Strand 4  Solutions and Solubility
SCH 3U Strand 4 Solutions and Solubility:  Solutions

Liquid Mixtures

     A mixture is a type of substance that is impure due to the presence of more than one type of particle; that is, it is a physical blending of two or more dissimilar substances. A major characteristic of mixtures is that the materials do not chemically combine. A mixture can be physically separated into pure compounds or elements. Just about everything that you can think of is probably a mixture. Even the purest of materials still contain other compounds as impurities. Mixtures may exhibit a changing set of physical properties. Mixtures can be divided into those that are evenly distributed (homogeneous) and those that aren't (heterogeneous which is really called a mechanical mixture since it must be forced to mix together). A mechanical or heterogeneous mixture contains very large and different particles that are very readily distinguishable and very easy to separate from one another by physical or mechanical means, such as settling by mass, evaporation, hand sortation, flotation, filtering, etc. For example a garden salad or chicken soup.  


     Liquid mixtures are quite unique and are generally called dispersions. These homogenous mixtures stay uniformly distributed for some time. Dispersions are more difficult to distinguish or separate that is, the particles remain dispersed within the medium or surroundings. Since they are homogenous, the dispersion’s composition and properties are uniform or consistent throughout the entire substance. Dispersions are actually classified according to the size of the dispersed particulate matter because this greatly affects its properties. Dispersions can be created with various combinations of solids, liquids and gases. Below is an example of a dispersion at the molecular level.


     There are three common dispersions chemists deal with regularly: suspensions, colloids or solutions. They can be compared based on the size of the dispersed particles.

Comparing Dispersios

     Suspension are not mechanical mixtures because, even though the dispersed particles can be huge, they must still be small enough to remain dispersed for some period of time; whereas, mechanical mixtures settle immediately.. For example chalk dust in air stays well mixed in the air for several minutes. The scattering of light by large dispersed particles is called the Tyndall effect. It might be surprising that blood is a suspension, but actually the liquid of blood is the yellow-coloured plasma which contains many cells such as red blood cells which will settle to the bottom over time. These red blood cells are so well suspended that most people think the fluid, blood, is red but it is actually the effect of the suspension.

     Colloids are quite interesting because they are really a range between the other types of dispersions and as a result show a range in properties. You might wonder then, what distinguishes an extremely large colloid particle from the smallest suspension particle or likewise at the lower boundary for colloids. The classification can be challenging but is due to a collection of properties. Milk is similar to blood in that the fat globules are dispersed so well, they colour is white, but in fact, the fluid portion of milk is more yellow-white due to other chemicals present. Milk is not a suspension though, unless you are looking at fresh cow's milk. The abundant larger fat particles in fresh milk settle quickly and can actually be seen on your cereal or in your glass. Store-bought milk is homogenized which involves removing the water, grinding the fat into fine small particles and then mixing the fat back into water by shooting them through a higher pressure nozzle and now you have a colloid that takes a long time to separate. Colloids are further classified based on the type of dispersed particle and medium.

Types of Colloids

     Even more interesting is the fact that how the particles in the dispersion behave can alter the type of dispersion produced. For example, vinegar and oil is a heterogeneous mechanical mixture when poured together since two liquid layers are present; however, if the mixture is violently shaken, the oil will separate into smaller particles making the dispersion a suspension until the oil drops reunite into large enough drops. Air with clouds is a suspension, as the clouds contain large droplets of liquid water allowing them to be distinguished but they remain dispersed locally; however, these water droplets may either clump to settle out as heterogeneous mechanical mixture (precipitation) or they water droplets may spread out to the point that the air becomes a solution.

     In the diagram above, fine sand (silica) was added to water and quickly settled, producing a heterogeneous mixture with water on top and silica on the bottom. However, on the right, the same proportion of silica, but specially prepared using a base produces a colloidal dispersion. The particles of hydrated silica, SiO2*xH2O, are much larger than atoms and ordinary molecules. However, the sodium hydroxide used in preparing the dispersion causes the silica particles to acquire a negative charge from adsorbed hydroxide ions. The similarly charged silica particles repel one another and stay suspended indefinitely.

     Solutions have particles which are the size of atoms or molecules - too small to be seen. A solution is a mixture of two or more substances in a single phase. At least two substances must be mixed in order to have a solution. Since solutions are mixtures, their compositions may vary over a very wide range. The concentrations may be expressed using a variety of measures. The non-specific terms concentrated and dilute are sometimes used. A concentrated solution has a relatively large (but non-specific) amount of solute dissolved in a solvent. A dilute solution has a smaller quantity of solute dissolved. One way to distinguish between colloids and solutions is to test them with a beam of light. Light passes through solutions, but reflects on colloid partials allowing the beam to become visible. This is referred to as the Tyndall effect. In the diagram below, light passes undetected through the solution on the right, but is visible in the colloid on the left.

What is a Solution

     Before examining solutions in great deal, we must first ask "What is a solution?"  A number of possible answers to this question might be proposed. Perhaps it is a mixture of two substances.  Perhaps it is a good mixture; one in there appears to only be one substance.  Perhaps there is more of one substance than there is of the other.  All of these considerations have elements of truth to them.  However, there is something more to consider.  In all matter, there is room between one particle and the next.  It is possible in some cases for particles of another substance to occupy some of this space, and this is what happens in a solution.  This state of matter has some unique properties, some of which will be explained in detail below.

     First, consider how the process of dissolving could come to pass.  Since in nature we find many solutions with water as the solvent, we will first direct our attention to aqueous solutions- solutions in which a substance, the solute is dissolved in water, the solvent.  Consider the following animation.

     As you probably remember, a water molecule is shaped like a V, and it is a polar molecule.  That is to say, the oxygen atom, being highly electronegative, attracts the electrons in each hydrogen-oxygen bond more strongly, and hence has negative charges surrounding it more often. The hydrogens then, have the shared electrons around them less often, and consequently carry a partial positive charge.  Because of its polarity, water molecules have the tendency to become attracted to any other charged or partially charged species in their vicinity.  In the animation, we see how the partially positive hydrogen end of the water molecule is attracted to the negatively charged chlorine ions.  We also see how the partially negative oxygen end of the water is attracted to the positively charged sodium ions.  Due to the random motion of water molecules in a liquid, water molecules attracted to the solid exert a force on it as they pass.  If enough water molecules exert such a force on the ions in solid, they can actually pull ions out of the solid.  With more surface area exposed, more water molecules cling to the ion until it is fully surrounded.  The ion is not said to be in solution.

     While the process is essentially the same for any polar substance dissolved in water, there are slight differences between the solvation of ionic, polar and non-polar covalent solids. Ionic solids are pulled apart into ions as described above.  Soluble ionic compounds are said to “dissociate” in water.  Polar covalent substances behave differently depending on the strength of their intramolecular bonds (the polar covalent bonds holding one atom of the molecule to the next). To explain this, recall that the water molecules exert a force on polar substances being dissolved as they pass.  In the case of some polar covalent molecules, this force is enough to break the covalent bonds, and create two ions.  The more electronegative atom from the original polar covalent substance becomes the negative ion (anion) while the lesser electronegative atom becomes the positive ion (cation).  This process is known as ionization. If the force applied by passing water molecules is insufficient to ionize the polar covalent molecule, it is often still strong enough to break up the intermolecular forces holding one molecule of the solute to another.  This is the case in a molecule such as glucose. The negative end of the polar glucose molecule is attracted to the positive end of water (and vice versa), and each individual glucose molecule is dissolved in its entirety.

Factors Affecting Solvation

     The solubility of a solute is the maximum quantity of solute that can dissolve in a certain quantity of solvent or quantity of solution at a specified temperature.  The polar nature of the solute molecules compared to the solvent will also alter the solubility.  For aqueous solutions where polar water is the solvent, the more polar the solute molecules are the more likely the solubility will be higher.  Non-Polar Hydrocarbons have a very low solubility in polar water. So do such things as carbon tetrachloride, oxygen gas, nitrogen gas, and other non-polar molecular substances.  On the other hand polar solutes like ionic salts and polar molecular substances will have higher solubilities in water.  There is an old adage "Like dissolves in Like". If we were to use a non-polar solvent instead of polar water then the non-polar solutes above would have higher solubilities in that solvent. There are three factors that explain why solubilities will vary even within the same solution. 

1. Solute/Solvent Contact 

     The molecular size of the solute molecules will affect the solubility. The larger the solute molecules the more likely the solubility will diminish.  This explains why crushing the solute improves solubility.  The smaller the solute particles the more surface area that can contact solvent.  Hence, solute/solvent contact increases and as a result, solubility improves.  Likewise, if you stir the solution, there will be a greater opportunity for solute/solvent contact and again solubility improves.

2. Temperature 

    Generally speaking, the water solubility of a liquid or solid will increase with increasing temperature.  However, there are some exceptions to this.  Some solutes like solid Ce2(SO4)3 will have a decreasing water solubility with increasing temperature.  This depends upon the thermodynamics of the solution process.  If dissolving is an endothermic process; that is, it requires heat, then adding heat will increase the solubility of that solute.  This is true for most solids and liquids.  In this case the original solid or liquid was in a more stable lattice arrangement than there are when solvated and so it takes heat to overcome this stability.  However, for some solids and liquids and all gases, dissolving is an exothermic process.  This is because the original solute was not as stable as the solvated lattice and so energy which was needed to hold the original lattice together is no longer needed and is released from the solution.  If more heat is applied, there is increased resistance to dissolve the solute and solubility decreases.  Consequently, the solubility of all gases in water decrease with increasing temperature. That is why carbonated drinks that have carbon dioxide gas dissolved in them will become "flat" tasting when heated (or left to warm up).  The sparkle of the drink will have disappeared along with the carbon dioxide gas. 

3. Pressure 

    Pressure changes above the solution do not affect the solubility limits of solids or liquids dissolved in water.  However gaseous solutes are affected.  If the pressure of the gas is increased above the gaseous solution then the solubility will be increased in a linear fashion.  This is because the gas molecules are now confined into a smaller space and the number of collisions with the solvent will increases.  Hence solute/solvent contact increases.  If the pressure is reduced, there will be less collisions and less solute/solvent contact and solubility decreases.  In fact, more gas escapes than is being dissolved.  This explains why pop makes that hissing sound when you first open it.  There is a sudden reduction in pressure following immediately by a sudden decrease in solubility of carbon dioxide gas. This escape of a gas from solution is called effervescence.

Consequences of Solvation 

1. Dissolving and Temperature

    In order for a solute to be dissolved in a solvent, the attractive forces between the solute and solvent particles must be great enough to overcome the attractive forces within the pure solvent and pure solute. The solute and the solvent molecules in a solution are expanded compared to their position within the pure substances. 

Energy and Dissolving

     The process of expansion, for both the solute and solvent, involves a change in the energy of the system: this process can be either exothermic or endothermic. After dissolving, the solute is said to be fully solvated (usually by dipole-dipole or ion-dipole forces), and when the solvent is water, the solute is said to be hydrated. The separation of the solute particles from one another prior to dissolving is an endothermic process for both solvent and solute (steps 1 and 2), but when the solute and solvent combine with each other, this is an exothermic process (step 3). If the energy released in step 3 is greater than the energy absorbed in steps 1 and 2, the solution forms and is stable.  In either case, once this reaction is initiated, it cannot be stopped providing there is some attraction between the solute and the solvent. 

     This can be put to good use.  If you require a cold pack for an injury use endothermic dissolving. When the internal container is broken, the chemical inside this inner package mixes with the solvent surrounding it, absorbing heat in the process.  In this case, more heat was used to separate the solute and the solvent than was released when they formed a new lattice.  As a result, the pack becomes colder and removes heat from the injured area. 

Cold Pack

     On the other hand, if you require a hot pack, simply use exothermic dissolving.  In this case, when the two substances are allowed to mix, the result will be a generation of heat.  In this case, more heat was released upon formation of the new combined lattice than was required to separate the solute and the solvent from their original lattice networks. 

Hot Pack

     Reversible hot packs work by boiling and gently cooling the pack.  This allows it to form a fragile aqueous lattice network.  Typically, you push some sort of disk inside the bag to fracture the delicate lattice network which allows the intermolecular bonds to reform in a more standard fashion.  This new alignment is more stable and releases heat.  To reverse the process, the bag is boiled and the addition of heat forces the lattice to reform into a delicate state.  The process is not really reversible since external heat must be continually applied to create this delicate lattice.  When the lattice is broken, the heat added by boiling is simply released. 

2. Melting Ice 

    Melting ice is actually more complicated than it looks. When the rate of freezing is the same as the rate of melting, the amount of ice and the amount of water won't change on average (although there are short-term fluctuations at the surface of the ice).  The ice and water are said to be in dynamic equilibrium with each other.  The balance between freezing and melting can be maintained at 0°C, the melting point of water, unless conditions change in a way that favors one of the processes over the other. There will be no change in the size of the ice cube. 

Ice Water Equilibrium

    The balance between freezing and melting processes can easily be upset. If the ice/water mixture is cooled, the molecules move slower. The slower-moving molecules are more easily captured by the ice, and freezing occurs at a greater rate than melting. The ice cube will grow larger. 


    Conversely, heating the mixture makes the molecules move faster on average, and melting is favored. The ice cube will grow smaller. 


    Adding salt disrupts the equilibrium.  The salt molecules dissolve in the water, but do not attach easily to the solid ice. There are fewer water molecules in the liquid because some of the water has been replaced by salt. This means that the number of water molecules able to be captured by the ice (frozen) goes down, so the rate of freezing goes down.  In a sense, the salt prevents liquid water molecules from joining the lattice of the solid.  The rate of melting of the ice is unchanged by the presence of the salt, so melting is now occurring faster than freezing.  So the ice eventually melts. Typically, salts used on icy roads also dissolve exothermically, so as the salt ions become solvated, they also release heat in the process which subsequently increases the rate of melting and so the ice cube melts more. In essence, salt extracts water from ice and then melts it.  The stronger the attraction of the salt ions to water, the better it will be at extracting the water and subsequently melting the ice.  For this reason, CalCl2 is more effective (but more expensive) at melting ice due to the 2+ charge on calcium. 

3. Boiling Point Elevation/Freezing Point Depression 

    In order to return the system to equilibrium, where the number of molecules of water that are freezing is equal to the number of ice molecules that are melting (this is the freezing/melting point, we must lower the temperature sufficiently to make the water molecules slow down enough so that more can attach themselves to the ice.  When the number of water molecules that are freezing equals the number of ice molecules that are melting, equilibrium will be reached.  The higher the concentration of salt, the lower the freezing point drops. As ice begins to freeze out of the salt water, the fraction of water in the solution becomes even lower, and the freezing point drops further!  However, this doesn't continue indefinitely.  At some point the solution will become saturated with salt. This happens for salt in water at -21.1°C, which; therefore, is the coldest a saturated solution of salt and water can get.  At that temperature, the salt begins to crystallize out of solution, along with the ice, until the solution completely freezes.  The frozen solution is a mixture of separate salt (NaCl·2H2O) crystals and ice crystals. This heterogeneous mixture is called a eutectic mixture.  Any foreign substance added to the water will cause a freezing point drop.  For every mole of foreign particles dissolved in a kilogram of water, the freezing point goes down by roughly 1.8°C. Sugar, alcohol, or any chemical salt will also lower the freezing point and melt ice.  Salt is used on roads and walkways because it is inexpensive and readily available.  You might suppose that larger molecules might inhibit the freezing of water molecules even more, and have a more dramatic effect on the freezing point.  However, that isn't the case.  Actual molecules are so tiny compared to the distance they move through the liquid that size is hardly a factor at all. 

    Both pure solid and pure liquid phases of the solvent can coexist with their vapor and therefore, both phases have a nonzero vapor pressure.  In solution, the solvent coexists with its vapor.  Thus, at the freezing point of a solution, the solvent in the solution and the solvent in the solid (which is composed only of solvent) must coexist. For the solid phase, we can consider only the pure solid solvent vapor pressure curve, while in solution, we need to consider the solution vapor pressure curve. Thus, freezing point corresponds to the point where the solution vapor pressure curve intersects that of the pure solid solvent. This is the triple point on the phase diagram, as this is the point at which all three phases coexist.  We can compare this curve to that of the pure liquid solvent as shown in the figure below: 

Freezing Point Depression

    The extent of the depression, ΔTf, = i Kfm.  Here i accounts for the number of individual ions formed by a compound in solution (1 if the solute does not ionize or dissociate), Kf is the Freezing Point Depression Constant which should be negative (for water = -1.86oC/m) and m represents the concentration in moles of solute per kilogram of solvent.  The actual math is beyond the scope of this course.
 Solutes have a similar effect on the boiling point of water.  Since the solute is surrounded by water molecules, there will be fewer free water molecules available to leave the solution to enter the vapour phase.  The normal boiling point of a liquid is the temperature at which the vapor pressure reaches 1 atm.  Since solutions exhibit vapor pressure lowering, the temperature at which the vapor pressure reaches 1 atm will be elevated, a phenomenon known as boiling point elevation.  Consider the following curve:

Boiling Point Elevation

    The extent of the elevation, ΔTb, = i Kbm.  Here i accounts for the number of individual ions formed by a compound in solution (1 if the solute does not ionize or dissociate), Kf is the Boiling Point Elevation Constant (for water = 0.512 oC/m) and m represents the concentration in moles of solute per kilogram of solvent.  The actual math is beyond the scope of this course.  Boiling point elevation are often used to calculate molar masses which is also beyond the scope of this course. 4. Conductivity 

    The presence of ions in a solution allows electrical current to pass through the solution.  Consider the diagram below in which NaCl and NaOCl is dissolved into water. Two electrodes are connected to a source of electricity and inserted into the solution.  Electrons flow to one of the electrodes making it negative.  This electrode is called the cathode.  It will now attract the positively charged ions in solution, the cations.  When the cations contact the cathode, the cations pick up an electron and become neutral.  The removal of electrons from the system draws electrons from the other electrode, the anode, making it positively charged.  As a result, the negative ions in solution, the anions, are drawn to the anode where they donate their electron to the anode.  In the process, the anions become neutral.  The process will stop when there are no more ions available.  The circuit is not really complete in the sense that electricity flows in a complete circuit, but it is uninterrupted if free ions are available.  Water dissociates or ionizes solutes to create these free ions.  The extent of the flow of electricity depends upon the type and surface area of the electrodes used.  Consider the example below.

Solution Conductivity

Acids and Bases 

There are several methods of defining acids and bases.  While these definitions don't contradict each other, they do vary in how inclusive they are.  Initial definitions were based solely on how acids and bases behaved or operated .  These definitions were mostly qualitative in nature, but were not really formal definitions.  Several scientists contributed to these definitions: 

    In 1661, Robert Boyle described acids and bases by very simple characteristics: 
  • Sour taste 
  • Corrosive 
  • Change litmus (dye extracted from lichens) from blue to red 
  • Become less acidic when combined with alkalies. 
Alkalies (Bases): 
  • Feel slippery 
  • Change litmus from red to blue 
  • Become less alkaline when combined with acids.
Robert Boyle
    Around the same time, Antoine Lavoisier stated that all acids contained oxygen after studying several acids such as sulfuric and nitric acid.  Antoine Lavoisier
    It was not until 1881, that Lavoisier's work was formally questioned.  Humphry Davy noted that hydrochloric acid did not contain oxygen yet is an acid.  Soon thereafter, several more acids without oxygen were found such as hydrobromic and hydroiodic acid. Humphrey Davy

    In 1838,  Justus von Liebig suggested that acids contain one or more hydrogen atoms which can be replaced by metal atoms to produce salts.  For example HSCN is an acid because the H atom can be replaced by a metal to form a salt, such as NaSCN. 

Justus von Liebig
Svante Arrhenius     Although Justus von Liebig was the first to suggest that acids and bases can be more formally defined, actual conceptual definitions were not formally presented until 1884 when Svante Arrhenius proposed that salts dissociate when they dissolve in water to give charged particles which he called ions and in 1887, he extended this his idea by defining acids and bases as the following: 
  • Arrhenius acid - Any substance that ionizes when it dissolves in water to give the H+ ion. 
  • e.g. Arrhenius Acid
  • Arrhenius base - Any substance that ionizes when it dissolves in water to give the OH- ion. 
  • e.g. Arrhenius Base


   Arrhenius's theory helped to explain why acids have similar characteristics, since they all give H+ ions when they dissolve in water.  It also explained why acids are neutralized by bases and why bases are neutralized by acids;  the H+ ions from acids combine with the OH- ions from bases to form water: 


   Though the Arrhenius theory helped to explain more about acids and bases, there were still several drawbacks to this theory. 

  • The theory can only classify substances when they are dissolved in water since the definitions are based upon the dissociation of compounds in water. 
  • It does not explain why some compounds containing hydrogen such as HCl dissolve in water to give acidic solutions and why others such as CH4 do not. 
  • The theory can only classify substances as bases if they contain the OH- ion and cannot explain why some compounds that don't contain the OH- such as Na2CO3 have base-like characteristics.
    To extend the Arrhenius theory a little further, consider the formation of water from the combination of an H+ ion and an OH- ion as a dynamic equilibrium:

     Based on the fact that the above reaction is reversible, we can conclude the following: 

  • and acid  is any substance that increases the concentration of the H+ ion when it dissolves in water. 
  • a base is any substance that increase the concentration of the OH- ion when it dissolves in water. 
    Now, substances not containing H+ or OH- ions can be classified as acids or bases if they alter the [H+] or [OH-] when they dissolve in water.   For example CO2 cannot dissociate to give H+ but it does increases [H+] when it is dissolves in water. 

Creating Acid

     Likewise CaO cannot dissociate to give OH- but it does increase [OH-] when it dissolves in water. 

Creating Base

    In 1923  Johannes Brønsted and Thomas Lowry extended the understanding of acids and bases as they separately proposed a new set of definitions for acids and bases which are known as either Brønsted acids and bases or Brønsted-Lowry acids and bases. 

Johannes BrønstedThomas Lowry

  • Brønsted Acid - Any substance that can donate a proton, H+ ion to a base.  That is, it is a hydrogen-ion donor or proton donor.
  • Brønsted Base - Any substance that can accept a proton, H+ ion from an acid.  That is, it is a hydrogen-ion acceptor or proton acceptor .

    In the above reaction, the H from HCl is donated to H2O which accepts the H to form H3O+, leaving a Cl- ion.  The Brønsted-Lowry model of acids and bases brings rise to the concept of conjugate acid-base pairs.  The part of the acid remaining when an acid donates a H+ ion is called the conjugate base.  The acid formed when a base accepts a H+ ion is called the conjugate acid. 

For the generic acid HA: Acid

For the generic base A-: Base

More examples of conjugate acid-base pairs:



    In the following reactions, it is shown how H2PO4- and H2O can act as both acids and bases.  Such compounds are said to be amphoteric


    Strong acids have weak conjugate bases and strong bases have weak conjugate acids.  Water has the tendency to equalize the strengths of all strong acids and strong bases, regardless of the strength of the acid itself.  This is known as the leveling effect.  Acids are limited to the strength of the H3O+ ions that they form when they lose H+ ions when they dissolve in water.  Likewise, bases are limited to the strength of the OH- ions that they form when they gain H+ ions when they dissolve in water. 

    There are numerous advantages to the Brønsted-Lowry model of acids and bases:

  • Acids and bases can now be ions or neutral molecules. 
  • Acids and bases can now be any molecule with at least one pair of nonbonding electrons. 
  • It explains the role of water in acid-base reactions; H2O accepts H+ ions from acids to form H3O+ ions.  The  H3O+ is present in all acid base reactions, but is usually represented by simply H+.  It is called the hydronium ion.
  • It can be applied to solutions with solvents other than water and even in reactions that occur in the gas or solid phases. 
  • It relates acids and bases to each other with conjugate acid-base pairs and can explain their relative strengths. 
  • It can explain the relative strengths of pairs of acids or pairs of bases. 
  • It can explain the leveling effect of water.
    Acidic and basic are two extremes that describe a chemical property chemicals.  Mixing acids and bases can cancel out or neutralize their extreme effects.  A substance that is neither acidic nor basic is neutral. The pH scale measures how acidic or basic a substance is.  The pH scale typically ranges from 0 to 14.  A pH of 7 is neutral.  A pH less than 7 is acidic.  A pH greater than 7 is basic.  

pH Scale

    The pH scale is logarithmic and as a result, each whole pH value below 7 is ten times more acidic than the next higher value. For example, pH 4 is ten times more acidic than pH 5 and 100 times (10 times 10) more acidic than pH 6.  The same holds true for pH values above 7, each of which is ten times more alkaline (another way to say basic) than the next lower whole value. For example, pH 10 is ten times more alkaline than pH 9 and 100 times (10 times 10) more alkaline than pH 8.

    Acid/base indicators (also known as pH indicators) are substances which change colour with pH.  They are usually weak acids or bases, which when dissolved in water dissociate slightly and form ions.  Consider an indicator which is a weak acid, with the formula HIn. At equilibrium, the following equilibrium equation is established with its conjugate base:

pH Indicator Equation

    The acid and its conjugate base have different colours.  At low pH values the concentration of H3O+ is high and so the equilibrium position lies to the left. The equilibrium solution has the colour A.  At high pH values, the concentration of H3O+ is low - the equilibrium position thus lies to the right and the equilibrium solution has colour B.  At a low pH, a weak acid indicator is almost entirely in the HIn form, the colour of which predominates. As the pH increases - the intensity of the colour of HIn decreases and the equilibrium is pushed to the right. Therefore the intensity of the colour of In- increases. An indicator is most effective if the colour change is distinct and over a low pH range.  A Universal Indicator is a mixture of indicators which give a gradual change in colour over a wide pH range - the pH of a solution can be approximately identified when a few drops of universal indicator are mixed with the solution.  Consider the following list of some indicators. Notice their distinct acid and base colours and the transition zone of each indicator. 

Indicator Chart

    The colours associated with a common universal indicatory is as follows:

Universal Indicator

    A pH meter is a device used for more precise pH measurements.  The device measures the potential difference between known reference electrode and the measuring pH electrode conducted by the presence of the H3O+ ions in the solutions and scaled in such a way that it displays not the measured potential, but ready pH value.

pH Meter

    One common source of confusion deals with the following pairs of words:  concentrated/dilute and strong/weak.  Lets consider an acid.  If the acid is concentrated, there is lots of it added to water.  However, this acid may still only be weak if only a small proportion of it ionizes to release the H+ into water.  Since lots of acid was added to water, the solution may still be quite acidic.  For example, lets say 1 million molecules of of acid were added to water.  If the acid only ionizes 5%, the acid is weak, but since 5% of 1 million is 50000, there is still a fair bit of H+ present.  Now, lets say that 100000 molecules (1/10 as much) acid is added to water so this acid is more dilute. However, this second acid is strong and ionizes 90% so there 90% of 100000 or 90000 molecules of H+ present.  To determine the acidity of a solution, you must consider both the strength and the concentration. 

Strong Versus Weak

    Typically, concentrated solutions must be diluted to make them more useful.  One way to do this, would be to remove some of the solute, but this is tricky to do.  It is usually easier to simply add more solvent to the solution.  However, if we wanted to dilute 100 mL of a 10 molar solution into a 0.1 molar solution, we would have to add 9.9 L of water which would be ridiculous especially considering we only likely need about 100 mL of the diluted solution.  To do this, we often set up a serial dilution or dilution series.  This allows us to keep volumes small and quickly achieve the desired concentration.

    Consider a ten times or ten fold dilution. 10 Fold Dilution This series has 4 consecutive dilutions and so the concentration of the final vial is now 10000 times less than the original vial. Notice that since this is a 10 fold dilution, the ratio of the sample to be dilution to the final volume of the solution it creates is 1:10 (1 mL of sample was added to 9 mL of water to make a total volume of 10 mL. 

Using Solubility

     Reactions where soluble compounds react to form insoluble ones are called precipitation reactions.  The reverse is a dissolution reaction where solid compounds dissolve upon addition into water.  There is no simple set of rules which can be used to predict which compounds will be soluble or not; however, some useful trends have been observed:
  • multiple-charged ions are less soluble than single-charged ions.
  • smaller ions are more soluble than larger ions.
     Explicit solubility rules exist and can be applied most of the time with certainty.  Something is considered insoluble if it forms a precipitate at a concentration of 0.1 mol/L or more.

Solubility Table

    This information can be put to great practical use.  It is often used to analyze the quality of drinking water by determining if any polluting ions are present.  For example, lead in drinking water can be verified by adding some chloride to the water (assuming there is no other ions of concern present).  This table can also be use to remove known ions from a solution to purify the water.  In each case the precise sequence of steps in critical to absolutely verify and/or remove a specific ion.

Separation of Ions

    Notice that to remove an ion, a soluble salt is added that leads to the desired precipitation following which the solid is filtered out of the remaining solution which is then further treated.

    When you are asked to predict whether a precipitation reaction takes place when two aqueous solutions of ionic compounds are mixed follow these steps:

Step 1 Determine the formulas for the possible products by identifying the type of reaction involved (such as double displacement, single displacement and etc.).

e.g. AB  +  CD       AD  +  CB
Step 2 Predict whether any of the possible products is water insoluble.  If a possible product is insoluble, a precipitation reaction takes place, and you will continue with step 3. If nothing is insoluble, write “No reaction”.
Step 3

Write the physical state for each formula.  The insoluble product will be followed by (s) and water-soluble ionic compounds will be followed by (aq).  Balance the equation.  Don't forget to use proper nomenclature to ensure you have the correct number of ions within a compound.

    For example:  Predict whether a precipitate will form when water solutions of silver nitrate, AgNO3(aq), and sodium sulfide, Na2S(aq), are mixed.  If there is a precipitation reaction, write the complete and net ionic equation that describes the reaction.

This is a double displacement reaction:
AgNO3(aq)  +   Na2S(aq)   Ag2S(?)  +   NaNO3(?)
Now determine if any precipitates will form:
Ag2S(?) is not soluble.
Complete the equation and balance
2 AgNO3(aq)  +   Na2S(aq)   Ag2S(s)  +   2 NaNO3(aq)

Ionic Equations

    The presence (or lack) of a driving force for reactions done in water can be evaluated using a net ionic equation. The total and net ionic equations show the real physical state of every component of the reaction. The compound is a strong electrolyte (a soluble ionic compound or a strong acid), then the compound is separated into the appropriate aqueous ions. Otherwise it is left in an undissociated (ionic compounds) state or un-ionized state (weak acids, bases and water). Consider some examples:

Example 1
Example 1

    Example 2 reveals a very unique type of ionic reaction; one in which there is simply a transfer of electrons from one species to another.  These reactions appear simple but actually have profound effects in our environment.  For example, rusting of metal is an example of such a reactions.  These reactions are called redox reactions, or oxidation-reduction reactions and they have a number of similarities to acid-base reactions.  Fundamentally, redox reactions are a family of reactions that are concerned with the transfer of electrons between species. Like acid-base reactions, redox reactions are a matched set in that you don't have an oxidation reaction without a reduction reaction happening at the same time. The compound that loses an electron is said to be oxidized. This comes from the observation that materials combine with oxygen in varying amounts. For instance, an iron bar oxidizes (combines with oxygen) to become rust. We say that the iron has oxidized. The substance which gains an electron is said to be reduced. Consider the following animation (please note that it starts over):

Redox Reaction

    Notice that the sodium atoms lose an electron (oxidation) while the chlorine atoms gain an electrons (reduction).  Consider another example:  

Fe(s) + Cu2+(aq) Fe2+(aq) + Cu(s)

    Fe(s) donates two electors (oxidation) to the Cu2+(aq) (reduction) to form Cu(s) . Since the Cu2+(aq) takes two electrons from the Fe(s) causing Fe(s) to be oxidized, Cu2+(aq)) is an oxidizing agent.   Likewise since Fe(s) forces Cu2+(aq) to accept two electrons, thus Fe(s) brings about reduction and is therefore a reducing agent. Try not to confuse the terms.


Redox Nemonic

Example 1

Example 1