Ch. 9: Solutions (Complete) Flashcards
defn: solutions
homogeneous mixtures of two or more substances that combine to form a single phase (usually the liquid phase)
NOTE: almost all reactions in living organisms take place in solutions
the MCAT focuses mostly on solids dissolved in aqueous solutions, but combinations of the three phases of matter can also form solutions, give 3 examples.
- gases dissolved in liquids (carbonating soda)
- liquids dissolved in other liquids (ethanol in water)
- solids dissolved in other solids (metal alloys)
why are gases “dissolved” in gases considered as mixtures more than as solutions?
gas molecules do not interact all that much chemically
what is the relationship between mixtures and solutions?
all solutions are mixtures
not all mixtures are solutions
defn: solute
the part that is dissolved or dispersed in a solvent
what are 4 common examples of a solute?
- NaCL
- NH3
- C6H12O6
- CO2
defn: solvent
the component of the solution that remains in the same phase after mixing
what are 3 common examples of a solvent?
- H2O
- Benzene
- Ethanol
if the two substances are already in the same phase, how do you determine which is the solvent and which is the solute?
the solvent is the component present in greater quantity
how do you determine which is the solvent and which is the solute if the two substances are already in the same phase and are in equal proportions in the solution?
the component that is more commonly used as a solvent in other contexts is considered the solvent
how do solute molecules interact with the solvent?
solute molecules move about freely in the solvent and interact with it by way of intermolecular forces such as ion-dipole, dipole-dipole, or hydrogen bonding
why do chemical reactions occur easily in solution?
because dissolved solute molecules are also relatively free to interact with other dissolved molecules of different chemical identities
defn + aka: solvation
the electrostatic interaction between solute and solvent molecules
involves breaking intermolecular interactions between solute molecules and between solvent molecules and forming new intermolecular interactions between solute and solvent molecules together
aka: dissolution
what is solvation referred to as when water is the solvent?
hydration
when is solvation exothermic?
is this favored at high or low temperatures?
example?
when the new interactions are stronger than the original ones
- favored at lower temperatures
- ex: dissolution of gases into liquids (such as Co2 in to water)
explain why the dissolution of gases into liquids is exothermic
because the only significant interactions that must be broken are those between water molecules
CO2, as a gas, demonstrates minimal intermolecular interaction
when is solvation endothermic? why?
is this favored at higher or lower temperatures?
2 examples?
when new interactions are weaker than the original ones
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 (as heat) must be supplied to facilitate the formation of these weaker, less stable interactions
favored at higher temperatures
ex: dissolving ammonium nitrate or sugar into water
are most dissolutions endothermic or exothermic?
endothermic
defn: ideal solution
approximated by: overall strength of the new interactions is approximately equal to the overall strength of the original interactions
what is the enthalpy of dissolution for an ideal solution?
0
what are the three factors that affect the spontaneity of dissolution?
- enthalpy change for solute and solvent
- entropy change for solute and solvent (at constant temperature and pressure, entropy always increases upon dissolution)
- change in Gibbs free energy (spontaneous (deltaG < 0); nonspontaneous (deltaG > 0))
at constant temperature and pressure, does entropy increase or decrease for dissolution?
increases
define entropy in terms of microstates (entropy as a measure of molecular disorder)
entropy = the number of energy microstates available to a system at a given temperature = the molecules’ ability to move around in different ways
defn: solubility
the maximum amount of that substance that can be dissolved in a particular solvent at a given temperature
defn: saturated
when the max amount of solute has been added, the dissolved solute is in equilibrium with its undissolved state, the solution is saturated
what happens if more solute is added to a saturated solution?
it will not dissolve (it will precipitate to the bottom of the container)
defn: dilute vs. concentrated solutions
DILUTE = the proportion of solute to solvent is small
CONCENTRATED = the proportion of solute to solvent is large
can both dilute and concentrated solutions be unsaturated?
yes! if the maximum equilibrium concentration (saturation) has not yet been reached
what does gibbs free energy tell us about dissolution reactions?
Gibbs free energy for dissolution rxn is negative at a given temperature –> process is spontaneous –> solute is soluble
Gibbs free energy for dissolution rxn is positive at a given temperature –> process is nonspontaneous –> solute is insoluble
above what molar solubility are solutes generally considered soluble? what range of changes gibbs free energy is this associated with?
0.1 M in solution
large magnitude negative changes in free energy
defn: sparingly soluble salts
what range of changes gibbs free energy is this associated with?
solutes that dissolve minimally in the solvent (molar solubility under 0.1 M)
slight negative changes in free energy, so the equilibrium position lies closer to the undissociated (reactants) side of the reaction
defn + char: aqueous solution
the most common type of solution
solvent = water
rely on interactions between water molecules and solutions in solutions
in some solutions, the formation of a complex called the hydronium ion (H3O+) can occur
how does a hydronium ion (H3O+) form in solution?
the transfer of a hydrogen ion (H+) from a molecule in solution to a water molecule (H2O)
how is H+ found in solution? how is it not? why?
NOT: never found alone in solution bc a free proton is difficult to isolate
FOUND: bonded to an electron pair donor/carrier molecule such as a water molecule
what are the 7 general solubility rules for aqueous solutions? (only familiarity required)
- 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, except those formed with Ag+, Pb2+, Hg22+
- all salts of the sulfate ion (SO42-) are water-soluble, except those formed with Ca2+, Sr2+, Ba2+, Pb2+
- all metal oxides are insoluble, except those formed with alkali metals, ammonium, and CaO, SrO, and BaO, all of which hydrolyze to form solutions of the corresponding metal hydroxides
- all hydroxides are insoluble, except those formed with the alkali metals, ammonium, and Ca2+, Sr2+, Ba2+
- all carbonates (CO32-), phosphates (PO43-), sulfides (S2-), and sulfites (SO32-) are insoluble, except those formed with alkali metals and ammonium
what are the 2 solubility rules for aqueous solutions that you should absolutely remember?
- all salts of group 1 metals are soluble
- all nitrate salts are soluble
what are sodium and nitrate ions generally used for?
as counterions to what is actually chemically important
example: how should you interpret if a pH problem gives a sodium formate concentration as 0.10M?
it is really indicating that the concentration of the formate ion is 0.10 M because the sodium ion concentration does not affect pH
what is the only situation in which one needs to worry about the nitrate ion concentration? why?
what do you do in all other cases?
an oxidation-reduction reaction bc the nitrate ion can function as an oxidizing agent
in all other cases with nitrate ions: only focus on the cation as the chemically reacting species
defn: complex ion // coordination compound
a molecule in which a cation is bonded to at least one electron pair donor (which could include the water molecule)
defn: ligands
the electron pair donor molecules that the cation is bonded to in a coordination compound
what type of bonds hold together coordination complexes? explain this type of bond
coordinate covalent bonds
an electron pair donor (a lewis base) and an electron pair acceptor (a Lewis acid) form very stable Lewis acid-base adducts
what is the corresponding reaction involved in the formation of a complex ion?
complexation reaction
defn: chelation
a complex where the central cation can be bonded to the same ligand in multiple places
what is a typical requirement for chelation?
large organic ligands that can double back to form a second (or even third) bond with the central cation
what is chelation therapy typically used for?
to sequester toxic metals (lead, arsenic, mercury, etc)
defn: concentration
the amount of solute dissolved in a solvent
what are the 5 common ways that concentration is expressed on the MCAT?
what is the MOST common?
- percent composition by mass
- mole fraction
- molarity (most common)
- molality
- normality
eqn + use: percent composition by mass
a way of expressing concentration
use for aqueous solutions, metal alloys, and other solid-in-solid solutions
eqn + use: mole fraction of a compound
what will the sum of mole fractions in a system always equal?
sum of mole fractions in a system always equals 1
used to calculate the vapor pressure depression of a solution, and the partial pressures of gases in a system
eqn + use + symbol: molarity of a solution
indicated by [X]
used for rate laws, the law of mass action, osmotic pressure, pH and pOH, the Nernst equation
eqn + use: molality of a solution
for dilute aq. solns at 25 deg C, molality is approx. = to molarity bc the density of water is 1 kg/L at this temp
required in boiling point elevation and freezing point depression
why is molality approximate to molarity for dilute solutions at 25 deg C?
because the density of water is 1 kg/L at this temp
what happens as aqueous solutions become more concentrated with their solute? (2)
- their densities become much different from that of pure water
- most water-soluble solutes have molar masses much greater than that of water, so the density of the solution increases as the concentration increases
defn + how do we think of this + what do we need to know to calculate it: normality of a solution
N = the number of equivalents of interest/liters of solution
THINK OF THIS AS THE MOLARITY OF THE STUFF OF INTEREST IN THE REACTION
we need to know what purpose the solution serves BC we are concerned with the concentration of the reactive species
so it is reaction dependent
defn: equivalent
a measure of the reactive capacity of a molecule
a mole of the species of interest = protons, hydroxide ions, electrons, or ions
defn + eqn: solution after dilution
when solvent is added to a solution of higher concentration to produce a solution of lower concentration
M = molarity
V = volume
i and f = initial and final values, respectively
defn: saturation point
the point where the solute concentration is at its maximum value for the given temperature and pressure
the equilibrium position in the process of creating a solution
explain the equilibrium flux of a saturation point
- when the solution is dilute, the thermodynamically favored process is dissolution. the rate of dissolution > rate of precipitation
- as solution becomes more concentrated and approaches saturation, rate of dissolution lessens, rate of precipitation increases
- eventually, saturation point is reached (solution now exists in dynamic equilibrium state for which the rates of dissolution and precipitation are equal), concentration of dissolved solute reaches a steady-state value
what is the first step for any solution stoichiometry or solution equilibrium question?
write out the balanced dissociation reaction for the ionic compound in question
what is true of most solubility problems on the MCAT?
they deal with sparingly soluble salts (ionic compounds that have very low solubility in aqueous solutions)
defn + eqn: solubility product constant (Ksp)
for a saturated solution of an ionic compound with the formula AaBb, this is the equilibrium constant for its solubility in aqueous solution
where the concentrations of the ionic constituents are equilibrium (saturation concentrations)
why do Ksp expressions never have denominators?
because dissociation reactions by definition have a solid salt as a reactant and solids do not appear in equilibrium constants
how does Ksp change with increasing temperature for non-gas solutes? for gas solutes?
NON-gas solutes: Ksp increases as temp increases
GAS solutes: Ksp decreases as temp increases
how is Ksp affected by higher pressures for gas solutes?
higher pressures favor dissolution of gas solutes
Ksp will be larger for gases at higher pressures than at lower pressures
defn + eqn + use: ion product (IP)
what concentrations do we use when calculating this?
the reaction quotient (Q) of solubility
we may not know whether the system has reached saturation. so to determine where the system is with respect to the equilibrium position, we calculate IP
the concentrations used here are the concentrations of the ionic constituents at that given moment in time, which may differ from equilibrium concentrations
does each salt have the same or distinct Ksp at a given temperature?
each salt has its own distinct Ksp at a given temperature
explain and describe: IP < Ksp, IP = Ksp, and IP > Ksp (at a given set of conditions)
IP < Ksp – solution not yet at equilibrium and is UNSATURATED
- dissolution is thermodynamically favored over precipitation
IP > Ksp – solution is beyond equilibrium and is SUPERSATURATED
- thermodynamically unstable, and any disturbance to the solution (addition of more solid, further cooling) will cause spontaneous precipitation of the excess dissolved solute
IP = Ksp – solution AT equilibrium (rate dissolution = rate precipitation) and is SATURATED
how is a supersaturated solution formed?
created by dissolving solute into a hot solvent and then slowly cooling the solution
defn: molar solubility of a substance
the molarity of a solute in a saturated solution
if X moles of AmBn (s) can be dissolved in one liter of solution to reach saturation, then the moral solubility of AmBn (s) is X molar
does the formation of complex ions increase or decrease the solubility of a salt in solution?
increases
explain why (in basic terms) complexes are more stable in solution than isolated ions
- 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 inereactions
- this stabilizes the dissolution of the complex ion
- end result: such complexes tend to have very high Ksp values
defn: Kf (formation or stability constant)
the subsequent formation of the complex ion in solution
why must there be a distinction between Kf and Ksp?
when you form a complex ion, you have to use a mixture of solutions
THUS a distinction must be made between Ksp of the SOLUTION and of the complex ion itself
which is larger, Kf or Ksp?
Kf of the complex ion is significantly larger than the Ksp of the compound providing the metal ion
defn: common ion effect
reduction in molar solubility that results from a salt being dissolved in a solution that already contains one of its constituent ions (as compared to its solubility in a pure solvent)
does the common ion effect have an effect on Ksp?
no, just results in a reduction of the molar solubility of the salt
explain why the common ion effect is le chatelier’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 shift left, reforming the solid salt
as a result, the molar solubility for the solid is reduced, and less of the solid dissolves in the solution, although the Ksp remains constant
how can the common ion effect be used to separate out specific compounds in a solution mixture?
by adding an appropriate counterion in excess, the dissociation reaction shifts to the left, forming the solid salt
defn: 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
usually associated with dilute solutions
what are the 4 colligative properties?
- vapor pressure depression
- boiling point elevation
- freezing point depression
- osmotic pressure
defn + func + eqn: Raoult’s law
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
explain why Raoult’s law is valid on a molecular level
the presence of the solute molecules can block the evaporation of solvent molecules but not their condensation
this reduces the vapor pressure of the solution compared to the pure solvent
Raoult’s law holds only under what situation?
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
defn: ideal solutions
solutions that obey Raoult’s law
defn: boiling point
the temperature at which the vapor pressure of the liquid equals the ambient (incident) pressure
what happens to the boiling point 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
eqn: boiling point elevation
Kb is a proportionality constant characteristic of a particular solvent
defn: van’t Hoff factor
corresponds to the number of particles into which a compound dissociates in solution
why do covalent molecules have i (van’t Hoff factor) values of 1?
because they don’t readily dissociate in water!
why must a greater amount of energy be removed from solution in order for the solution to solidify when there are solute particles present?
the presence of solute particles in a solution interferes with the formation of the lattice arrangement of solvent molecules associated with the solid state
eqn: freezing point depression
Kf is a proportionality constant characteristic of a particular solvent
defn (formal + layman’s) + eqn: osmotic pressure
a “sucking” pressure generated by solutions in which water is drawn into a solution
the amount of pressure that must be applied to counteract this attraction of water molecules for the solution
does water move in the direction of higher or lower solute concentration?
higher
