Dinosaur chemistry
Unnecessary concepts
It is all too easy to think that in the past chemists got some of it wrong and that these days we know exactly what we are doing. There are obvious examples of theories that were once prevalent but are now discarded, even discredited. Examples include the phlogiston theory of the seventeenth century and the theory of ‘molecules’ such as sodium chloride which transfer electrons when they conduct - a theory which was overthrown by Arrhenius who showed that strong electrolytes such as sodium chloride are completely dissociated into their ions in solution. But even if some of the theories used in chemistry remain intact it is interesting to see how the way in which they are taught has changed.
In the past all chemists learned the ‘correct’ way of approaching problems. Certain facts were extremely important and it seemed as though chemistry could not be taught or understood without using these facts. In reality these were more about a particular approach to solving problems. Without changing the underlying theory we now solve these problems in a different way. I have coined the term ‘Dinosaur Chemistry’ to describe these obsolete approaches. Why is this worth looking at? My answer is that if we can see that approaches we adopted in the past and once seemed indispensible are now no longer necessary then perhaps we can look at today’s approaches and predict what may seem equally archaic and unnecessary to chemists in the future.
1. Equivalent weight and normality
When you buy ampoules of hydrochloric acid to make volumetric solutions you will see that they are often labelled still using N as well as M. N stands for a normal solution whereas M of course is a molar solution measured in mol dm-3.
When I was at school we had to learn about equivalent weights and normality in order to solve titration problems.
The equivalent weight of an element was (is) defined as the number of parts by weight of the element which combine with (or displace) one part by weight of hydrogen or eight parts by weight of oxygen. It was a legacy from the early years of chemistry when chemical formulas and atomic weights were often arrived at by combining elements with oxygen or hydrogen.
For example: 0.180 g of magnesium burns in air to form 0,300 g of magnesium oxide.
From this we can calculate the equivalent weight of magnesium.
Weight of oxygen = 0.300 – 0.180 = 0.120 g
1 g of oxygen combines with 0.180 / 0.120 = 1.50 g of magnesium
So 8 g of oxygen combine with 8 x 1.50 = 12.0 g of Mg
Hence the equivalent weight of magnesium = 12.0 g
Chemists related equivalent weight to valency by the formula
Valency = atomic weight ÷ equivalent weight
The atomic weight (we now call it atomic mass) of magnesium is approximately 24 g mol-1 so the valency of magnesium is 2.
When solutions were made up their concentration were expressed as a normality rather than as a molarity.
Normality:
A normal solution (N) of a substance contains the equivalent weight of the substance in 1.0 dm3 of solution.
So that for any balanced equation:
n1 x v1 = n2 x v2
This now seems very strange so why was it necessary?
Let’s consider the redox reaction between Fe2+(aq) ion and an acidified solution of manganate(VII) ions
A typical question might have been:
If 25.0 cm3 of 0.1N Fe2+(aq) react with 25.0 cm3 of KMnO4(aq) in acid solution what is the concentration of the KMnO4 solution?
Using the equation n1 x v1 = n2 x v2 very quickly leads us to the answer 0.1N
Nowadays students will not understand what is meant by a 0.1N solution and solve such problems using molarities .
If we use molarities to solve the problem we will need to convert 0.1N into a concentration in mol dm-3.
The half equation is:
Fe2+(aq) → Fe3+(aq) + e–
So the equivalent weight is simply equal to the formula mass
So 0.1N is the same as 0.1 mol dm-3.
From the equation: MnO4-(aq) + 8H+(aq) + 5Fe2+(aq) → Mn2+(aq) + 4H2O(l) + Fe3+(aq)
we can see that 5Fe2+(aq) reacts with one mole of MnO4-(aq)
Since 25.0 cm3 of 0.1 mol dm-3 Fe2+(aq) react with 25.0 cm3 KMnO4
Concentration of KMnO4(aq) = 0.02 mol dm-3
This means that a molar solution of KMnO4(aq) is actually five times stronger than a normal solution of potassium manganate(VII) (see image).
Fifty years ago it seemed impossible to solve these sorts of titration problems without using the concepts of equivalent weight and normality. Now they are completely redundant although normality is still used in many catalogues and even now I still occasionally come across students still using normality when I am marking extended essays.
2. Relative density
Another type of dinosaur chemistry is the use of Relative Density. In the past molecular masses of gases were often determined by comparing equal volumes of two gases and Relative Density usually meant a comparison with hydrogen although any known gas could be used. By comparing the mass of a volume of the gas with the mass of the same volume of hydrogen at the same temperature and pressure the Relative Density simply needed to be multiplied by two (the molar mass of hydrogen gas) to give the molecular mass of the gas. Now we use other ways such as a mass spectrometer to determine molecular masses so the term relative density, although still valid, has fallen into misuse.
Candidates for future redundancy?
So can you think of any concepts that we use now which may become redundant in the future? One of my favourites is oxidation states (see How useful are oxidation states?).
Oxidation states do not exist in reality (atoms do not go around with their oxidation state stamped on them) and actually make little sense when it comes to organic chemistry particularly as you have to assume a covalent compound is ionic in order to arrive at the oxidation state. I was amused to see that the new Cambridge Pre-U qualification now talks about lower and higher order carbonyl compounds depending on the oxidation number (unlike the IB, they still call it oxidation number rather than oxidation state) of the carbon atom in the carbonyl compounds. I wonder how long this will last as higher and lower order carbonyl compounds do not appear on any other programme as far as I am aware.
Another concept that seems to me totally pointless is the use of the cell convention. It really doesn’t matter which is the left hand cell and which is the right hand cell and whether you write the oxidized form first or the reduced form first. You can calculate all you need so long as you know the values for the standard electrode potentials of the two half-cells.
Do you agree with my two choices or do you have another favourite candidate for a redundant concept in the future? How about getting rid of the use of anode and cathode and just using (+) and (−) then it wouldn't matter whether we are talking about an electrolytic cell where the anode is (+) or a voltaic cell where the anode is (−)?