Since this is a 2:1 ratio;
n(H2SO4) = ½n(NaOH) = ½(2.002x10-2 mol)
= 1.001x10-2 mol
Using the molarity equation (again);
V = n/C
V = (1.001x10-2mol)/(1.0707mol.dm-3)
V = 0.009349 dm3 ± 0.05 %
V2(H2SO4) = 9.349 cm3 ± 0.05 %
This is the volume of sulfuric acid required to induce the reaction in (II), producing Na2SO4, it will result in the presence of half the number of mole NaOH in H2SO4. Note that as would be expected mathematically this volume is half that required in (I).
Results:
Fig. 1
A Table to Show Initial Observations of Solutions
Fig. 2
A Table to Show Burette Volumes and Relevant Values
*Note: ± 0.10 due to addition of component errors for initial and final volume
†Note: This is the initial burette reading plus the amount of sulfuric acid required for the reaction. It serves as a useful calculation when using a burette, as it gives an end-point value.
Fig. 3
A Table to Show Observations During Addition of Nominated Volumes Sulfuric Acid
Fig. 4
A Table to Show Observations During Evaporation and Preparation of Stock Solution
Initial Observations of Sodium Carbonate: Fine White Powder
Fig. 5
A Table to Show Results of Litmus Test and Addition of Sodium Carbonate for Both Solutions
Fig. 6
A Table to Show Results of Litmus Test for Sodium Hydrogen Carbonate (NaHCO3)
Fig. 7
A Table to Show Key Errors
Collating, Interpreting and Analysing Results:
Responding to the method
7. The litmus and carbonate tests (fig. 5) were crucial in highlighting the properties of the products for each reaction. Clearly the 1:1 reaction produces an acidic entity, since it turned litmus red, where the 2:1 solution turned litmus blue, indicating the presence of a base. This evidence is corroborated by each reaction with sodium carbonate; the 1:1 reacted, where the 2:1 did not. This is because any acid will react with a carbonate to produce salt, water and carbon dioxide, and a base will not generally react at all. A simple reaction between an acid and a carbonate is exemplified below:
CaCO3(s) + 2HCl(aq) → CaCl2(aq) + H2O(l) + CO2(g)
The reaction between the acid salt produced in reaction (I) and sodium carbonate would be governed similarly:
2NaHSO4(aq) + Na2CO3(s) → 2Na2SO4(aq) + H2O(l) + CO2(g)
Thus, if the evolved gas was tested with a flaming splint, it would be extinguished – CO2 gas is evolved.
-------------------------------------------------------//--------------------------------------------------------------
8. Although these equations have already been explored during preliminary calculations, they shall be restated here:
(I) NaOH(aq) + H2SO4(aq) → NaHSO4(aq) + H2O(l)
(II) 2NaOH(aq) + H2SO4(aq) → Na2SO4(aq) + 2H2O(l)
In summary; when the number of moles for each substance (NaOH and H2SO4) are equivalent the stoichiometry dictates the formation of sodium hydrogen carbonate (an acid salt), and when sodium hydroxide is present in twice the molar quantity sodium sulfate (a basic compound) is formed.
Clearly then, since (I) yielded an acidic product, its 1:1 ratio is self-evident (since the volumes were carefully designed to preserve this ratio).
Data Processing and Presentation:
Responding to Processing Questions
- Defining mono/di/polyprotic acids:
- monoprotic;
An acid capable of donating only one proton (H+ ion) during a reaction. Examples of monoprotic acids include:
- hydrochloric acid (HCl);
-
nitric acid (HNO3);
-
ethanoic acid (CH3COOH)
These acids only react in one stage in donating their proton – for example:
HCl(aq) + H2O(l) → H3O+(aq) + Cl-(aq)
A B A B
In this example the acid (HCl) donates a single proton to the water molecule, forming hydronium and anionic chlorine.
- diprotic;
An acid capable of donating two protons (H+ ions) during a reaction. Examples of diprotic acids include:
-
sulfuric acid (H2SO4);
-
carbonic acid (H2CO3);
-
sulfurous acid (H2SO3)
These acids can react in two stages, donating both their protons – for example:
H2SO4(aq) + H2O(l) → HSO4-(aq) + H3O+(aq)
A B B A
HSO4-(aq) + H2O(l) → SO42-(aq) + H3O+(aq)
A B B A
In this example the acid (H2SO4) donates two protons to the water molecules, forming hydronium and anionic sulfate.
- polyprotic;
this is a general term for an acid capable of donating two or more protons (H+ ions) during a reaction. Examples of polyprotic acids range in protenicity:
-
diprotic – sulfuric acid (H2SO4);
-
triprotic – tri hydrogen phosphoric acid (H3PO4);
Polyprotic acids can react in quite a number of stages, donating however many feasible protons they have – for example:
H3PO4(aq) + H2O(l) → H2PO4-(aq) + H3O+(aq)
A B B A
H2PO4-(aq) + H2O(l) → HPO42-(aq) + H3O+(aq)
A B B A
HPO42-(aq) + H2O(l) → PO43-(aq) + H3O+(aq)
A B B A
In this example the acid (H3PO4) donates three protons to the water molecules, forming hydronium and anionic phosphate.
- When evaporating the prepared solutions it is essential that the acid be volatile while the salt is non-volatile so that the fluid content may evaporate readily (due to its volatility) without removing the salt in the process. This will ensure that there will be a solid salt remainder, which can then be used to produce a stock solution for testing of the nature of the salt.
- An acid salt is produced through the reaction in equation (I):
NaOH(aq) + H2SO4(aq) → NaHSO4(aq) + H2O(l)
NaHSO4 – the acid salt component – is referred to as such due (obviously) to its acidic properties. These properties were demonstrated through litmus testing (turned litmus red) and through reaction of a carbonate with a solution of this salt (produced by reaction (I)). For these reason sodium hydrogen sulfate is commonly referred to as an acid salt. The acidic properties themselves are due to the presence of hydrogen in the molecule, which can be donated during a reaction; the salt is therefore – according to the Brønsted-Lowry definition – an acid.
-
Not all acids react to form salts that produce acidic solutions, a fact specifically illustrated by the tests performed on the corresponding acid salt of diprotic carbonic acid, NaHCO3(aq). When a solution of sodium hydrogen carbonate was tested with litmus, a deep purple colour was observed – the solution was extremely basic. This is a clear indication that not all acid salts give acidic solutions.
Conclusion:
By carefully calculating the correct volumes of diprotic sulfuric acid required to satisfy each of the stoichiometric ratios necessary for the production of two different salts (NaHSO4 and Na2SO4), we were able to successfully dictate which salt evolved. The equations governing each reaction were found to be as follows:
(I) NaOH(aq) + H2SO4(aq) → NaHSO4(aq) + H2O(l)
-
2NaOH(aq) + H2SO4(aq) → Na2SO4(aq) + 2H2O(l)
From this it is clear that when the number of moles of both reactants are equivalent, the stoichiometry dictates the formation of sodium hydrogen carbonate (an acid salt), and when sodium hydroxide is present in twice the molar quantity of sulfuric acid, sodium sulfate (a basic compound) is formed. Given the molarity of each reactant, it was a simple matter to calculate the volumes of sulfuric acid required to react with the sodium hydroxide to produce each of the products (18.70 cm3 ± 0.10 cm3 for equation (I) and for equation (II) 9.349 cm3 ± 0.10 cm3). These volumes were calculated based on an assessment of the correct molar quantities of acid when n(NaOH) = 2.002x10-2 mol; being n(I)(H2SO4) = 2.002x10-2 mol and n(II)(H2SO4) = 1.001x10-2 mol. All this was done to preserve the molar ratios dictated by each equation.
Evaluation:
The experiment was entirely successful in confirming the predictions of the products for each reaction based on the equation stoichiometry, but the quantitative aspect of this experiment is minimal.
Results of others who perform this experiment may vary (minutely) for several reasons, most of them the result of gross error. The most obvious source of potential error is in the initial calculation or application of the determined volumes required for the production of each salt. Gross inaccuracies in measurement or calculation of the volumes could result in a violation of the stoichiometric ratios, and the formation of a salt other than that intended. Several other obvious errors may have had a significant impact on results, and caused a variation in results between two standardisations of the same acid:
Sources of Error:
Gross/Random – no method of controlling this factor
-
Inaccuracies during the standardisation of NaOH and H2SO4 could have resulted in flawed estimation of molar quantities, in turn affecting the validity of the stoichiometry preservation
Random/Systematic – ensure all measurements are taken at eye level
- Parallax error causes inaccuracies in reading measurements, though it is largely accounted for by individual calibration errors for most apparatus (though not for apparatus with a single calibration line)
Random – no method of controlling this factor
- Burette, pipette and volumetric flask are all calibrated for specific temperatures and pressures. While the experiment was carried out in approximately the correct order of magnitude for all of these, it would not have been carried out precisely to these calibration requirements
Gross/Random – both solutions agitated thoroughly
- Failing to produce a homogenous solution of each solution prior to evaporation could result in tainted concentrations for each stock solution, which may have affected the reaction with sodium hydrogen carbonate or the pH reading using the litmus
Gross – take care to remove funnel from burette
- Failing to remove the funnel from the burette could result in drops of acid falling into the burette, altering measured volumes
Random/Gross – care taken to right clamp vertically
- Failing to clamp burette vertically could result in mis-reading volumes and/or parallax
Random – rinse sides of conical flask down with de-ionized water
- Failing to react all the acid with all the base could result in obvious errors in finding the accurate neutralisation point
Systematic
-
There is a ± 0.05 cm3 error associated with the burette and a ± 0.01 cm3 error with the pipette. These uncertainties have been included in calculations above without working.
*NOTE: Since relatively small masses are involved, even a minute error of ± 0.01 cm3 becomes significant
It is likely that a culmination of these errors – with an emphasis on the possibility of gross error in a miscalculation/measurement – could cause a variation to these results. Regardless, the results confirmed theoretical predictions quite confidently; the ratio’s determined and measured yielded theoretically predictable products.
Improvements:
Several aspects of this experiment could be developed to yield more definitive results. The influences of random and systematic errors in this experiment were essentially insignificant. Key improvements to this experiment would, however, include:
- Use equipment consistently calibrated for a specific temperature and pressure (i.e. SLC) and maintain conditions at their calibration specifications
- Agitating solutions directly prior to evaporation and preparation of stock solutions
- Taking greater care to maintain burette vertically
- Ensure that all apparatus is rinsed with distilled water when appropriate to ensure that all the acid and base react
- Collation of multiple results, or repetition of experiment to ensure a more accurate average volume for each titre
- Pooling and compiling results to achieve a more precise average
- Allow more time for evaporation and use a larger quantity the acids to produce a more concentrated stock solution for testing
-
Test gas evolved in reaction between NaHSO4 and Na2CO3, ensuring that it is CO2 to experimentally confirm theoretical reaction.
If these improvements were implemented, the results would prove more definitive – there would be further proof of the theoretical expectations.
Chris Bolton