1
Q
Mass Unit
A
kilogram (kg)
2
Q
Length Unit
A
meter (m)
3
Q
time unit
A
second (s)
4
Q
temperature unit
A
kelvin (K)
5
Q
luminous intensity
A
Candela (cd)
6
Q
amount of substance
A
mole (mol)
7
Q
mega
A
M… 10^6
8
Q
kilo
A
k… 10^3
9
Q
Deci
A
d… 10^-1
10
Q
centi
A
c… 10^-2
11
Q
milli
A
m… 10^-3
12
Q
micro
A
u… 10^-6
13
Q
nano
A
n… 10^-9
14
Q
pico
A
p… 10^-12
15
Q
femto
A
f… 10^-15
16
Q
all matter is composed of tiny particles called
A
atoms
17
Q
proton’s mass
A
about 1
18
Q
neutron’s mass
A
about 1
19
Q
electron’s mass
A
0
20
Q
the BLANK of an atom refers to the number of protons contained in the nucleus. The number of protons is the sole determining factor for what element the atom is
A
atomic number (Z)
21
Q
The BLANK of an atom denotes the number of protons plus the number of neutrons in the nucleus of an atom
A
mass number (A)
m= p + n
22
Q
BLANK are atoms of the same element but differ in the number of neutrons. BLANK behave chemically identically in almost every way, but sometimes differences can be detected by smell
A
isotopes
23
Q
BLANK describe where electrons are likely to exist
- a maximum of 2 electrons can fit in each atomic orbitals
A
Atomic orbitals
24
Q
Which of the following sets of elements are most likely to react in similar ways?
A
A. N, O
B. O, S
C. Ca, O
D. Mg, S
B. which are in the same column
25
BLANK represent areas of space where you are likely to find an electron
atomic orbitals
26
BLANK is always spherical
s-orbital
27
Blank has a dumbbell shape. Remember that there are three of these within a given subshell, each in different spatial interactions
p-orbitals
*peanut*
28
x p-orbital
29
z p-orbital
30
y p-orbital
31
A BLANK will always be a "clover leaf," with the exception of the one "bow tie shape."
d-orbital
32
energy level and average distance from nucleus
principal quantum number (n)
n= 1,2,3,...
ex. Ni
n=2
33
orbital shape
(s, p, d, f)
angular momentum quantum number (l)
l=0, to n-1
l=0 is the s-subshell
l=1 is the p-subshell
l=2 is the d-subshell
ex. Ni
l=0 (2-1) = 0,1... l=1
34
orbital orientation
(px, py, pz)
magnetic quantum number (ml)
ml= -1 to +1
for l=0 there is only one orientation
for l=1 there are three orientations along each of the x, y, and z axes
35
describes spin of electron
spin quantum number (ms)
ms= +1/2 and -1/2
either up or down
36
This principle states that when multiple electrons are present in an atom, electrons will always occupy the lowest available energy level.
Aufbau principle
37
This principle states that electrons will fill each subshell singly before spin-pairing. As a result, atoms will have as many unpaired electrons as possible
Hund's rule
38
This principle states that no two electrons in an atom can occupy the same space; therefore, two electrons cannot have the same four quantum numbers. Even electrons in the same orbital must have opposing spins (ms= -1/2 and +1/2)
pauli exclusion principle
39
what is the shell
same n values
40
what is the subshell
same n and l values
41
what are the orbital
same n, l, and ml values
42
although subshell energies will be different, orbitals within a subshell are usually
degenerate (equivalent energy)
43
atoms and ions are called BLANK when they have the same electron configuration
isoelectronic
- atleast one this is an ion
- for instance, adding an electron to bromine ([Ar]4s23d104p5) would form an anion (Br-) making it isoelectronic with [Kr]
44
lowest-energy state
ground state
45
highest-energy state
excited state
46
BLANK will have additional electron(s). Add them in the same order as usual
anions
47
BLANK will lose the outermost (highest n) electrons first. These are not necessarily the last electrons listed in the configuration
cations
48
if all electrons are spin-paired, the magnetic fields cancel-out, and the element is...
diamagnetic
49
if there are unpaired electrons present, the element has a net magnetic field and is...
paramagnetic
50
The BLANK model describes the possible electron excitations and relaxations for single-electron species.
Bohr model
51
electron moves to higher energy as light is
absorbed
52
electron moves to lower energy as light is
emitted
53
metals are atoms that tend to lose electrons and form positive ions (cations) to obtain a noble gas confguration
- metal atoms bond with other metal atoms or form ionic bonds with negative ions that are usually non-metals
metallic character
54
non-metals tend to form negative ions (anions) to obtain a noble gas configuration
non-metallic character
55
non-metals form BLANK with other non-metals and form ionic bonds with positive ions (cations) that are usually metals
covalent bonds
56
non-metals are BLANK of heat and electricity, non-malleable, brittle (when solid), and possess low to moderate melting points
poor conductors
57
BLANK have both metal and non-metal characteristics
metalloids
58
is the ability of an atom to attract the shared electrons in a covalent bond
electronegativity
-it increases right and up the periodic table
-remember the "FONCI" elements
59
is the net positive charge felt by an orbiting electron. It increases down and right on the periodic table. The horizontal trend is the most important. In fact, the MCAT will expect you to use the approx. zeff = #etotal - #escreening
effective nuclear charge (zeff)
**look at the column number
60
is the distance from the nucleus to the outermost electrons
atomic radius
*it increases from top to bottom of the periodic table
61
is the amount of energy required to break apart an ionic lattice. It combines two factors: change strength and atomic radius
1. the greater the charge, the greater this is
2. the closer the ions, the greater this is as well
lattice energy
E= kQ1Q2 / r .... as r decreases E increases
62
is the minimum energy required to completely remove and electron from a neutral gaseous atom
ionization energy
**byeonization E
-it generally increases right and up on the periodic table
63
Is the second ionization energy for an atom is always greater than the first
true.
-removing an electron from something already +
64
is the energy released when an electron is added to a neutral gaseous atom. Generally, it increases to the right and up the periodic table
electron affinity
- how much an atom wants to gain electrons
65
BLANK are defined as an oscillating disturbance that travels through space and time, usually accompanied by the transfer of energy
Waves
66
BLANK is inversely proportional to frequency
wavelength (nm)
67
is the number of crests that pass through a given point per second. measured in hertz (Hz = s-1)
frequency
68
is the distance a wave travels through space per unit of time. Often measured in meters per second
velocity (m/s)
69
is a type of energy. as this energy travels forward, electric and magnetic fields are created perpendicular to each other and to the direction of travel
electromagnetic radiation (light)
70
When electromagnetic radiation travel in a vacuum, its velocity is equal to the speed of light
c= f(wavelength)
where c= speed of light (3x10^8 m/s)
71
Plancks constant
6.62607 x 10^-34 Jxs
72
tells us that the exact position and momentum of such as electron cannot be measured simultaneously. In other words, the more you know about the momentum of an electron, the less you know about its exact position
Heisenberg Uncertainty Principle
p=mv
73
BLANK bonds occur within a molecule
intramolecular bonds
ex.
ionic bonds:
- metal + non-metal
metallic bonds:
74
metals are ductile (easily stretched), malleable (hammered into thin sheets), conductive (able to conduct electrons and heat), and possess a characteristic luster
metallic bonds
75
- non-metal and non-metal (similar electronegativity)
-sharing of electron pairs
- ex. O2
Covalent bonds
76
solids are composed of atoms continuously covalently bonded through the material (eg. diamond). There are no distinct molecules.
covalent networks
77
occur between identical atoms. The difference in electronegativity is exactly zero so the electrons in the covalent bond are shared equally
pure covalent bonds
ex. Cl2, N2, F2, H2
78
when
polar covalent bonds
79
2 atoms of different electronegativity form a covalent bond, the electrons are not shared equally. The results in this bond contains a dipole
polar covalent bonds
80
involve the weak interactions between molecules. When these forces are present, they determine physical properties.
intermolecular interactions
*weaker
81
what are the three main types of intermolecular forces
1. dipole-dipole forces
2. hydrogen bonding
3. dispersion (london dispersion) forces
82
when two polar molecules are near each other, they align with the positive end of one molecule close to the negative end of the other. This repeating pattern of head-to-tail alignment gives rise to attractive forces among the molecules
dipole-dipole forces
83
occurs when a highly electronegative atom with a lone pair of electrons share its nonbonding electrons with a positively polarized hydrogen atom. BLANK are simply very strong dipole-dipole interactions
hydrogen bonding
84
Two requirements for h-bond formation:
1. there must be an electron-deficient hydrogen atom to act as a hydrogen donor. Hydrogen atoms in O-H, F-H, and N-H bonds meet this requirement
2. there must be a highly electronegative atom to act as a hydrogen acceptor. O, N, and F meet this requirement
85
dispersion forces arise from the formation of an instantaneous dipole in the electron cloud of a given molecule
london dispersion forces
86
As the number of BLANK between molecules increase, the viscosity (how slowly something flows), surface tension, boiling point and melting point of that chemical compound also increase
intermolecular forces
87
what are the strengths of bonds?
ionic = metallic = covalent >> h-bonding >_ dipole-dipole > London dispersion
88
When water boils, which of the following bonds break?
O-H hydrogen bonds
*if we have intermolecular forces, then those are the ones that will break with phase changes
89
Decay: involves a nucleus becoming more stable by emitting an alpha particle
- a particle containing 2 protons and 2 neutrons
- sometimes denotes 4/2 He 2+
down 4 on atomic mass and 2 on atomic number
alpha decay
90
involves beta particles:
a neutron breaks down into a proton and electron. The electron is emitted
add 1 to proton number/ atomic number
beta decay
91
occurs when an excited nucleus relaxes. Gamma rays (high-energy photons) are emitted
*travels far; requires thick lead walls to shield
gamma decay (y)
92
Uranium 238 (238 U 92 U) undergoes alpha decay once and negative beta decay twice. What is the resulting material?
234, 92 U
93
is the amount of time it takes for half of a given sample to decay. For radioactive decay, the half-life is independent of time
half life
94
A fossil is found to have 20% of the concentration of 14C as the atmosphere (while living, things contain roughly the same concentration of 14C as the atmosphere). How old if the fossil? The half-life of 14C is 5730 years.
13305 years
94
95
Kinetic Molecular Theory of Ideal Gases
The kinetic molecular theory of ideal gases has four assumptions and a consequence: 1. Gas particles have no volume. 2. A gas particle exhibits constant, random, straight-line motion. 3. All gas collisions are elastic, meaning momentum and kinetic energy are conserved. 4. Gas particles do not experience intermolecular forces with other gas particles. Consequence: Gas particles span a distribution of speeds, where their average kinetic energy is directly proportional to the temperature of the entire sample: Ek avg oc T.
96
Measuring Pressure by a Simple Barometer or Manometer
A barometer is used to measure atmospheric pressure. It consists of a long tube that is closed at one end and inverted into a dish of mercury. The tube itself is partially filled with mercury as well. The closed end of the tube is a vacuum, so its pressure can be assumed to be zero. < P= 0 atmospheric pressure Potm = pgh, where p is the density of mercury and h is the height of the mercury in the column.
97
98
99
The Ideal Gas Law
The ideal gas law describes the relationship between volume, temperature, and pressure of an ideal gas.
100
PV = nRT
Where: • P: Pressure of the gas (atm, Torr, kPa, mmHg) • V: Volume of the gas/container (L) • n: Number of moles of the gas (mol) • T: Temperature of the gas (K) • R: Gas Constant
101
Five values of the gas constant R
• 8.314 J-mol-•K- • 0.082 L•atm-mol-K- • 8.314 L-kPa-mol K-1 • 62.3 L• Torr-mol-K- • 1.98 cal-mol-•K
102
Select which R constant to use based on the units given for pressure!
103
MCAT Tip
You must know how to convert between common units of pressure: • 1 atm = 760 Torr = 760 mmHg 1 atm = 101 kPa • 1 atm = 1 bar
104
MCAT Tip
STP = Standard Temperature and Pressure = 0°C (273.15K) and 1 bar. At STP, 1 mol of any ideal gas occupies volume of 22.4 L.
105
Examples:
How many moles of an ideal gas exist in 44.8 L container at STP? What is the volume of a container containing 1 mol of gas at 0.5 atm and 0°C?
106
total pressure
The total pressure is equal to the sum of partial pressures for each gas within the mixture.
107
partial pressure
A partial pressure of a gas is the same pressure it would exert if it were alone in the container.
108
Partial pressure is proportional to the mole fraction
Partial pressure is proportional to the mole fraction of the gas in the mixture.
109
mole fraction of gas A
To determine mole fraction of gas A: ПА XA = - ntotal Where: Ntotal = na + n +•.•.
110
partial pressure of gas A
The partial pressure of gas A would then be the mole fraction of A multiplied by the total pressure of the mixture: PA = XA • Ptotal.
111
Real Gases
Gases mostly behave ideally at high temperature and low pressure.
112
Ideal Gas
For an ideal gas, the ratio of PV/nRT (the compressibility factor, Z) is 1 at any pressure or temperature.
113
Deviation from Ideal Gas Behavior
A plot of PV/nRT for 1 mol of molecular hydrogen against increasing pressure shows deviation from ideal gas behavior.
114
Cause of Deviation from Ideality
This deviation from ideality is caused by PV being greater than nT.
115
Ideal Case Pressure-Volume Relationship
In an ideal case, increases in pressure would perfectly correspond to decreases in volume.
116
Real Gases Volume Resistance
In real gases, there is a little more resistance to decreases in volume: the molecules themselves take up.
117
ideal gases
For ideal gases, we assume that individual gas particles take up no volume.
118
pressure
Real gas particles take up space. As a result, real gases occupy volumes greater than predicted at high pressures.
119
PV = nRT
can be > 1 or < 1 depending on which is the predominant factor— intermolecular forces or molecular volume.
120
Positive deviation
due to volume
121
Negative deviation
due to attraction
122
Figure adapted from Chemistry 2e
by Paul Flowers et al and published by Rice University via OpenStax. It has a CC BY 4.0 license. Original text available for free at openstax.org
123
Lesson 3: Gases, Kinetics, and Equilibrium
For ideal gases, we assume that individual gas particles take up no volume.
124
pressure
high pressure
125
V Prep101
Real gas particles take up space. As a result, real gases occupy volumes greater than predicted at high pressures.
126
P-
can be > 1 or < 1 depending on which is the predominant factor- intermolecular forces or molecular volume.
127
Consider the following graph:
2.0 Positive deviaton N2 CHa due to volume - H2 СОг 1.5- PV nRT 1.0 0.5 Negative deviaton due to attraction 200 400 600 800 1000 Pressure (atm)
128
Figure adapted from Chemistry 2e by Paul Flowers et al and published by Rice University via OpenStax.
It has a CC BY 4.0 license. Original text available for free at openstax.org.
129
What is the value of **ART** when the volume is larger than predicted?
Greater than 1
## Footnote
This indicates that the gas is deviating from ideal behavior.
130
True or false: All gases deviate from **ideal behavior** as pressure increases.
TRUE
## Footnote
This is an important consideration in gas behavior under varying pressures.
131
At **low pressures**, how does the volume of a real gas compare to that predicted by the ideal gas law?
Less volume
## Footnote
This is represented as PVreal < PVideal.
132
At **high pressures**, how does the volume of a real gas compare to that predicted by the ideal gas law?
Greater volume
## Footnote
This is represented as PVreal > PVideal.
133
What factors influence the **deviation from ideal behavior** of gases?
* Pressure
* Temperature
## Footnote
Both pressure and temperature affect how closely a gas behaves like an ideal gas.
134
As **temperature increases**, what happens to the attractive forces between gas molecules?
Become negligible
## Footnote
At high temperatures, gas molecules move too fast to experience significant attractive forces.
135
What is the equation for the **ideal gas law**?
PV = nRT
## Footnote
This equation relates pressure, volume, number of moles, the ideal gas constant, and temperature.
136
As the concentration of **reactants decreases**, what happens to the **rate** of the reaction?
The rate usually decreases
## Footnote
This occurs because there are fewer reactant molecules present, leading to fewer opportunities for collisions.
137
In the reaction **H202(aq) → H20(l) + ½ O2(g)**, how does the average rate of formation of the products change over time?
* Decreases over time
## Footnote
Given a certain concentration of H202(aq), the average rate of formation decreases as time progresses.
138
What is the **rate law** or **rate equation**?
A mathematical expression predicting the rate of a reaction based on concentrations
## Footnote
It relates the overall rate of a reaction to initial concentrations of reactants through a rate constant (k).
139
True or false: The **orders** in the rate law can be determined from the stoichiometric coefficients.
FALSE
## Footnote
The orders must be determined experimentally.
140
For a reaction **aA + bB → cC + dD**, what is the general form of the **rate law**?
Rate = K[A]^m[B]^n
## Footnote
A and B are chemical species whose concentrations affect the rate of the reaction.
141
In the rate law, what do the **exponents m and n** represent?
* The orders of the reactants
* How large of an effect [A] and [B] have on the rate
## Footnote
The value of m gives the order with respect to A, and n gives the order with respect to B.
142
If m = 1 and n = 2, what is the **overall order** of the reaction?
3
## Footnote
The overall order is the sum of m and n.
143
For **zero-order reactions**, what is the linear plot?
Concentration vs. time
## Footnote
This indicates that the rate of reaction is constant and does not depend on the concentration of the reactants.
144
For **first-order reactions**, what is the linear plot?
In(concentration) vs. time
## Footnote
The slope of this plot is equal to -K, where K is the rate constant.
145
For **second-order reactions**, what is the linear plot?
1/(concentration) vs. time
## Footnote
The slope of this plot is equal to K, indicating a relationship between the rate of reaction and the concentration of reactants.
146
What is the slope for **zero-order reactions**?
-K
## Footnote
This reflects the constant rate of reaction irrespective of reactant concentration.
147
What is the slope for **first-order reactions**?
-K
## Footnote
This indicates that the rate of reaction decreases exponentially as the concentration decreases.
148
What is the slope for **second-order reactions**?
K
## Footnote
This indicates that the rate of reaction increases with the square of the concentration of the reactants.
149
What do **rate laws** predict?
* The rate of the reaction from initial concentrations of reactants
* Reactant (or product) concentration as a function of time
## Footnote
Rate laws are fundamental in understanding the kinetics of chemical reactions.
150
What does the **Arrhenius equation** summarize about the rate constant (k) for a chemical reaction?
k is dependent on:
* Collision frequency (z)
* Fraction of collisions with proper alignment (p)
* Fraction of collisions with sufficient kinetic energy (f)
## Footnote
The Arrhenius equation relates the rate constant to temperature and activation energy.
151
In the **collision theory**, what does 'z' represent?
Collision frequency
## Footnote
It refers to the number of molecular collisions per unit time.
152
In the **collision theory**, what does 'p' represent?
Fraction of collisions with proper alignment
## Footnote
This is necessary for breaking and forming new bonds.
153
In the **collision theory**, what does 'f' represent?
Fraction of collisions with sufficient kinetic energy
## Footnote
KE must be greater than the activation energy (Ea) for a reaction to occur.
154
What is the **activated complex** in transition state theory?
An unstable, high-energy species
## Footnote
It exists between reactants and products and corresponds to a peak on the energy graph.
155
What can the **activated complex** do?
* React to form products
* Revert to reactants
## Footnote
This indicates the dual nature of the transition state.
156
Name the factors that affect **reaction rates**.
* Reactant concentrations
* Temperature
* Catalysts
## Footnote
Each factor influences the collision frequency and the energy of the particles involved.
157
How does **temperature** affect reaction rates?
* Increases collision frequency (z)
* Increases fraction of particles colliding with E > Ea (f)
## Footnote
Higher temperatures lead to increased speed of particles.
158
What role do **catalysts** play in chemical reactions?
They raise the reaction rate without being permanently consumed
## Footnote
Catalysts provide an alternative pathway with a lower activation energy.
159
160
161
What is a **mechanism** in chemistry?
A mechanism is composed of one or more elementary reactions that describe the sequence of steps in a chemical reaction
## Footnote
The sum of the individual steps must give the overall balanced chemical reaction.
162
What is the **rate-determining step**?
The slowest step in a chemical reaction mechanism
## Footnote
It determines the rate of the overall reaction.
163
In a reaction mechanism, what does the **molecularity** refer to?
The number of reactant molecules involved in the balanced equation
## Footnote
It helps classify elementary reactions as unimolecular, bimolecular, etc.
164
True or false: The same **elementary step** may occur more than once in a mechanism.
TRUE
## Footnote
This allows for complex reactions to be described in terms of simpler steps.
165
What is the overall reaction for the mechanism: Br + H2 → HBr + H and H + Br2 → HBr + Br?
H2 + Br2 → 2 HBr
## Footnote
This is derived from summing the individual steps of the mechanism.
166
For elementary reactions, how can the **stoichiometric coefficients** be used?
To determine the order of the reaction with respect to each reactant
## Footnote
This applies only to elementary reactions.
167
What is the rate law for the first (slow) elementary step: NO3(g) + CO(g) → NO2(g) + CO2(g)?
Rate = k[NO3(g)]
## Footnote
The rate is determined by the concentration of the reactants in the rate-determining step.
168
Fill in the blank: The **overall reaction** for the elementary reaction NO2(g) + CO(g) → NO2(g) + CO2(g) is __________.
NO2(g) + CO(g) → NO2(g) + CO2(g)
## Footnote
This illustrates the process of elementary reactions.
169
Define **molecularity** in the context of a chemical reaction.
The number of molecules involved in a given elementary reaction
## Footnote
Molecularity can be unimolecular, bimolecular, or termolecular.
170
What is **unimolecular** molecularity?
One molecule dissociates
## Footnote
This type of reaction involves a single reactant molecule.
171
What is **bimolecular** molecularity?
Two molecules collide
## Footnote
This type of reaction involves the interaction of two reactant molecules.
172
What is **termolecular** molecularity?
Three molecules collide
## Footnote
This type of reaction is rare due to the low probability of three molecules colliding simultaneously.
173
What is the general form of a **differential rate law** for a reaction involving one reactant?
Rate = K[A]
## Footnote
This indicates that the rate is proportional to the concentration of a single reactant.
174
What is the general form of a **differential rate law** for a reaction involving two reactants?
Rate = k[A][B]
## Footnote
This indicates that the rate depends on the concentrations of two reactants.
175
What is the general form of a **differential rate law** for a reaction involving three reactants?
Rate = k[A][B][C]
## Footnote
This indicates that the rate depends on the concentrations of three reactants.
176
What is a key **MCAT Tip** regarding elementary reactions in a mechanism?
Elementary reactions must add up to give the overall chemical equation
## Footnote
This ensures that the proposed mechanism is consistent with the overall reaction.
177
What must not appear in the overall chemical equation or final rate law?
Intermediates
## Footnote
Intermediates are species that are formed and consumed during the reaction mechanism.
178
The reaction mechanism must explain the experimentally determined **rate law**. True or False?
TRUE
## Footnote
The mechanism should align with observed reaction rates.
179
The reaction mechanism and overall rate law must be feasible: usually _______ or _______.
uni- or bimolecular
## Footnote
This reflects the common occurrence of reactions involving one or two molecules.
180
What are **catalysts** in chemistry?
Chemical species that increase the rate of a reaction by providing an alternative pathway with lower activation energy
## Footnote
Catalysts are not consumed in the reaction and are not part of the overall stoichiometry.
181
True or false: A **catalyst** is consumed during the reaction.
FALSE
## Footnote
Catalysts change the mechanism but remain unchanged after the reaction.
182
What is an example of an **uncatalyzed reaction**?
O3 → O2 + O
## Footnote
In this reaction, there is no catalyst, and atomic O is an intermediate.
183
What is an example of a **catalyzed reaction**?
NO + O3 → NO2 + O2
## Footnote
NO is the catalyst in this reaction.
184
What are the two types of **catalysts** based on their phases?
* Homogeneous
* Heterogeneous
## Footnote
Homogeneous catalysts are in the same phase as the reactants, while heterogeneous catalysts are in a different phase.
185
How do **catalysts** affect the activation energy of reactions?
They decrease the activation energy of both the forward and reverse reactions
## Footnote
This allows the reaction to reach equilibrium faster without changing the final state.
186
In a reaction diagram, what do the labels **Ea (uncatalyzed)** and **Ea (catalyzed)** represent?
Activation energy for uncatalyzed and catalyzed reactions
## Footnote
Catalysts lower the activation energy, allowing for faster reactions.
187
What does the **equilibrium constant** provide?
A quantitative ratio of products to reactants
## Footnote
It is affected by pressure, changes in volume, or the presence of a catalyst.
188
Is the **equilibrium constant** influenced by temperature?
Yes, it is temperature-dependent
## Footnote
Changes in temperature can affect the value of the equilibrium constant.
189
For the reaction aA + bB = CC + dD, what is the expression for **Kc**?
Kc = [C]^c [D]^d / [A]^a [B]^b
## Footnote
This expression represents the concentrations of the reactants and products.
190
What types of species are **not included** in equilibrium expressions?
* Pure liquids
* Solids
## Footnote
Their concentrations do not change significantly during a chemical reaction.
191
For gases, the equilibrium constant can be expressed in **partial pressures**. What is the symbol used?
Kp
## Footnote
Kp is used instead of Kc when dealing with partial pressures (atm or bar).
192
What is the expression for **Kp**?
Kp = (P_C^c * P_D^d) / (P_A^a * P_B^b)
## Footnote
This expression is based on the partial pressures of the gases involved.
193
Name two types of **equilibrium constants** encountered in MCAT preparation.
* Kp: gas-phase species based on partial pressures
* Kc: aqueous-phase or gas-phase species based on concentrations
## Footnote
These constants help in understanding different chemical equilibria.
194
What do **Ka/Kb** represent?
Acid/base dissociation constants
## Footnote
They indicate the strength of acids and bases in solution.
195
What is **Kw**?
Dissociation constant for the self-ionization of water
## Footnote
Kw is a crucial constant in acid-base chemistry.
196
What does **Ksp** stand for?
Solubility product for solutes in solution
## Footnote
Ksp helps determine the solubility of ionic compounds.
197
True or false: Gases can only be defined by **partial pressure**.
FALSE
## Footnote
Gases can also be defined by concentration in moles/L, and Kc is often used unless partial pressure is specified.
198
What is the **reaction quotient (Q)**?
An expression that combines all the instantaneous concentrations or partial pressures in a system, regardless of equilibrium
## Footnote
Q is calculated in the same way as K.
199
How is **Q** calculated for the reaction 2NO(g) + O2(g) = 2NO2(g)?
Q = [NO2]^2 / ([NO]^2 * [O2])
## Footnote
This formula represents the concentrations of the reactants and products.
200
What does it indicate if **Q < K**?
Q must increase as the system moves toward equilibrium
## Footnote
This means the reaction will proceed in the forward direction.
201
What does it indicate if **Q = K**?
The system is at equilibrium
## Footnote
No net change occurs in the concentrations of reactants and products.
202
What does it indicate if **Q > K**?
Q must decrease as the system moves toward equilibrium
## Footnote
This means the reaction will proceed in the reverse direction.
203
What is the equilibrium constant (Kp) for the reaction 2NOCI(g) = 2NO(g) + Cl2(g) at 500°C?
4.4 × 10^4
## Footnote
This value is used to determine the direction of the reaction based on partial pressures.
204
If the partial pressure of each substance in the reacting mixture is 0.100 atm, in which direction will the reaction proceed?
Towards the reactants
## Footnote
This conclusion is drawn from comparing Q to K.
205
What does **Le Chatelier's Principle** state?
When a system is disturbed from a state of equilibrium, it will shift the position of equilibrium in the direction that reduces the effect of the disturbance
## Footnote
This principle applies to changes in concentration, temperature, and pressure.
206
In the reaction **A(g) + B(g) = C(g) + heat**, what happens if the concentration of **A** is increased?
Shift to the right
## Footnote
Increasing the concentration of a reactant drives the reaction forward to produce more products.
207
In the reaction **A(g) + B(g) = C(g) + heat**, what happens if the concentration of **C** is decreased?
Shift to the right
## Footnote
Decreasing the concentration of a product drives the reaction forward to produce more products.
208
In the reaction **A(g) + B(g) = C(g) + heat**, how does increasing the **temperature** affect the equilibrium?
Shift to the left
## Footnote
Increasing temperature favors the endothermic direction, which is the reverse in this case.
209
In the reaction **A(g) + B(g) = C(g) + heat**, what happens if the **volume** of the container is increased?
Shift to the left
## Footnote
Increasing volume decreases pressure, favoring the side with more moles of gas.
210
What is the effect of **temperature** on the value of **K**?
Temperature is the only variable that causes a change in the value of K
## Footnote
The effect of temperature on K depends on the enthalpy change of the reaction.
211
What happens to the equilibrium position if the **volume decreases**?
Shift towards the side with fewer moles of gas
## Footnote
This shift occurs to decrease pressure.
212
True or false: **Changes in pressure or volume** affect the value of K.
FALSE
## Footnote
Changes in pressure or volume do not change the value of K.
213
In endothermic reactions, how should heat be viewed?
As a reactant
## Footnote
This perspective is crucial for applying Le Chatelier's Principle.
214
In exothermic reactions, how should heat be viewed?
As a product
## Footnote
This perspective is crucial for applying Le Chatelier's Principle.
215
Solubility
The solubility of a species describes the maximum amount that will dissolve into a given volume of solvent at a given temperature.
216
Solubility expression
Solubility is expressed as either moles per volume or as mass per volume.
217
Example of solubility
For instance, if the solubility of silver phosphate (AgPOq) in water at 298 K was 0.0075 gL'; this implies that 0.0075 g or 1.8 × 10-5 moles of AggPO completely dissolve in 1 L of water at 298 K.
218
Equilibrium in dissolution
When a solid dissolves in a liquid, equilibrium is reached between the aqueous ions in solution and the solid.
219
Equilibrium constant
As with any equilibrium, an equilibrium constant for the dissociation reaction can be derived (NOTE: Because solids do not appear in equilibrium constant expressions, this K value will only have a numerator).
220
Solubility product (Ksp)
The solubility product, Ks, describes the equilibrium between a solid solute and its dissociated ions in a saturated solution at a given temperature.
221
General form of Ksp
Generally, for ABs) = aA(ag) + bB(aq) Ksp = [A]°[B]
222
Ksp and solubility relationship
Solubility products can vary over many orders of magnitude. In general, the larger the value of Ksp, the greater the concentration of ions at equilibrium (i.e., greater solubility).
223
Ksp of AgI
At 298 K: Ksp = 8.5 × 10-17
224
Ksp of CaCO3
At 298 K: Ksp = 9.8 × 10-9
225
Ksp of NaCl
At 298 K: Ksp = 6.2
226
Ksp expression setup
When setting up a Ksp expression, remember that solids and liquids are never included.
227
What principle is the **common ion effect** based on?
Le Chatelier's principle
## Footnote
The common ion effect can only be used for saturated (at equilibrium) solutions.
228
The common ion effect causes the equilibrium to shift towards the **reactants** when _______ ions are already present in solution.
I⁻
## Footnote
This shift occurs more than expected from the Ksp value.
229
What is the **Ksp** value for the reaction involving Pb at 298 K?
8.7 × 10⁻⁸
## Footnote
This value indicates the solubility product constant for lead.
230
Name a source of **cations** that can be added as salts in the context of the common ion effect.
* NH₄⁺
* H⁺
## Footnote
These cations can influence the solubility equilibrium.
231
Name a source of **anions** that can be added as salts in the context of the common ion effect.
* Cl⁻
* CN⁻
* I⁻
## Footnote
These anions can also affect the solubility equilibrium.
232
True or false: The common ion effect can be applied to solutions that are not at equilibrium.
FALSE
## Footnote
The common ion effect is applicable only to saturated solutions at equilibrium.
233
What happens to ions with **similar characteristics** in the context of the common ion effect?
They don't want to dissolve
## Footnote
This is due to the shift in equilibrium caused by the presence of common ions.
234
What do ions with **unlike properties** do in the context of the common ion effect?
They dissolve
## Footnote
Ions with unlike properties can still dissolve despite the presence of common ions.
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