respectively.
b) Each series specifies the wavelengths emitted when an electron drops from higher energy levels to the first, second, and third energy level, respectively.
c) Each series specifies the wavelength emitted when an electron drops down a specific number of orbitals, decreasing by exactly two, three, or four orbitals,
respectively.
d) Each series specifies the wavelengths emitted when an electron drops from a higher energy level to the second, third, and fourth energy level, respectively.
b) Each series specifies the wavelengths emitted when an electron drops from higher energy levels to the first, second, and third energy level, respectively.
141
a) n=2 to n=4
b) n=4 to n=2
c) n=3 to n=4
d) n=4 to n=3
b) n=4 to n=2
142
a) energy of electron is greater than the energy of the photon
b) energy of the electron is less than the energy of the photon
c) energy of electron is zero
d) energy of electron is the same as the energy of the photon
b) energy of the electron is less than the energy of the photon
143
a) n=1 to n=2
b) n=1 to n=0
c) n=infinity to n=1
d) n=1 to n=infinity
d) n=1 to n=infinity
144
a) 286 nm
b) 867 nm
c) 367 nm
d) 486 nm
145
a) 3.4 eV
b) -1.5 eV
c) -13.6 eV
d) 8 eV
b) -1.5 eV
146
a) electron energy dec since E1 is neg
b) electron energy inc since E1 is neg
c) electron energy inc since E1 is pos
d) electron energy dec since E1 is pos
b) electron energy inc since E1 is neg
147
a) KE is neg, PE is pos
b) KE is neg, PE is neg
c) KE is pos, PE is pos
d) KE is pos, PE is neg
d) KE is pos, PE is neg
148
a) 3/2
b) 4/9
c) 9/4
d) 2/3
b) 4/9
149
a) nitrogen
b) carbon
c) hydrogen
d) sulfur
d) sulfur
150
a) (mass x velocity) / (magnetic force)
b) (mass x acceleration) / (magnetic force)
c) (mass) x (velocity)^2 / (magnetic force)
d) (mass) x (acceleration)^2) / (magnetic force)
c) (mass) x (velocity)^2 / (magnetic force)
151
a) distance the ion traveled through the electric field
b) radius of the ion's path in magnetic field
c) charge of the ion
d) size of the ion
b) radius of the ion's path in magnetic field
152
heisenberg uncertainty principle
it is impossible to simultaneously determine, with perfect accuracy, the momentum and position of an electron
153
uncertainty principle eq (2)
p=mv, p = momentum
ΔxΔp ≥ h/4π
154
uncertainty principle
dec uncertain in position...
inc uncertainty in momentum
155
the uncertainty principle has nothing to do with...
the precision of the instrument we are using to measure
156
wave particle duality
matter on a subatomic level can act as particles and as waves
157
energy eq with lambda
Ephoton = hc/lambda
158
as principal quantum number (n) inc, the avg distance of electron \_\_\_
inc, atom is larger (radius), and energy inc
159
principal quantum number (n)
n = any positive integer
main energy level/shell
160
azithmuthal (angular momentum) quantum number (l)
designates the subshell where the electron is located
shape of the orbital --\> most probable location of electrons
important implications for chemical bonding and bond angles
l = 0, 1, 2... n-1
l = 0 --\> s orbital
161
magnetic quantum number (ml)
designates the exact orbital in which our electron is in -\> direction of angular momentum
orientation
ml = -l to +l
162
spin quantum number (ms)
ms = +1/2 up
ms = -1/2 down
because orbitals can have a max of 2 electrons
163
pauli exclusion principle
any two electrons in a given atom cannot have the same 4 quantum numbers
164
number of total orbitals within a shell level (n) can be found using
n2
165
eq to calculate energy of electron using nth principle quantum number
En = Ei / n2
166
eq to calculate az quantum number (l) and angular momentum
L2 = _h_2l(l+1)
L = angular momentum
167
space quantization
any orbital has a certain number of orientations in space
168
where do you lose the electron from?
from the orbital that is highest in energy
169
hard part of writing electronic configurations
metals at Cr -\> jumps to [Ar] 4s1 2d5 -\> fills up d until
Cu -\> [Ar] 4s1 3d10 -\> fills up 4 and stays at 3d10
170
Aufbau Principle
electrons will fill orbitals from lowest to highest energy
171
hund's rule
electrons are going to fill orbitals one at a time before doubling up
172
electron configuration periodic table
173
What is the electron configuration of osmium (Z=76)?
[Xe] 6s2 4f14 5d6
174
What is the electron configuration of Fe3+?
[Ar] 3d5
electrons are removed from the 4s subshell before the 3d because it has a higher principal quantum number
175
write out and compare an orbital diagram for a neutral oxygen atom and an O2- atom
* both have fully filled 1s and 2s orbitals
* O has 4 electrons in 2p
* 2 paired, other 2 have own orbital
* O2- has 6 electrons in 2p
* all are paired
176
where did the Bohr model fail?
* did not take into account the repulsion between multiple electrons surrounding the nucleus
* electrons do NOT follow a clearly defined circular pathway or orbit at a fixed distance from the nucleus
177
what do we know today about orbitals?
electrons move rapidly and are localized within regions of space around the nucleus called orbitals
178
if we want to assess the position of an electron...
the electron has to stop (removing its momentum)
179
energy state
the position and energy of an electron described by its quantum numbers
180
eq max number of electrons within a shell
2n2
181
eq max number of electrons within a subshell
4*l* + 2
182
as *l* values inc.... the energies of the subshell
increase
183
as atomic number inc, the number of electrons
also increases
184
parallel spins
different orbitals with the same ms values
185
n+*l* rule
the lower sum of the values of the first and second quantum numbers, *n+l*, the lower the energy of the subshell -\> fills up first
186
*n+1* rule
if 2 subshells possess the same *n+l* value, the subshell with ______________ has a lower energy and will fill with electrons first
lower n value
187
anions fill orbitals...
lowest to highest energy (same rules)
188
cations fill orbitals...
* start with neutral atom, remove electrons from the subshell with the highest n value first
* if multiple subshell are tied for the highest n value, then electrons are removed from the subshell with the highest *l* value among these
189
basis for hund's rule
electron repulsion: electrons in the same orbital tend to be closer to each other and thus repel each other more than electrons placed in different orbitals
190
according to hund's rule, which types of orbitals have lower energy/high stability than other states?
half filled and fully filled orbitals
191
exceptions to electron configuration
chromium (and other elements in its group) and copper (and other elements in its group)
192
electron configuration
chromium exception
[Ar] 4s1 3d5
moving one electron from the 4s to the 3d allows the 3d to be half filled --\> even though energetically unfavorable, making the 3d subshell half-filled outweighs the cost
193
electron configuration
copper exception
[Ar] 4s13d10
a full d subshell outweighs the cost of moving an electron out of the 4s subshell
194
paramagnetic
materials composed of atoms with unpaired electrons that will orient their spins in alignment with a magnetic field
the magnetic fields of the electrons add together
magnetic field causes parallel spins --\> attraction
195
diamagnetic
materials consisting of atoms that have only paired electrons --\> slightly repelled by magnetic field
magnetic fields of the electrons cancel each other out
(sufficiently strong magnetic fields can still make diamagnetic substances levitate tho)
196
label the groups on periodic table
197
label how many valence electrons in each
198
valence electrons
elements in period 3 (starting w Na) and below
may accept electrons into their d subshell
which allows them to hold more than 8 electrons in their valence shell
199
valence electrons
langthanide and actinide series
highest s and f subshells
200
valence electrons
groups IIIA - VIIIA (groups 13-18)
highest s and p subshells
201
IR spectroscopy
purpose and how
determine chemical structure because different bonds will absorb different wavelengths
202
UV-Vis spectroscopy
uses
absorption of light in visible and ultraviolet range
203
absorption spectra may be represented as...
1. a color bar with peak areas of absorption represented by black lines
2. graph with absolute absorption as a function of wavelength
204
fluorescence
1. excite fluorescent substance with UV radiation
1. UV light photons have high freq (short wavelengths)
2. electron excited to higher energy state and returns to OG state in 2 or more steps
1. each step involves less energy
2. at each step, a photon is emitted with a lower frequency (longer wavelength) than the absorbed UV photon
205
the farther away from nucleus, the electron is the ______ the potential energy
the higher the potential energy
206
what equation can be used to determine the radius of any electron in terms of its principal quantum number?
rn = n2 r1
```
rn = Radius of electron at n
n = Principle quantum number
r1 = Radius of electron at n of 1 (5.3 ⋅ 10^-11)
```
207
valence electrons
groups IA and IIA
highest s subshell
208
Why doesn't Hydrogen produce a continuous spectrum of light?
Because its orbitals have specific energy differences between them, resulting in light waves with specific energy levels and thus specific wavelengths. It cannot produce every different wavelength level.
209
How does the Heisenberg uncertainty principle apply to the Bohr model of the electron?
It is impossible to know the location and momentum of an electron at the same time, which proves the Bohr model to be an inaccurate representation of the atom.
210
Which quantum number is also called the azimuthal quantum number, and includes integer values up to n-1?
(A) l
(B) n
(C) m(l)
(D) m(s)
A) l
The angular momentum quantum number, or azimuthal quantum number, may have integer values as large as n-1.
211
What are the possible values of n, l, m(l), and m(s) for when n = 2? How many electrons are in this shell?
```
n = 2
l = 0 or 1 (s or p)
m(l) = -1, 0, or +1
m(s) = +1/2 (up) or -1/2 (down)
# of electrons = 8
```
212
Draw the shape of the d orbital(s) and f orbital(s).
d has 5 orientations
f has 7 orientations
213
Which of these principles requires that an electron in the same orbital as another electron must spin in the opposite orientation?
(A) Aufbau principle
(B) Pauli exclusion principle
(C) Hund's rule
(D) Electron configuration principle
(B) Pauli exclusion principle
214
What is the shorthand electron configuration for Sc+?
(A) [Ar] 3d1 4s1
(B) [Ar] 3d2
(C) [Ar] 4d1 4s1
(D) [Ar] 4d2
(A) [Ar] 3d1 4s1
Note that the Scandium loses one of its 4s electrons and not its 3d electron when obtaining a +1 charge.
215
what is the max number of electrons that can be found in the 3rd energy level?
a) 10
b) 8
c) 32
d) 18
d) 18
216
What is the max number of electrons in an atom that have the following set of quantum numbers?
n=5 l=3 ml=+2 ms=+1/2
a) 2
b) 14
c) 1
d) 6
c) 1
217
Which of the following subshells do electrons enter first?
a) 4p
b) 2d
c) 3d
d) 4s
d) 4s
218
What is the total number of orbitals for the subshell with quantum numbers n=5 and l=1?
a) 6
b) 5
c) 3
d) 7
c) 3
219
What is the azimuthal quantum number (l) for the orbital from which an electron is lost to form Na+1?
a) 3
b) 1
c) 0
d) 2
c) 0
220
chemical families of periodic table
221
a) I, II, III
b) I and III only
c) II and III only
d) II only
a) I, II, III
222
a) Hund's rule
b) Aufbau Principle
c) Markovnikov's rule
d) Pauli exclusions principle
d) Pauli exclusions principle
223
Which can hold more electrons, the n=2 shell or the l=2 subshell, and by how large a difference?
a) n=2 can hold 4 more electrons than l=2
b) l=2 can hold 2 more electrons than n=2
c) n=2 can hold 2 more electrons than l=2
d) l=2 can hold 4 more electrons than n=2
b) l=2 can hold 2 more electrons than n=2
224
how many electrons are in the n=3 energy level?
a) 8
b) 16
c) 18
d) 10
c) 18
225
An astrologer is observing the emission spectrum of a star to see what gases it contains. What does an emission spectrum look like?
a) a noncontinuous spectrum with bright lines at wavelengths where a gas emits light
b) a continuous spectrum with dark lines at wavelengths where a gas absorbs light
c) a noncontinuous spectrum with bright lines at wavelengths where a gas absorbs light
d) a continuous spectrum with dark lines at wavelengths where a gas emits light
a) a noncontinuous spectrum with bright lines at wavelengths where a gas emits light
226
Potassium reacts with chlorine to form KCl, an ionic compound. What is the electronic configuration of potassium in KCl?
a) [Ar] 4s2
b) [Ar] 4s13d2
c) [Ar] 4s1
d) [Ar]
d) [Ar]
227
Yttrium in its ground state has 39 electrons. Which set of quantum numbers is not possible for Yttrium?
a) 4, 3, -3, -1/2
b) 2, 2, 1, 1/2
c) 3, 1, 0, 1/2
d) 5, 0, 0, -1/2
b) 2, 2, 1, 1/2
228
a) has uncertainty of 2.6 x 10^-10 m
b) has uncertainty of at least 2.6 x 10^-11 m
c) has average value of 2.6 x 10^-11 m
d) has avg value of 2.6 x 10^-10 m
a) has uncertainty of 2.6 x 10^-10 m
229
a) 656 nm
b) 80 nm
c) 322 nm
d) 122 nm
d) 122 nm
230
a) light w a wavelength of 656 nm is emitted
b) light w a wavelength of less than 656 nm is emitted
c) light w a wavelength greater than 656 nm is emitted
d) light w a wavelength of 656 nm is absorbed
b) light w a wavelength of less than 656 nm is emitted
231
does each orbital have to have one unpaired electron to be paramagnetic?
no, as long as there are only unpaired electrons
232
electronegativity
the power of an atom in a molecule to attract or accept electrons to itself
233
if an element is more electronegative...
it has a greater power to attract electrons to itself
234
nonpolar covalent bond
no difference in electronegativity between the two atoms --\> no pull
(difference of less than 0.5)
235
polar covalent bond
difference in electronegativity between two atoms
(greater than 0.5)
236
ionic bond
* more electronegative atom steals the electrons
* (about more than 1.7 difference)
* atoms with low ionization energy will form bonds with atoms that have high electron affinity
237
alkali metals
* very reactive
* react readily with nonmetals, especially halogens
* react with water
* found in nature with other elements
* have only one loosely bound valence electron -\> easily lose to form univalent cations
238
alkaline earth metals
* reactive (less than alkali metals)
* found in nature with other elements
* have 2 valence electrons -\> easily removed to form divalent cations
239
metals
* solids at room temp except for mercury
* lustrous, malleable, ductile
* good conductors of heat and electricity
* high melting points and densities (some exceptions)
240
nonmetals
* poor conductors of heat and electricity
* solid ones - brittle, little/no luster
* include halogens, noble gases
241
halogens
* very reactive non metals
* 7 valence electrons
* really want to complete their octets by gaining an electron
* corrosive
* especially reactive toward alkali and alkaline earth metals
* found in nature as ions (halides) or diatomic molecules
242
noble gases
* aka inert gases
* colorless
* generally unreactive because stable from having 8 electrons in valence shell
* highest ionization energy
* extremely low BP
* exist as gases at room temp
243
metalloids
* aka semimetals
* in between metals and nonmetals
* reactivities dependent on the elements with which they are reacting
244
transition metals
* other def: element whose atom has an incomplete d subshell or which can give rise to cations with an incomplete d subshell
* more general def: d orbital ones
* very hard
* have high melting and boiling points
* malleable
* good conductors due to loosely held electrons in d orbitals
* can have different oxidation states
* bc capable of losing different numbers of electrons from s and d valence electrons
* form many different ionic compounds
* many act as cofactors for enzymes
245
as you go down the groups in periodic table... atomic radius \_\_\_\_
atomic radius increases
(bc adding electrons of higher energy levels that are farther away from nucleus)
246
as you go across periods to the right... atomic radius \_\_\_\_
decreases
more protons as you go across -\> more electrons in outer shell also increases -\> nucleus pulls outer shell electron in more than inner shell electrons shield
Zeft inc
247
electron shielding
* inner electrons shielding outer electron from positively charged nucleus
* protons in nucleus pull outer shell electron in -\> inner shell electrons repels outer shell electron
* (outer electrons don't shield)
248
as you go down a group in periodic table, ionization energy \_\_\_\_\_
decreases
249
factors that affect ionization energy
* effective nuclear charge
* nuclear charge: inc nuclear charge, inc ionization energy
* electron shielding: inc electron shielding, dec ionization energy
* distance of outer electron from nucleus: distance inc, dec ionization energy
250
nuclear charge
the more positive charge you have in nucleus, the more of an attractive force the electrons field -\> the harder it would be to pull the electron away
increase in nuclear charge -\> increase in ionization energy
251
effective nuclear charge
simple way to calculate
~ Zeft = Z - S
Z = nuclear charge
S = shielding electrons
252
electron affinity
* energy released when a gaseous species gains an electron
* adding electron to neutral atom
* gives off energy to form anion
* negative energy
* exothermic process
* ΔHrxn = negative
* but reported as positive number
253
adding an electron but EA ≥ 0
* adding energy in to force the electron to form the anion
* positive energy
* because atom does not have electron affinity
254
as you go across periods to the right... nuclear charge \_\_\_\_
increases
255
as you go across periods to the right... effective nuclear charge \_\_\_\_
increases
256
as you go across periods to the right... electron affinity \_\_\_\_
+exceptions
increases
because effective nuclear charge increases -\> electron adding would feel more of an attractive force
exceptions: Be, B ish, N, Ne
257
as you go down a group, the atomic radius \_\_\_
increases
because addition of a new energy shell
258
as you go across the periodic table, ionization energy \_\_\_
increrases
inc Zeft -\> greater force
AR decreases
259
the higher your ionization energy is, the ___ likely you are to give up electrons
less likely
electrons harder to remove
260
as you go across periods to the right, electronegativity \_\_\_\_
increases
Zeft increases --\> higher force pulling in electrons
261
most electronegative atom
Flourine
262
As you go down the groups. electronegativity \_\_\_\_
decreases
AR increases -\> force of nuclear protons decreases
263
for noble gases, electronegativity is \_\_\_\_
undefined
noble gases can't accept any more electrons
264
the lower your ionization energy is, the ___ likely you are to give up electrons
more
electrons easily removed
265
high electron affinity means that
atoms will easily accept electrons
266
low electron affinity means that
atoms will not easily accept electrons
267
ionic bonds
periodic trends
atoms with low ionization energy will form bonds with atoms that have high electron affinity
(atoms of the left and right side of period table, noble gases not included)
268
periodic law
the chemical and physical properties of the elements are dependent, in a periodic way, upon their atomic numbers
269
representative elements
A elements in periodic table
have their valence electrons in orbitals of s or p subshells
270
nonrepresentative elements
include transition elements, lanthanide, and actinade series
271
transition metals
periodic trends
* 2 or more oxidation states
* bc valence electrons are loosely held
272
metals
periodic trends
* easily give up electrons
273
oxidation state
charges when forming bonds with other atoms
274
nonmetals
periodic t5rends
inability to easily give up electrons
275
groups in periodic table that are likely to be lustrous
metals
276
groups in periodic table that are likely to be be poor conductors of heat and electricity
nonmetals
277
groups in periodic table that are likely to be good conductors but brittle
metalloids
278
effective nuclear charge
electrostatic attraction between the valence shell electrons and the nucleus
279
atomic radius
equal to 1/2 of the distance between the centers of 2 atoms of an element that are briefly in contact with each other
(akin to diameter)
(opposite of all other periodic trends)
280
can atomic radius be measured using single atom?
cannot be measured by using a single atom because electrons are constantly moving around
281
ionic radii
assumptions
1. metals lose electrons and become positive; non-metal gain electrons and become negative
2. metalloids can go in either direction, but tend to follow the trend based on which side of the metalloid line they fall on
282
nonmetals close to metalloid line
require more electrons than other nonmetals to achieve the most stable electronic configuration -\> possess larger ionic radius than other nonmetals
283
metals closer to metalloid line
have more electrons to lose to achieve most stable electronic configuration -\> ionic radius is smaller than other metals
284
ionization energy
endo or exothermic?
endothermic because removing an electron from an atom requires an input of heat
285
why is the removal of a 2nd or 3rd electron require more energy?
remove of more than 1 electron means that electrons are being removed from an increasing catatonic (positive) species
286
active metals
* include the alkali and alkaline earth methals
* low ionization energy
* reactive
* do not exist naturally in neutral forms -\> found in ionic compounds, minerals, or ores
287
electron affinity
endo or exothermic?
exothermic because it expels energy in form of heat to acquire an electron
288
periodic trends
289
which has larger radius?
F or F-
F-
290
which has larger radius?
K or K+
K
291
chalcogens
* crucial for normal biological functions
* at high concentrations, many can be toxic or damaging (usually the heavier ones)
292
groups in periodic table that exhibit:
high reactivity to water
groups 1 and 2
293
groups in periodic table that exhibit:
negative oxidation states
almost all groups but most notably transition metals
294
groups in periodic table that exhibit:
multiple oxidation states
all groups but most notably transition metals
295
groups in periodic table that exhibit:
possess a full octet in the neutral state
noble gases
296
When Dmitri Mendeleev published the first periodic table, it was based on the periodic law. Which of the following best describes the periodic law?
(A) Elements' physical properties depend on their atomic numbers in a periodic way.
(B) Elements' chemical properties depend on their atomic numbers in a periodic way.
(C) Elements' chemical and physical properties depend on their atomic numbers in a periodic way.
(D) Elements' chemical, physical, and nuclear properties depend on their atomic numbers in a periodic way.
(C) Elements' chemical and physical properties depend on their atomic numbers in a periodic way.
297
Sodium (electronegativity of .9) bound to Chlorine (electronegativity of 3.0) is what type of bond? How did you know this based on electronegativity values?
(A) Nonpolar covalent
(B) Polar covalent
(C) Ionic
(D) Hydrogen
(C) Ionic
Sodium and Chlorine have an electronegativity difference of 2.1, which is greater than 1.7, making their bond an ionic bond.
298
Periodic table columns are referred to as _________ while periodic table rows are referred to as \_\_\_\_\_\_\_\_\_.
Fill in the blanks using the following options:
- periods
- halogens
- metalloids
- groups
Periodic table columns are referred to as groups while periodic table rows are referred to as periods.
299
Which of the following best explains why elements in the same group share similar chemical properties?
(A) Elements in the same group have a similar number of nucleons, giving them similar chemical properties.
(B) Elements in the same group have the same number of valence electrons, giving them similar chemical properties.
(C) Elements in the same group have the same number of electrons, giving them similar chemical properties.
(D) Elements in the same group have various multiples of the first element's valence electrons, giving them similar chemical properties.
(B) Elements in the same group have the same number of valence electrons, giving them similar chemical properties.
Having the same number of valence electrons allows the different elements to interact with their environments in similar ways.
300
True or false? Alkaline earth metals are more reactive than Alkali metals because Alkaline earth metals have more electrons.
False. Alkali metals are more reactive than Alkaline earth metals because Alkali metals have an unpaired electron.
301
True or false? Metals are considered to be good conductors because they can exist in multiple oxidation states. This means that valence electrons are only loosely held by an atom and are free to move.
True. Metals are considered to be good conductors because they can exist in multiple oxidation states. This means that valence electrons are only loosely held by an atom and are free to move.
This is also a description of the term "sea of electrons".
302
Halogens belong to which group as compared to Noble Gases?
Halogens belong to group 7A as compared to Noble Gases, which belong to group 8A.
303
True or false? Noble gases are more reactive than halogens because noble gases are more mobile and mixable as gases.
False. Halogens are more reactive than noble gases because halogens have an unpaired electron in their outermost orbital.
304
The octet rule states that an element will gain or lose electrons to achieve a stable octet formation, like the noble gases have. However, the octet rule has many exceptions. Which of the following is NOT one of the exceptions experimentally seen?
(A) Noble gases like Xenon can form covalent bonds using more than eight electrons like XeF6.
(B) Helium is a noble gas and relatively inert, even though it only has two valence electrons.
(C) Nitrogen can form covalent bonds using more than eight electrons, like in the compound NO3 (with a central Nitrogen and double bonds between 2 O and N)
(D) Silicon can form covalent bonds using more than eight electrons, like in SiO4 (with a central Silicon and double bonds between 2 O and Si).
(C) Nitrogen can form covalent bonds using more than eight electrons, like in the compound NO3 (with a central Nitrogen and double bonds between 2 O and N)
The octet rule does not have exceptions in period 2, where Nitrogen is found. Also, that description of NO3 in the answer is incorrect.
305
True or false? Metalloids can also be called semimetals, because they can alternatively act as metals or nonmetals in the same environment.
False. Metalloids (also called semimetals) can act as either metals or nonmetals DEPENDING on the environment. For example, Boron will act like a nonmetal when surrounded by Sodium but like a metal when surrounded by Fluorine.
306
Why does IUPAC not consider Zn(s) a transition metal?
It has a complete d-subshell, even in its +2 cation form.
307
A cation is bigger or smaller than its original atom? Why?
Smaller because you are taking an electron away, potentially reducing electron repulsion, making it smaller.
308
An anion is bigger or smaller than its original atom? Why?
Bigger because you are adding an electron, potentially increasing electron repulsion, making it bigger.
309
If Hydrogen and Lithium both have the same Zeff, why does Hydrogen have a higher ionization energy?
Because Lithium's outer electron is farther away from the nucleus.
310
The acidity of a compound depends on how well an atom can accommodate a negative charge. How does this relate to electronegativity's trends of increasing moving up and to the right?
(A) Acidity will have the same trends because electronegativity is how much an element wants a negative charge, and donating the proton gives it a negative charge.
(B) Acidity will have the same trends because electronegativity increases as atomic radius decreases, and a smaller radius will allow a more delocalized negative charge, providing more stability.
(C) Acidity will not have the same trends because electronegativity is how much an element wants a negative charge, and donating the proton gives it a negative charge.
(D) Acidity will not have the same trends because electronegativity increases as atomic radius decreases, and a larger radius will allow a more delocalized negative charge, providing more stability.
(D) Acidity will not have the same trends because electronegativity increases as atomic radius decreases, and a larger radius will allow a more delocalized negative charge, providing more stability.
Acidity increases moving to the right and moving down the periodic table.
311
Which has a greater electron affinity: Lithium or Berylium? Why?
Lithium. This breaks the normal electron affinity trend because Berylium's new electron will be added to a far away p orbital, which will experience more shielding than Lithium completing its s orbital.
312
Which has a greater electron affinity: Nitrogen or Carbon? Why?
Carbon. This breaks the normal electron affinity trend because Nitrogen's new electron will be added to a p orbital with another electron, which increases the amount of repulsion that the new electron will experience.
313
general rule for balancing complex chemical equations
1. balance molecules with multiple elements in them first
1. (leave single element molecules for last)
2. treat groups as one thing
314
balance
315
combination reactions
A + B -\> AB
two or more reactants combine to form new compound
316
decomposition reactions
AB -\> A + B
reactant dissociates into 2 or more new compounds
usually using heat, high frequency radiation, or electrolysis
317
single displacement reaction
aka substitution reactions
C kicks out B
318
double displacement reaction
aka metathesis reaction
one of products is removed from the solution as a precipitate or gas or when two of the of species combine to form a weak electrolyte that remains undissociated in solution
319
combustion reactions
hydrocarbons are burned in presence of diatomic gas to form CO2 and H2O molecules
320
redox reactions
oxidation states of atoms change
chemical reactions in which at least one chemical species is gaining electrons (being reduced) and another chemical species is losing electrons (being oxidized)
321
how to manually calculate oxidation state
valence electrons - non bonding electrons - electrons gained from bonds
* more electronegative element gets all the electrons in a bond
322
assigning ox numbers
1. if atom is made of only **one kind** (elemental state) of atom, ox = 0
2. atoms in monoatomic (single atom) ions, ox = their charge
3. assign ox to FHO
1. F = -1
2. H = +1
3. O = -2
1. except peroxides = -1
2. more electroneg elements = +2
4. Group IA, ox = +1
5. Group IIA, ox = +2
6. Group VIIA, ox = -1
1. except when combined with element of higher electronegativity
323
Rules for balancing redox equations
1. Assign oxidation numbers to see what is being reduced and what is being oxidized.
2. Divide reaction into two half reactions
3. Balance each half reaction
1. a) balance all elements except O and H
2. b) add H2O to balance O
3. c) add H+ to balance H
4. Balance the charges by adding electrons and multiplying to make sure -\> each ½ reaction has the same number of electrons
5. Add the two ½ reactions together.
6. IF BASIC, add OH- to both sides to neutralize (cancel) the H+’s
1. cancel
2. OH- combine with H+ to make water
7. Check atoms and charges
324
oxidation number def
effective ionic charge obtained by assuming 100% ionicity
guide to the type of reactions that a substance can undergo because an element in its highest oxidation state may only be reduced whereas one in an intermediate oxidation state may either be oxidized or reduced
325
reduction
gain of e- (OIL RIG)
ox # gets smaller
326
reducing agent
gains/accepts e-
is oxidized
* low electronegativity
* low IE
* easily loses electrons
327
oxidation
loss of e- (OIL RIG)
ox # gets bigger
328
oxidizing agent
loses/donates e-
is reduced
* high electronegativity
* high IE
* easily gains electrons
329
gain of e-
ox # gets smaller
reduction
330
gains/accepts e-
is oxidized
reducing gaent
331
loss of e-
ox # gets bigger
oxidation
332
loses/donates e-
is reduced
oxidizing agent
333
balance
redox
334
assign ox number
335
a) all of the above
b) 3rd ionization energy
c) 2nd ionization energy
d) 1st ionization energy
b) 3rd ionization energy
336
Which of the following has the highest ionic radii?
a) N3-
b) F-
c) Li+
d) C
a) N3-
337
Which of the following is TRUE about transition elements?
a) they cannot conduct electricity
b) they have varying oxidation states
c) they all have an incomplete d orbital
d) they are more electronegative than halogens
b) they have varying oxidation states
338
a) electron affinity, since both describe an energy change when the number of electrons is changed in an atom
b) electronegativity, since both describe qualitative changes in electron density
c) electronegativity, since both describe an energy change when the number of electrons is changed in an atom
d) electron affinity, since both describe qualitative changes in electron density
b) electronegativity, since both describe qualitative changes in electron density
339
a) no, iron is never seen at oxidation states higher than +3 due to hund's rule
b) yes, iron is occasionally seen with a charge of +8, giving it the same electron configuration as Argon
c) yes, iron is occasionally seen with a charge of +6, giving it the same electron configuration as calcium
d) no, iron is only able to lose its highest s orbital electrons
c) yes, iron is occasionally seen with a charge of +6, giving it the same electron configuration as calcium
340
a) silicon can form covalent bonds using more than 8 electrons, like in the compound SiO4
b) helium is a noble gas and relatively inert, even though it only has 2 valence electrons
c) xenon can form covalent bonds using more than 8 electrons, like in the compound XeF6
d) nitrogen can form covalent bonds using more than 8 electrons, like in the compound NO3
d) nitrogen can form covalent bonds using more than 8 electrons, like in the compound NO3
341
a) add an ethanol subunit
b) reduce the concentration of crown ether
c) increase the concentration of crown ether
d) remove an ethanol subunit
d) remove an ethanol subunit
342
a) it would decrease by about 1
b) it would increase by about 1
c) it would stay the same
d) it cannot be determined
b) it would increase by about 1
343
a) At
b) F
c) Li
d) Cs
d) Cs
344
a) He
b) Br
c) Mg
d) Rb
a) He
345
a) K+ \> Na+ \> Cl- \> Ar
b) K+ \> Cl- \> Ar \> Na+
c) Cl- \> Ar \> K+ \> Na+
d) Na+ \> K+ \> Ar \> Cl-
c) Cl- \> Ar \> K+ \> Na+
346
which of the following classifications of elements is least conductive?
a) transition metals
b) metalloids
c) alkali metals
d) alkaline earth metals
b) metalloids
347
the elements fall into 3 broad categories: metals, metalloids, and nonmetals. which of the following groups contains exactly one nonmetal?
a) 1
b) 5
c) 9
d) 2
a) 1
348
a) they are all alkaline earth metals
b) they are all alkali metals
c) they are all transition metals
d) they are all metalloids
b) they are all alkali metals
349
a) Ca
b) Sr
c) Ba
d) Mg
c) Ba
350
bond polarity is important for understanding the interactions of molecules. which of the following would have the least polar bond?
a) F-N
b) Na-Cl
c) B-H
d) Cl-P
c) B-H
351
a) with the hydrogens closer to the ion, bc oxygen is more electronegative than hydrogen
b) with the hydrogens closer to the ion, bc hydrogen is more electronegative than hydrogen
c) with the oxygen closer to the ion, bc oxygen is more electronegative than oxygen
d) with the oxygen closer to the ion, bc hydrogen is more electronegative than oxygen
c) with the oxygen closer to the ion, bc oxygen is more electronegative than oxygen
352
a student wants to make a permanent magnet, and is learning about paramagnetic and diamagnetic materials. which material would make a better magnet?
a) neither
b) both
c) diagmagnetic
d) paramagnetic
a) neither
353
a student is surprised to learn that the acidity of a compound depends on how well an atom can accommodate a negative charge. how does this relate to electronegativity's trends of increasing moving up and to the right?
a) Acidity will have the same trends because electronegativity depends on Z(eff), which increases as you move up and to the right on the periodic table.
b) Acidity will have the same trends because electronegativity is how much an element wants a negative charge, and donating the proton gives it a negative charge.
c) Acidity will not have the same trends as electronegativity.
d) Acidity will have the same trends because electronegativity increases as atomic radius decreases, and a smaller radius will allow a more delocalized negative charge, providing more stability.
c) Acidity will not have the same trends as electronegativity.
354
Reactions between metals and nonmetals are often
extremely exothermic, since both elements gain stability
by moving closer to the octet rule. Which of the following
would not be an example of this scenario?
a) boron and nitrogen
b) aluminum and phosphorus
c) radium and polonium
d) rubidium and astatine
a) boron and nitrogen
355
law of conservation of mass and charge
reactants consumed must equal the mass of products generated
ensure that the number of atoms of each element on the reactant side equals the number of atoms of that element on the product side
356
stoichiometric coefficients
used to indicate the relative number of moles in a given species
357
3 step process for stoichiometry
1. convert from the given units to moles
2. use mole ratio
3. convert from moles to the desired units
358
1 mole of any ideal gas at STP = ___ L
22.4 L
359
1 mole of any substance = ___ particles
6.022 x 1023
360
1 mole of any substance = ____ grams
its molar mass in grams
361
when given quantities of 2 reactants are given on MCAT....
expect to have to figure out limiting reagent
362
limiting reagent
limits the amount of product that can be formed in the reaction
363
determining limiting reagent
1. balance eq
2. covert mass to moles first
3. use mole ratio to convert how much of product is made from mole of each product
1. reactant that produces smaller amount of product is LR
4. find amount remaining excess reactant by subtracting mass of excess reagent consumed from total mass of excess reagent given
364
How many grams of calcium chloride are needed to prepare 71.7 g of silver chloride according to the following equation?
CaCl + 2 AgNO3 --\> Ca(NO3)2 + 2 AgCl
365
Example: If 27.9 g of Fe react with 24.1 g of S to produce FeS, what would be the limiting reagent? How many grams of excess reagent would be present in the vessel at the end of the reaction?
Fe + S --\> FeS
366
theoretical yield
max amount of product that can be generated as predicted from the balanced eq, assuming that all of the limiting reactant is consumed
367
actual yield
amount of product one actually obtains during the reaction
368
percent yield eq
percent yield = actual yield/theoretical yield x 100%
369
What is the percent yield for a reaction in which 28 g of Cu is produced by reacting 32.7 g of n in excess CuSO, solution?
The balanced equation is as follows:
Zn (s) + CuSO4 (aq) --\> Cu (s) + ZnSO4 (aq)
370
ionic compounds are made up of...
* cations --\> usually metals
* anions --\> usually nonmetals
371
coulombs law
ionic compounds
* distance between nuclei in ionic bonds is inversely proportional to the force
* ionic compounds with long bond distances are much more weakly held together
372
ionic compounds with long bond distances are \_\_\_strongly/weakly\_\_\_ held together
weakly
373
nomenclature ionic compounds
elements that can form more than one positive ion
(usually metals)
charge - roman numeral in parenthesis
followed by name of element
ex: Fe3+ Iron(III)
374
nomenclature ionic compounds
-ous
ions with lesser charge
ex: Fe2+ Ferrous
375
nomenclature ionic compounds
-ic
ions with greater charge
ex: Fe3+ Ferric
376
nomenclature ionic compounds
-ide
monoatomic anions
ex: H- hydride, S2- sulfide
377
nomenclature ionic compounds
-ite
when an element forms 2 oxyanions, name of one with less oxygen
(l**ite**st anions have the fewest oxygens)
ex: NO2- nitrite
378
oxyanion
polyatomic anions that contain oxygen
379
nomenclature ionic compounds
-ate
when an element forms 2 oxyanions, the name of the one with more oxygen
(heaviest anions ate the most oxygens)
ex: NO3- nitrate
380
nomenclature ionic compounds
hypo-
extended series of oxyanions
less oxygen
ex: ClO- hypochlorite
381
nomenclature ionic compounds
per-
(hyper)
extended series of oxyanions
more oxygen
ex: Perchlorate ClO4-
382
alkali metals are not typically found in nature in uncharged state because...
they are highly reactive with moisture
instead they are found as cations in salts (NaCl)
383
ion charges
alkali metals (group 1)
+1
384
ion charges
alkaline earth metals (group 2)
+2
385
ion charges
halogens (group 17)
form monoatomic anions
-1
386
nonmetals generally form \_\_cations/anions\_\_\_
anions
387
how does color of a solution indicate ox state of a given element in the solution?
same element in different ox states can undergo different electron transitions and absorb different frequencies of light
388
ionic compounds make \_\_good/bad\_\_ electrolytes because....
good
they dissolve most readily
389
nonpolar covalent compounds make \_\_good/bad\_\_ electrolytes because....
bad
they do not form current carrying ions
390
solid ionic compounds are \_\_good/bad\_\_ conductors of electricity because...
bad
charged particles are rigidly set in place by the lattice arrangement of the crystalline solid
391
ionic compounds in aqueous solutions are \_\_good/bad\_\_ conductors of electricity because...
good
lattice arrangement is disrupted by ion dipole interactions -\> cations and anions free to move --\> ions can conduct electricity
392
electrolytes
* solutes that enable solutions to carry currents
* electrical conductivity of solutions is governed by presence and concentration of ions
393
a solute is considered a strong electrolyte if...
ex
it dissociates completely into its constituent ions
ex: CaCl and KI -\> highly polar covalent bonds
394
a solute is considered a weak electrolyte if...
ex
it ionizes or hydrolyzes incompletely in aqueous solution
only some of the solute is dissolved into its ionic constiutents
395
a solute is considered a nonelectrolyte if...
ex
they do not ionizat at all inf water, retaining their molecular structure and solution (may also limit solubility)
ex: nonpolar gasses and organic compounds - CO2, glucose
396
identify the following ions as cations or anions, then provide the formula
phosphate
anion
PO43-
397
identify the following ions as cations or anions, then provide the formula
hypochlorite
anion
ClO2-
398
identify the following ions as cations or anions, then provide the formula
ammonium
cation
NH4+
399
identify the following ions as cations or anions, then provide the formula
bicarbonate
anion
HCO3-
400
identify the following ions as cations or anions, then provide the formula
nitrite
anion
NO2-
401
identify the following ions as cations or anions, then provide the formula
chromium(II)
cation
Cr2+
402
almost all oxidizing agents contain...
oxygen or another strongly electronegative element (such as a halogen)
403
reducing agents often contain...
metal ions or hydrides (H-)
404
complete ionic equation
various species split into all ions present + spectator ions
405
conventions of formula writing
cation and anion
cation 1st and anion 2nd
406
spectator ions
ions not taking part in overall reaction but simply remaining in solution unchanged
407
net ionic equation
* all aqueous compounds split into constituent ions
* solid slats should be kept together
* no spectator ions
408
disproportionation reactions
aka dismutation
type of redox reaction
element undergoes both oxidation and reduction in producing its products
409
redox titrations
* follow the transfer of charge (as electrons) to reach the equivalence point
* utilize indicators that change color at a particular voltage (emf) value
410
iodometric titration
* type of redox titration
* relies on titration of free iodine radicals
* uses starch indicators to identify iodine complexes
* presence of iodine is initially determined by a dark solution in the presence of starch
* at the endpoint of the titration, a colorless solution develops
411
412
A sample is assayed for lead by a redox titration with I3- (ag). A 10.00 g sample is crushed, dissolved in sulfuric acid, and passed over a reducing agent so that all the lead is in the form Pb2+. The Pb2+ (aq) is completely oxidized to Pb4+ by 32.60 mL of a 0.7 M solution of NaI3.
The balanced equation for the reaction is:
I3- + (aq) + Pb2+ (aq) --\> Pb4+ (aq) + 3 I- (ag)
Calculate the mass of lead in the sample.
413
The law of conservation of matter could also be applied to electrons, so that electrons are not created nor destroyed in a reaction, only transferred from one element to another. If this is true, which of the following can be changed in a chemical reaction?
(A) Number of nucleons
(B) Oxidation state
(C) Electron affinity
(D) None of the above can change
(B) Oxidation state
The oxidation state of an atom is similar to an atom's "ownership" of valence electrons, and could change when interacting with other compounds, like in a chemical reaction.
414
How is the formal charge of Chlorine changing in this hypothetical reaction: Li(s) + Cu(s) + OCl- -\> CuO + LiCl?
(A) Chlorine goes from a -1 to a +1 formal charge
(B) Chlorine goes from a +1 to a -1 formal charge
(C) Chlorine goes from a 0 to a -1 formal charge
(D) Chlorine does not change, staying at a -1 formal charge
(C) Chlorine goes from a 0 to a -1 formal charge
Li(s) + Cu(s) + OCl- -\> CuO + LiCl
Chlorine's formal charge goes from 0 to -1.
We know that Chlorine starts out at a neutral formal charge based on the Lewis dot structure of Hypochlorite and how we calculate formal charges.
We also know that Chlorine has a -1 formal charge in LiCl because of the nature of that ionic bond.
415
Define electrolytes, and explain why they are much better at conducting than solid salts.
Electrolytes are solutes that enable solutions (often aqueous) to carry currents. Electrolytes are free to move around in the solution, so the charges can move, unlike in solid salts where charged particles are rigidly set in place by lattice structures.
416
from the electronic configurations given below, which element is most likely to act as an oxidizing agent?
a) [Ne]3s1
b) [Ar]4s23d104p6
c) [Ar]4s23d10
d) [Ne]3d9
d) [Ne]3d9
417
Some bacteria can oxidize sulfur to sulfuric acid and sulfite depending on the environmental conditions. What is the role of oxygen in these reactions?
a) it acts as a reducing agent
b) it loses electrons
c) it makes the reactions go faster
d) it acts as an oxidizing agent
d) it acts as an oxidizing agent
418
What is the ox number of chlorine in sodium hypochlorite (NaClO)?
a) +2
b) +1
c) -1
d) 0
b) +1
419
a) none of the above
b) I only
c) II and III only
d) III only
a) none of the above
420
MnO4- reacts with I- to form I2 and Mn2+ in a basic solution. What is the balanced eq for this reaction?
a) 2 MnO4- + 10 I- + 16 OH- --\> 5 I2 + 2 Mn2+ + 8 H2O
b) 2 MnO4- + 10 I- + 16 H+ --\> 5 I2 + 2 Mn2+ + 8 H2O
c) 2 MnO4- + 10 I- + 8 H2O --\> 5 I2 + 2 Mn2+ + 16 H+
d) 2 MnO4- + 10 I- + 8 H2O --\> 5 I2 + 2 Mn2+ + 16 OH-
d) 2 MnO4- + 10 I- + 8 H2O --\> 5 I2 + 2 Mn2+ + 16 OH-
421
Acid-base reactions are often also redox reactions. MnO4- reacts with I- in an acidic solution to form I2, and Mn2+. What is the ratio of H+ to I- used up by this reaction?
a) 1.6
b) 0.6
c) 0
d) 1
a) 1.6
422
Two aldehydes can react with each other to form an alcohol and a carboxylic acid. This is an example of a:
a) acid base rxn
b) elimination rxn
c) combustion rxn
d) disproportionate rxn
d) disproportionate rxn
423
What are the ox number of oxygen in OF2 and ClO2-?
a) -2 and -2
b) +2 and -2
c) -2 and +1
d) +1 and -2
b) +2 and -2
424
A student is trying to test the conductivity of different solvents, preparing for a big experiment. They are testing deionized water and pure diethyl ether as their two solvents. Which solution(s) will work well as a conductor in this experiment?
a) Both diethyl ether and deionized water will be good
conductors, since they are both polar.
b) Deionized water will be a better conductor, since it is more polar than diethyl ether.
c) Neither deionized water nor diethyl ether will be good conductors, since they both lack electrolytes.
d) Diethyl ether will be a better conductor, since it is less polar than deionized water.
c) Neither deionized water nor diethyl ether will be good conductors, since they both lack electrolytes.
425
Although the law of conservation of matter says that mass remains constant in a chemical reaction, those masses are free to move. Which of the following can be changed in a chemical reaction?
a) atom's mass number
b) atom's electron affinity
c) atom's polarizability
d) atom's oxidation state
d) atom's oxidation state
426
molecule
combination of 2 or more atoms held together by covalent bonds
smallest units of compounds that display their identifying properties
427
ionic compounds do not form true molecules because...
of the way the oppositely charged ions arrange themselves in the solid state
428
ionic compounds form combinations of elements with ____ electronegativity differences
large
429
molecular compounds form combinations of elements with ____ electronegativity differences
similar
430
formula unit
empirical formula of an ionic compound
431
molecular weight
sum of the atomic weights of all of the atoms in the molecule
units: amu/molecule
432
formula weight of an ionic compound is found by
adding up the atomic weights of the constituent ions according to its empirical formula
433
molar mass
mass of one mole of a compounds
g/mol
434
equivalents
how many moles of the thing we are interested in (protons, hydroxide ions, electrons, ions) will one mole of a given compound produce?
435
gram equivalent weight
eq
produces one equivalent of the particle of interest
gram equivalent weight = molar mass / n
n = number of particles of interest produced or consumed per molecule
436
equivalents and gram equivalent weight eq
equivalents = mass of compounds / gram equivalent weight
437
conversion for normality to molarity
molarity = normality / n
n = number of particles of interest produced or consumed per molecule
438
law of constant composition
any pure sample of a given compounds will contain the same elements in an identical mass ratio
439
empirical formula
simplest whole number ratio of the elements in the compound
440
molecular formula
exact number of atoms in each element in the compound
is a multiple of empirical formula
441
percent composition of an element
percent of sa specific compound that is made up of a given element
442
percent composition eq
mass of element in formula / molar mass x 100%
443
What are the empirical and molecular formulas of a carbohydrate that contains 40.9% carbon, 4.58% hydrogen, and 54.52% oxygen and has a molar mass of 264 g/mol?
444
neutralization reaction
type of double displacement reaction
acid reacts with a base to produce a salt (and usually water)
HCl + NaOH -\> NaCl + H2O
445
Describe in words what occurs when Zn(NO3)2 is dissolved in (NH4)2S.
Ammonium cations swap places with (or displace) zinc cations yielding
ammonium nitrate and zinc(II) sulfide. Zinc(II) sulfide then precipitates out of solution as a solid salt.
446
classify the most likely reactions
447
You have a substance that is composed of 20.8% Iron and 79.2% Chlorine by mass. What is the chemical formula of this compound?
(A) FeCl3
(B) Fe2Cl3
(C) FeCl6
(D) Fe3Cl6
(C) FeCl6
Pretend you have 20.8 g of Fe. That would be 20.8 g / 55.845 g -\> approx. .33 moles Fe
Pretend you have 79.2 g Cl. That would be 79.2 g / 35.45 g -\> approx. 2 moles Cl.
So, for every 1 Fe atom, you must have 6 atoms of Cl, making the formula FeCl6.
448
How would you find the empirical formula when given only the molecular formula?
You can find the empirical formula by finding the ratio of atoms in the molecular formula and then dividing by their greatest common factor.
449
To determine the number of moles in a given mass of a compound, you need to convert mass to moles. Which of the following terms do you divide the mass by to get moles?
(A) Atomic mass units
(B) Formula weight
(C) Avogadro's number
(D) Percent composition
(B) Formula weight
The formula weight (i.e. molecular weight or molar mass) of a compound is the sum of the atomic weights of each atom in the compound.
450
When the mass of a sample of a given compound is half of its formula weight, how many moles of the compound do you have?
(A) 1/4 of a mole
(B) 1/2 of a mole
(C) 1 mole
(D) 2 moles
(B) 1/2 of a mole
Recall that the formula weight of a compound is equal to the mass of one mole of the compound.
451
You have 8.32 g of P4(s) and 3.22 g of Cl2. Considering the following unbalanced equation, how much PCl3 will be produced: P4 + Cl2 -\> PCl3
(A) 4.12 g
(B) 1.82 g
(C) 0.27 g
(D) 15.78 g
(A) 4.12 g
The equation is balanced as follows: P4 + 6Cl2 -\> 4PCl3
8.32 g P4 × (1 mol P4/[30.97×4] g P4) × 4 moles PCl3/mol P4 = approx. 0.25 mol PCl3 (actual: 0.269 mol PCl3)
3.22 g Cl2 × (1 mol Cl2/[35.45×2] g Cl2) × (4 moles PCl3/6 mol Cl2) = approx. 0.025 mol PCl3 (actual: 0.030 mol PCl3)
Thus, Cl2 is the limiting reagent.
0.03 mol PCl3 × 137.33 g/mol PCl3 = about 4g PCl3 (actual: 4.12 g PCl3)
Since Cl2 is the limiting reagent, and my answer will be close to 4g , making (A) 4.12 g the correct answer.
452
Equivalents are related to limiting reagents, because using a different compound means a different mass needs to be used to have the same effect (whether in buffering, electrons to donate or protons to donate). How many grams of Magnesium would need to be used to donate as many electrons as 50g of Sodium?
(A) 57.6 g
(B) 82.1 g
(C) 15.2 g
(D) 26.4 g
(D) 26.4 g
Notice that Magnesium could donate two electrons per atom, whereas Sodium could only donate one.
(50g Na)/(22.99g/mol Na) × (I mol e- / mol Na)= approx 2 mol e- (actual: 2.17 mol e-)
(2.17 mol e-) × (1 mol Mg/ 2 mol e-) × (24.3g Mg/ mol) = approx 25g Mg (actual 26.4g Mg)
453
how many sulfur atoms exist in 3 moles of sulfur dioxide?
a) 6 x 10^23 atoms
b) 1.2 x 10^24 atoms
c) 1.8 x 10^23 atoms
d) 1.8 x 10^24 toms
d) 1.8 x 10^24 toms
454
how many moles are equal to 4.084 x 10^26 molecules?
a) 67.8 x 10^24 moles
b) 678 moles
c) 67.8 moles
d) 6.78 x 10^23 moles
b) 678 moles
455
How many sodium ions are present in 316 g of sodium phosphate, Na3PO4?
a) 3.6 x 10^23
b) 3.5 x 10^24
c) 6.6 x 10^23
d) 6.6 x 10^24
b) 3.5 x 10^24
456
a) 15.2 g
b) 26.4 g
c) 82.1 g
d) 57.6 g
b) 26.4 g
457
a) I and III only
b) I and II only
c) II and III only
d) I, II, and III only
c) II and III only
458
a) II and III only
b) I and II only
c) I, II, and III only
d) I and III only
c) I, II, and III only
459
a) 2 mol
b) 5 mol
c) 6 mol
d) 3 mol
c) 6 mol
