Chapter 9: Solutions
Colligative properties
physical properties of solutions that are dependent on the concentration of dissolved particles but not on the chemical identity of the dissolved particles These properties—vapor pressure depression, boiling point elevation, freezing point depression, and osmotic pressure—are usually associated with dilute solutions.
Complex ion (coordination compound) -ligands
refers to a molecule in which a cation is bonded to at least one electron pair donor (which could include a water molecule) the electron pair donor molecules are called ligands complexes are held together by coordinate covalent bonds. an electron pair donor (Lewis base) and an electron pair acceptor (Lewis acid) form very stable Lewis acid-base adducts
Ideal Solution
Sometimes the overall strength of the new interactions is approximately equal to the overall strength of the original interactions In this case the overall enthalpy change for the dissolution is close to zero These types of solutions approximate the formation of an ideal solution for which the enthalpy of dissolution is equal to zero
mole fraction (X)
sum of mole fractions in a system will always equal 0 The mole fraction is used to calculate the vapor pressure depression of a solution, as well as the partial pressures of gases in a system
Solvation (dissolution)
the electrostatic interaction between solute and solvent involved breaking intermolecular interactions between solute molecules and between solvent molecules and forming new intermolecular interactions between solute and solvent molecules together
Molar solubility
the molar solubility of a compound is its concentration (in moles per liter) at equilibrium at a given temperature. If X moles of AmBn (s) can be dissolved in one liter of solution to reach saturation, then the molar solubility of AmBn (s) is X molar.
Dilute
the proportion of solute to solvent is small
Concentrated
the proportion of solvent to solute is small can still be unsaturated if max saturation has not been reached
R value
.0812 L x atm/molxK
Kb of water
.512 K x kg/mol
Which of the following will cause the greatest increase in the boiling point of water when it is dissolved in 1.00 kg H2O?
0.5 mol iron(III) nitrate The equation to determine the change in boiling point of a solution is as follows: ΔTb = iKbm. m is the molality of the solution, and Kb is the boiling point elevation constant. In this case, the solvent is always water, so Kb will be the same for each solution. What is needed is the number of dissociated particles from each of the original species. This is referred to as the van't Hoff factor (i) and is multiplied by molality to give a normality (the concentration of the species of interest—in this case, all particles). The normality values determine which species causes the greatest change in boiling point. 2x.4 = .8 for CaSo4 .5 x 4 = 2.0 for Fe(NO3)3 1 x 1 = 1 for sucrose
Solubility Rules
1. All salts containing ammonium NH4+ and alkali metal (Group 1) cations are watersoluble 2. All salts containing nitrate (NO3-) and acetate (CH3COO-) anions are water soluble 3. Halides (Cl-, Br-, I-), excluding fluorides are water soluble with the exceptions of those formed with Ag2+, Pb2+, and Hg22+ 4. All salts of the sulfate ion (SO42-) are water soluble, with the exceptions of those formed with Ca2+, Sr2+, Ba2+, Pb2+ 5. 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 6. All metal hydroxides are insoluble, with the exception of those formed with the alkali metals, ammonium and Ca2+, Sr2+ and Ba2+ 7. All carbonates (CO32-), phosphates (PO43-), Sulfides (S2-) and sulfites (SO32-) are insoluble with the exception of those formed with the alkali metals and ammonium
Kf of water
1.86 K x kg/mol
Two organic liquids, pictured in the figure below, are combined to form a solution. Based on their structures, will the solution closely obey Raoult's law? benzene and toluene
Benzene and toluene are both organic liquids and have very similar properties. They are both nonpolar and are almost exactly the same size. Raoult's law states that ideal solution behavior is observed when solute-solute, solvent-solvent, and solute-solvent interactions are all very similar. Therefore, benzene and toluene in solution will be predicted to behave as a nearly ideal solution.
A saturated solution of cobalt(III) hydroxide (Ksp = 1.6 × 10−44) is added to a saturated solution of thallium(III) hydroxide (Ksp = 6.3 × 10−46). What is likely to occur?
Both thallium(III) hydroxide and cobalt(III) hydroxide precipitate. Since both salts have a formula MX3 (one of one particle, three of another), it is possible to directly compare the molar solubilities of each. When the solutions are mixed, [OH−] is above saturation levels for both the cobalt and the thallium in the solution. Since thallium(III) hydroxide has a smaller Ksp than that of cobalt(III) hydroxide, it will react first. The ion product of the mixed solution is higher than the Ksp for thallium(III) hydroxide, and the system will shift left to precipitate solid thallium(III) hydroxide. After the thallium(III) hydroxide precipitates, a small excess of OH- will remain, which gives an ion product slightly above the K of sp cobalt (III) hydroxide. This will cause a small amount (1%-3%) of cobalt(III) hydroxide to also precipitate.
What is a colligative property?
Colligative properties are those that depend on the amount of solute present, but not the actual identity of the solute particles. Examples include vapor pressure depression, boiling point elevation, freezing point depression, and osmotic pressure.
The entropy change when a solution forms can be expressed by the term ΔS°soln. When water molecules become ordered around an ion as it dissolves, the ordering would be expected to make a negative contribution to ΔS°soln. An ion that has more charge density will have a greater hydration effect, or ordering of water molecules. Based on this information, which of the following compounds will have the most negative contribution to ΔS°soln?
CaS CaS will cause the most negative contribution to ΔS°soln through hydration effects because the Ca2+ and S2− ions have the highest charge density compared to the other ions. All of the other ions have charges of +1 or -1,whereas Ca2+ and S2− each have charges with a magnitude of 2.
The process of formation of a salt solution can be better understood by breaking the process into three steps: Breaking the solute into its individual components Making room for the solute in the solvent by overcoming intermolecular forces in the solvent Allowing solute-solvent interactions to occur to form the solution
Endothermic, endothermic, exothermic
Which of the following combinations of liquids would be expected to have a vapor pressure higher than the vapor pressure that would be predicted by Raoult's law?
Ethanol and hexane Mixtures that have a higher vapor pressure than predicted by Raoult's law have stronger solvent-solvent and solute-solute interactions than solvent-solute interactions. Therefore, particles do not want to stay in solution and more readily evaporate, creating a higher vapor pressure than an ideal solution. Two liquids that have different properties, like hexane (hydrophobic) and ethanol (hydrophilic, small) in (A), would not have many interactions with each other and would cause positive deviation; i.e. higher vapor pressure. (B) and (C) are composed of liquids that are similar to one another and would not show significant deviation from Raoult's law. (D) contains two liquids that would interact very well with each other, which would actually cause a negative deviation from Raoult's law—when attracted to one other, solutes and solvents prefer to stay in liquid form and have a lower vapor pressure than predicted by Raoult's law.
Molality (m)
For dilute aqueous solutions at 25°C, the molality is approximately equal to molarity because the density of water at this temperature is 1 kilogram per liter. However, note that this is an approximation and true only for dilute aqueous solutions. As aqueous solutions become more concentrated with solute, their densities become significantly different from that of pure water; most water-soluble solutes have molar masses significantly greater than that of water, so the density of the solution increases as the concentration increases. used in BP elevation and FP depression
Name two ions that form salts that are always soluble:
Group I metals, ammonium, nitrate, and acetate salts are always soluble.
Unsaturated
IP < Ksp For unsaturated solutions, dissolution is thermodynamically favored over precipitation
Saturated
IP = Ksp 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
Supersaturated
IP > Ksp solution is beyond equilibrium, and the solution is considered supersaturated. It is possible to create a supersaturated solution by dissolving solute into a hot solvent and then slowly cooling the solution. A supersaturated solution is thermodynamically unstable, and any disturbance to the solution, such as the addition of more solid solute or other solid particles, or further cooling of the solution, will cause spontaneous precipitation of the excess dissolved solute
Coenzymes and cofactors
Many coenzymes (vitamins) and cofactors also contain complexes of transition metals such as cobalamin (vitamin B) The presence of a transition metal allows coenzymes and cofactors to bind other ligands or assist with electron transfer
Which of the following explanations best describes the mechanism by which solute particles affect the melting point of ice?
Melting point is depressed because solute particles interfere with lattice formation. Melting point depresses upon solute addition, making (A) and (B) incorrect. Solute particles interfere with lattice formation, the highly organized state in which solid molecules align themselves. Colder-than- normal conditions are necessary to create the solid structure.
How are molality and molarity related for water? How are they related for other solvents?
Molarity (M) and molality (m) are nearly equal at room temperature. This is only because 1 L solution is approximately equal to 1 kg solvent for dilute solutions (the denominators of the molarity and molality equations, respectively). For other solvents, molarity and molality differ significantly because their densities are not 1 g/mL like water.
all salts of Group I metals and all nitrate salts (NO3-) are soluble
Sodium and nitrate ions are generally used as counterions to what is actually chemically important for example if a pH problem gives a sodium formate concentration of 00 M it is really indicating that the concentration of the formate ion is 00 M because the sodium ion concentration does not affect pH The only time one needs to worry about the nitrate ion concentration is in an oxidationreduction reaction for the nitrate ion can functionalthough only weaklyas an oxidizing agent In all other cases with nitrate ions only focus on the cation as the chemically reacting species
Describe the differences between solubility and saturation:
Solubility is the amount of solute contained in a solvent. Saturation refers to 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.
What is one way in which solubility of a compound can be increased?
Solubility of solids can be increased by increasing temperature. Solubility of gases can be increased by decreasing temperature or increasing the partial pressure of the gas above the solvent (Henry's law).
Ksp is temp dependent (sometimes pressure) and never have a denominator
Solubility product constants, like all other equilibrium constants (Keq, Ka, Kb, and Kw) are temperature dependent. When the solution consists of a gas dissolved into a liquid, the value of the equilibrium constant, and hence the position of equilibrium (saturation), will also depend on pressure
Describe the process of solvation.
Solvation refers to the breaking of intermolecular forces between solute particles and between solvent particles, with formation of intermolecular forces between solute and solvent particles. In an aqueous solution, water is the solvent.
Ion Product (IP)
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 the ion product (IP), which is analogous to the reaction quotient, Q, for other chemical reactions. The difference is that the concentrations used in the ion product equation are the concentrations of the ionic constituents at that given moment in time, which may differ from equilibrium concentrations. As with the reaction quotient Q, the utility of the ion product lies in comparing its value to that attained at equilibrium, Ksp. Each salt has its own distinct Ksp at a given temperature
Normality (N)
The normality (N) of a solution is equal to the number of equivalents of interest per liter of solution. An equivalent is a measure of the reactive capacity of a molecule. Most simply, an equivalent is equal to a mole of the species of interest—protons, hydroxide ions, electrons, or ions. To calculate the normality of a solution, we need to know what purpose the solution serves because it is the concentration of the reactive species with which we are concerned. For example, in acid-base reactions, we are most concerned with the concentration of hydrogen ions; in oxidation- reduction reactions, we are most concerned with the concentration of electrons. Normality is unique among concentration units in that it is reaction dependent. For example, in acidic solution, 1 mole of the permanganate ion (MnO4−) will readily accept 5 moles of electrons, so a 1 M solution would be 5 N. However, in alkaline solution, 1 mole of permanganate will accept only 1 mole of electrons, so in alkaline solution, a 1 M permanganate solution would be 1 N.
Freezing point depression
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 greater amount of energy must be removed from the solution (resulting in a lower temperature) in order for the solution to solidify.
When ammonia, NH3, is used as a solvent, it can form complex ions. For example, dissolving AgCl in NH3 will result in the complex ion [Ag(NH3)]+2 . What effect would the formation of complex ions have on the solubility of a compound like AgCl in NH3?
The solubility of AgCl will increase because complex ion formation will consume Ag+ ions and cause the equilibrium to shi away from solid AgCl. Formation of complex ions between silver ions and ammonia will cause more molecules of solid AgCl to dissociate. The equilibrium is driven toward dissociation because the Ag+ ions are essentially being removed from solution when they complex with ammonia. This rationale is based upon Le Châtelier's principle, stating that when a chemical equilibrium experiences a change in concentration, the system will shi to counteract that change.
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. This reduction in molar solubility is called the common ion effect. Pay attention to the effect of the common ion: its presence results in a reduction in the molar solubility of the salt. Note, however, that the presence of the common ion has no effect on the value of the solubility product constant itself. For example, if a salt such as CaF2 is dissolved into water already containing Ca2+ ions (from some other salt, perhaps CaCl2), the solution will dissolve less CaF2 than would an equal amount of pure water. The common ion effect is really Le Châtelier's principle in action. Because the solution already contains one of the constituent ions from the products side of the dissociation equilibrium, the system will shi toward the le side, reforming the solid salt. As a result, molar solubility for the solid is reduced, and less of the solid dissolves in the solution—although the Ksp remains constant.
Vant hoff factor
The van't Hoff factor corresponds to the number of particles into which a compound dissociates in solution. For example, i = 2 for NaCl because each formula unit of sodium chloride dissociates into two particles—a sodium ion and a chloride ion—when it dissolves. Covalent molecules such as glucose do not readily dissociate in water and thus have i values of 1.
Boiling point elevation
When a nonvolatile solute is dissolved into a solvent to create a solution, the boiling point of the solution will be greater than that of the pure solvent. The boiling point is the temperature at which the vapor pressure of the liquid equals the ambient (incident) pressure. We've just seen that adding solute to a solvent results in a decrease in the vapor pressure of the solvent in the solution. If the vapor pressure of a solution is lower than that of the pure solvent, then more energy (and consequently a higher temperature) will be required before its vapor pressure equals the ambient pressure. Kb values vary between each solvent
Exothermic salvations
When the new interactions are stronger than the original, solvation is exothermic and the process is favored at low temp The dissolution of gases into liquids such as CO into water is an exothermic process because the only significant interactions that must be broken are those between water moleculesCO as a gas demonstrates minimal intermolecular interaction Le Châteliers principle tells us this is the reason that lowering the temperature of a liquid favors solubility of a gas in the liquid (doesn't need as much energy to break bonds)
Molarity (M)
[A] used for rate laws, Law of mass action, osmotic pressure, pH and pOH, and the Nernst equation
Raoult's Law
accounts for vapor pressure depression caused by solutes in solution as solute is added to a solvent the vapor pressure decreases proportionately On a molecular level, the presence of the solute molecules can block the evaporation of solvent molecules but not their condensation.
Which phases of solvent and solute can form a solution? I. solid solvent, gaseous solute II. solid solvent, solid solute III. gaseous solvent, gaseous solute
all three All three choices can make a solution as long as the two components create a mixture that is of uniform appearance (homogeneous). Hydrogen in platinum is an example of a gas in a solid. Brass and steel are examples of homogeneous mixtures of solids. The air we breathe is an example of a homogeneous mixture of gases; while these are more commonly simply referred to as mixtures, they still fit the criteria of a solution.
Biological complexes
complex ions have profound biological applications in macromolecules such as proteins For instance many active sites of proteins utilize complex ion binding and transition metal complexes to carry out their function One classic example is the iron cation in hemoglobin which can carry oxygen carbon dioxide and carbon monoxide as ligands The iron in hemoglobin can bind various gases leading to the formation of oxyhemoglobin O2 carbaminohemoglobin CO2 and carboxyhemoglobin CO
Concentration
denotes the amount of solute dissolved in a solvent
Osmotic Pressure
efers to a "sucking" pressure generated by solutions in which water is drawn into a solution. Formally, the osmotic pressure is the amount of pressure that must be applied to counteract this attraction of water molecules for the solution. Water moves in the direction of higher solute concentration. For instance, pure water (no solute concentration) will traverse a semipermeable membrane to a solution containing solute particles (such as NaCl) and increase the level of the water as a result
Entropy and micro states
entropy can be thought of as the degree to which energy is dispersed throughout a system or the amount of energy distributed from the system to the surroundings at a given temperature Another way to understand entropy is the measure of molecular disorder or the number of energy microstates available to a system at a given temperature When solid sodium chloride dissolves into water the rigidly ordered arrangement of the sodium and chloride ions is broken up as the ionion interactions are disrupted and new iondipole interactions with the water molecules are formed The ions freed from their lattice arrangement have a greater number of energy microstates available to them in simpler terms they are freer to move around in different ways and consequently their energy is more distributed and their entropy increases The water however becomes more restricted in its movement because it is now interacting with the ions The number of energy microstates available to it that is the water molecules ability to move around in different ways is reduced so the entropy of the water decreases In the end the increase in the entropy experienced by the dissolved sodium chloride is greater than the decrease in the entropy experienced by the water so the overall entropy change is positive energy is overall dispersed by the dissolution of sodium chloride in water
Solvent
ex. H2O, benzene, ethanol the component of the solution that remains in the same phase after mixing if two same phase components are mixed, the solvent is the component present In greater quantity. if equal amount of both, which is which will be context dependent
Solute
ex. NaCl, NH3, C6H12O6, CO2 solutes are dissolved (dispersed) in a solvent
Kf (formation or stability constant)
formation of the complex ion in solution
Solutions
homogenous (the same throughout) mixtures of two or more substances that combine to form a single phase ; usually the liquid phase
Chelation
in some complexes, the central cation can be bonded to the same ligand in multiple places this is called chelation and generally requires large organic ligands that can double back to form a second (or even third) bond with the central cation Chelation therapy is oen used to sequester toxic metals; lead arsenic mercury and so on
Molar Solubility
molarity of a solute in a saturated solution is called the molar solubility
Solubility
solubility of a substance is the maximum amount of that substance than can be dissolved in a particular solvent at a given temperature
Ksp
solubility product constant increases with increasing temp for non gas solutes and decreases for gas solutes Higher pressures favor dissolution of gas solutes, and therefore the Ksp will be larger for gases at higher pressures than at lower ones.
What is considered soluble
solutes are considered soluble if they have a molar solubility above .1 M in solution
Sparingly soluble salts
solutes that dissolve minimally in the solvent (molar solubility under .1 M)
Hydration
solvation when water is the solvent
Aqueous solution
solvent is water (aq)
Structure of the tretraaquadioxouranyl complex cation
water and oxygen act as ligands with a U+6 cation
Endothermic solvations
when new interactions are weaker than the original ones, solvation is endothermic and the process is favored at high temp 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
Gibbs free energy and solubility
when non-spontaneous, solute Is insoluble when spontaneous, solute is soluble
Saturated -precipitate
when the maximum amount of solid has been added, the dissolved solute is in equilibrium with its undissolved state if more solute were to be added it will not dissolve and would precipitate
Saturation point
where the solute concentration is at its maximum value for the given temperature and pressure When the solution is dilute (unsaturated), the thermodynamically favored process is dissolution, and initially, the rate of dissolution will be greater than the rate of precipitation. As the solution becomes more concentrated and approaches saturation, the rate of dissolution lessens, while the rate of precipitation increases.
The salt KCl is dissolved in a beaker. To an observer holding the beaker, the solution begins to feel colder as the KCl dissolves. From this observation, one could conclude that:
ΔS°soln is large enough to overcome the unfavorable ΔH°soln. Dissolution is governed by enthalpy and entropy, which are related by the equation ΔG°soln = ΔH°soln − TΔS°soln. The cooling of the solution indicates that heat is used up in this bond-breaking reaction. In other words, dissolution is endothermic, and ΔH is positive. The reaction is occurring spontaneously, so ΔG must be negative. The only way that a positive ΔH can result in a negative ΔG is if entropy, ΔS, is a large, positive value as in (A). Conceptually, that means that the only way the solid can dissolve is if the increase in entropy is great enough to overcome the increase in enthalpy. (B) is incorrect because it is clearly stated in the question stem that KCl dissolves; further, all salts of Group 1 metals are soluble. (C) is incorrect because ΔS°soln must be positive in order for KCl to dissolve. Finally, (D) is incorrect because solute dissolution would cause the boiling point to elevate, not depress. It is also not a piece of evidence that could be found simply by observing the beaker's temperature change.
