Erroneous IB Syllabus a Great Opportunity To Shift Solubility Away From Acid-Base Education

Acid-base indicators incorporate many levels of acid-base theory.  Usually what an indicator is follows the analysis of a titration curve to assess the understanding of end point vs. equivalence point.  By this time, the student is expected to have some familiarity with pH, strong vs. weak, salt hydrolysis, buffers, acid-base equilibrium calculations as well as the stoichiometry calculations that accompany a titration.  

Indicators can be a fantastic opportunity to reinforce complex ideas from acids and bases because the selection of an appropriate indicator can accompany analyzing why the pH changes most rapidly near the equivalence point and this can provide clarity to the logarithmic scale of pH that is often masked by the even more troubling scales of concentration.  Beyond the titration curve, an indicator is a great chance to teach or reinforce the notion of weak acids and bases.  

An acid-base indicator is typically a weak acid that changes color when it changes form to the conjugate base (or vice-versa).  When a small amount of indicator is put into a solution, the indicator will give information about the solution via a very simple observation, the color it becomes.  A solution with higher concentrations of hydrogen ions will form mostly the protonated version of the indicator while a solution with higher concentrations of hydroxide ions will form mostly the deprotonated version.  There will always be some of both forms present.  When the pH of the solution matches the pKa of the indicator, there will be equal quantities of basic and acidic form of the indicator.  These explanations of indicators are very well summarized by the 2016 IB chemistry syllabus on page 94 (copied below).  

The indicator in the acidic form will occur in acidic solutions where pH is less than the pKa, the basic form of the indicator will be present more in solutions where pH is greater than the pKa.  When the pH is near the pKa we will see some of both colors present.  For example, bromothymol blue is an indicator that is yellow in the acidic form and blue in the basic.  When put in a solution with a pH near the pKa of bromothymol blue (7.1) then both forms of indicator are present and the solution will be green.  At a pH of 6.1 there will be 10x as much of the yellow form and at a pH of 8.1 there will be 10x as much of the blue form.  

There is a small problem, however, with the language used by the IB syllabus above.  The indicator is really not a weak acid in isolation, rather it is a chemical that can exist in an acidic or basic form.  It can be both a weak acid and a weak base and which it is largely depends on what solution it is put into.  The reaction could easily have been rewritten in reverse and nothing about the chemical would be altered.  This could stem from an early introduction to acids and bases where they are treated in isolation.  Many are taught that acids are chemicals with a pH less than 7 and bases have a pH greater than 7.  Instead it should be taught that an acid is something that transfers a hydrogen ion to something else and a base takes that hydrogen ion from something else (or electron pair interactions if we expand to include Lewis acids/bases).  This is a very critical difference that often is overlooked.  The key here is that an acid does not exist until a base is present.  Much like oxidation and reduction being reliant upon one another, so are acids and bases.  Hydrochloric acid is not an acid until it has something capable or taking away its proton.  Once you put hydrochloric acid into water, the hydrochloric acid reacts with that water to transfer a proton to it and the water acts as the base accepting that proton.  For an indicator, the indicator is acidic when it is transferring a proton to something else and a base when taking a proton away from something else.  To label an indicator as a weak acid is both incomplete and it also reinforces the weaker approach to acids and bases.  From here the syllabus gets worse.

Below the explanation of weak acid indicators, weak base indicators are introduced.  There are several issues with this.  First of all, there are no indicators that react in this manner.  Every indicator presented in the data booklet is of the type that changes color via a proton transfer (the indicators presented in section 22 of the data booklet can be found at the end of this post).  Some have a pKa less than 7 indicating that they change color in an acidic environment and some have a pKa greater than 7 indicating that they change color in a basic environment.  But as explained above, these are all chemicals that act as both weak acids and weak bases.  If the IB syllabus had stated that the indicators that are weak bases react as: In-(aq) + H+(aq) → HIn(aq) it would be misleading and in poor form educationally, but not as incorrect.  Instead they chose to show a very odd solubility reaction that in no way is an acid-base reaction.  Why would they choose this generic reaction?  

The purpose of this reaction is to accompany a poor way of defining acid strength, solubility.  A strong acid is nearly always defined as an acid that completely ionizes when dissolved in water.  Strong bases are treated the same.  This is a problem because it is often interpreted that the action that makes something a strong acid is that it dissociated in water.  What makes something a strong acid is that it has a weak bonding interaction with the hydrogen ion attached to it.  This allows most bases to easily remove that hydrogen ion.  In other words, a strong acid is very good at transferring a proton to something else.  The ability to transfer the hydrogen ion makes something a strong acid.  A strong base is something that has a strong affinity for hydrogen ions and can easily take them from something else.  When you put a strong acid into water, the strong acid is good at donating a proton to the water and so it does so very frequently with fast kinetics and very soon all of the hydrochloric acid has reacted with water to form chloride ions and hydronium ions.  

HCl(g) + H2O (l) → Cl-(aq) + H3O+ (aq)

Sometimes this is written as HCl(g) → H+(aq) + Cl-(aq) which is not nearly as accurate because the hydrochloric does not split apart.  A water molecule or other base is needed for this separation to occur.  

Some strong bases are not even used in water.  NaH is a strong base using commonly in organic chemistry and is delivered in an organic solvent.  The NaH is delivered using an organic solvent to the reaction mixture.  If water is present the NaH will react with the water to produce hydrogen gas and NaOH.  NaH is a strong base because the hydride has a very large amount of negative charge that can easily attract a proton to become hydrogen gas H2.  

Historically we have easily been able to detect the presence of these ions through conductivity tests that show that the strong acids like HCl produce ions in solution.  There are a lot of limitations that arise when we try to make this our definition of strength.  The act of dissolving is not actually what produces the ions.  When ammonia (NH3) dissolves in water it forms aqueous ammonia.  A very small amount of this aqueous ammonia reacts with water to form hydroxide ions.  The reason why ammonia is a weak base is because it is not as good at pulling an H+ away from a water molecule as water is at pulling it back.  Very few collisions result in the transfer of a proton from the water to the ammonia.  But ammonia is highly soluble in water (for a gas).  It is the reaction with water, the acid-base reaction that produces the ions.  

For strong bases, some of them dissolve in water to produce ions that directly impact the equilibrium affecting pH.  Sodium hydroxide dissolves in water to produce sodium ions and hydroxide ions.  

NaOH(s) → Na+(aq) + OH-(aq)

This is where a lot of problems arise.  This dissolving is directly producing ions that will affect the equilibrium in the solution.  These ions still react.  A hydroxide ion will react with a water molecule, but it produces a different hydroxide ion and a different water molecule.  If we dissolved NaOH into water with an oxygen isotope we would see reactions of OH-(aq) + H2O18(l) → H2O(l) + O18H-(aq) where a proton would transfer from the water to the dissolved hydroxide.  Another problem here is how to label bases such as Mg(OH)2 and Al(OH)3.  These bases have very limited solubility in water, but the dissolved components are ionized.  The large amount of charge on the metal cations causes the electron density on the oxygen atom in the hydroxide to shift further towards the cations and is less available to pull hydrogen ions away from water molecules.  Some call these bases insoluble bases, some call them strong bases and some call them weak bases.  

When we try and expand this definition of strong acids and bases to apply towards weak acids and bases things continue to falter.  A weak acid or weak base is defined as an acid or base that partially ionizes when dissolved in water.  I once heard a teacher explain that weak and strong are bad adjectives to use to describe acids and bases because it is always relative to what it will react with.  They recommended using weaker and stronger instead.  This makes a lot of sense on a particulate level.  A generic reaction for this would be:

HA + B → A- + BH+

If HA is a stronger acid than BH+ or B is a stronger base than A- the reaction will be product favored at equilibrium.  A- could be a very strong base and B could still be stronger and that would be all that matters.  

Defining strength for acids and bases by acid-base reactions relative to the conjugates formed allows a much better and deeper understanding of acids and bases.  It can be used to reinforce the notion that an acid and base are dependent upon one another to react.  It also can easily be summarized in the minds of a student since a strong acid is something good at being an acid.  A strong base is something good at being a base.  This can be tied into bonding and thermochemistry quite easily.  Using solubility to define acid-base strength is confusing.  It muddles solubility and acid-base reactions and produces students that do not fully understand acid-base chemistry.  The IB chemistry syllabus should never have had the definition of a weak base indicator.  It is rife with misconception and poor concept of acid-base chemistry and will mislead teachers and students in exchange for an algorithmic exchange of chemistry related statements (IE when I say weak acid, you say partially ionized).  Even reading through the email responses I see flaws in the analysis of my challenges to the syllabus.  The full text of emails are presented below with names removed.  

Text of email exchanges with IB support:

Stakeholder By Web Form (Scott Milam)

07/15/2015 02.51 AM

Hello, I am working on the acid-base 18.5 section in the 2016 syllabus and it says: For an indicator which is a weak base: – BOH(aq) --> B+(aq) + OH-(aq) Colour A Colour B This makes no sense to me, is this an error that is being revised or is there some example of some odd organic indicator that fits this criteria? This is not an acid-base reaction and I would teach a basic indicator as In-(aq) + H+(aq) --> HIn(aq) exactly the same as it indicates for a weak acid indicator since the conjugate is merely another weak acid or base. I would prefer not to teach what it says in the syllabus as it appears to me based on the application of the limited definition of completely ionized in water for strength.

Response Via Email (Administrator)

07/17/2015 07.33 PM

Dear Scott, Thank you for your e-mail. I believe that this example has been taken to draw attention to the fact that indicators can also be weak bases although, you are right, they are not really common. These reactions are not acid base reactions but relate to dissociation of weak acids and bases. Teaching the following reaction, In-(aq) + H+(aq) --> HIn(aq), to explain indicators is perfectly fine. Best regards,

Stakeholder Via Email	08/10/2015 06.10 PM
Is there one example of a weak base indicator you could provide that fits the syllabus definition?

Discussion Thread
Response Via Email (Administrator)	08/14/2015 05.58 PM
Dear Scott, I agree that the equation should have been a salt reacting with water forming hydroxide ions but the authors of the syllabus decided to simplify the reaction to stick to the rather general definition of a (weak) base. Best regards,

I can’t format a double arrow, so I just used single arrows even when equilibrium for a reaction would make a double arrow more appropriate.  

I also ignored the dissolving of a cation as a Lewis acid-base reaction in parts of this analysis for simplicity, but complexation is a type of acid-base reaction as well.