NaOH + CH3COOH → CH3COONa + H2O
mol
= 0.0020 mol
Mole ratios
mol
n(CH3COOH)= 0.0020 mol
Convert n(CH3COOH) to c(CH3COOH)
Convert diluted c(CH3COOH) to undiluted c(CH3COOH)
Molar Concentration
*To one significant figure:
Concentration Percentage
mL
m/v
*To one significant figure: 5% m/v
Discussion:
Many different kinds of vinegar can be bought to be used in the kitchen for cooking and preservation and for other purposes as well. The chemical compound the produces its sour taste and overpowering smell is ethanoic (acetic) acid.
In this investigation, titration was used to measure the unknown amount of the titrant (CH3COOH) by adding an amount of the titrating solution with a known concentration (NaOH). The endpoint of the reaction was then monitored when the sodium hydroxide reacted with the acetic acid. The values obtained from the titration and the chemical equation was used to determine the concentration of ethanoic acid in vinegar through stoichiometric calculations.
When some of the water molecules lose a hydrogen atom and become hydroxyl ions (OH−), the lost hydrogen ions combine with water molecules to form hydronium ions (H3O+). In pure water, there are an equal number of hydrogen ions and hydroxyl ions hence, the solution is neither acidic nor basic. An acid, like acetic acid, is a substance that donates hydrogen ions, hence, when acetic acid is dissolved in water, the balance between hydrogen ions and hydroxyl ions is shifted. Now there are more hydrogen ions than hydroxyl ions in the solution, therefore this solution is acidic. A base like sodium hydroxide is a substance that accepts hydrogen ions so that when a base is dissolved in water, the balance between hydrogen ions and hydroxyl ions shifts the opposite way. Since sodium hydroxide is a strong base, it dissociates almost completely in water. So for every NaOH molecule that was added to the solution, it is expected to produce a hydroxyl ion. Since phenolphthalein is colourless when the solution is acidic or neutral, when the solution becomes slightly basic, phenolphthalein turns pinkish, and then purple as the solution becomes more basic.
In any practical investigation, there will be factors which may influence the final results. These errors include random, systematic and even human mistakes.
Random errors would have affected the final results. Although the measuring instruments have a calibrated error, the chances of a person reading the result in the same way is quite small, hence the volume might have been misread on the measuring instruments. This might have been a parallax error, when I might have read the volume looking at an angle, or error in counting graduation marks. When volumes are misread, the accuracy in volume is affected. In general- if volume of any of the titrant or the standard solution is affected, and this reading is used in the calculation, the number of mols is calculated in a volume different to the actual one, thus affecting the accuracy and reliability of results.When reading the volume on the burette scale it is not uncommon for someone to read both upper and lower values in different lighting conditions, thus influencing the final recorded results. If the volume recorded was higher than the actual value then it is thought that more sodium hydroxide was used for the same amount of acetic acid, and then the concentration of the acetic acid is much stronger than it actually is. The same thing applies if the volume recorded was lower than the actual value, and then the concentration of the acetic acid is thought weaker than it actually is. This will give a false representation of what the actual concentration of a substance is.
We also pipetted straight from a volumetric flask instead of pouring a volume of the solution into a clean, dry beaker and pipetting from that. This may have resulted in cross-contamination which will affect the titration and its endpoint. Another random error might have been a misjudgment of the colour change not only was the colour change very delicate and slow, but different people have different sensitivity to colours. Therefore the actual endpoint of the titration might have occurred before what we thought; hence more NaOH was added, thus increasing the concentration of the acetic acid than what it actually is. This will in turn give us inaccurate results in the end.
The titration might have taken place at a wrong temperature as most equipment and methods have an ideal temperature where the readings are more accurate. Since the titration equipment is being calibrated to be only accurate at a certain temperature. The liquid will expand slightly if it is hotter than this, causing the reading to be incorrect as the volume recorded is larger than what it actually is. This will affect the rest of the stoichiometric calculations regarding the moles, molar concentration and concentration percentage of acetic acid. This may be one of the reasons why the molarity of the acetic acid we obtained (0.79M), was higher than the actual molarity which was 0.75M.Since pH indicators changes with temperature, its effect on the indicator colour also changes which will affect the endpoint of the titration. I might have not filled the burette properly; therefore I might have been oblivious to the air bubbles that were trapped around the stopcock of the burette. This would have resulted in inaccurate volume readings as the air bubbles would have flowed with the titrant; therefore we would have no idea of the real volume of solution used. This makes it look like more titrant was used than actual, which makes molarity of unknown appear higher. This may be another reason why the molarity we got for acetic acid (0.79M) was higher than its actual molarity which was 0.75M.
Some acetic acid solution may have been left in the pipette when transferring into the conical flask. This causes systematic error in the amount of moles of acetic acid present in the conical flask. The amount of moles of acetic acid in the conical flask is consistently lesser than what is expected; hence, the amount of moles of sodium hydroxide is lesser too. This affects the final stoichiometric calculation as incorrect values will be recorded, i.e., lesser moles of the solution(s) than it actually is, which in turn influences the rest of the calculations such as the molar concentration and concentration percentage which will be lower than it actually is. Therefore, it will affect accuracy of the results obtained.
Also affecting the results were the systematic errors, which were made by the calibration of measuring equipment that were used, such as; the burette, pipette and volumetric flask. Obviously, human mistakes are a factor which may have affected the results of the titration. Although they did not happen during this titration, they include; not rinsing the conical flasks with demineralised water (to ensure that everything is collected), leaving the funnel on top of the burette, recording a wrong result, or a miscalculation during the stoichiometry.
After the titration and stoichiometric calculations were carried out, it was found that this particular commercial vinegar only had the ethanoic concentration of 5% m/v, despite their published standards of 6%. This is considered as “false claims” under the Competition and Consumer Act 2010, it breaches the law and the Australian Competition & Consumer Commission (ACCC) can take legal action. Businesses and manufacturers are prohibited by ‘The Competition and Consumer Act’ from deceiving and misleading consumers by making false claims. Failure to comply with these requirements will result in fines and injunctions. This act clearly states that a business is not to falsely represent the standard, quality and grade of goods. In this case, the manufacturer has given a misleading impression about the concentration of the product which may result in fines up to $1.1 million for businesses and $220,000 for individuals.
The original method was effective and valid, as it achieved the aim of determining the concentration of ethanoic acid. However, it could have been improved by including a detailed list of materials, more steps rearranged in a chronological order (e.g. rinsing all the equipment required in the beginning) so that the chance of human mistakes made is reduced. Other improvements to the actual titrations practical would be to perform the titration in a temperature controlled room to ensure that the results were more accurate. A working bench friendlier to short people would improve the precision of the results because it would decrease the random errors like misreading the instruments. However, this can also be done by adjusting the burette clamp etc.
In order to get a uniform colour, it is important to swirl the conical flask vigorously enough to ensure that all parts of the solution react to present a uniform colour; otherwise it is difficult to detect the end point of the titration. To eliminate air bubbles around the stopcock of the burette, tap it; otherwise perform that step again ensuring there are no air bubbles present.
The accuracy of glassware is limited, and not all volumetric glassware is created equally. Hence, another improvement to increase the chance of the class finding similar results would be to ensure that everyone uses similar equipment – in particular, the burette. By ensuring that everyone uses a strategic approach to the swirling and inversion to give a homogenous solution, a consistent result can be produced across the class. Some procedural changes would include arranging the titration to have bigger titres (reducing errors), and using larger aliquots (reducing errors).
The titration can also be repeated more times to ensure that the endpoint is more accurate and consistent. By attaching a white paper slide with a broad line to burette so that the line is vertical, as well as taking care to avoid parallax error, precise readings of the volume can be made.
Conclusion:
The concentration of ethanoic acid in commercial vinegar was determined by titrating vinegar with sodium hydroxide. Though the results may be subjected to human mistake, systematic errors and random errors, the findings of the titration were that the concentration of ethanoic acid in commercial vinegar was 5% m/v.