Inorganic Chemistry Flashcards
1
Q
Atomic mass
A
= sum of masses of protons and neutrons
2
Q
Isotopes
A
Different number of neutrons (i.e. same element with a different atomic mass)
3
Q
Atomic weight
A
The weighted average of the naturally occurring isotopes of an element
4
Q
Planck relation (frequency) (E)
A
= h f
5
Q
Angular momentum of an electron (Bohr model) (L)
A
= (n h) / (2 pi)
6
Q
Energy of an electron (Bohr model) (E)
A
= - Rh / n^2
7
Q
Planck relation (wavelength) (E)
A
= h c / lambda
8
Q
Energy of electron transition (Bohr model) (E)
A
= - Rh ( 1 / n0^2 - 1 / n^2)
9
Q
Maximum number of electrons within a shell
A
= 2 n^2
10
Q
Maximum number of electrons within a sub-shell
A
= 4 l + 2
11
Q
Mass of proton
A
1 amu
12
Q
Mass of neutron
A
1 amu
13
Q
Atomic number
A
The number of protons
14
Q
Mass number
A
The number of protons and neutrons
15
Q
The isotopes of hydrogen
A
Protium, deuterium, tritium
16
Q
Quantum
A
The energy difference between energy levels
17
Q
Quantization
A
There is not an infinite range of energy levels available to an electron; electrons can exist only at certain energy levels
18
Q
Does the energy level of an electron increase or decrease when it is farther from the nucleus?
A
Increase
19
Q
Heisenberg uncertainty principle
A
It is impossible to know both an electrons position and its momentum exactly at the same time
20
Q
Quantum numbers
A
Describe an electron in an atom
21
Q
What are the quantum numbers?
A
Principal quantum number (n), azimuthal quantum number (l), magnetic quantum number (ml), spin quantum number (ms)
22
Q
Principal quantum number (n)
A
Describes the average energy of a shell
23
Q
Azimuthal quantum number (l)
A
Describes the subshells within a given principal energy level (s, p, d and f)
must be < n - 1
24
Q
Magnetic quantum number (ml)
A
Specifies the particular orbital within a sub-shell where an electron is likely to be found at a given moment in time
must be between - l and + l
25
Spin quantum number (ms)
Indicates the spin orientation (+/- 1/2) of an electron in an orbital
26
n + l rule
Electrons fill the principal energy levels and subshells according to increasing energy
27
Hund's rule
Sub-shells with multiple orbitals (p, d and f) fill electrons so that every orbital in a sub-shell gets one electron before any of them gets a second
28
Paramagnetic materials
Have unpaired electrons that align with magnetic fields, attracting the material to a magnet
29
Diamagnetic materials
Have all paired electrons, which cannot easily be realigned; they are repelled by magnets
30
Where are the valence electrons in representative elements (Groups 1, 2, 13 to 18)?
s and p orbitals
31
Where are the valence electrons in transition elements?
s and either d or f orbitals
32
Effective nuclear charge (Zeff)
The strength with which protons in the nucleus can pull on electrons
33
In which direction does Zeff increase on the periodic table?
Left to right
34
What happens to valence electrons as the principal energy level, n, increases from top to bottom in a group?
They become increasingly separated from the nucleus
35
In which direction does the principal energy level, n, increase on the periodic table?
Top to bottom
36
In which direction does the atomic number increase on the periodic table?
Left to right and top to bottom
37
Ionic radius
The size of charged species
38
Are cations smaller or larger than their corresponding neutral atom?
Smaller
39
Are anions smaller or larger than their corresponding neural atom?
Larger
40
Ionization energy
The amount of energy necessary to remove an electron from the valence shell of a gaseous species
41
In which direction does ionization energy increase on the periodic table?
Left to right and bottom to top
42
Electron affinity
The amount of energy released when a gaseous species gains an electron in its valence shell
43
In which direction does electron affinity increase on the periodic table?
Left to right and bottom to top
44
Electronegativity
A measure of the attractive force of the nucleus for electrons within a bond
45
In which direction does electronegativity increase on the periodic table?
Left to right and bottom to top
46
What is the oxidation state of alkali metals?
+1
47
Do alkali metals prefer to gain or lose electrons?
Lose
48
Which metals are the most reactive on the periodic table?
Alkali metals and alkaline earth metals
49
What is the oxidation state of alkaline earth metals?
+2
50
Do alkaline earth metals prefer to gain or lose electrons?
Lose
51
What is the oxidation state of chalcogens?
-2 or +6 (depending on whether they are non-metals or metals, respectively)
52
What is the oxidation state of halogens?
-1
53
Do halogens prefer to gain or lose electrons?
Gain
54
Which group has the highest electronegativity on the periodic table?
Halogens
55
Noble gases
Have fully filled valence shell in their standard state and prefer not to give up or take on additional electrons
They have very high ionization energies and virtually nonexistent electronegativities and electron affinities
56
Transition metal features
Take on multiple oxidation states
| Can form hydration complexes with water, increasing their solubility in water
57
Elements that form incomplete octets
H, He, Li, Be, B
58
Elements that form expanded octets
Any element in period 3 or greater
59
Characteristics of ionic bonds
- High melting and boiling points due to electrostatic attractions
- Solubility of ions in water and other polar solvents due to interactions with polar solvents
- Good conductors of heat and electricity
- Have crystal lattice arrangements to minimize repulsive forces
- Large electronegativity differences between ions (delta EN > 1.7)
60
What is bond strength defined by?
The electrostatic attraction between nuclei and electrons
61
Do multiple bonds (higher bond order) increase or decrease bond strength?
Increase
62
Does a longer bond length result in a stronger bond or a weaker bond?
Weaker bond
63
Bond energy
The minimum amount of energy needed to break a bond
64
Does a stronger bond have higher bond energy or lower bond energy?
Higher
65
delta EN in non-polar covalent bonds
= between 0 and 0.5
66
delta EN in polar covalent bonds
= between 0.5 and 1.7
67
delta EN in ionic bonds
= 1.7 or higher
68
Intermolecular forces
Hydrogen bonding, dipole-dipole interactions and London dispersion forces
They are electrostatic attractions between molecules
They are weaker than covalent bonds (which are weaker than ionic bonds)
69
What is the strongest intermolecular force?
Hydrogen bonding
70
What is the weakest intermolecular force?
London dispersion forces
71
Dipole-dipole interactions
A dipole consists of a segment of a molecule with partial positive and partial negative regions. The positive need of one molecule is attracted to the negative end of another molecule, and vice-versa.
72
When does hydrogen bonding occur?
A molecule must have a hydrogen bonded to either N, O or F
73
Octet rule
Elements will be most stable with eight valence electrons
74
Bond order
Refers to whether a covalent bond is a single bond, double bond or triple bond
75
Where do non-polar covalent bonds form?
In molecules in which both arms have exactly the same electronegativity
76
When do polar covalent bonds form?
When there is a significant difference in electronegativity, but not enough to transfer electrons and form an ionic bond
77
When do coordinate covalent bonds form?
When a single atom provides both bonding electrons while the other atom does not contribute any
Common in Lewis acid-base chemistry
78
When do resonance structures appear?
When a molecule has a pi system of electrons
| Represent all the possible configurations of electrons (stable and unstable) that contribute to the overall structure
79
Valence shell electron pair repulsion (VESPER) theory
Predicts the three-dimensional molecular geometry of covalently bonded molecules
Electrons arrange themselves to be as far apart as possible from each other in three-dimensional space
80
Do non-bonding electrons exert more or less repulsion on each other than bonding electrons?
More repulsion because they reside closer to the nucleus
81
Electronic geometry
Refers to the position of all electrons in a molecule, whether bonding or nonbonding
82
Molecular geometry
Refers to the position of only the bonding pairs of electrons in a molecule
83
What kind of bonds do polar molecules have?
Polar bonds
84
What kind of bonds do non-polar molecules have?
Polar bonds or non-polar bonds
85
What are sigma bonds results of?
The result of head-to-head overlap
86
What are pi bonds results of?
The result of the overlap of two parallel electron cloud densities
87
London dispersion forces
Present in all atoms and molecules
88
What happens to London dispersion forces when an atom or molecule's size increases?
Increase
89
Where are dipole-dipole interactions most prominent?
Liquids and solids
90
Dipole moment (p)
= q d
```
q = charge
d = distance between charges
```
91
Formal charge
= N valence electrons - N non-bonding - (1/2 N bonding)
92
Molarity (M)
= moles of solute / liters of solution (mol/L)
Used for rate laws, the law of mass action, osmotic pressure, pH and pOH and the Nernst equation
93
How are molecular formulas and empirical formulas similar?
Both contain the same elements in the same ratios
94
How are molecular formulas and empirical formulas different?
The molecular formula gives the actual number of atoms of each element in the compound
The empirical formula gives only the ratio and therefore may or may not give the actual number of atoms
95
Molecular formula
Gives the actual number of atoms of each element in the compound
96
Empirical formula
Gives only the ratio and therefore may or may not give the actual number of atoms
97
How to find the empirical formula when we only have the percent mass of the elements in the molecule?
1. Assume the percentages are the total grams of each element (e.g. 40.9% C = 40.9 g C)
2. Convert grams to moles (e.g. 40.9 g C = 3.4 mol C)
3. Once that is done for all the compounds, divide all the moles by the smallest mole value obtained (e.g. C 3.4 mol C, 4.6 mol H, 3.4 mol O => 1 mol C, 1.33 mol H, 1 mol O)
4. Multiply the mole values obtained by whole numbers, until all the mole values are whole numbers (e.g. multiply by 2 => 2 mol C, 2.66 mol H, 2 mol O // multiply by 3 => 3 mol C, 3.99 mol H [~ 4 mol H], 3 mol O)
Answer C3H4O3
98
How to find the empirical formula when both the percent mass of the elements in the molecule and the molar mass of the whole molecule are given?
1. Multiply the molar mass by the percentages given to find the mass of each element present in 1 mole of the molecule (e.g. [40.9% C * molar mass of whole molecule] / 12 g/mol C = 9 mol)
2. Divide all the mole values you get by the common denominator (e.g. 9 mol C, 12 mol H, 9 mol O => 3 mol C, 4 mol H, 3 mol O)
Answer C3H4O3
99
What happens when Zn(NO3)2 dissolves in (NH4)2S?
Ammonium swaps 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.
100
2 H2 + O2 -->
2 H2O
101
What type of reaction is 2 H2 + O2 --> 2 H2O
Combination
102
Al(OH)3 + H3PO4 -->
3 H2O + AlPO4
103
What type of reaction is Al(OH)3 + H3PO4 --> 3 H2O + AlPO4
Neutralization (a type of double-displacement)
104
2 H2O --(electricity)-->
2 H2 + O2
105
What type of reaction is 2 H2O --(electricity)--> 2 H2 + O2
Decomposition
106
NaNO3 + CuOH -->
NaOH + CuNO3
107
What type of reaction is NaNO3 + CuOH --> NaOH + CuNO3
Double-displacement (metathesis)
108
Zn + AgCl -->
ZnCl + Ag
109
What type of reaction is Zn + AgCl --> ZnCl + Ag
Single-displacement
110
Is HCl an electrolyte or a non-electrolyte?
Electrolyte
111
Is sucrose an electrolyte or a non-electrolyte?
Non-electrolyte
112
Is MgBr2 an electrolyte or a non-electrolyte?
Electrolyte
113
Is CH4 an electrolyte or a non-electrolyte?
Non-electrolyte
114
Phosphate
PO4^3-
115
Hypochlorite
ClO-
116
Ammonium
NH4+
117
Phosphide
P^3-
118
Bicarbonate
HCO3-
119
Nitrite
NO2-
120
Chromium (II)
Cr^2+
121
Moles from mass
= mass of sample / molar mass
122
Gram equivalent weight (GEW)
= molar mass / n
A measure of the mass of a substance that can donate one equivalent of the species of interest
123
Equivalents from mass
= mass of compound / gram equivalent weight
124
Molarity from normality
= normality / n
125
Percent composition
= (mass of element in formula / molar mass) 100
Determine the mass of the individual element and divide by the molar mass of the compound
126
Percent yield
= (actual yield / theoretical yield) 100
127
Molecular weight
The mass (in amu) of the constituent atoms in a compound as indicated by the molecular formula
128
Molar mass
The mass of one mole of a compound (g/mol)
129
Avogadro's number
6.022e23
130
Normality
The ratio of equivalents per liter
The molarity of the species of interest
= molarity * n
n = equivalents of H+ or OH-
Used for acid-base and oxidation-reduction reactions
131
How is normality related to molarity?
By multiplying the molarity by the number of equivalents present per mole of compound
132
Equivalents
Moles of the species of interest
133
Law of constant composition
Any pure sample of a compound will contain the same elements in the same mass ratio
134
Combination reaction
Occurs when two or more reactants combine to form one product
135
Decomposition reaction
Occurs when one reactant is chemically broken down into two or more products
136
Combustion reaction
Occurs when a fuel and an oxidant (typically oxygen) react, forming the products water and carbon dioxide (if the fuel is a hydrocarbon)
137
Displacement reaction
Occurs when one or more atoms or ions of one compound are replaced with one or more atoms or ions of another compound
138
Single-displacement reaction
Occurs when an ion of one compound is replaced with another element
139
Double-displacement reaction
Occurs when elements from two different compound trade places with each other to form two new compounds
140
Neutralization reaction
Occurs when an acid reacts with a base to form a salt and usually water
141
Limiting reagent
The reactant that will be consumed first in a chemical reaction
142
Theoretical yield
The amount of product generated if all of the limiting reactant is consumed with no side reactions
143
What are Roman numerals in ion nomenclature used for?
Non-representative elements to denote ionic charge
144
What is -ous in ion nomenclature used for?
Indicate lesser charge
145
What is -ic in ion nomenclature used for?
Indicate greater charge
146
What is -ide in ion nomenclature used for?
The ending of monatomic anions
147
Oxyanions suffixes
-ite and -ate
148
What is -ite in ion nomenclature used for?
Contain a lesser amount of oxygen
149
What is -ate in ion nomenclature used for?
Contain a greater amount of oxygen
150
Oxyanions prefixes
Hypo- and per-
151
What is hypo- in ion nomenclature used for?
Contain the fewest amount of oxygen
152
What is per- in ion nomenclature used for?
Contain the highest amount of oxygen
153
Hydrogen-donating polyatomic ions
Hydrogen/bi- and dihydrogen
154
What is hydrogen/bi- in ion nomenclature used for?
Polyatomic ion that donates one hydrogen
155
What is dihydrogen in ion nomenclature used for?
Polyatomic ion that donates two hydrogens
156
Electrolytes
Contain equivalents of ions from molecules that dissociate in solution
157
What determines the strength of an electrolyte?
Its degree of dissociation (i.e. solvation)
158
Ionic compounds
1. They are composed of atoms held together by ionic bonds
2. They associate charged particles with large differences in electronegativity
3. Form large arrays of ions in crystalline solids
4. Measured with formula weights
5. Electrons are donated from the less electronegative atom to the more electronegative atom
159
Chalcogens
Group 16 elements
160
Do noble gases have high or low ionization energies?
Very high
161
Do noble gases have high or low electronegativities?
Virtually nonexistent
162
Do noble gases have high or low electron affinities?
Virtually nonexistent
163
Do ionic bonds have high or low melting points?
High
164
Do ionic bonds have high or low boiling points?
High
165
Why do ionic bonds have high melting and boiling points?
Due to their electrostatic attractions
166
Do ionic compounds interact with polar solvents?
Yes
167
Are ionic compound good conductors of heat and electricity?
Yes
168
How do ionic bonds minimize repulsive forces?
Because they have crystal lattice arrangements
169
Do ionic bonds have large electronegativity differences between ions?
Yes, EN > 1.7
170
Rate-determining step
The slowest step of a reaction
Determines the overall rate of the reaction because the reaction can only proceed as fast as the rate at which this step occurs
171
Activation energy
The minimum energy needed for a chemical reaction to occur
172
Transition state theory of chemical kinetics
States that molecules form a transition state or activated complex during a reaction in which the old bonds are partially dissociated and the new bonds are partially formed
Requires a certain activation energy to be overcome in order for a reaction to occur (therefore not all reactions will occur)
Focuses on forming a high-energy activated complex that can then proceed forward or backward, forming the products or reverting to the reactants, respectively
173
Collision theory of chemical kinetics
States that a reaction rate is proportional to the number of effective collisions between the reacting molecules
Requires a certain activation energy to be overcome in order for a reaction to occur (therefore not all reactions will occur)
Focuses on the energy and orientation of reactants, and considers each potential reaction to be "all-or-nothing" (either there is enough energy to form the products, to there is not).
174
What happens to a zero-order reaction if the temperature is lowered?
The reaction rate decreases
175
What happens to a first-order reaction if the temperature is lowered?
The reaction rate decreases
176
What happens to a second-order reaction if the temperature is lowered?
The reaction rate decreases
177
What happens to a zero-order reaction if all the reactants' concentrations doubled?
The reaction rate is unaffected
178
What happens to a first-order reaction if all the reactants' concentrations doubled?
The reaction rate is doubled
179
What happens to a second-order reaction if all the reactants' concentrations doubled?
The reaction rate is quadrupled
180
What happens to a zero-order reaction if a catalyst was added?
The reaction rate increases
181
What happens to a first-order reaction if a catalyst was added?
The reaction rate increases
182
What happens to a second-order reaction if a catalyst was added?
The reaction rate increases
183
How to determine the rate law of a reaction?
1. Compare two sets of data in which only one of the reactant's concentration has changed
2. Rate after concentration change / rate before concentration change = (rate of second concentration / rate of first concentration) ^ x
3. Solve for x
Answer = x is the order of that reactant
Repeat for the rest of the reactants and then put it all in rate = k [A]^x [B]^y ...
184
Gibbs free energy (delta G)
Determines whether or not a reaction is spontaneous
| Is derived from both enthalpy and entropy values for a given system
185
Intermediates
Molecules that exist within the course of a reaction, but are neither reactants nor products overall
186
How can a collision be effective?
Molecules must be in the proper orientation and have sufficient kinetic energy to excel the activation energy
187
Arrhenius equation
k = A e^(-Er / [RT])
188
What is the highest point on a free energy reaction diagram?
Transition state
189
How does changing the medium affect the rate of a reaction?
It can increase or decrease the rate of the reaction, depending on how the reactants interact with the medium
190
Homogeneous catalysts
Are in the same phase as the reactants
191
Heterogenous catalysts
Are in a different phase from the reactions
192
Reaction rate
Measured in terms of the rate of disappearance of a reactant or appearance of a product
193
Rate law
rate = k [A]^x [B]^y
194
Do rate orders match stoichiometric coefficients?
Not necessarily
195
How are rate laws determined?
Through experimentation
196
Rate order
The sum of all individual rate orders in the rate law
197
Zero-order reactions
Have a constant rate that does not depend on the concentration of reactant
198
How can the rate of a zero-order reaction be affected?
Changing the temperature or adding a catalyst
199
What is the shape of a concentration v. time curve of a zero-order reaction?
Linear
200
What is the slope of a concentration v. time curve of a zero-order reaction?
-k
201
First-order reactions
Have a non constant rate that depends on the concentration of a reactant
202
What is the shape of a concentration v. time curve of a first-order reaction?
Nonlinear
203
What is the slope of a concentration v. time curve of a first-order reaction?
ln [A] v. time
| slope = -k
204
What does a first-order reaction depend on?
Concentration of a reactant
205
Second-order reactions
Have a non constant rate that depends on the concentration of a reactant
206
What does a first-order reaction depend on?
Concentration of a reactant
207
What is the shape of a concentration v. time curve of a second-order reaction?
Nonlinear
208
What is the slope of a concentration v. time curve of a second-order reaction?
1/[A] v. time
| Slope = k
209
Broken-order reactions
Reactions with nonifnteger orders
210
Mixed-order reactions
Reactions that have a rate order that changes over time
211
Collision theory equation
Rate = Z x f
212
Definition of rate
Rate = - (delta [A]) / (a delta t) = - (delta [B]) / (b delta t) = (delta [C]) / (c delta t) = (delta [D]) / (d delta t)
213
Radioactive decay [A]t
= [A]o e^(-k t)
214
Is a non-spontaneous reaction endergonic or exergonic?
Endergonic
215
Is an endergonic reaction reaction spontaneous?
No
216
Endergonic (non-spontaneous) reaction
The activation energy of the forward reaction is greater than the activation energy of the reverse reaction
The products have a higher free energy than the reactants
217
Are catalysts used up in a reaction?
No
218
How do catalysts increase the rate of the reaction?
By lowering the activation energy
219
Do catalysts alter the thermodynamics of the reaction?
No
220
What is the direction of the reaction when Q = Keq?
The reaction is at equilibrium, no net reaction
221
Q = Keq
Reaction is at dynamic equilibrium
222
What is the Gibbs free energy of a reaction when Q = Keq?
0
223
What is the direction of the reaction when Q > Keq?
Reaction proceeds towards the reactants
224
What is the Gibbs free energy of a reaction when Q > Keq?
Positive
225
What is the direction of the reaction when Q < Keq?
Reaction proceeds towards the products
226
What is the Gibbs free energy of a reaction when Q < Keq?
Negative
227
Kp
= Partial pressure of products in gas phase / Partial pressure of reactants in gas phase
228
Kc
= Concentration of products in aqueous phase / Concentration of reactants in aqueous phase
229
What conditions favor the formation of a kinetic product?
Low temperatures with low heat transfer
230
What conditions favor the formation of a thermodynamic product?
High temperatures with high heat transfer
231
On a reaction coordinate diagram, how would the kinetic pathway appear as compared to the thermodynamic pathway?
Kinetic pathways require a smaller gain in free energy to reach the transition state. They also have a higher free energy of the products, with a smaller difference in free energy between the transition state and the products.
232
Equilibrium constant (Keq)
= Concentrations of products / concentrations of reactants = ([C]^c [D]^d) / ([A]^a [B]^b)
233
Reaction quotient (Qc)
= Concentrations of products / concentrations of reactants = ([C]^c [D]^d) / ([A]^a [B]^b)
234
Reversible reactions
Eventually reach a state in which energy is minimized and entropy is maximized
235
Dynamic chemical equilibria
The reactions are still occurring, just at a constant rate
The concentrations of reactants and products remain constant because the rate of the forward reaction equals the rate of the reverse reaction
236
Why are the concentrations of reactants and products constant in dynamic chemical equilibria?
Because the rate of the forward reaction equals the rate of the reverse reaction
237
Law of mass action
Gives the expression for the equilibrium constant Keq
238
What is the difference between the reaction quotient (Q) and the equilibrium constant (Keq)?
They use the same formula, but the reaction quotient can be calculated at any concentrations of reactions and products (i.e. at any point in the reaction). The equilibrium constant can only be calculated at equilibrium.
239
Does the equilibrium constant (Keq) change when the temperature changes?
Yes
240
Do solids appear in the law of mass action?
No
241
Do liquids appear in the law of mass action?
No
242
Do gases appear in the law of mass action?
Yes
243
Do aqueous species appear in the law of mass action?
Yes
244
Le Châtelier's principle
When a chemical system experiences a stress, it will react so as to restore equilibrium
245
Types of stresses applied to a system
Changes in concentration, pressure, volume and temperature
246
Will increasing the concentration of reactants shift the reaction to the right or to the left?
Right
247
Will decreasing the concentration of products shift the reaction to the right or to the left?
Right
248
Will increasing the concentration of products shift the reaction to the right or to the left?
Left
249
Will decreasing the concentration of reactants shift the reaction to the right or to the left?
Left
250
Will Increasing the pressure on a gaseous system increase or decrease its volume?
Decrease
251
Will increasing the pressure on a gaseous system (decreasing its volume) shift the reaction towards the side with the fewer moles of gas or the side of the more moles or gas?
Fewer
252
Will decreasing the pressure on a gaseous system increase or decrease its volume?
Increase
253
Will decreasing the pressure on a gaseous system (increasing its volume) shift the reaction towards the side with the fewer moles of gas or the side of the more moles or gas?
More
254
Will increasing the temperature of an endothermic reaction shift the reaction to the right or to the left?
Right
255
Will decreasing the temperature of an exothermic reaction shift the reaction to the right or to the left?
Right
256
Will decreasing the temperature of an endothermic reaction shift the reaction to the right or to the left?
Left
257
Will increasing the temperature of an exothermic reaction shift the reaction to the right or to the left?
Left
258
Kinetic products
Higher in free energy than thermodynamic products
Can form in lower temperatures
Form more quickly than thermodynamic products
259
Which product is higher in free energy: kinetic or thermodynamic?
Kinetic
260
Which product can form in low temperatures: kinetic or thermodynamic?
Kinetic
261
Which product can form faster: kinetic or thermodynamic?
Kinetic
262
Thermodynamic products
Lower in free energy than kinetic products
More stable than kinetic products
Slower in formation than kinetic products
More spontaneous than kinetic products
Have a more negative Gibbs free energy than kinetic products
263
Which product is lower in free energy: kinetic or thermodynamic?
Thermodynamic
264
Which product is more stable: kinetic or thermodynamic?
Thermodynamic
265
Which product forms slower: kinetic or thermodynamic?
Thermodynamic
266
Which product forms more spontaneously: kinetic or thermodynamic?
Thermodynamic
267
Which product has a more negative Gibbs free energy: kinetic or thermodynamic?
Thermodynamic
268
Does a reaction in equilibrium have equal forward and reverse reactions?
Yes
269
Kc >> 1
The equilibrium mixture will favor products over reactants
270
Ka
The ratio of products to reactants with each species raised to its stoichiometric coefficient
271
Ka > 10^-7
The compound contains more H+ than HA- at equilibrium, which makes it an acid
272
Keq of the reverse reaction
= (Keq of the forward reaction)^-1
273
Negative delta H
Exothermic reaction
274
A person snaps an ice pack and places it on his or her leg. In terms of energy transfer, what would be considered the system and what would be the surroundings in this scenario?
System: icepack
Surroundings: the person
275
Isothermal process
Temperature is constant
276
Adiabatic process
No heat is flowing into or out of the system
277
Isobaric process
Pressure is constant
278
Isovolumetric (isochoric)
Volume is constant
279
Delta U in an isothermal process
= 0
280
Q in an isothermal process
= W
281
W in an isothermal process
= Q
282
Q in an adiabatic process
= 0
283
Delta U in an adiabatic process
= - W
284
P-V graph in an isobaric process
Flat line
285
W in an isochoric process
= 0
286
Delta U in an isochoric process
= Q
287
Standard conditions
```
T = 25 C (298 K)
P = 1 atm
Concentration = 1 M
```
288
When are standard conditions used for calculation?
Kinetics, equilibrium and thermodynamics calculations use standard conditions
289
State function
Properties of a system at equilibrium
Independent of the path taken to achieve equilibrium
State functions may depend on one another
290
Process function
Define the path between equilibrium states
291
Common state functions
Pressure (P), density (p), temperature (T), volume (V), enthalpy (H), internal energy (U), Gibbs free energy (G) and entropy (S)
292
Triple point
The combination of temperature and pressure at which all three phases are in equilibrium
293
Critical point
The temperature and pressure above which liquid and gas phases are indistinguishable and the heat of vaporization is zero
294
Common process functions
Heat (Q) and work (W)
295
Temperature
Indirect measure of the thermal content of a system that looks at average kinetic energy of particles in a sample
296
Heat
The thermal energy transferred between objects as a result of differences in their temperatures
297
Specific heat
The energy required to raise the temperature of one gram of a substance by one degree Celsius
298
Heat capacity
The product of mass and specific heat and is the energy required to raise any given amount of a substance one degree Celsius
299
Constant-volume calorimetry (bomb calorimeter)
Heats of certain reactions (like combustion) can be measured indirectly by assessing temperature change in a water bath around the reaction vessel
300
Constant-pressure calorimetry (coffee cup calorimeter)
Exposed to constant (atmospheric) pressure. As the reaction proceeds, the temperature of the contents is measured to determine the heat of the reaction
301
What is the specific heat of water (in calories)?
1 cal / g dot K
302
Endothermic process
An increase in heat content of a system from the surroundings
303
Exothermic process
A release of heat content from a system
304
delta H in an endothermic process
Positive
305
delta H in an exothermic process
Negative
306
Enthalpy of a reaction
= bonds broken - bonds formed
307
Rank the phases of matter from lowest to highest entropy
Solids < liquids < gases
308
Entropy in terms of energy dispersal and disorder
Entropy increases as a system has more disorder or freedom of movements, and energy is dispersed in a spontaneous system
Entropy of the universe can never be decreased spontaneously
A measure of the degree to which energy has been spread throughout a system or between a system and its surrounds
A ratio of heat transferred per mole per unit kelvin
309
Liquid water --> solid water : increase or decrease in entropy?
Decrease (freezing)
310
Dry ice sublimates into carbon dioxide : increase or decrease in entropy?
Increase (sublimation)
311
NaCl (s) --> NaCl (aq) : increase or decrease in entropy?
Increase (dissolution)
312
N2 (g) + 3 H2 (g) --> 2 NH3 (g) : increase or decrease in entropy?
Decrease (fewer moles of gas)
313
An ice pack is placed on a wound : increase or decrease in entropy?
Increase (heat is transferred)
314
Gibbs free energy (delta G)
= delta H - T delta S
315
Gibbs free energy at equilibirum
= 0
316
First law of thermodynamics
delta U = Q - W
317
Heat transfer (no phase change) (q)
m c delta T
318
Heat transfer (duding phase change) (q)
m L
319
Generalized enthalpy of reaction (delta H reaction)
= H products - H reactants
320
Standard enthalpy of reaction (delta H reaction)
= sum of delta Hf, products - sum of delta Hf, reactants
321
Bond enthalpy (delta H reaction)
= sum of delta H bonds broken - sum of delta H bonds formed
| = total energy absorbed - total energy released
322
Entropy (delta S)
= Q rev / T
323
Second law of thermodynamics
delta S universe = delta S system + delta S surroundings > 0
324
Standard entropy of reaction
delta S reaction = sum of delta Sf, products - sum of delta Sf, reactants
325
Standard Gibbs free energy of reaction (delta G reaction)
= sum of delta Gf, products - sum of delta Gf, reactants
326
Standard Gibbs free energy from equilibrium constant (delta G reaction)
= - R T ln (Keq)
327
Gibbs free energy from reaction quotient (delta G reaction)
= delta G reaction + R T ln (Q) = R T ln (Q / Keq)
328
Isolated systems
Exchange neither matter nor energy with the environment
329
Closed systems
Exchange energy, but not matter, with the environment
330
Open systems
Exchange both energy and matter with the environment
331
When do phase changes exist?
At characteristic temperatures and pressures
332
Where do fusion (melting) and freezing (crystallization or solidification) occur?
At the boundary between the solid and the liquid phases
333
Where do vaporization (evaporation or boiling) and condensation occur?
At the boundary between the liquid and the gas phases
334
Where do sublimation and deposition occur?
At the boundary between the solid and the gas phases
335
Phase diagram
Graphs the phases and phase equilibria as a function of temperature and pressure
336
Enthalpy
A measure of potential energy of a system found in intermolecular attractions and chemical bonds
337
Hess's law
The total change in potential energy of a system is equal to the changes of potential energies of the individual steps of the process
338
When is entropy maximized?
At equilibrium
339
Negative delta G
Reaction proceeds in the forward direction (spontaneous)
Endothermic reaction (conclusion, not a comment)
340
Positive delta G
Reaction proceeds in the reverse direction (nonspontaneous)
Exothermic reaction (conclusion, not a comment)
341
Does Gibbs free energy depend on temperature?
Yes
342
Temperature-dependent processes
Change between spontaneous and nonspontaneous, depending on the temperature
343
What must be the value of Keq if ln (Keq) is positive?
Keq > 1
344
Combustion
Hydrocarbon reacting with oxygen to produce carbon dioxide and water
345
Do longer hydrocarbon chains yield greater or less amounts of combustion products?
Greater
346
Do longer hydrocarbon chains release more or less heat in a combustion process?
More
347
Spontaneous forward reaction
Keq > Q
348
Characteristics of gas phase
Compressible, have rapid molecular motion, have large intermolecular distances, have weak intermolecular forces
349
What will happen to the level of mercury in a mercury barometer if the barometer was moved to the top of a mountain?
Decreases
350
What will happen to the level of mercury in a mercury barometer if the barometer was placed ten meters under water?
Increases
351
Standard temperature and pressure (STP) conditions
T = 273 K (0 C)
P = 1 atm
One mole of ideal gas occupies 22.4 L
352
Standard conditions
```
T = 298 K (25 C)
P = 1 atm
Concentration = 1 M
```
353
How can the concentration of carbon dioxide in sodas or other carbonated beverages be so much higher than that of atmospheric carbon dioxide?
High pressures of carbon dioxide are forced on top of the liquid in sodas, increasing its concentration in the liquid.
354
Assumptions made by the kinetic molecular theory
Negligible volume of gas particles, no intermolecular forces (neither attraction nor repulsion), random motion, elastic collisions among gas particles, and the average kinetic energy of the gas particles is recruit proportional to temperature
355
Hydrogen sulfide (H2S) has a very strong rotten egg odor. Methyl salicylate (C8H8O) has a wintergreen odor, and benzaldehyde (C7H6O) has a pleasant almond odor. If the vapors for these three substances were released at the same time from across the room, in which order would one smell the orders?
The rotten egg odor first, the almond next and the wintergreen last. Because all of the gases have the same temperature, they have the same kinetic energy; thus, the lightest molecules travel the fastest.
356
In what ways do real gases differ from ideal gases?
Real gas molecules have nonnegligible volume and attractive forces. Real gases deviate from ideal gases at high pressure (low volume) and low temperature.
357
If the attractive forces between gas molecules were to increase while the actual volumes of the molecules remained negligible, what would occur?
According to the van der Waals equation, if a is increased while b remains negligible, the correction term (n^2 a) / V^2 gets larger, and the pressure or volume must drop to compensate
358
If the volumes of gas molecules were to increase while the attractive forces between the molecules remained negligible, what would occur?
Increasing the volume of gas molecules while keeping attraction negligible makes the term V - nb smaller; thus, the pressure or volume must rise to compensate
359
Ideal gas law
P V = n R T
| Describes the relationship between the four variables of the gas state for an ideal gas
360
Density of a gas
p = m / V = (P M) / (R T)
361
Combined gas law
(P1 V1) / T1 = (P2 V2) / T2
A combination of Boyle's, Charles's and Gay-Lussac's laws; it shows an inverse relationship between pressure and volume along with direct relationships between pressure and volume with temperature
362
Avogadro's principle
n / V = k
n1 / V1 = n2 / V2
A special case of the ideal gas law for which the pressure and temperature are held constant; it shows a direct relationship between the number of moles of gas and volume
363
Boyle's law
P V = k
P1 V1 = P2 V2
A special case of the ideal gas law for which temperature and the tumble of moles are held constant; it shows an increase relationship between pressure and volume
364
Charles's law
V / T = k
V1 / T1 = V2 / T2
A special case of the ideal gas law for which pressure and number of moles are held constant; it shows a direct relationship between temperature and volume
365
Gay-Lussac's law
P / T = k
P1 / T1 = P2 / T2
A special case of the ideal gas law for which volume and number of moles are held constant; it shows a direct relationship between temperature and pressure
366
Dalton's law (total pressure from partial pressures) (PT)
= PA + PB + ...
367
Dalton's law (partial pressure from total pressure) (PA)
= XA PT
| XA = moles of A / total number of moles
368
Henry's law
[A] = kH x PA
[A1] / P1 = [A2] / P2 = kH
States that the amount of gas dissolved in solution is directly proportional to the partial pressure of that gas at the surface of a solution
369
Average kinetic energy of a gas (KE)
= 1/2 m v^2 = 1.5 kB T
370
Root-mean-square speed (u rms)
= square root of ([3 R T] / M)
371
Graham's law
r1 / r2 = square root of (M2 / M1)
Describes the behavior of has diffusion or effusion, stating that gases with lower molar masses will diffuse or effuse faster than gases with higher molar masses at the same temperature
372
van der Waals equation of state
(P + [n^2 a] / v^2) (V - n b) = n R T
| Is used to correct the ideal gas law for intermolecular attractions (a) and molecular volume (b)
373
What is the least dense phase of matter?
Gases
374
Do gases conform to the shape of the container?
Yes
375
Are gases easily compressible?
Yes
376
Variables that describe gas systems
Temperature, pressure, volume and moles
377
1 atm in mmHg
760
378
760 mmHg in atm
1
379
1 atm in torr
760
380
760 mmHg in torr
760
381
760 torr in mmHg
760
382
760 torr in atm
1
383
1 atm in Pa
101.325 kPa
384
760 mmHg in Pa
101.325 kPa
385
760 torr in Pa
101.325 kPa
386
101.325 kPa in torr
760
387
101.325 kPa in mmHg
760
388
101.325 kPa in atm
1
389
Simple mercury barometer
Measures incident (usually atmospheric) pressure. As pressure increases, more mercury is forced into the column, increasing its height. As pressure decreases, mercury glows out of the column under its own weight, decreasing its height.
390
Dalton's law of partial pressures
States theta individual gas components of a mixture of gases will exert individual pressures in proportion to their mole fractions. The total pressure of a mixture of gases is equal to the sum of the partial pressures of the component gases.
391
Kinetic molecular theory
The average kinetic energy of a molecule of gas is directly proportional to the temperature of the gas in kelvins
392
Diffusion
The spreading out of particles from high to low concentrations
393
Effusion
The movement of gas from one compartment to another through a small opening under pressure
394
How do real gases deviate from ideal gases at modernly high pressure, low volumes or low temperatures?
They will occupy less volume than predicted by the ideal gas law because the particles have intermolecular attractions
395
How do real gases deviate from ideal gases at extremely high pressures, low volumes or low temperatures?
They will occupy more volume than predicted by the ideal gas law because the particles occupy physical space
396
Graham's law of effusion
States that the relative rates of effusion of two gases at the same temperature and pressure are given by the inverse ratio of the square roots of the masses of the gas particles
397
Which factors affect a gas's pressure?
Temperature and volume
| If one of those two values is kept constant, the pressure will always be found
398
What does heating a gas at constant volume do?
Increase pressure
399
What does cooling a gas at constant volume do?
Decrease pressure
400
How can pressure be calculated when both volume and temperature are changing?
If both changes in volume and temperature have the same effect on pressure, then pressure can be predicted
e.g.
Cooling the gas and increasing its volume will both decrease pressure
401
What would increasing volume do to a gas if temperature is constant?
Decreases its pressure
402
What would decreasing volume do to a gas if temperature is constant?
Increases its pressure
403
Solvation
The breaking of intermolecular forces between solute particles and between solvent particles, with formation of intermolecular forces between solute and solvent particles
404
Solubility
The amount of solute contained in a solvent
405
How can the solubility of solids be increased?
Increasing the temperature
406
Which ions can form salts that are always soluble?
Group I metals, ammonium, nitrate and acetate salts
407
What is the solvent in aqueous solution?
Water
408
Saturation
The maximum solubility of a compound at a given temperature; one cannot dissolve any more of the solute just by adding more at this temperature
409
How can the solubility of gases be increased?
Decreasing temperature or increasing the partial pressure of the gas above the solvent (Henry's law)
410
Colligative property
Physical properties of solutions that depend on the amount of solute present, but not the actual identity of the solute particles
411
How are molality and molarity related for water?
Molarity and molality are nearly equal at room temperature because 1 L solution if approximately equal to 1 kg solvent for dilute solutions (like the denominators of the molarity and molality equations, respectively).
412
How are molality and molarity related for solvents other than water?
Molarity and molality differ significantly because their densities are not 1 g/mL like water
413
Examples of colligative properties
Vapor pressure depression, boiling point elevation, freezing point depression and osmotic pressure
414
Percent composition by mass
= (mass of solute / mass of solution) x 100%
Used for aqueous solutions and solid-in-solid solutions
415
Mole fraction (Xa)
= moles of a / total moles of all species
Used for calculating vapor pressure depression and partial pressures of gases in a system
416
Molality (m)
= moles of solute / kilograms of solvent (mol/kg)
Used for boiling point elevation and freezing point depression
417
Dilution formula
Mi Vi = Mf Vf
418
Solubility product constant (Ksp)
= [A^n+]^m [B^m-]^n
The equilibrium constant for a dissociation reaction
419
Ion product (IP)
= [A^n+]^m [B^m-]^n
Determines the level of saturation and behavior of the solution
420
Raoult's law (vapor pressure depression)
Pa = Xa Poa
The presence of other solutes decreases the evaporation rate of a solvent without affecting its condensation rate, thus decreasing its vapor pressure
421
Boiling point elevation (delta Tb)
= i Kb m
A shift in the phase equilibrium dependent on the molality of the solution
422
Freezing point depression (delta Tf)
= i Kf m
A shift in the phase equilibrium dependent on the molality of the solution
423
Osmotic pressure (pi)
= i M R T
Primarily depends on the molarity of the solution
424
Solution
Homogenous mixtures composed of two or more substances
| Combine to form a single phase, generally the liquid phase
425
Solvent
Surrounds solute particles via electrostatic interactions in a process called salvation or dissolution
426
Hydration
Salvation in water
427
Are dissolutions endothermic or exothermic?
Endothermic
428
Is the dissolution of gas into liquid endothermic or exothermic?
Exothermic
429
How is solubility expressed?
As molar solubility
430
Molar solubility
The molarity of the solute at saturation
431
Complex ions (coordination compounds)
Composed of metallic ions bonded to various neutral compounds and anions (ligands)
Increase the solubility of otherwise insoluble ions (the opposite of the common ion effect)
432
How are complex ions formed?
The process involves electron pair donors and electron pair acceptors such as those seen in coordinate covalent bonding
433
IP < Ksp
The solution is unsaturated
| If more solute is added, it will dissolve
434
IP = Ksp
The solution is saturated (at equilibrium)
| There will be no change in concentrations
435
IP > Ksp
The solution is supersaturated
| A precipitate will form
436
Formation (stability constant, Kf)
The equilibrium constant for complex formation
| Its value is usually much greater than Ksp
437
Is Kf greater than or less than Ksp?
Greater than
438
Why does the formation of a complex ion increase the solubility of other salts containing the same ion?
Because it uses up the products of those dissolution reactions, shifting the equilibrium to the right (the opposite of the common ion effect)
439
Common ion effect
Decreases the solubility of a compound in a solution that already contains one of the ions in the compound. The pretense of that ion in solution shifts the dissolution reaction to the left, decreasing its dissociation.
440
As vapor pressure decreases, does the temperature (energy) required to boil the liquid increase or decrease?
Increase
441
van't Hoff factor (i)
It is used in freezing point depression, boiling point elevation and osmotic pressure calculations for solutes that dissociate
442
What is the van't Hoff factor (i) for a solute that does not dissociate into smaller components?
1
443
Raoult's law
The ideal solution behavior is observed when solute-solute, solvent-solvent and solute-solvent interactions are all very similar
444
When is melting point depressed?
Upon solute addition
| Solute particles interfere with lattice formation
445
Lattice formation
The highly organized state in which solid molecules align themselves
446
Is breaking the solute into its individual components an endothermic or exothermic process?
Endothermic
| Energy is required to break molecules apart
447
Is making room for the solute in the solvent by overcoming intermolecular forces in the solvent an endothermic or exothermic process?
Endothermic
| The intermolecular forces in the solvent must be overcome to allow incorporation of solute particles
448
Is allowing solute-solvent interactions to occur to form the solution an endothermic or exothermic process?
Exothermic
| Polar water molecules will interact with the dissolved ions, creating a stable solution and releasing energy
449
Is a spontaneous reaction endothermic or exothermic?
Endothermic
450
Arrhenius acid
Dissociates to form excess H+ in solution
451
Arrhenius base
Dissociates to form excess OH- in solution
452
Bronsted-Lowry acid
H+ donor
453
Bronsted-Lowry base
H+ acceptor
454
Lewis acid
Electron pair acceptor
455
Lewis base
Electron pair donor
456
Amphoteric species
A species that can act as an acid or a base
457
If a compound has a Ka value >> water, what does it mean about its behavior in solution? How does this compare with a solution that has only a slightly higher Ka than water?
High Ka indicates a strong acid, which will dissociate completely in solution. Having a Ka slightly greater than water means the acid is a weak acid with minimal dissociation.
458
If a compound has a Kb value >> water, what does it mean about its behavior in solution? How does this compare with a solution that has only a slightly higher Kb than water?
High Kb indicates a strong base, which will dissociate completely in solution. Having a Kb slightly greater than water means the base is a weak base with minimal dissociation.
459
What is the mathematical relationship between Ka, Kb and Kw?
Ka x Kb = Kw
460
What species are considered the equivalents of acids and bases?
Acids use moles of H+ (H3O+) as an equivalent
| Bases use moles of OH- as an equivalent
461
Buffering region in a titration curve
Occurs when [HA] ~= [A-]
| The flattest portion of the titration curve (i.e. resistant to changes in pH)
462
Half-equivalence point in a titration curve
The center of the buffering region, where [HA] = [A-]
463
Equivalence point in a titration curve
The steepest point of the titration curve
| Occurs when the equivalents of acid present equal the equivalents of base added (or vice-versa)
464
Endpoint in a titration curve
The pH at which an indicator turns its final color
465
In which part f the pH range will the equivalence point fall for titration with a strong acid with a weak base?
In the acidic range
466
In which part f the pH range will the equivalence point fall for titration with a strong base with a weak acid?
In the basic range
467
In which part f the pH range will the equivalence point fall for titration with a strong acid with a strong base?
pH = 7 (neutral)
468
In which part f the pH range will the equivalence point fall for titration with a weak acid with a weak base?
In the acidic range, basic range or neutral range, depending on the relative strengths of the acid and base
469
Buffer solution
Resists changes in pH and has optimal buffering capacity within 1 pH point from its pKa
Consist of a mixture of a weak acid and its conjugate salt or a weak base and its conjugate salt
470
Autoionization constant for water (Kw)
= [H3O+] [OH-] = 10^-14 (at 25 C, 298 K)
471
Definition of pH
= - log [H+] = log (1 / [H+])
472
Definition of pOH
= - log [OH-] = log (1 / [OH-])
473
Relationship of pH and pOH at 298 K
pH + pOH = 14
474
p scale value approximation
p value ~= m - 0.n
475
Acid dissociation constant (Ka)
= ([H3O+] [A-]) / [HA]
476
Base dissociation constant (Kb)
= ([B+] [OH-]) / [BOH]
477
Equivalence point
Na Va = Nb Vb
478
Henderson-Hasselbalch equation (acid buffer)
pH = pKa + log ([A-] / [HA])
479
Henderson-Hasselbalch equation (base buffer)
pOH = pKb + log ([B+] / [BOH])
480
Are all Arrhenius acids and bases also Bronsted-Lowry acids and bases?
Yes
481
Are all Bronsted-Lowry acids and bases also Lewis acids and bases?
Yes
482
Are all Lewis acids and bases also Bronsted-Lowry acids and bases?
No
483
Are all Bronsted-Lowry acids and bases also Arrhenius acids and bases?
No
484
Amphoteric species
Can behave as either acids or bases
485
Amphiprotic species
Amphoteric species that specifically can behave as a Bronsted-Lowry acid or base
486
Examples of amphoteric and amphiprotic species
Water and conjugate species of polyvalent acids and bases
487
What is the water dissociation constant (Kw, auto ionization constant of water) affected by?
Changes in temperature
488
Strong acids and bases
Completely dissociate in solution
489
Weak acids and bases
Do not completely dissociate in solution and have corresponding dissociation constants (Ka and Kb)
490
Bronsted-Lowry definition
Acids have conjugate bases that are formed when the acid is deprotonated. Bases have conjugate acids that are formed when the base is protonated. Strong acids and bases have very weak (inert) conjugates. Weak acids and bases have weak conjugates.
491
Neutralization reactions
Form salts and (sometimes) water
492
Equivalent
One mole of the species of interest
493
Normality
The concentration of acid or base equivalents in solution
494
Polyvalent acids and bases
Can donate or accept multiple electrons
495
What is the normality of a solution containing polyvalent species?
It is the molarity of the acid or base times the number of protons it can donate or accept
496
Titrations
Used to determine the concentration of a known reactant in a solution
497
Titrant
Has known concentration and is added slowly to the titrand to reach the equivalence point
498
Titrand
Has an unknown concentration but a known volume
499
Indicators
Weak acids or bases that display different colors in their protonated and deprotonated forms
500
How should the indicator for a titration be chosen?
The indicator should have a pKa close to the pH of the expected equivalence point
501
Titration curve or polyvalent species
Has multiple buffering regions and equivalence points
502
Buffering capacity
Refers to the ability of a buffer to resist changes in pH
503
When is maximal buffering capacity seen?
Within 1 pH point of the pKa of the acid in the buffer solution
504
Henderson-Hasselbalch equation
Quantifies the relationship between pH and pKa for weak acids and between pPH and pKb for weak bases
When a solution is optimally buffered, pH = pKa and pOH = pKb
505
Acids ending in -ic
Derivatives of acids ending in -ate
506
Acids ending in -ous
Derivatives of acids ending in -ite
507
Strong bases
Soluble hydroxides of Group IA and IIA metals
508
Ka = 1
Strong acid
509
Oxidation
Loss of electrons
510
Reduction
Gain of electrons
511
Oxidizing agent
Facilitates the oxidation of another compound and is reduced itself in the process
512
Reducing agent
Facilitates the reduction of another compound and is itself oxidized in the process
513
Common oxidizing agents
Contain oxygen or a similarly electronegative element
514
Common reducing agent
Contain metal ions or hydrides
515
Half-reaction method (ion-electron method)
1. Separate the two half-reactions
2. Balance the atoms of each half-reaction
3. Balance the charges of each half-reaction by adding electrons
4. Multiply the half-reactions as necessary to obtain the same number of electrons in both half-reactions
5. Add the half-reactions, cancelling out terms on both rides of the reaction arrow
6. Confirm the mass and charge are balanced
516
Complete ionic equation
Accounts for all of the ions present in a reaction
517
How to write a complete ionic equation?
Split all aqueous compounds into their relevant ions. Keep solid salts intact
518
Net ionic equation
Removes spectator ions to focus only on the species that actually participate in the reaction
519
How to obtain the net ionic reaction?
Subtract the ions appearing on both sides of the reaction
520
Spectator ions
Ions appearing on both sides of the reaction
521
For reactions that contain no aqueous salts, what does the net ionic equation look like?
It is generally the same as the overall balanced reaction
522
For double displacement (metathesis) reactions that do not form a solid salt, what does the net ionic equation look like?
There is no net ionic equation because all ions remain in solution and do not change oxidation numbers
523
Disproportionation (dismutation) reaction
A type of reduce reaction in which one element is both oxidized and reduced forming at least two molecules containing the element with different oxidation states
524
Oxidation-reduction titrations
Follow the transfer of charge
| Use indicators to indicate when certain voltages of solutions are achieved
525
Potentiometric titration
A form of redox titration in which a voltmeter or external cell measures the electromotive force (emf) of a solution
No indicator is used
The equivalence point is determined by a sharp change in voltage
526
What is the site of oxidation in a galvanic (voltaic) cell?
Anode
527
In a galvanic (voltaic) cell, what to electrons flow towards?
Cathode
528
In a galvanic (voltaic) cell, what does current flow towards?
Anode
529
In a galvanic (voltaic) cell, what has a (-) designation?
Anode
530
In a galvanic (voltaic) cell, what attracts cations?
Cathode
531
What is the site of oxidation in an electrolytic cell?
Anode
532
In an electrolytic cell, what to electrons flow towards?
Cathode
533
In an electrolytic cell, what does current flow towards?
Anode
534
In an electrolytic cell, what has a (-) designation?
Cathode
535
In an electrolytic cell, what attracts cations?
Cathode
536
Cell diagram for the discharging state of a lead-acid battery
Pb (s) | H2SO4 (4 M) || H2SO4 (4 M) | PbO2 (s)
537
Which type of cell has a positive delta G?
Electrolytic cells (because they are non spontaneous)
538
Which type of cell has a positive Ecell?
Galvanic (voltaic) cells (because they are spontaneous and have a negative delta G and therefore a positive Ecell)
539
Is Ni-Cd (discharging) galvanic or electrolytic?
Galvanic
540
Is Ni-Cd (charging) galvanic or electrolytic?
Electrolytic
541
Is molten NaCl (discharging) galvanic or electrolytic?
Electrolytic
542
Is Daniell cell (discharging) galvanic or electrolytic?
Galvanic
543
Is lead-acid (charging) galvanic or electrolytic?
Electrolytic
544
Is lead-acid (discharging) galvanic or electrolytic?
Galvanic
545
What is the anode material in Ni-Cd (discharging)?
Cd
546
What is the anode material in Ni-Cd (charging)?
Cd(OH)2
547
What is the anode material in molten NaCl (discharging)?
Anything
548
What is the anode material in Daniell cell (discharging)?
Zn
549
What is the anode material in lead-acid (charging)?
PbSO4
550
What is the anode material in lead-acid (discharging)?
Pb
551
What is the anode charge in Ni-Cd (discharging)?
Negative
552
What is the anode charge in Ni-Cd (charging)?
Positive
553
What is the anode charge in molten NaCl (discharging)?
Positive
554
What is the anode charge in Daniell cell (discharging)?
Negative
555
What is the anode charge in lead-acid (charging)?
Positive
556
What is the anode charge in lead-acid (discharging)?
Negative
557
What is the cathode material in Ni-Cd (discharging)?
NiO(OH)
558
What is the cathode material in Ni-Cd (charging)?
Ni(OH)2
559
What is the cathode material in molten NaCl (discharging)?
Anything
560
What is the cathode material in Daniell cell (discharging)?
Cu
561
What is the cathode material in lead-acid (charging)?
PbSO4
562
What is the cathode material in lead-acid (discharging)?
PbO2
563
What is the cathode charge in Ni-Cd (discharging)?
Positive
564
What is the cathode charge in Ni-Cd (charging)?
Negative
565
What is the cathode charge in molten NaCl (discharging)?
Negative
566
What is the cathode charge in Daniell cell (discharging)?
Positive
567
What is the cathode charge in lead-acid (charging)?
Negative
568
What is the cathode charge in lead-acid (discharging)?
Positive
569
How are standard reduction potentials measured?
Setting up a cell relative to a standard hydrogen electrode, with is given a reduction potential of 0 V by convention
570
Positive emf
The cell is spontaneous (galvanic)
571
Negative emf
The cell is non spontaneous (electrolytic)
572
Is:
2 Fe^3+ (aq) + 2 Cl- (aq) --> 2 Fe^2+ (aq) + Cl2 (g) || Ecell = -0.59 V
galvanic or electrolytic?
Electrolytic, because emf is negative
573
Is:
2 Fe^3+ (aq) + 2 I- (aq) --> 2 Fe^2+ (aq) + I2 (aq) || Ecell = +0.25 V
galvanic or electrolytic?
Galvanic, because emf is positive
574
If Keq = 1.2e-2, is delta G positive or negative?
Positive
575
If Keq = 2e2, is delta G positive or negative?
Negative
576
If Keq = 1 is delta G positive or negative?
0
577
If Keq = 1.2e-2, is the reaction spontaneous or non spontaneous?
Nonspontaneous
578
If Keq = 2e2, is the reaction spontaneous or non spontaneous?
Spontaneous
579
If Keq = 1, is the reaction spontaneous or non spontaneous?
The cell is in equilibrium
580
If Keq = 1.2e-2, is Ecell positive or negative?
Negative
581
If Keq = 2e2, is Ecell positive or negative?
Positive
582
If Keq = 1, is Ecell positive or negative?
0
583
What is the direction of the reaction if Q = 10^-3 and Keq = 10^-2?
Forward
584
What is the direction of the reaction if Q = 10^2 and Keq = 1.1?
Backward
585
What is the direction of the reaction if Q = 1 and Keq = 1?
The reaction is in equilibrium
586
If Q = 10^-3 and Keq = 10^-2, is Ecell positive or negative?
Positive
587
If Q = 10^2 and Keq = 1.1, is Ecell positive or negative?
Negative
588
If Q = 1 and Keq = 1, is Ecell positive or negative?
0
589
How is the equation:
delta G = - R T ln Keq
affected if Keq < 1?
ln Keq < 0 and delta G > 0
590
How is the equation:
delta G = - R T ln Keq
affected if Keq > 1?
ln Keq > 0 and delta G < 0
591
How is the equation:
delta G = - R T ln Keq
affected if Keq = 1?
ln Keq = 0 and delta G = 0
592
Moles of electrons transferred during reduction
M^n+ + n e- --> M (s)
593
Electrodeposition equation
mol M = (I t) / (n F)
594
Standard electromotive force of a cell (Ecell)
= E red, cathode - E red, anode
The difference in standard reduction potential between the two half-cells
595
Standard change in free energy from standard emf (delta G)
= - n F Ecell
596
Nernst equation (full)
Ecell = Ecell - ([{R T} / {n F}] ln Q)
Describes the relationship between the concentration of species in a solution under nonstandard conditions and the electromotive force
597
Nernst equation (simplified)
Ecell = Ecell - (0.0592 / n) log Q
598
Reaction quotient (Q)
= ([C]^c [D]^d) / ([A]^a [B]^b)
599
Standard change in free energy from equilibrium constant (delta G)
= - R T ln Keq
600
Free energy change (nonstandard conditions) (delta G)
= delta G standard + R T ln Q
601
Electrochemical cell
Any cell in which oxidation-reduction reactions take place
602
Characteristics of electrochemical cells
Electrodes, anodes, cathodes, electrons flow from the anode to the cathode, and current flowed from the cathode to the anode
603
Electrodes
Strips of metal or other conductive materials placed in an electrolyte solution
604
Anode
The site of oxidation
| Attracts anions
605
Cathode
The site of reduction
| Attracts cations
606
Cell diagrams
Shorthand notation that represent the reactions taking place in an electrochemical cell
Written from anode to cathode with electrolytes (the solution) in between
A cortical line represents a phase boundary, and a double vertical line represents a salt bridge or other physical boundary
607
Galvanic (voltaic) cell
Spontaneous (delta G < 0)
| Positive electromotive force
608
Electrolytic cell
Nonspontaneous (delta G > 0)
| Negative electromotive force
609
Concentration cell
Specialized form of galvanic cell in which both electrodes are made of the same material
Rather than a potential difference causing the movement of charge, it is the concentration gradient between the two solutions
610
Rechargeable batteries
Electrochemical cells that can experience charging (electrolytic) and discharging (galvanic) states
Ranked by energy density
611
Energy density of a rechargeable battery
The amount of energy a cell can produce relative to the mass of battery material
612
Lead-acid batteries (discharging)
Consist of a Pb anode and a PbO2 cathode in a concentrated sulfuric acid solution
613
Lead-acid batteries (charging)
PbSO4-plated electrodes are dissociated to restore the original Pb and PbO2 electrodes and concentrate the electrolyte
614
Do lead-acid batteries have high or low energy density?
Low
615
Are lead-acid batteries rechargeable?
Yes
616
Nickle-cadmium batteries (Ni-Cd) (discharging)
Consist of a Cd anode and a NiO(OH) cathode in a concentrated KOH solution
617
Nickle-cadmium batteries (Ni-Cd) (charging)
Ni(OH)2 and Cd(OH)2-plated electrodes are dissociated to restore the original Cd and NiO(OH) electrodes and concentrate the electrolyte
618
Do Ni-Cd batteries have high or low energy density?
Low, but higher than that of lead-acid batteries
619
Are Ni-Cd batteries rechargeable?
Yes
620
Nickel-metal hydride (NiMH) batteries
Replace Ni-Cd batteries because they have higher energy density, are more cost effective and are less toxic
621
Do NiMH batteries have high or low energy density?
High, higher than those of lead-acid batteries and Ni-Cd batteries
622
Are NiMH batteries rechargeable?
Yes
623
Surge current
An above-average current transiently released at the beginning of the discharge phase, it eases rapidly until a stable current is achieved
624
Reduction potential
Quantifies the tendency for a species to gain electrons and be reduced
625
What does a higher reduction potential mean?
The more the species wants to be reduced
626
Standard reduction potential (Ered)
Calculated by comparison to the standard hydrogen electrode (SHE) under the standard conditions of 298 K, 1 atm pressure and 1 M concentrations
627
Standard hydrogen electrode (SHE)
Has a standard reduction potential of 0 V
628
In a galvanic cell, what is the sign of the difference of the reduction potentials of the two half reactions?
Positive
629
In an electrolytic cell, what is the sign of the difference of the reduction potentials of the two half reactions?
Negative
630
How are the signs of electromotive force and change in free energy relate to one another?
They are always opposites of each other
631
In a galvanic cell, can the reduction potential of the cathode be less than that of the anode?
No
632
Keq in an electrolytic cell
< 1
