Ch 9,10,11- Solutions, AcidBase, Redox Flashcards
1
Q
“Incidentally, gases “dissolved” into other gases can be thought of as solutions, but are more properly defined only as ______because gas molecules do not interact all that much chemically, as described by the kinetic molecular theory of gases.
A
Mixtures
As a point of clarification: all solutions are considered mixtures, but not all mixtures are considered solutions.”
2
Q
“is the component of the solution that remains in the same phase after mixing. If the two substances are already in the same phase (for example, a solution of two liquids), it is the component present in greater quantity”
A
Solvent
3
Q
“is the electrostatic interaction between solute and solvent molecules. This is also known as dissolution, and when water is the solvent, it can be called hydration”
A
Solvation
4
Q
“The dissolution of gases into liquids, such as CO2 into water, is an ______ process because the only significant interactions that must be broken are those between water molecultes
A
Exothermic
“CO2, as a gas, demonstrates minimal intermolecular interaction. Le Châtelier’s principle tells us this is the reason that lowering the temperature of a liquid favors solubility of a gas in the liquid.”
5
Q
“When the new interactions are weaker than the original ones, solvation is_______ and the process is favored at ______ temperatures.
A
Endothermic; high
Most dissolutions are of this type. Two such examples have already been given: dissolving ammonium nitrate or sugar into water. Because the new interactions between the solute and solvent are weaker than the original interactions between the solute molecules and between the solvent molecules, energy (heat) must be supplied to facilitate the formation of these weaker, less stable interactions.”
6
Q
“At constant temperature and pressure, entropy always ____ upon dissolution”
A
Increases
7
Q
“the overall dissolution of table salt into water is ________ and favored at high temperatures.”
A
Endothermic
“The formation of these ion–dipole bonds is exothermic, but the magnitude is slightly less than the energy required to break the ionic bonds and hydrogen bonds
8
Q
“ The ________ of a substance is the maximum amount of that substance that can be dissolved in a particular solvent at a given temperature. ”
A
Solubility
“When this maximum amount of solute has been added, the dissolved solute is in equilibrium with its undissolved state, and we say that the solution is saturated.”
9
Q
“When this maximum amount of solute has been added, the dissolved solute is in equilibrium with its undissolved state, and we say that the solution is
A
saturated.”
10
Q
“A solution in which the proportion of solute to solvent is small is said to be _______, and one in which the proportion is large is said to be ________”
A
Small; Concentrated
11
Q
“In general, solutes are considered soluble if they have a molar solubility above _____ in solution.”
A
0.1 M
12
Q
“In aqueous solutions, there are ____ general solubility rules:”
A
7
- All salts containing ammonium (NH4+) and alkali metal (Group 1) cations are water-soluble.
- All salts containing nitrate (NO3−) and acetate (CH3COO−) anions are water-soluble.
- Halides (Cl−, Br−, I−), excluding fluorides, are water-soluble, with the exceptions of those formed with Ag+, Pb2+, and HG22+.
- All salts of the sulfate ion (SO42−) are water-soluble, with the exceptions of those formed with Ca2+, Sr2+, Ba2+, and Pb2+.
- All metal oxides are insoluble, with the exception of those formed with the alkali metals, ammonium, and CaO, SrO, and BaO, all of which hydrolyze to form solutions of the corresponding metal hydroxides.
- All hydroxides are insoluble, with the exception of those formed with the alkali metals, ammonium, and Ca2+, Sr2+, and Ba2+.
- All carbonates (CO32−), phosphates (PO43−), sulfides (S2–), and sulfites (SO32−) are insoluble, with the exception of those formed with the alkali metals and ammonium.”
13
Q
- All salts containing ammonium (NH4+) and alkali metal (Group 1) cations are water ______
A
Soluble
“All salts containing nitrate (NO3−) and acetate (CH3COO−) anions are also water-soluble.”
14
Q
“Halides (Cl−, Br−, I−), excluding fluorides, are water-______, with the exceptions of those formed with
A
Soluble;
Ag+, Pb2+, and HG22+.”
15
Q
“All salts of the sulfate ion (SO42−) are water-______, with the exceptions of those formed with”
A
Soluble
Ca2+, Sr2+, Ba2+, and Pb2+
16
Q
“All metal oxides are _____, with the exception of those formed with
A
Insoluble; the alkali metals, ammonium, and CaO, SrO, and BaO”
17
Q
“All hydroxides are ________, with the exception of those formed with
A
insoluble; the alkali metals, ammonium, and Ca2+, Sr2+, and Ba2+.”
18
Q
All carbonates (CO32−), phosphates (PO43−), sulfides (S2–), and sulfites (SO32−) are ______, with the exception of those formed with
A
Insoluble; the alkali metals and ammonium.”
19
Q
—refers to a molecule in which a cation is bonded to at least one electron pair donor (which could include the water molecule). The electron pair donor molecules are called _________.”
A
“complex ion—or coordination compound; Ligand
“Complexes are held together with coordinate covalent bonds, in which an electron pair donor (a Lewis base) and an electron pair acceptor (a Lewis acid) form very stable Lewis acid–base adducts”
20
Q
“In some complexes (complex ion), the central cation can be bonded to the same ligand in multiple places. This is called _______, and it generally requires large organic ligands that can double back to form a second (or even third) bond with the central cation”
A
Chelation
“Chelation therapy is often used to sequester toxic metals (lead, arsenic, mercury, and so on)”
21
Q
What is percent composition by mass?
A
Mass of solute/mass of solution
“Percent composition is used not only for aqueous solutions, but also for metal alloys and other solid-in-solid solutions.”
22
Q
“The mole fraction (X) of a compound is given by the equation”
A
Moles of A/Total moles of all species
23
Q
Define molarity
A
Moles of solute/Liters of solution
24
Q
Define molality
A
Moles of solute/kg of solvent
“For dilute aqueous solutions at 25°C, the molality is approximately equal to molarity because the density of water at this temperature is 1 kg/L
“You won’t use molality very often, so be mindful of the special situations when it is required: boiling point elevation and freezing point depression.”
25
“The ______ of a solution is equal to the number of equivalents of interest per liter of solution”
normality (N)
26
What is the dilution equation?
MiVi=MfVf
27
“The _______ can be applied to a solution at equilibrium; that is to say, when the solution is saturated and the solute concentration is at a maximum and is dynamically stable”
Law of mass action
“For a saturated solution of an ionic compound with the formula AmBn, the equilibrium constant for its solubility in aqueous solution, called the solubility product constant (Ksp), can be expressed by
Ksp = [Ag+][Cl−]
“On the MCAT, if you remember that Ksp is just a specialized form of Keq, then you can simplify a lot of problems by using the same concepts that you do for all equilibria, including Le Châtelier’s principle.”
28
“Solubility product constants, like all other equilibrium constants (Keq, Ka, Kb, and Kw) are ______ dependent.”
Temperature
29
“Generally speaking, the solubility product constant ______ with increasing temperature for non-gas solutes and ______ for gas solutes.”
Increases; Decrease
30
“Higher pressures favor ______ of gas solutes, and therefore the Ksp will be larger for gases at higher pressures than at lower ones.”
Dissolution
31
“We may not know whether the solution has reached saturation, and so to determine where the system is with respect to the equilibrium position, we can calculate a value called ______, which is analogous to the reaction quotient, Q, for other chemical reactions. The ion product equation has the same form as the equation for the solubility product constant”
the ion product (IP)
32
“If, at a given set of conditions, a salt’s IP is less than the salt’s Ksp, then the solution is not yet at equilibrium and is considered _______. For unsaturated solutions, dissolution is thermodynamically favored over precipitation. If the IP is greater than the Ksp, then the solution is beyond equilibrium, and the solution is considered _______”
Unsaturated; Supersaturated
“If the calculated IP is equal to the known Ksp, then the solution is at equilibrium—the rates of dissolution and precipitation are equal—and the solution is considered saturated. The molarity of a solute in a saturated solution is called the molar solubility of that substance.”
33
“If the calculated IP is equal to the known Ksp, then the solution is at _______—the rates of dissolution and precipitation are equal—and the solution is considered saturated. The molarity of a solute in a saturated solution is called the molar solubility of that substance.”
Equilibrium
34
The molarity of a solute in a saturated solution is called
the molar solubility of that substance.”
35
“IP
unsaturated, solute will continue to dissolve”
36
IP = Ksp
“saturated, solution is at equilibrium”
37
“IP > Ksp:
supersaturated, precipitation will occur”
38
“Every sparingly soluble salt of general formula MX will have Ksp = _____, where x is the molar solubility (assuming no common ion effect).”
x^2
“Every sparingly soluble salt of general formula MX2 will have Ksp = 4x^3, where x is the molar solubility (assuming no common ion effect).”
“Every sparingly soluble salt of general formula MX3 will have Ksp = 27x^4, where x is the molar solubility (assuming no common ion effect).”
39
“Every sparingly soluble salt of general formula MX2 will have Ksp = ______, where x is the molar solubility (assuming no common ion effect).”
4x^3
“Every sparingly soluble salt of general formula MX3 will have Ksp = 27x^4, where x is the molar solubility (assuming no common ion effect).”
“Every sparingly soluble salt of general formula MX3 will have Ksp = 27x^4, where x is the molar solubility (assuming no common ion effect).”
40
“Every sparingly soluble salt of general formula MX3 will have Ksp =_____, where x is the molar solubility (assuming no common ion effect).”
27x^4
Every sparingly soluble salt of general formula MX2 will have Ksp = 4x^3, where x is the molar solubility (assuming no common ion effect).”
“Every sparingly soluble salt of general formula MX3 will have Ksp = 27x^4, where x is the molar solubility (assuming no common ion effect).”
41
“The formation of complex ions ______ the solubility of a salt in solution.”
Increases
“it should make sense that, if a complex ion contains multiple polar bonds between the ligands and the central metal ion, it should be able to engage in a very large amount of dipole–dipole interactions. ”
“The end result is that such complexes tend to have very high Ksp values.”
42
“The dissolution of the original solution is termed Ksp, and the subsequent formation of the complex ion in solution is termed
Kf (the formation or stability constant of the complex in solution).”
43
What is the common ion effect.”
“The solubility of a salt is considerably reduced when it is dissolved in a solution that already contains one of its constituent ions as compared to its solubility in a pure solvent.
Note presence of the common ion has no effect on the value of the solubility product constant itself.”
44
“The ____ _____ are physical properties of solutions that are dependent on the concentration of dissolved particles but not on the chemical identity of the dissolved particles.”
Colligative Properties
“These properties—vapor pressure depression, boiling point elevation, freezing point depression, and osmotic pressure—are usually associated with dilute solutions.”
45
______ accounts for vapor pressure depression caused by solutes in solution. As solute is added to a solvent, the vapor pressure of the solvent decreases proportionately.”
Raoul’s Law
“PA = XAPA°
Equation 9.8
where PA is the vapor pressure of solvent A when solutes are present, XA is the mole fraction of the solvent A in the solution, and PA° is the vapor pressure of solvent A in its pure state.”
“Raoult’s law holds only when the attraction between the molecules of the different components of the mixture is equal to the attraction between the molecules of any one component in its pure state. When this condition does not hold, the relationship between mole fraction and vapor pressure will deviate from Raoult’s law. Solutions that obey Raoult’s law are called ideal solutions.”
46
“Vapor pressure depression goes hand in hand with boiling point elevation. The lowering of a solution’s vapor pressure would mean that
a higher temperature is required to match atmospheric pressure, thereby raising the boiling point.”
47
“When a nonvolatile solute is dissolved into a solvent to create a solution, the boiling point of the solution will be ______ than that of the pure solvent.
Greater
The boiling point is the temperature at which the vapor pressure of the liquid equals the ambient (incident) pressure”
48
The extent to which the boiling point of a solution is raised relative to that of the pure solvent is given by the formula
ΔTb = iKbm
“where ΔTb is the increase in boiling point, i is the van ’t Hoff factor, Kb is a proportionality constant characteristic of a particular solvent (which will be provided on Test Day), and m is the molality of the solution. ”
“The boiling point elevation formula calculates the amount that the normal boiling point is raised. The value calculated is not the boiling point itself.”
49
“The presence of solute particles in a solution interferes with the formation of the lattice arrangement of solvent molecules associated with the solid state. Thus, a _____ amount of energy must be removed from the solution (resulting in a lower temperature) in order for the solution to solidify. ”
Greater
“The formula for calculating the freezing point depression for a solution is
ΔTf = iKfm”
50
“The formula for calculating the freezing point depression for a solution is
ΔTf = iKfm”
51
refers to a “sucking” pressure generated by solutions in which water is drawn into a solution.”
“Osmotic pressure
52
“The equation for osmotic pressure is
∏ = iMRT
where ∏ is the osmotic pressure, i is the van ’t Hoff factor, M is the molarity of the solution, R is the ideal gas constant, and T is the temperature.”
53
“Most dissolutions are endothermic, although the dissolution of gas into liquid is
exothermic.”
54
“is the maximum amount of a solute that can be dissolved in a given solvent at a given temperature; it is often expressed as molar solubility—the molarity of the solute at saturation.”
Solubility
55
“are composed of metallic ions bound to various neutral compounds and anions, referred to as ligands.
“Complex ions or coordination compounds”
“Formation of complex ions increases the solubility of otherwise insoluble ions (the opposite of the common ion effect).
The process of forming a complex ion involves electron pair donors and electron pair acceptors such as those seen in coordinate covalent bonding.”
56
“is simply the equilibrium constant for a dissolution reaction.”
The solubility product constant (Ksp)
57
are in equilibrium at that particular temperature.”
“Saturated solutions
58
“is the equilibrium constant for complex formation. Its value is usually much greater than Ksp.”
The formation or stability constant (Kf)
59
“The formation of a complex does what to solubility
increases the solubility of other salts containing the same ions because it uses up the products of those dissolution reactions, shifting the equilibrium to the right (the opposite of the common ion effect).”
60
“The common ion effect does what to solubility
decreases the solubility of a compound in a solution that already contains one of the ions in the compound. The presence of that ion in solution shifts the dissolution reaction to the left, decreasing its dissociation.”
61
“An ______ will dissociate to form an excess of H+ in solution, and an _______ will dissociate to form an excess of OH− in solution.”
Arrhenius acid; Arrhenius base
“The advantage of this definition over Arrhenius’s is that it is not limited to aqueous solutions. For example, OH−, NH3, and F− are all Brønsted–Lowry bases because each has the ability to accept a hydrogen ion. However, neither NH3 nor F− can be classified as Arrhenius bases because they do not dissociate to produce an excess of OH− ions in aqueous solutions”
62
“Every Brønsted–Lowry acid (or base) can also be classified as
a Lewis acid (or base)”
“Every Arrhenius acid (or base) can also be classified as a Brønsted–Lowry acid (or base)”
63
“is one that reacts like an acid in a basic environment and like a base in an acidic environment.”
An amphoteric species
64
“For pure water at 298 K, the water dissociation constant, Kw, has been experimentally determined:
Kw = [H3O+][OH−] = 10−14”
“Each mole of water that autoionizes produces one mole each of hydrogen (or hydronium) ions and hydroxide ions, so the concentrations of the hydrogen ions and hydroxide ions are always equal in pure water at equilibrium. Thus, the concentration of each of the ions in pure water at equilibrium at 298 K is 10−7 M.
65
“However, the concentrations of the two ions will not always be equal. In fact, they will only be equal when the solution is neutral. Nevertheless, the product of their respective concentrations will always equal
10^14 when the temperature of the solution is 298 K”
66
“Kw is an equilibrium constant; unless the temperature of the water is changed, the value for Kw cannot be changed. Thus, the product of the concentrations of the hydrogen ions and the hydroxide ions in an aqueous solution at 298 K must always equal
10^14”
67
“ A p scale is defined as
the negative logarithm of the number of items.”
68
How to estimate P value (pH or pKa)
where n × 10^-m
=m-0.n
69
“Strong acids commonly encountered on the MCAT include _____ Strong bases commonly encountered include
HCl (hydrochloric acid), HBr (hydrobromic acid), HI (hydroiodic acid), H2SO4 (sulfuric acid), HNO3 (nitric acid), and HClO4 (perchloric acid) ;
NaOH (sodium hydroxide), KOH (potassium hydroxide), and other soluble hydroxides of Group IA metals. ”
70
“The smaller the Kb, the
weaker the base, and consequently, the less it will dissociate”
Generally speaking, we can characterize a species as a weak acid if its Ka is less than 1.0 and as a weak base if its Kb is less than 1.0. On the MCAT, molecular (nonionic) weak bases are almost exclusively amines.”
71
““The conjugate of a strong acid or base is sometimes termed inert because it is almost completely
unreactive.”
72
“An acid equivalent is equal to _______; a base equivalent is equal to
one mole of H+ (or, more properly, H3O+) ions; one mole of OH− ions. ”
73
“Some acids and bases are polyvalent; that is
each mole of the acid or base liberates more than one acid or base equivalent. ”
“Under the Brønsted–Lowry definition, such acids or bases could also be termed polyprotic”
74
“Titration is
a procedure used to determine the concentration of a known reactant in a solution.”
“There are different types of titrations, including acid–base, oxidation–reduction, and complexometric (metal ion). The MCAT frequently tests the first two types.”
75
“Titrations are performed by
adding small volumes of a solution of known concentration (the titrant) to a known volume of a solution of unknown concentration (the titrand) until completion of the reaction is achieved at the equivalence point.”
76
“In acid–base titrations, the equivalence point is reached when
the number of acid equivalents present in the original solution equals the number of base equivalents added, or vice-versa.”
77
“At the equivalence point, the number of equivalents of acid and base are equal. This fact allows us to calculate the unknown concentration of the titrand through the equation:
NaVa = NbVb ”
“where Na and Nb are the acid and base normalities, respectively, and Va and Vb are the volumes of acid and base solutions, respectively. Note that, as long as both volumes use the same units, the units used do not have to be liters.”
78
“The equivalence point in an acid–base titration is determined in two common ways:
”
evaluated by using a graphical method, plotting the pH of the unknown solution as a function of added titrant by using a pH meter, or estimated by watching for a color change of an added indicator.
79
“The indicator must always be
a weaker acid or base than the acid or base being titrated; otherwise, the indicator would be titrated first! ”
80
“The point at which the indicator changes to its final color is not the equivalence point but rather the
endpoint. If the indicator is chosen correctly and the titration is performed well the volume difference between the endpoint and the equivalence point is negligible and may be corrected for or simply ignored.
81
“Any question involving the selection of an ideal indicator will require you to know what the pH of the reaction at it equivalence point will be, whether graphically or mathematically. Once you have determined where the equivalence point is,
select the indicator that has the closest pKa value to it.”
82
“the equivalence point for strong acid/strong base titrations is always at pH
7
83
“To identify which type of titration is being shown in a graph
identify the starting position in the graph (pH ≫ 7 = titrant is a strong base, > 7 (slightly) = weak base,
84
“A buffer solution consists of
a mixture of a weak acid and its salt (which is composed of its conjugate base and a cation) or a mixture of a weak base and its salt (which is composed of its conjugate acid and an anion).”
85
“when [conjugate base] = [weak acid], the pH = pKa because log(1) = 0. This occurs at
the half-equivalence points in a titration, and buffering capacity is optimal at this pH.”
86
“What would happen if the concentrations of both the acid and its conjugate base were doubled?
While the pH would not change, the buffering capacity—the ability to which the system can resist changes in pH—has doubled.”
87
“The ______ has a known concentration and is added slowly to the titrand to reach the equivalence point.”
“The _______ has an unknown concentration but a known volume.”
Titrant; Titrand
88
“The half-equivalence point is
the midpoint of the buffering region, in which half of the titrant has been protonated (or deprotonated); thus, [HA] = [A−] and a buffer is formed.”
89
“Strong acid and strong base titrations have equivalence points at _____
Weak acid and strong base titrations have equivalence points at ______
Weak base and strong acid titrations have equivalence points at ______
pH = 7.
pH > 7.
pH
90
“Indicators are
weak acids or bases that display different colors in their protonated and deprotonated forms.”
“The indicator chosen for a titration should have a pKa close to the pH of the expected equivalence point.”
91
“The endpoint of a titration is
when the indicator reaches its final color.”
92
“Buffer solutions consist of
a mixture of a weak acid and its conjugate salt or a weak base and its conjugate salt; they resist large fluctuations in pH.”
93
“Buffering capacity refers to
the ability of a buffer to resist changes in pH; maximal buffering capacity is seen within 1 pH point of the pKa of the acid in the buffer solution.”
94
“The oxidation number of a free element is
zero. For example, the atoms in N2, P4, S8, and He all have oxidation numbers of zero.”
95
“The oxidation number of each Group IA element in a compound is ____
The oxidation number of each Group IIA element in a compound is _____
The oxidation number of each Group VIIA element in a compound is_____
+1; +2; -1
96
“The oxidation number of hydrogen is usually _____; however, its oxidation number is _____ in compounds with less electronegative elements
+1; -1
(Groups IA and IIA). Hydrogen is +1 in HCl, but –1 in NaH.”
97
“The sum of the oxidation numbers of all the atoms present in a neutral compound is. The sum of the oxidation numbers of the atoms present in a polyatomic ion is equal to
zero; the charge of the ion.”
98
“ The most common method for balancing redox equations is the _____________________, in which the equation is separated into two half-reactions—the oxidation part and the reduction part. Each half-reaction is balanced separately, and they are then added to give a balanced overall reaction.”
half-reaction method, also known as the ion–electron method
99
How to write net ionic equations
all aqueous compounds should be split into their constituent ions. Solid salts, on the other hand, should be kept together as a single entity.”
100
“Common oxidizing agents almost all contain ____
| Common reducing agents often contain ______
oxygen or a similarly electronegative element; metal ions or hydrides (H–).”
101
“A complete ionic equation accounts for all of the ions present in a reaction. To write a complete ionic reaction,
split all aqueous compounds into their relevant ions. Keep solid salts intact.”
102
“Net ionic equations ignore spectator ions to focus only on the species that actually participate in the reaction. To obtain a net ionic reaction,
subtract the ions appearing on both sides of the reaction, which are called spectator ions.”
“For reactions that contain no aqueous salts, the net ionic equation is generally the same as the overall balanced reaction
“For double displacement (metathesis) reactions that do not form a solid salt, there is no net ionic reaction because all ions remain in solution and do not change oxidation number.”
103
“are a type of redox reaction in which one element is both oxidized and reduced, forming at least two molecules containing the element with different oxidation states.”
Disproportionation (dismutation) reactions
104
How to determine oxidation state?
MFHOR
Monday-Friday Hor
```
Metals (+1 and +2)
Fluorine (-1)
Hydrogen (-1, 0 w carbon, +1 w electroneg)
Oxygen (-2)
Rest of Halogen and chalcogens (-1, -2)
```
