Lab Report 2 – Titration
CHEM1903 – Chemistry 1A (SSP) Michael West (305159240)

1. Experiment 2.2 – Titrimetric determination of the molecular mass of an organic acid
Method An unknown organic acid was supplied in solid form. The acid was known to be diprotic and had the reference number 19. Using an analytical balance, 1.5397 g of the acid were weighed out, and made up with deionised water into 250 mL of solution. 25 mL of the acid solution was added to a conical flask with phenolphthalein indicator and titrated against standardized 0.0983 M NaOH solution. Three titrations were performed and the results averaged. The molar mass of the acid was then calculated and compared to a list of given possibilities. Results and Calculations The three titres were 26.30 mL, 26.50 mL and 26.30 mL. The mean titre volume was hence 26.37 mL. The number of moles of NaOH was thus moles.

Because the acid was diprotic, reaction stoichiometry dictates that there was one mole of acid for every two moles of NaOH. Accordingly, in 25 mL of the acid solution, there were moles of acid. The molar mass of the acid is then g⋅mol-1. This matches most closely with

succinic acid, for which the given molar mass was 118.1 g⋅mol-1. Although this represents a 0.6% discrepancy, the error is small enough to identify the acid as succinic acid with a high degree of certainty, given the possibilities listed.

2. Experiment 2.3 – Determination of the carbon dioxide and hydrogencarbonate contents of soda water by indirect titration
Method A known excess of standardised sodium hydroxide was added to a known amount of chilled soda water. The remaining sodium hydroxide after reaction with the soda water was determined by titration against standardised HCl, to two separate endpoints. A fuller description of the method is available in the lab book. Results Define titre A as the volume of HCl added in the back-titration to the phenolphthalein endpoint, and titre B as the volume of HCl added between the phenolphthalein and methyl orange endpoints. Volume (mL) Titration 2 Titration 3 0.00 0.00 18.00 47.95 29.95 18.10 48.25 30.15

Point Origin Phenolphthalein endpoint (volume of titre A) Methyl orange endpoint Volume of titre B

Titration 1 0.00 17.85 47.60 29.75

Mean n/a 17.98 n/a 29.95

The titrations were not exactly reproducible, with the volume of both titre 1 and titre 2 increasing from titration 1 to titration 3. This might be because the titrations were not performed simultaneously, with the soda water aliquots of titrations 2 and 3 allowed to stand exposed to the atmosphere while titration 1 was being carried out. As it warmed, this soda water would have undergone degassing, with HCO3- and H2CO3 being converted to gaseous CO2, released to the atmosphere. Thus, less NaOH reacted with the later aliquots, and more was left to react with the HCl, so the titres of HCl were larger.

3.
Equilibria between carbonic acid, hydrogen carbonate and carbon dioxide, and ions in water:

Titration of HCl with NaOH:

4.
If the density of water at 20°C is 0.998 g/mL, then the solubility of CO2 in water at 20°C of 0.145 g/100 g is equivalent to . The molar mass of CO2 is . .

Hence, the concentration of the gas at saturation is

As a check, we can derive a similar result using Henry’s Law for the concentration of gases in solution. Details of this are given in an appendix, with the result of matching the above result to only 1 significant figure, but at least giving confirmation of the order of magnitude. The numbers for saturation point are hence much lower than the experimentally determined concentration of CO2 in the soda water, . Since the experimental concentration was determined after the cap was taken off, we conclude that even with the cap off, the soda water was still super-saturated in CO2 (at least initially; over time, degassing will cause the soda water to become unsaturated). Before the cap was removed and any degassing occurred, the concentration must have been higher, so it would also have been super-saturated then. Two key general principles are operating. The first is related to Henry’s Law, . As mentioned, more details are available in the appendix, but we note that concentration of CO2 in the aqueous phase is directly proportional to the partial pressure of CO2 in the gaseous phase above the liquid, . Hence, when the seal is broken and the pressure drops, the concentration of aqueous CO2 falls as it is released to the atmosphere as gas. Secondly, we can observe the effect of temperature on the equilibrium constant of the reaction (leaving aside for a moment other species such as and ). As the soda water warms, the change in temperature will shift the equilibrium position of the reaction, favouring the production of gaseous CO2 by dissolution.

5.
The solubility of the CO2 in solution is positively correlated with partial pressure of CO2 outside and negatively correlated with the temperature of the solution. Accordingly, degassing could be prevented by opening the bottle of soda water in a high-pressure CO2 environment or at a very low-temperature environment. Of these two, the second is probably more feasible, and the motivation behind using chilled soda water in the experiment.

6.
To determine the carbonate content of beverages where titration is not possible, we may take advantage of the degassing effect. Firstly, we weigh the beverage, then encourage the release of all CO2 contained within the solution. This may be achieved by heating, by the reduction of the outside pressure, by the addition of nucleation sites such as salt crystals (although any added mass must be recorded and accounted for), and by the addition of basic substances to convert the carbonate contained in carbonic acid and hydrogencarbonate ions into CO2. When the beverage’s mass becomes constant, the change in mass from the initial value indicates the amount of dissolved CO2 and carbonate that was originally present.

Appendix – concentration of CO2 in water using Henry’s Law
The concentration of gases in solution is given by Henry’s Law, the gas in the aqueous phase (in mol/L of solution) and atm). , where is the concentration of is the partial pressure in the gaseous phase (in where is the enthalpy of solution and is the

is Henry’s constant, which can be determined using ),

the symbol denotes standard conditions ( gas constant. From reference material, for carbon dioxide, We aim to find the concentration at

(1) and , so we have:

(2).

Since the given CO2 head pressure is 101.3 kPa, equivalent to 1 atm, we can write simply:

1

R. Sander (1999), Compilation of Henry's Law Constants for Inorganic and Organic Species of Potential Importance in Environmental Chemistry (3rd ver.), [www.henrys-law.org]. 2 J.F. Liebman (1997), Some thoughts on the solubility of carbon dioxide and silicon dioxide in water, Struct. Chem. 8(5), 379-381.