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Aqueous Reactions

Properties of Solutes in Aqueous Solutions

Solvent versus Solute

Water has the ability to dissolve many different types of substances, resulting in a homogeneous

mixture.

In homogeneous mixtures involving water, water is considered to be the solvent:

The typical molar concentrations of substances dissolved in water would be on the order of 10-6

to

101 molar, thus, they are present at far lower molar concentration and are considered to be the

solute.

How does water "dissolve" a solute?

The polar nature of the water molecule

The Lewis Structure of water:

The central Oxygen has a tetrahedral geometry for the valence electron pairs

Thus the H2O molecule will adopt a bent molecular geometry:

The Oxygen (3.5) is more electronegative than the Hydrogen (2.1), thus the O-H bond is polar

covalent

The bent geometry results in an overall dipole for the water molecule:

Thus, H2O can participate in the following types of non-covalent interactions:

o London Dispersion Forces

o Dipole-dipole interactions

o Ion-dipole interactions (with ions)

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o Hydrogen bonds (between other water molecules or with an appropriate solute)

It is the ability of water to participate in these diverse non-covalent interactions that allows water to

"dissolve" a variety of solutes

Ionic Compounds in Water

Water can participate in ion-dipole interactions.

Water molecules will organize around an ion to orient the appropriate opposite partial charge of

the water dipole:

Water molecules will separate, surround and disperse ions in an ionic solid:

Although H2O is a poor conductor of electricity, dissolved ions in an aqueous solution can

conduct electricity. Thus, ionic aqueous solutions are known as electrolytes.

An electrolyte solution (a solution of ions) can be described by either the concentration of the ionic

compound that was dissolved, or by the relative concentrations of the anion and cation components

Molecular compounds in an aqueous solution

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Generally speaking, interaction with H2O will not break any covalent bonds. Thus, the ability of

water to dissolve a molecular compound is based on non-covalent interactions of H2O with the

molecular compound

Example: Methanol (CH3OH) mixed with water

The alcohol (-OH) group of methanol is similar to water in that the oxygen valence electron

geometry is tetrahedral with two non-bonding pairs of electrons

There is a polar bond between the hydrogen and oxygen of the alcohol group in methanol, similar

to that in water:

Thus, water can form Hydrogen bonds with the alcohol group of methanol, and in this way the

water molecules can separate, surround and disperse the molecules of methanol

Some molecular compounds interact so strongly with H2O that covalent bonds of the compound

may be broken

o Although the molecular compound in question may be neutral, the molecular fragments

produced by the bond breakage may be oppositely charged ions. An example of this

would be the uncharged molecular compounds H-Cl. Interaction with water is so strong

that it results in the breakage of the H-Cl bond. This produces an H+ ion, and a Cl

- ion.

o Thus, although these types of molecular compounds may be neutral, in water they result

in the production of ions and therefore, are electrolytes

o Acids are one example of neutral compounds that ionize in aqueous solution (i.e. interact

so strongly that a covalent bond is broken and ions are produced)

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Strong and weak electrolytes

Some compounds in aqueous solution dissociate completely into ions. This would include most

ionic compounds, and some molecular compounds (like H-Cl)

Other compounds only have a slight tendency to ionize in aqueous solutions. In other words, only

a few of the molecules in solution will ionize, and most will remain as neutral compounds

(although the compound is completely dissolved in water)

Compounds that ionize completely are known as strong electrolytes

Compounds that ionize only partially are known as weak electrolytes

For example, acetic acid only partially ionizes (i.e. is a weak electrolyte) when dissolved in H2O. This

ionization involves the breaking of a covalent bond between an oxygen and hydrogen atom in the acid:

The double arrow means that the reaction is significant in both directions

o At any given moment, some of the anion form of acetic acid is combining with H+ cation

to form a covalent bond and produce the neutral acetic acid

o Likewise, at any given moment, some of the neutral acetic acid in aqueous solution is

dissociating (i.e. ionizing) to form the anion and H+ cation

o The balance (i.e. the chemical equilibrium) between these two processes determines the

relative concentration of the neutral molecule and ion forms

Chemists use a double arrow to indicate the ionization of weak electrolytes, and a single arrow to indicate

the ionization of strong electrolytes (i.e. in strong electrolytes, the ions have essentially no tendency to

recombine to form the neutral compound)

Acids, Bases and Salts

Acids

Acids are substances that are able to ionize in aqueous solutions to form H+ ions (and an associated anion)

A Hydrogen atom consists of a single proton and a single electron (no neutron)

Thus, an H+ ion is just a proton

Acids are often referred to as "proton donors"

Different types of acids can ionize to release one or more protons

A monoprotic acid releases a single proton when it ionizes (e.g. hydrochloric acid):

A diprotic acid can release two protons when it ionizes (e.g. sulfuric acid):

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o For sulfuric acid, the first ionization is complete (as indicated by the single arrow),

therefore it is a strong electrolyte

o However, only some of the molecules of HSO4- undergo a second ionization. Thus, the

deprotonization of HSO4-, and the protonization of SO4

2- are significant reactions (as

indicated by the double arrows)

o Aqueous solutions of sulfuric acid will therefore contain some mixture of protons (H+),

HSO4- and SO4

2-ions

Bases

Bases are substances that react with (or accept) H+ ions

A common example of a base is the hydroxide ion (OH-):

(Note that this represents the formation of a chemical bond between the ions to produce a water molecule)

Any substance that increases the concentration of OH-(aq) is a base

Some of the most common bases are metal hydroxides (e.g. NaOH, KOH, Ca(OH)2)

o These are ionic compounds, that when dissolved in H2O release the metal ion and one or

more hydroxide (OH-) ions

Other compounds can react with H2O in such a way that they are considered bases (even though they do not

directly contribute a OH- ion)

Ammonia (NH3) reacts chemically with a H2O. It has a high affinity for the proton that can be

obtained from a water molecule. This leaves a hydroxide ion (OH-) left over in solution:

Ammonia has "accepted a proton" from the water molecule, and the concentration of OH- ions has

increased in solution, thus, ammonia is a base

Note that the double arrows indicate that only some of the ammonia molecules will react with

water to accept a proton. Thus, ammonia is a weak electrolyte.

Strong and Weak Acids and Bases

Acids and bases that ionize completely in solution are strong electrolytes, and are therefore known

as strong acids or strong bases

Likewise, acids and bases that only partially ionize in solution are weak electrolytes, and are

known as weak acids or weak bases

o If a chemical reaction depends upon the concentration of H+ ions, then strong acids will

be more chemically reactive than weak acids

o However, the anion that is also released in the ionization of an acid can also be reactive.

If a weak acid (e.g. HF) releases an anion that is highly reactive (e.g. F-), then the weak

acid can also be chemically highly reactive

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Observations about acids and bases:

1. The commonly used acids, hydrochloric, nitric and sulfuric, are all strong acids

2. Several of the strong acids are a combination of hydrogen and a halogen (the exception is HF,

which is a weak acid)

3. There are not a lot of strong acids, most acids are weak acids

4. There are not a lot of strong bases; the strong bases are metal hydroxides (group 1A and heavy

group 2A metals)

5. Ammonia (NH3) is a weak base

Based on the above discussion, here is a flow chart to help you decide if a compound is a strong or weak

electrolyte (or a nonelectrolyte):

Neutralization reactions and salts

When an acid solution and a base solution are mixed, a neutralization reaction occurs.

In general, a neutralization reaction between an acid and a metal hydroxide (strong base) produces

water and a salt

The term salt has come to mean any ionic compound whose cation comes from a base and whose

anion comes from an acid

Example: aqueous hydrochloric acid mixed with aqueous sodium hydroxide

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Ionic Equations

For strong electrolytes, where the molecules in a reaction completely dissociate to ionic forms in solution,

there are two different ways we could think of writing the chemical reaction

In an example of the neutralization reaction of hydrochloric acid and the metal hydroxide of

sodium (i.e. sodium hydroxide) the molecular equation for the reaction would be:

Since HCl and NaOH are strong electrolytes and will dissociate completely to their ionic forms,

we can also write the reaction in terms of a complete ionic equation:

As with any equation, we can algebraically manipulate the complete ionic equation to cancel the

same terms that are present as both reactants and products:

which yields the overall reaction of:

This is known as the net ionic equation for this particular reaction of strong electrolytes. It has the

following characteristics

o The net ionic equation only shows the ions and molecules directly involved in reaction.

The other ions (Na+ and Cl

- in this case) are called spectator ions.

o Like any reaction, the overall charge on the reactant and products must be equal. In the

above case the net charge on the reactants (H+ and OH

-) is zero, and the net charge on the

product (H2O) is zero.

This equation represents the net reaction of any strong acid with any strong base (i.e. a proton and

hydroxide ion react to produce H2O)

Only reactions involving soluble, strong electrolytes can be written in ionic form

Metathesis Reactions

In many aqueous reactions it seems that the reaction involves the ionic compounds swapping their ionic

partners. For example, in the reaction involving the ionic compounds silver nitrate and potassium chloride

we have:

The silver cation exchanges its nitrate anion partner for the chloride anion. Likewise, the potassium cation

exchanges chloride anion for the nitrate anion:

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This swapping of ions in aqueous reactions can be symbolically represented as follows:

This type of reaction is known as a Metathesis reaction

Note: metathesis is not pronounced "meta-thesis", but rather "meh-TATH-eh-sis" (apparently the Greeks

prefer to pronounce it that way)

There is something subtle in the above example that is important to note.

Take a look at the state of the silver chloride ion pair.

The AgCl(s) is a solid, and therefore the ions are not dispersed in solution. In fact, they precipitate

out of solution.

The precipitation effectively removes the AgCl ions from the solution, and this is the driving

force for the observed metathesis reaction

The driving force for metathesis reactions is the removal of ions from solution

What are the ways in which ions can be removed from solution and thus drive a metathesis reaction?

1. Certain ions can associate to form an insoluble precipitate (as with the formation of AgCl(s))

2. Certain ions can chemically combine to form a neutral molecular compound (resulting in

either a non-electrolyte, or a weak electrolyte).

o Acid/base neutralization reactions that produce water from H+ and OH

- ions are an

example of this.

o Being a non-electrolyte (or weak electrolyte) the formation of the molecular compound

from the constituent ions is essentially an irreversible process

3. Certain ions can chemically combine to form a gas, and the gas physically escapes from the

solution

Precipitation Reactions

Metathesis reactions that result in an insoluble precipitate are called precipitation reactions

Solubility refers to the amount of a substance that can be dissolved in a given quantity of water

The solubility of an ionic compound determines whether it will precipitate or not

Any substance with a solubility less than 0.01 moles/L will be considered as essentially insoluble

The reaction of KI and Pb(NO3)2:

The PbI2 is an insoluble ionic compound that will precipitate and drive the metathesis reaction:

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Introduction to Oxidation-Reduction Reactions

Many metals, when placed in an acid solution, will react chemically to produce hydrogen gas. For

example, zinc metal in an aqueous solution of hydrochloric acid will react as follows:

What is going on in this type of reaction?

The zinc starts out as a solid metal, i.e. the neutral (uncharged) elemental form of zinc. The zinc

ends up as an ion (Zn2+

)in an ionic compound involving chloride ion

The hydrogen starts out as a hydrogen ion (H+) from the acid, and ends up in the neutral

(covalently bonded) form of diatomic hydrogen

Chloride ion (Cl-) is a spectator ion in this reaction

In order for these ionic changes to occur, electrons leave the metal (producing zinc cations) and join the

hydrogen ions (producing neutral diatomic hydrogen).

When an atom has become more positively charged (by LOSING electrons) chemists say that it has been

OXIDIZED

In the above reaction, the zinc metal has become oxidized (producing Zn2+

cations)

When an atom becomes more negatively charged (by GAINING electrons) chemists say that is has been

REDUCED

In the above reaction, the hydrogen ions have been reduced (to produce neutral diatomic hydrogen

gas)

Oxidation of the metal refers to the fact that this type of reaction was actually first characterized by

studying the reaction of metals with oxygen:

Many metals react with oxygen in the air to form metal oxides

In this reaction the neutral metal loses electrons to the oxygen, forming a cation (Ca2+

in this case).

The metal cation participates in forming an ionic solid with the oxide (O2-

) anion. Thus, the metal

becomes oxidized

When one substance in a reaction loses electrons, another substance in the reaction must gain them.

The oxidation of a reactant is always accompanied by the reduction of another reactant in the reaction

These types of reactions are therefore called oxidation/reduction (or REDOX) reactions

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In predicting oxidation reactions between different metals, compare the relative ease with which the two

metals can be oxidized

Of the two metals, the one that is relatively more difficult to oxidize will preferentially exist in

the neutral elemental form (i.e. be reduced)

The other metal will preferentially exist in the oxidized (cation) form (i.e. will give its electrons up

to the other metal)

A ranking of the metals by the relative ease with which they can be oxidized, is known as an Activity

Series. Hydrogen is also included in such lists so as to include the behavior of acids along with the metals:

At the top (the easiest to oxidize) are the alkali metals (group 1A) and the alkaline earth metals

(group 2A)

At the bottom (the least likely to oxidize) are the transition metals, and in particular, the metals

that are typically found in jewelry or coins are at the very bottom

The Activity Series can be use to predict the outcome of reactions between metals and either metal

salts or acids

Any metal in the list can be oxidized by the ions of any metal below it (or hydrogen ions if they

fall below the metal in question). The corresponding ions of the other metal (or hydrogen) will

become reduced to the elemental form

For example, silver is below copper in the Activity Series, therefore, copper metal will be oxidized by

silver cations (resulting in the formation of copper cations and elemental silver)

The Solution Process

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Important characteristics of solutions:

They are homogenous mixtures

Solutions may be gasses, liquids or solids

Each substance in a solution is a component of the solution. Usually, the component with the

highest concentration is termed the solvent (other components are termed solutes)

Most solutions we will deal with are those in a liquid state, where the solvent is H2O (i.e. aqueous

solutions)

The liquid state, and the solid state, are known as condensed states

In condensed states, the attractive forces between molecules are strong enough (in comparison to

the temperature-induce kinetic energy) to hold neighboring molecules together.

o In solids, the neighbors are held rigid

o In liquids, the neighbor molecules can slide past each other

Homogenous mixtures (solutions) can form only when the following attractive forces are approximately

equivalent:

Attraction between solvent and solute molecules

Attraction of solute molecules for other solute molecules

Attraction of solvent molecules for other solvent molecules

If the attractive forces of solute molecules for other solute molecules are greater than the attractive forces of

solute molecules for water, then the solute will not dissolve

For ionic solids, the lattice energy describes the attractive forces between the solute molecules (i.e.

ions)

For an ionic solid to dissolve in water, the water-solute attractive forces has to be strong enough to

overcome the lattice energy

The process known as solvation is where the solute-solvent interactions are strong enough separate,

surround and disperse a solute

If the solvent is H2O, then solvation is referred to as hydration

Energy Changes and Solution Formation

We have previously studied enthalpy changes associated with chemical reactions (e.g. combustion

reactions, ΔHrxn) and with physical processes (e.g. changes of state, ΔHfusion, and the heating of matter in a

specific state, Molar Heat Capacity)

The process of solvation involves energy changes also, known as the enthalpy of solvation (ΔHsolv)

It is a physical process, not chemical

What are the different processes that contribute to the enthalpy of solvation ΔHsolv?

There is the energy (ΔH1) associated with dispersing the solutes. It is the lattice energy (here's an

example for an ionic solid):

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In order to accommodate the dispersed solutes within the H2O solution, the H2O molecules have to

separate from one another to provide the necessary space (in other words, we have to disrupt

solvent-solvent interactions). There is an energy (ΔH2) associated with this process:

And, finally, there is the energy (ΔH3) associated with the formation of solvent-solute interactions

for the solvated solute molecules:

What is the overall energy associated with these three distinct energetic steps involved in solvation of a

solute?

ΔH1, separating the solute molecules from each other, requires an input of energy to overcome the

attractive forces holding the solute molecules together. Thus, ΔH1 will be positive in sign

(endothermic).

ΔH2, disrupting solvent-solvent interactions, to allow space for dispersed solute molecules, will

also require the input of energy. Thus, ΔH2 will be positive in sign (endothermic).

ΔH3, formation of solvent-solute interactions, will release energy. Thus, ΔH3 will be negative in

sign (exothermic).

If the energy released by the formation of solvent-solute interactions (ΔH3) is greater than the sum of the

energies required to disrupt solute-solute interactions and solvent-solvent interactions (ΔH1 + ΔH2), then

the overall enthalpy of solvation (ΔHsolv) will be exothermic and energetically favorable and spontaneous:

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Some solvation enthalpies are actually positive (i.e. endothermic) and absorb heat from their surroundings

(this process is used in cold packs for muscle sprains):

Although some heats of solvation are positive (e.g. the hydration of ammonium nitrate), their reactions

proceed spontaneously.

Up to this point, when trying to predict whether a reaction or process is spontaneous we have

considered the overall change in heat energy (enthalpy)

Processes in which the overall heat energy of the system decreases tend to be spontaneous (i.e. a negative

value for the overall ΔH indicates spontaneity)

However, the hydration of Ammonium Nitrate is spontaneous, but has a slightly positive value for

the overall enthalpy (it has absorbed heat energy). Clearly, something else is going on that forces

the hydration of Ammonium Nitrate to occur

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The answer to this riddle lies in the behavior of collections of objects. Consider a field with a herd of sheep,

some are black and some are white.

If we let the sheep roam around and then take an aerial picture of the field, the black and white

sheep would be randomly distributed:

Suppose the farmer wants to separate the sheep. A sheepdog will have to do work to separate

them:

Not only that, but the dog will have to keep working to keep them separated, because their natural

tendency would be to randomly distribute throughout the field

We can postulate that if energy must be expended to keep the sheep from being randomly

distributed, then if we reverse the process (i.e. start with sheep in an ordered arrangement and

let them become randomly distributed) then energy is somehow released

The same situation is true with collections of molecules. In the case of the solvation of ammonium nitrate

we start with a crystal of the ionic solid being placed in a container with water:

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When the crystal initially dissolves all the ions, though hydrated, are not randomly distributed

throughout the water in the container - they are concentrated in the vicinity of the original crystal

The solvated ions are therefore initially not in a random distribution throughout the container.

They are initially in a more organized or ordered state (i.e. concentrated in one location in the

container) and will naturally want to become more disordered or randomly distributed

Processes in which the disorder of the system increases tend to occur spontaneously

Thus the increase in disorder is a driving force that can overcome the slight positive enthalpy

associated with the hydration of Ammonium Nitrate

The term for the degree of disorder in a system is ENTROPY

A system with high entropy is disordered

A system with low entropy is ordered

Systems tend to go from low to high entropy

The entropy of the universe is increasing (just check out your kitchen or bathroom for evidence of

this)

Ways of Expressing Concentration

Dilute versus Concentrated Solutions:

Some liquid cleaning solutions are sold in "concentrated" form, and the instructions require you to

"dilute" them prior to use

A concentrated solution is one where the solute is present in high concentration

A dilute solution is one where the solute is present in low concentration

"high" and "low" are relative terms and do not give us any precise information regarding the actual

concentration of solute in a solution.

We need a formal description of the concentration of a solute in a solution

Mass Percentage

A formal description of concentration that refers to the ratio of the mass of the component of

interest (e.g. solute) with the entire mass of the sample:

For example, a 100g aqueous solution containing NaCl is evaporated to dryness, leaving only the NaCl.

The NaCl is weighed and contains 5g. What was the mass percentage of NaCl in the original solution?

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For very dilute solutions (again, a relative term) the concentrations are often expressed as "parts per

million" (abbreviated as ppm)

ppm of component is given as:

A mass percentage of 1% would be equal to 10,000 ppm

1 ppm would be equal to .0001%

For really, really dilute solutions, the concentrations of solute are sometimes expressed as parts per billion

(abbreviated as ppb)

ppb of component is given as:

Mole Fraction, Molarity and Molality

It is quite common to refer to concentration of a solute in terms of the number of moles of solute

Mole Fraction:

The symbol X is used to refer to the mole fraction of a component in a sample

A subscript is sometimes added to indicate the component being referred to. For example, the

mole fraction of lead ion in a sample might be indicated by

The sum of the moles fractions of all components in a solution (including the solution itself) must equal 1.0

Molarity:

The symbol M is used to refer to the molarity of a solute in a solution

Molarity defines the concentration of a solute in terms of the number of moles of solute and the

total volume of solution

Molality:

The symbol m is used to refer to the molality of a solute in solution

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Molality defines the concentration of a solute in terms of the number of moles of solute and the

mass of the solvent component

Saturated Solutions and Solubility

When a salt (NaCl) crystal is initially place in a sample of H2O the solution is devoid of hydrated

Na+ and Cl

- ions:

As the water molecules surround, separate and disperse the Na+ and Cl

- ions, the solution becomes

populated by the hydrated ions:

The dispersed ions in the solution will collide with water molecules, the surfaces of the container

and potentially other ions as well.

If the original crystal has not completely dissolved, then dispersed ions can also collide with the

remaining crystal. These collisions can result in the incorporation of the ions back into the NaCl

crystal lattice:

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Thus, there are two opposing processes that can potentially occur:

1. Dissolving of the crystal, resulting in hydration of the individual Na+ and Cl

- ions, and

2. Collision of dispersed ions resulting in an increase in the crystal mass (a process also

known as Crystallization)

Note that when a hydrated ion collides with a crystal and is incorporated into the crystal lattice, that the

waters that are hydrating the ion are released (i.e. the exact reverse of the hydration process)

These two opposing processes of dissolving and crystallization can be represented as follows:

If the rate of dissolution is greater than the rate of crystallization, then the crystals of NaCl in the

solvent will get smaller

If the rate of crystallization is greater than the rate of dissolution, then the crystals of NaCl in the

solvent will get larger

If the rates of the two opposing processes are equal, then the size of the NaCl crystals will remain

unchanged and the system is said to be in a dynamic equilibrium

A solution that is in dynamic equilibrium with undissolved solute is said to be saturated (i.e. no more solute

will dissolve into the solvent under the current conditions)

The concentration of solute present in the solution under conditions of saturation is known as the

solubility of that solute

At higher temperatures, usually more solute can be dissolved, and the solubility is higher.

Under some conditions it is possible to produce a supersaturated solution of a solute

Solute is dissolved to saturation at a high temperature

The solution is carefully and slowly cooled to a lower temperature (the idea is not to induce the

formation of tiny crystals that can serve to nucleate crystal growth)

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At the lower temperature, the concentration of solute is higher than the equilibrium concentration

at that temperature. The introduction of a "seed" crystal will stimulate rapid crystal formation

Factors Affecting Solubility

Factors that can affect solubility:

Properties of solute

Properties of solvent

Temperature

Pressure (Gases)

Solute-Solvent interactions

Gases

Gases are gases because the attractive forces are typically weak - involve primarily London

dispersion forces

Consequently, attractive forces (basis of solubility) between gas molecules and solvent are also

primarily London dispersion forces

London dispersion forces increase with increasing size and mass of molecules involved

Thus, the solubility of gas molecules typically increases with increasing mass of the gas molecules

Some gases appear to be a lot more soluble than what their mass would predict

Higher than normal solubilites of gases are an indication that a chemical reaction (in addition to the

physical process of dissolution) is occuring

Polar Solutes in Polar Solvents

Polar solutes tend to dissolve readily in polar solvents

Interactions between polar solutes are typically dipole-dipole (or Hydrogen-bonds)

Interactions between molecules of a polar solvent are also dipole-dipole (or Hydrogen-bonds)

Thus, the energies associated with disrupting solute-solute interactions and solvent-solvent

interactions are approximately equivalent

Entropic forces can subsequently drive the dissolution process

Polar liquids tend to dissolve readily in polar solvents

Pairs of liquids that mix in any proportion are termed miscible. Liquids that do not mix are termed

immiscible.

Ethanol contains a hydroxyl (OH) functional group that is similar in structure to water. Attractive

forces between ethanol molecules include Hydrogen bonds, like the attractive forces between

water molecules. Ethanol is miscible in H2O

Octane (gasoline) molecules contain only C-H and C-C bonds and are essentially non-polar

molecules. Attractive forces between octane molecules include primarily London dispersion

forces. Octane is not miscible in H2O

Ethanol is an alcohol (contains an OH functional group). Octanol is also an alcohol (it contains a

single OH functional group). However, it contains a string of 8 carbon groups compared to the two

in ethanol. The carbon groups cannot participate in Hydrogen bonding (only dispersion forces).

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The single OH group in octanol is not enough to provide solubility for octanol, and octanol is

essentially immiscible in H2O.

The observation that if similar attractive forces exits between solute-solute molecules and solvent-solvent

molecules results in miscible solutions (i.e. the ability of the solvent to dissolve the solute), has led to the

following generalization:

"Like dissolves like"

(in other words, substances with similar intermolecular attractive forces tend to be soluble in one another)

Pressure Effects on Gases and Solubility

Increasing the pressure (at constant T) results in more collisions of the gas molecules, per unit

time, with the surface of the solvent. This results in greater solubility.

Temperature effects on Gases and Solubility

Kinetic energy increases with increasing temperature

Thus, increasing the temperature reduces the solubility of gas molecules in a solvent

Temperature effects on Solid Solutes and Solubility

Thus, increasing temperature increases the solubility of solid solutes

Colligative Properties

Various kinds of solutes (e.g. NaCl, ethylene glycol) added to H2O result in a decrease in the freezing

temperature, as well as an increase in the boiling temperature, of H2O.

Solutes added to H2O are a useful mechanism with which to prevent water-cooled machines from

freezing in winter and boiling over in summer

Salt added to the roads can prevent the freezing of water in winter

These effects upon the physical properties of water are termed "colligative properties" and are primarily

dependent upon the amount of solute added, and are relatively independent of the type of solute

Colligative properties depend upon the collective effect of the number of solute particles in a solution

Raoult's Law

The vapor pressure of the solute-containing solution (PA) is equal to the vapor pressure of the pure solvent

(PA0) times the mole fraction of the solvent (XA)

An ideal solution is a solution that obeys Raoult's Law

Variations in the strength of the solvent-solute and solvent-solvent interactions can result in vapor

pressures that deviate from ideal solution behavior

o If solvent-solute interactions are really strong (e.g. involve Hydrogen bonding) then the

vapor pressure of a solution might be lower than ideal

o If the solvent-solvent interactions are really strong then the vapor pressure of a solution

might be higher than ideal

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Boiling Point Elevation

Here is a typical vapor pressure diagram for H2O:

The addition of a solute results in a lowering of the vapor pressure. At any given temperature, the

vapor pressure of a solution of H2O and a non-volatile solute will have a lower vapor pressure.

What does this mean with regard to the liquid/gas phase transition?

o At any given temperature, the vapor pressure of the liquid is lower

o This means that bubbles that form will have a lower vapor pressure and will now

collapse, where previously they remained (i.e. the sample boiled)

o The lower vapor pressure means the sample will boil at this temperature, but only if the

pressure is lowered:

Thus, at a constant pressure, the boiling point of an aqueous solution of a non-volatile solute will

be higher than pure H2O:

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The increase in the boiling point (ΔTb) relative to that of the pure solvent is directly proportional

to the number of solute particles per mole of solvent molecules

o Molality (m) represents the number of moles of solute per kg of the solvent component

(whereas molarity, M, relates to volume of total solution). Thus, molality gives

information about the number of moles of solute per mole of solvent

o Therefore, ΔTb is proportional to molality (and not molarity):

ΔTb = Kb * m

o Kb is called the molal boiling-point-elevation constant, and is dependent only upon the

choice of solvent

Freezing Point Depression

If an aqueous solution of ethanol is frozen, the water selectively freezes as a pure substance and

the ethanol (with a lower freezing point) is squeezed out as pure ethanol. Consequently the vapor

pressure diagram representing the vapor/solid equilibrium (vapor pressure of solid) for an aqueous

salt solution will be identical to that of pure H2O:

Thus, at a constant pressure, the melting point for a aqueous salt solution will be lower than pure

H2O:

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Thus, the addition of a solute will both raise the boiling point, and lower the melting point of an aqueous

solution

The decrease in the freezing point temperature, ΔTf, is directly proportional to the molality of the

solute:

ΔTf = -Kf * m

The value of Kf, the molal freezing-point-depression constant is a characteristic only of the solvent

o For water, the value of Kf is 1.86°C/m

o Thus, a 1 molal solution of solute will decrease the freezing point of water by 1.86°C (i.e.

ΔTf = -1.86°)

o A 0.5 molal solution of NaCl will produce 1.0 molal of solute (remember, it dissociates

into two ions)

Osmosis

Some materials, like cellophane and some biological membranes, have tiny pores. These pores are large

enough to allow solvent, like H2O, to freely pass across the membrane, but may prevent the passage of

larger solute molecules

Such materials are termed semi-permeable membranes

If two solutions, with different solute concentrations, are separated by a semi-permeable membrane, there

can be a net flow of solvent across the membrane

Like effusion with gas molecules, the rate of movement of solvent across the membrane is a

function of the concentration of solvent and the kinetic energy

Both solutions are at the same temperature, thus have the same kinetic energy. However, solvation

of solute molecules means there are less free solvent molecules to pass through the membrane (i.e.

those solvent molecules involved in hydrating solute are not free to pass through the membrane)

On the side with the pure solvent, more molecules of solvent per unit time can pass through the

membrane

Therefore, there will be a net flow of solvent molecules from the side with pure solvent to the side

containing solute

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Here is how such an effect can be demonstrated:

At the beginning of the experiment, here is how things might look in a 'U' shaped glass tube with

two solutions separated by a semi-permeable membrane:

After equilibrium is reached, here is how the apparatus and solutions might look:

There will be a net flow of solvent across the membrane, into the salt solution, until the height differential

(π) results in a pressure that offsets the flow of solvent. This pressure, π, is termed the osmotic pressure

The osmotic pressure behaves very much like pressure in the ideal gas equation:

πV = nRT

π = (n/V)RT

Where n = moles of solute, T is temperature (K), R is the gas constant and V is the volume of

solution

(n/V) is the number of moles of solute per volume of solution. This is also the molarity of the

solution (M). Thus:

π = MRT

The osmotic pressure is just a function of the molar concentration of solute

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Relative osmotic pressures

Two solutions with equal osmotic pressures are termed isotonic

If a solution has a lower osmotic pressure than another, it is termed hypotonic

If a solution has a higher osmotic pressure than another it is termed hypertonic

Summary of Colligative Properties

1. Raoult's Law

The vapor pressure of the solute-containing solution (PA) is equal to the vapor pressure of the pure solvent

(PA0) times the mole fraction of the solvent (XA)

2. Boiling Point Elevation

ΔTb = Kb * m

The increase in boiling point is equal to the molality of the solute times the molal boiling point elevation

constant, Kb, (unique for each solvent)

3. Freezing Point Depression

ΔTf = Kf * m

The decrease in the freezing point temperature is equal to the molality of the solute times the molal

freezing-point-depression constant, Kf, (unique for each solvent)

4. Osmotic Pressure

π = MRT

The osmotic pressure is equal to the molar concentration of solute times the gas constant

times the temperature in Kelvin

Determination of Molar Mass Using Colligative Properties

Colligative properties relate either the mole fraction, molal concentration, or molar concentration of a

solute to a measurable change in a physical property of a solution

Thus, we have a way to quantitate the number of moles of a solute based upon one of the four

colligative properties

If we know the quantity in grams of a solute added to a solution (to produce the measurable

colligative effect), then we have both the number of moles of the solute and the associated mass

Molar mass has the units of grams/mole; it describes the mass associated with 1 mole of the

compound (i.e. solute)

Colloids

The solutes in solutions that we have been considering up to this point are ions or small molecules

They form homogeneous solutions with the solvent

They do not slowly settle out of solution, or sink to the bottom of the solution, over a period of

time

Gravitational forces are small compared to the kinetic energy of the molecules in solution

As solutes get larger, at some point they my start to settle out, or sink to the bottom of the solvent

Gravitational forces are greater in comparison to the kinetic energy of the solution

In this case, we no longer have a homogeneous solution, but rather, a heterogeneous mixture

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Another property of "large" solute molecules, or high-molecular weight components in a mixture, is the

interaction of such molecules with visible light.

Particles can scatter light when their physical dimensions are similar to the wavelength of the light

Light scattering of this type will be manifest as the sample appearing "cloudy"

Visible light has wavelengths of from ~400nm to ~750nm

Particles with physical dimensions (i.e. diameter) of at least 400nm to 750nm can scatter visible

light and will therefore appear cloudy

A Carbon-Carbon bond has dimensions of approximately 0.15nm, therefore, a carbon-containing

molecule would need to have on the order of several thousand carbons in a chain to be of a size to

scatter visible light

Large solute molecules that are still small enough not to settle out

Between the tiny solutes we have been considering up to this point, and solutes that are so large that they

settle out of solution, are homogenous mixtures involving "big" solutes

These solutions are termed "colloidal dispersions", or just "colloids"

Colloids are somewhere between a homogenous solution and a heterogenous mixture

o They are small enough to where random collisions keep them dispersed throughout the

solution, and the won't settle out due to the effects of gravity, but they are not really

dissolved by the solvent

o They can be detected by their light scattering properties (i.e. their large molecular size).

The scattering of light indicates that dispersed throughout the solvent is a "solute" that is

actually comprised of large chunky bits. Although too small to settle out, their presence is

considered to represent a heterogenous mixture, rather than a homogenous solution

Colloidal dispersions can be gasses, liquids or solids. Here are some examples:

Phase of Colloidal

Dispersion

(i.e. "solvent" phase)

Colloidal "Solute" phase Official Name Example

Gas Liquid Aerosol Fog

Gas Solid Aerosol Smoke

Liquid Gas Foam Whipped cream

Liquid Liquid Emulsion Homogenized

Milk

Liquid Solid Sol Paint

Solid Gas Solid foam Marshmallow

Solid Liquid Solid emulsion Butter

Solid Solid Solid Sol Ruby glass

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Hydrophilic Colloids

Hydrophilic means "water loving"

Hydrophilic colloids have solutes with structural groups exposed on their surface that are able to

hydrogen bond with water (electronegative groups with or without hydrogen covalently attached)

Since "like dissolves like" hydrophilic colloids form aqueous colloidal dispersions

Hydrophobic Colloids

Hydrophobic means "water fearing"

Hydrophobic colloids have solutes with surface groups that cannot hydrogen bond, and typically

involving groups that can only interact via dispersion forces

Hydrophobic colloids cannot be prepared in water - unless some type of chemical alteration is

done to the solute.

o If the solute can adsorb ions onto its surface, then it may be able to interact strong enough

with water. (Note: adsorb means to stick to the surface).

o Another strategy is to combine the solute with another molecule that has two distinct ends

to it. One end is hydrophobic in nature, the other hydrophilic. The hydrophobic end binds

to the hydrophobic solute, and the hydrophilic end can interact with water (and solubilize

the hydrophobic solute). Examples of this are soap molecules, and the digestive juices

called bile (helps to dissolve fats in the diet in the aqueous environment of our bodies).

Removing colloidal solutes (or colloidal particles)

This can be a little tough to do because you cannot centrifuge them out, and often cannot filter them out.

Some times heat or addition of electrolytes can cause colloid particle to clump together - or coagulate

Heat increases the kinetic energy and rate of intermolecular collisions between colloidal particles.

Sometimes they have a natural tendency to stick together. Thus, larger aggregates are built up that

eventually settle out of solution or can be filtered out.

If some colloidal solutes are solubilized in aqueous solution by surface ions, then the addition of

counter ions can cause such particles to stick together (by eliminating the charge repulsion of like-

charges)

As seen with the section on osmosis, semi-permeable membranes can also separate colloids. This

is the basis of kidney dialysis.

© 2000 Dr. Michael Blaber