General Chemistry MCAT- Chapter 9: Solutions

Osmotic pressure

"sucking pressure" generated by solutions in which water is drawn into a solution. Formally: The amount of pressure that must be applied to counteract the attraction of water molecules for the solution. i = van't Hoff, M is molarity. Water moves in direction of higher solute concentration.

Mole fraction (X)

(Moles of solute per total moles) is used for calculating vapor pressure depression and partial pressure of gases in a system Xa = moles of A/moles of all species Sum of mole fractions equal 1

Molarity (M)

(moles of solute per liters of solution) is the most common unit for concentration and is used for rate laws, the law of mass action, osmotic pressure, pH and pOH, and the Nernst equation. -Concentration of molarity is sometimes presented using brackets

Molality (m)

(moles solute/kilograms solvent) used for boiling point elevation and freezing point depression. -The molality is equal to molarity when a dilute aqueous solutions is @ 25°C, b/c at that point, the density of water at this temp. is 1Kg/L

Normality

(number of equivalents per liters of solution) is the molarity of the species of interest and is used for acid-base and oxidation-reduction reactions. -it is a reaction dependent Ex: 1 mole of permanganate ion (MnO4^-) will accept 5 mol of electrons, so 1 M solution would be 5 N. But, in alkaline solution, permanganate will accept 1 mole of electrons, so 1 M of permanganate solution would be equivalent to 1 N. An equivalent= its a measure of the reactive capacity of a molecule; its equal to a mole of the species of interest (protons, hydroxide ions, electrons, ions)

When solutes have a slightly -ve changes in free energy, will that dissolve or not?

those kind of solutions are closer to the undissociated (reactants) side of the reaction, so those solutes will have minimal dissolve in the solvent (meaning that their solubility is under 0.1 M)

Which one is thermodynamically favored at equilibrium, dissolution or precipitation?

Neither dissolution nor precipitation is more thermodynamically favored at equilibrium b/c favoring either would result in a solution that is no longer in a state of equilibrium

Complex ion (coordination compound)

refers to a molecule in which a cation is bonded to at last one electron pair donor (ligands). This electron pair donor could include a water molecule. Complexes held together with COORDINATE COVALENT BONDS they have a variety of physical & chemical properties they have a wide range of solubilities and chemical reactions Often used in protein active sites to carry out their functions. e.g. Iron cation in hemoglobin can carry O2, CO2 and CO (carbon monoxide) as ligands. -the solubility of complex ions is determined by the Ksp -the formation of complex ions causes an increase in the solubility of a salt in solution -complex ion that contains multiple polar bonds btw the ligands and the central metal ion, it should be able to engage in a very large amount of dipole-dipole interactions; this stabilizes the dissolution of the complex ion. The end result is that such complexes tend to have very high Ksp values. -they increase the solubility of a substrate

Oxidation-reduction reactions concerned about the concentration of...

electrons

What can effect the solubility product constants (Keq,Ka,Kb,Kw)?

temp. & pressure

What happens to the boiling point when a nonvolatile solute is dissolved into a solvent?

the boiling point of the solution will be greater than that of the pure solvent

ion product (IP) vs. Ksp

the difference is the 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 -IP compare its values that was reached to the Ksp -each salt has a distinct Ksp at a given temp.

Molar solubility

the molarity of the solute at saturation

What effects can temperature have on the "solubility product constant"?

the solubility product constant increases w/ increasing temperature for non-gas solutes, and decreases for gas solutes

IP = Ksp

the solution is at equilibrium (the rate of dissolution and precipitation are equal) and the solution is considered saturated

IP > Ksp

the solution is beyond equilibrium and the solution is considered supersaturated For a supersaturated solutions: they're thermodynamically unstable b/c any disturbance to the solution will cause spontaneous precipitation of the excess dissolved solute -precipitation will occur

IP < Ksp

the solution is not yet at equilibrium and the solution is considered unsaturated For unsaturated solutions: dissolution is thermodynamically favored over precipitation -solute will continue to dissolve

Boiling point

the temperature at which the vapor pressure of the liquid equals the ambient (incident) pressure

Which one is considered to be the solvent if the solvent and the solute is present in equal amount in the solution?

then the component that is more commonly used as a solvent in other contexts is considered the solvent

boiling point elevation formula:

- the formula is used when: the boiling point of a solution is raised relative to that of the pure solvent △Tb= increase in boiling point i= van't Hoff factor Kb= proportionality constant of a solvent (given on the MCAT) m= molality

The application of common ion effect:

-A common ion decrease solubility of compounds. Hence it helps in the formation of precipitation of a particular salt whose ion is common. -We can take the advantage of this property to separate (or precipitate) our desired salt from a solution of different salts. -Ex: -We can separate AgCl(s) form a solution containing different salt solutions by adding NaCl that contains common ion Cl- -As we add NaCl, AgCl(s) separates out as precipitate by turning into solids NOTE: NaCl remains in the aqueous solution. Only AgCl(s) precipitates into solids.

Coordinate covalent bonds

-Complexes are held together with these bonds, in which an electron pair donor (a Lewis base) and an electron pair acceptor (a Lewis acid) form a stable Lewis acid-base product

What is significant about IP (ion product)?

-Its analogous to the reaction quotient (Q) -Its the same equation used for "solubility product constant"

Aqueous solutions (aq)

-Solvent is water. -relay on the interactions btw water molecules and solutes in solution -In an acidic solutions, Hydronium ion (H3O+) can be formed by transfer of a hydrogen proton ion (H+) from a molecule in solution to a water molecule. H+ never found on its own in solution b/c a free proton is difficult to isolate. So for that reason, it is found bounded to an electron pair donor (carier). Ex: H3O+

Solvation

-always wants to achieve equilibrium -equilibrium= its the lowest energy state of a system under a given set of temperature and pressure conditions spontaneous movement= when a system move towards equilibrium nonspontanous movement= when a system move away from equilibrium

Colligative properties

-are physical properties of solutions that depend on the concentration of dissolved particles but NOT on the chemical identity of the dissolved particles. -physical properties of solution that are affected by the # of solute particles. Ex: Boiling point, osmotic pressure, freezing pint, vapor pressure -usually associated with dilute solutions

Whats another example of somethings that would depend complex ions to function?

-coenzymes (vitamins) and cofactors they contain transition metal complexes which assist the coenzymes and cofactors to bind other ligands OR they could assist them with electron transfer.

Freezing point depression (decrease):

-its a colligative property (meaning that its depends on the concentration of particles, not no their identities. Presence of solute particles in a solution interferes with the formation of the solid lattice arrangement of solvent particles. Therefore a greater amount of energy must be removed from the solution (resulting in a lower temperature) in order for the solution to solidify. Ex: pure water freezes @ 0°C, but for very mole of solute dissolved in 1 Kg of water, the freezing point is lowered by 1.86°C Kf for water = 1.86 K.kg/mol -its given on the test △Tf= freezing point depression i= van't Hoff factor Kf= proportionality constant of a solvent m= molarity of the solution

what must happen when forming a complex ion?

-one must use a mixture of solutions; and you must distinct the Ksp solution from the complex ion. This is how we label them: Ksp= its the dissolution of the original solution Kf (the formation/stability constant)= its when the complex ion is formed in a solution

what happens after we dissolve the solute in the saturated solution?

-precipitation of the solute will begin to occur -this usually happens after the solution becomes more concentrated and approaches saturation, causing the rate of dissolution to decrease and the rate of precipitation to increase

solute

-their molecules move freely in solvent and interact with the solvent by intermolecular forces (ion-dipole, dipole-dipole, hydrogen bonding) -they interact freely w/ other dissolved molecules of different chemical identities, resulting in a chemical reaction in solution

Dissociation reactions

-they have a solid salt as a reactant -this causes Ksp to never have a denominator (reactant)

Density of water...

1 g/mL

What is the density of water at room temperature?

1 g/ml

Two factors in spontaneity of dissolution:

1) Enthalpy change AND entropy - at constant temperature and pressure entropy always increases upon dissolution. 2) Gibbs free energy is the key here: -spontaneous process associated w/ a decrease in free energy -nonspontaneous process associated w/ increase in free energy NOTE: solutions may form spontaneously for an endothermic & exothermic dissolutions

Concentration expressions

1) Percent composition by mass 2) Mole fraction 3) Molarity 4) Molality 5) Normality

What can disturb a supersaturated solution?

1) addition of more solid solute or other solid particles 2) or doing further cooling of the solution

What happens to a solution that contains a solute?

1) vapor pressure lowers 2) boiling point rises 3) freezing point lowers 4) osmotic pressure increase *They all depend on the molality of a solution (moles of solute/Kg of solvent) = m *molality is going to show how much all of those are effected (1-4)

Solubility is dependent on:

1)temperature od the solution 2)the solvent 3)pressure 4)the addition of other substances to the solution

Concentrated solution

A solution in which the proportion of solute to solvent is LARGE -considered unsaturated if the maximum equilibrium conc. (saturation) has not been reached

Dilute soultion

A solution in which the proportion of solute to solvent is SMALL -considered unsaturated if the maximum equilibrium conc. (saturation) has not been reached

Raoult's law

Accounts for "vapor pressure depression" caused by solutes in solution. -As solute is added to solvent, vapor pressure of solvent decreases proportionately. -Vapor pressure depression also explains boiling point elevation -- as the vapor pressure decreases, the temperature (energy) required to boil the liquid must be raised. -this law states that ideal solution behavior is observed when solute-solute, solvent-solvent, and solute-solvent interactions are very similar -mixtures that have a higher vapor pressure than presidctee by Raoult's law have stronger solvent-solvent and solute-solute interactions than solvent-solute interactions. Therefore, particles don't want to stay in solution and more readily evaporate, creating a higher vapor pressure than ideal solution pg. 321 -Initially we started with pure solvent A. -For pure A, the mole fraction Xa = 1, and Xb = 0. -and vapor pressure for pure A = P^oA -Now when we completely replace compound A by compound B, there will be only pure compound B on extreme right. -This graph simply shows how vapor pressure changes when two different solvents of different vapor pressure are mixed at differet proportion. -The vapor pressure of a mixture always lies in between the vapor pressure of ther respective pure form. -They want to tell us that when we add a solute to a solvent, the vapor pressure decreases. -the solute should be non-volatile. -From left to right, the mole fraction of B increases. -Also from left to right the vapor pressure of the mixture decreases. -Hence as the concentration of B in A increases, the vapor pressure decreases. -This conforms that addition of a non-volatile solute decreases vapor pressure.

Mixtures vs. solutions

All solutions are mixtures, but not all mixtures are solutions.

This is how ionic solid introduced into polar solvent would dissociate:

AmBn(s) --> mA^n+ + nB^m-

Differences in solubility

Based on Gibbs free energy: -when △G for dissolution is negative at given temp., will happen spontaneously (the solute is soluble) solutes are soluble if they have a molar solubility above 0.1 M in solution -when △G is positive, at given temp., will happen nonspontanously (the solute is insoluble) Solutes are insoluble if they have molar solubility under 0.1M in solution.

Formation of complex ions in regards to solubility?

Complex ions increase solubility of otherwise insoluble ions (the opposite of the common ion effect).

van 't Hoff factor (i)

For solutes that can dissociate, (i) can be used in freezing point depression and boiling point elevation, and osmotic pressure calculations.

Mixtures

Gas dissolved into another gas Since don't interact chemically -Not all mixtures are solutions

Solutions

Homogenous mixtures composed of two or more substances. They combine to form a single phase, generally a liquid phase. -all solutions are mixtures

What type of solutions obey Raoult's law?

Ideal solutions

Solvent

In a solution, the substance in the most abundance and the one in which the solute dissolves in. 1) Dissolves solute particles via electrostatic interactions in a process called solvation or dissolution.

Solubility rules

In aqueous solutions, there are seven solubility rules: pg.304&305 Bunch of rules. Know these: 1. All salts of group 1 metals are soluble 2. All nitrate salts are soluble -Also know that Pb2+ and Ag+ are insoluble -Sodium and nitrate ions are used as counter ions. Ex: Sodium formate has a concentration of 0.10 M, but what its really saying that the concentration of formate ion is 0.10 M, b/c sodium ion conc. doesn't effect pH -The only time to worry about nitrate ion conc. is during oxidation & reduction reactions

Stability constant (Kf)

Is the equilibrium constant for complex ion formation. Its value is usually much greater than Ksp. The formation constant increases the solubility of other salts containing the same ions because it uses up the products of those dissociation reactions, shifting the equilibrium to the right. In another words: products are being used up to form the complex ion, and as a result, the reactants would shift to the right to increase the product (thus increase complex ion formation)

How does the common ion effect the solubility product constant?

It does not effect the Ksp, only the molar solubility.

How does the addition of solute molecules effect evaporation and condensation?

It lowers evaporation, but does not affect condensation.

Saturation (equilibrium concentration)

Maximum amount of solute is added, and the dissolved solute is in equilibrium with undissolved state. Any more solute added will not dissolve (precipitate out), so if we add more solute, its will not dissolve.

Which formula is used to find the concentration of solution after dilution?

MiVi = MfVf M = molarity V= volume NOTE: the MCAT might use the term "part-per" to indicate concentration of a dissolved substance in a solution (such as water) "parts-per-million" is equal to ppm 10^-6 1 ppm (10^-6) equals to 1 mg/L of water -parts per million means 1 part of a substance (solute) is present in 1 million part (10^6 ) of solvent(usually water) -ppm unit is driven by mass/volume of water. -Density of water is 1 g/mL -So mass of 1L water = 1000 g -So mass of 1L water is 10^6 mg -i.e the mass of 1L water is 1 million (10^6) times of of an mg. Mass of 1L water = 1000g * [1000mg/1g] = 1000000mg = 10^6mg -So if 10 mg of a substance is present in 1L of water, it will be 10 ppm -e.g A water sample contains 100 ppm lead means there are 100 mg of lead in 1L of water -So 100 ppm lead = 100 mg lead / 1L water -mg/L is quivale to ppm

When does Raoult's law applied?

Only when attraction between molecules of different components of the mixture is equal to the attraction between the molecules of any one component in its pure state.

Ideal solution

Overall strength of new interactions approximately equal to overall strength of original interactions. Enthalpy change for dissolution is close to zero.

Raoul's law (vapor pressure depression)

PA= vapor pressure of solvent A when solutes are present XA= is the mole fraction of the solvent A in the solution P°A= the vapor pressure of the solvent A its pure state.

What is the equilibrium constant and position of equilibrium dependent on when gas is dissolved into a liquid?

Pressure

Why do we salt icy roads in winter?

Salt mixes and dissolves into the small amount of liquid water in equilibrium with the solid phase (snow and ice). Solute causes disturbance, so the rate of freezing decreases. Causes more ice to melt than water to freeze. Melting is endothermic, so heat is absorbed from liquid solution. This decreases solution temperature, which causes heat gradient causing flow of heat from warmer air to cooler aqueous solution - facilitates more melting. -melting point depresses upon solute addition. Solute particles interfere with lattice formation (the highly organized state in which solid molecules align themselves). Colder-than-normal conditions are necessary to create solid structure.

Common ion effect

Solubility of a salt is 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. Decreases the molar solubility (in moles per liter) of the salt. The presence of that ion in solution shifts the dissolution reaction to the left.

Solute and solvent

Solute dissolves in solvent. Solvent remains in the same phase after mixing, or is the component present in greater quantity (i.e. if two liquids are mixed)

e.g. NaCl dissolving

The enthalpy change: NaCl - component ions have to dissociate. Na+ and Cl- ionic bonds must break, hydrogen bonds in water must break. This is all endothermic. Now the water and Na+ and Cl- have to interact with each other (ion-dipole interactions). This is exothermic B/c the magnitude of the exothermic process is slightly less than then energy required to break those ionic bonds/hydrogen bonds, the overall process is endothermic and requires some energy (favored at high temperatures) The entropy change: NaCl -> Na+ + Cl- releases ions from their lattice, and gives them more energy microstates. Entropy increases. However, water becomes more restricted because of its interactions with the ions, so its microstates are decreased. However, the entropy increase for NaCl is greater than the decrease in entropy for the water, so overall positive entropy change. -b/c of the relatively low endothermicity & relatively large positive change in entropy, NaCl will spontaneously dissolve in liquid water (△G=△H-T△S)

Solubility product constant (Ksp)

The equilibrium constant for a dissociation reaction. -shows the concentrations of the ionic constituents when they're at an equilibrium (saturation) concentrations

Saturation point

The equilibrium, where solute concentration is at its maximum value for the given temperature and pressure. Movement away from this point is nonspontaneous. -this is when the change in free energy is ZERO (equilibrium) -Rates of dissolution and precipitation are equal at this point

What can increase the density of a solution?

The increase of solute concentration.

Solubility

The maximum amount of solute (substance) that can be dissolved in a particular solvent at a given temperature; it is often expressed in molar solubility- the molarity of the solute at saturation

Molar solubility

The molar amount of a solute that can dissolve in 1L of solvent until equilibrium - also called saturation - is reached.

molar solubility

The molarity of a solute in a saturated solution.

What does the van't Hoff factor correspond to?

The number of particles into which a compound dissociates in solution. Ex:for NaCl i = 2 -glucose will have i = 1 b/c it doesn't dissociate in water

What is molality of a dilute aqueous solution at 25 degrees C equal to?

The solution's molarity because the density of water at this temperature is 1 kilogram per liter

Hydration

The solvation of a solute molecule by water. -its when water act as the solvent -its the process through dissolution occurs

Percent composition by mass

This formulas used for: *aqueous solutions *solid-in-solid solutions *metal alloys (a mixture of metals or a mixture of a metal and another element)

What is the solubility of salts in complex ions?

Very high, because complex ion has multiple polar bonds between ligands and central metal ion - able to interact in tons of dipole-dipole interactions. Stabilizes dissolution, therefore very high Ksp values.

Solvation - exothermic vs. endothermic

When new interactions are STRONGER than old ones, solvation is exothermic and favored at lower temp. Ex: CO2 turning into water (gas to liquid); CO2 has minimal intermolecular interaction. When new interactions are WEAKER than old ones, solvation is endothermic and favored at high temperatures. Most dissolutions are of this type. Energy (heat) must be put in to facilitate these shitty interactions. Ex: dissolving ammonium nitrate or sugar into water (heat energy is supplied to facilite the formation of these water, less stable interactions)

Sparingly soluble salts

a kind of solutes that dissolve minimally in the solvent (molar solubility under 0.1 M) -they're ionic compounds that have very low solubility in aqueous solutions

What happens when the solute dissolves into the solvent, and they system approaches saturation?

a point is reached to where no more solute can be dissolved and any excess will precipitate to the bottom of the container

Solvation

aka dissolution - electrostatic interaction between solute and solvent molecules. -involves breaking intermolecular interactions between solute molecules and between solvent molecules, and forming new interactions between the solute and solvent molecules.

How can a supersaturated solution be formed?

by dissolving solute into a hot solvent and then slowly cooling the solution

How does proteins carry out their functions?

by utilizing: 1) complex ion binding 2) transition metal complexes

What effects can pressure have in the "solubility product constant"?

ex: lets consider what happens when the solution consists of gas that becomes dissolved into a liquid. This will effect the value of equilibrium constant, hence equilibrium (saturation) depend on pressure. -so the solubility product constant decrease for a gas solute

What happens to Ksp when pressure is high?

high pressures will favor the dissolution of gas solutes -so Ksp will be larger for gases when pressure is high than when the pressure is low

Acid-base reactions concerned about the concentration of...

hydrogen ions

When can we apply "the law of mass action"?

it can be applied to a solution at equilibrium (saturated) b/c thats when the solute concentration is at a maximum and is dynamically stable

what is the vapor pressure of water at 100°C

it is equal to 1 atm (atmospheric pressure)

What effects can lowering the solution's vapor pressure have on temperates?

it means that a higher temperature is required to match atmospheric pressure, thereby raising the boiling point

How can adding a solute to a solvent effect the vapor pressure of the solvent in the solution?

it will cause a decrease in vapor pressure

What is the function of a nitrate ion in an oxidation-reduction reaction?

its a weak oxidizing agent (it's weak at causing other electrons to lose their electrons) -in other cases w/ nitrate, focus on the cation as a chemically reacting species

What determines the solubility of a compound in water?

its determined by the relative changes in enthalpy and entropy associated with the dissolution of the ionic solute at a given temperature and pressure

Entropy

its 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 -its the measure of molecular disorder, or the number of energy microstates available to a system at a given temperature

What is a chelation therapy?

its used to sequester toxic metals (like lead, arsenic, mercury) Note: metals such as iron can be toxic if we have too much of it, so it would sometimes needs to be sequested.

Chelation

its when the central cation can be bonded to same ligand in multiple places. Requires large organic ligands that can double back to form 2nd or even 3rd bond with central cation. Used to sequester toxic metals.

Freezing point depression

m = molality

Boiling point elevation

m=molality

If the vapor pressure of a solution is lower than that of the pure solvent, then what must happen in order for its vapor pressure to equal that ambient pressure?

more energy (higher temperatures) is required

KEEP IN MIND:

solids and liquids don't show up in the equilibrium constant

Which aqueous solution is more concentrated with solute, water OR water-double solutes? and what effects might this have?

water-double solutes have molar mass greater than water. So as the density of a solution increase, the concentration of solute increases as well.

What maximizes (increase) entropy during dissolution of the proteins?

when proteins dissolve in solution with their most hydrophilic amino acids on the outside and hydrophobic amino acids on the inside (this maximizes the increase in entropy during dissolution)

When will the dynamic equilibrium be reached in a solution?

when the rates of dissolution and precipitation are equal (this occurs during the saturation point) -this is when the concentration of dissolved solute reaches a steady-state (constant) value.

When will the solution be unsaturated?

when the solution is dilute (dissolution) which is thermodynamically favored.

What is the 1st step that you must do when you are doing a "stoichiometry" OR "solution equilibrium" problem?

you must write out the balanced dissociation reaction for the ionic compound in the question this first step will help us calculate: 1) the solubility product constant 2) ion product 3) molar solubility 4) find out the common ion effect

TUTOR Explanation of Common ion effect:

•We have a beaker containing AgCl(s). •We poured water to dissolve it. •The salt NaCl has Cl^-(aq) which is common to both AgCl(s) and NaCl. •If we pour NaCl to a solution containing AgCl, then the solubility of AgCl(s) will decrease and the ions Ag+(aq) and Cl^-(aq) will convert back to AgCl(s) in presene of NaCl. •Now look at the equation that I wrote in the book (pg. 319) •Before the addition of NaCl, AgCl(s) was in equilibrium with Ag^+(aq) and Cl^-(aq). •Now when we add NaCl(aq), the concentration of Cl^-(aq) in the solution increases. •Since Cl^-(aq) is a product, according to Lechatelier's principle when we add a product, the reaction will move in such a way that the concentration of product decreases. •We added NaCl, that contains a common ion Cl to the AgCl solution and now watching what effect it is imparting to AgCl solution according to Lechatelier's principle. •Since the common ion Cl^-(aq) is present on product side of the dissociation equation of AgCl, when we added NaCl, the concentration of product increases. •According to Lechatelier's principle, when we increase the concentration of product, the reaction goes in such a direction that the concentration of products decreases. •In order to decrease the concentration of products Ag^+(aq) and Cl^-(aq), they must combine back to form AgCl(s). •i.e ions are combined back to form the solid precipitate when we added a common ion. •Since more solid precipitate is formed when we add a common ion, the solubility of the salt decreases. •The effect of decrease in solubility of a salt solution when a common ion is added is called common ion effect. •Although the solubility decreases in presence of common ion, but the value of Ksp remains the same for a particular compound and always remain constant. •Because the concentration of the product decreased only to keep the equilibrium constant(Ksp) unchanged according to Lechatelier's principle •Hence value of Ksp always remains constant whether or not we add a common ion.