Preparation and Analysis of Benzoic Acid


The purpose of this experiment is to prepare and analyze benzoic acid. It is also meant to demonstrate some general preparation methods, such as using a reflux apparatus, vacuum filter, and a hot gravity filter, plus determining melting point and titrating.


Benzoic acid was first synthesized in the 16th century [W2009]. It is an organic acid that naturally forms monoclinic leaflets and is found as a colourless solid. The molar mass of benzoic acid is 122.12 g and its melting point is 122°C. In water at 1°C, benzoic acid has a solubility of 1.7 g/liter, but at 95°C it has a solubility of 68 g/liter [ND2008]. Due to its natural ability to impede the growth of yeast, mold, and some bacteria, benzoic acid has been used to treat fungal skin diseases such as athlete's foot and as a food preservative.

Oxidation Theory and Calculations

In general, a primary alcohol can be oxidized to lend either aldehydes or carboxylic acids [JC2006]. The reaction goes through two processes. Note that R is shorthand for an arbitrary hydrocarbon counterpart. For example, RCH2OH would be ethanol if R was CH3.

  1. RCH2OH + O- ⇋ RCOH + H2O

  2. RCOH + O- ⇋ RCOOH

In the first stage, an available oxygen atom binds with two of the hydrogen atoms at the end of the alcohol (RCH2OH) to produce an aldehyde (RCOH). The oxygen within the alcohol then assumes a covalent configuration with the carbon and the result is a molecule known as an aldehyde. In the second stage, the hydrogen atom at the end of the aldehyde incorporates another available oxygen to produce a carboxylic acid (RCOOH). This reaction has several uses. For example, potassium dichromate (an oxidant) will produce ethanal or ethanoic acid when mixed with ethanol. This particular reaction used to be used as part of the ‘breathalyser’ procedure to test if drivers had consumed too much alcohol. In our particular experiment, we are oxidizing benzyl alcohol (C6H5CH2OH) with permanganate (MnO4-). The goal is to produce benzoic acid (C6H5COOH). Thus, in our case the R will be C6H5, a phenyl ring. The permanganate is created by dissolving potassium permanganate in sulphuric acid (H2SO4) (the reason for using a highly acidic solvent will be explained later). To understand how the reaction works, we must first understand how the oxidation process works. Oxidation processes are described using redox equation. Redox equations describe the requirements for changing the oxidation level of a chemical. First, consider the permanganate. It reacts in two steps: first to MnO2 then Mn2+. Their redox equation [VB1986] is written as:

  1. MnO4- + 4 H+ + 3 e- ⇋ MnO2 + 2 H2O

  2. MnO2 + 4 H+ + 2 e- ⇋ Mn2+ + 2 H2O

And for the two states of the alcohol to acid reaction:

  1. RCH2OH ⇋ RCOH + 2 H+ + 2 e-

  2. RCOH + H2O ⇋ RCOOH + 2 H+ + 2 e-

These equations are not complete since they have spare electrons on one side (because of this they are often called half-equations). Note, however, how equations 3 and 4, and 5 and 6 have spare electrons on the opposite sides. This is because 3 and 4 are reduction while 5 and 6 are oxidation. These equations can be combined to produce the full equations for the oxidation. For example, combine 3, 4, and 5 to determine the alcohol to aldehyde step:

  1. 3 RCH2OH + 2 MnO4- + 2 H+ ⇋ 3 RCOH + 2 MnO2 + 4 H2O

  2. RCH2OH + MnO2 + 2 H+ ⇋ RCOH + Mn2+ + 2 H2O

  3. 5 RCH2OH + 2 MnO4- + 6 H+ ⇋ 5 RCOH + 2 Mn2+ + 8 H2O

Equation 7 is the combination of 3 and 5, the first step of manganese's oxidation; 7 is from 4 and 5, the second step of the oxidation; and 9 is the combination of 7 and 8. The same can be done to determine the aldehyde to alcohol step:

  1. 3 RCOH + 2 MnO4- + 2 H+ ⇋ 3 RCOOH + 2 MnO2 + H2O

  2. RCOH + MnO2 + 2 H+ ⇋ RCOOH + Mn2+ + H2O

  3. 5 RCOH + 2 MnO4- + 6 H+ ⇋ 5 RCOOH + 2 Mn2+ + 3 H2O

Finally, equations 9 and 12 can be combined to produce the equation for the overall reation (assuming it goes to completion):

  1. 5 RCH2OH + 4 MnO4- + 12 H+ ⇋ 5 RCOOH + 4 Mn2+ + 11 H2O

Replacing R with C6H5 gives the reaction for the benzyl chemicals.

  1. 5 C6H5CH2OH + 4 MnO4- + 12 H+ ⇋ 5 C6H5COOH + 4 Mn2+ + 11 H2O

Experimental Procedure

Benzyl alcohol (4 ml) and KMnO4 (about 9 g) are placed into 150 ml of a highly acidic solution (H2SO4). Once in the solution, the KMnO4 will dissolve into K+ and MnO4-. There are about twice as many moles of KMnO4 to encourage the alcohol to react as far as possible. As the amount of benzyl alcohol approaches or exceeds the 5/4 of the amount of KMnO4, there will be significant amounts of benzylhyde left over. Also, the solution needs to be acidic so that it will supply the free hydrogen ions that appear on the left-hand side of equation 14.

In order to push the equilibrium to the right as quickly as possible, the solution is heated in a ball flask and stirred over a magnetic stirring plate. Unfortunately, as the solution heats up, the solution will start to evaporate. Benzaldehyde has the lowest boiling point (178.1°C) and will evaporate first. Benzaldehyde has a strong smell of vanilla, so it should be relatively noticeable when it starts to evaporate. The benzaldehyde is an intermediate chemical in the reaction and if it evaporates, the final yield will be reduced. To prevent this, the solution is placed in a reflux apparatus while the heating is taking place.

Apart from simply increasing the activity of the molecules, increasing the reaction. It also helps the oxidation process since KMnO4 decomposition is sensitive to heat.

The reaction takes a long time to complete, so the procedure is stopped part way through. As a result, the solution will still contain MnO4- and MnO2. These are reduced by adding sodium bisulphite (the MnO4- and MnO2 oxidize the bisulphite reacts with these to produce manganese, water, and sulphate ions, see equations 15 and 16). The result of this should be noticeable as the brownish color of manganese dioxide will fade. Not all of the MnO4- and MnO2 will be removed at this point, so some color will persist.

  1. 2 MnO4- + 4 HSO3- + 2 H+ ⇋ 2 Mn2+ + 5 SO42- + 3 H2O

  2. MnO2 + HSO3- + H+ ⇋ Mn2+ + SO42- + H2O

The benzoic acid is slowly cooled. It freezes at 122.4°C, creating crystals. The crystals are removed by vacuum filtration, and tests are run on them.

After this, the benzoic acid crystals are dissolved in water for further purification. Since MnO2 and Mn2+ are not soluble in water, they are easily removed via gravity filtration. Then the acid is split into two beakers: one hot, and one cool. These are then cooled so that the acid will recrystallize. Note that unless the solution is cooled ahead of time, the part of the solution that is place in the cool beaker will cool down very rapidly. This hinders the crystallization process because it produces large amounts of very small crystals.

The crystals are separated from the solution again and the final tests are made.

Testing Procedures

The effectiveness of the procedure is measured by 3 different tests. First, the crude acid crystals are weighed and compared against the expected yield to determine the % conversion. Second, some of the acid crystals are melted to test the accuracy of the melting point (ideally 122°C). Third, some of the purified benzoic acid is redissolved in water and titrated against sodium hydroxide (NaOH). The moles of NaOH needed for the titration equals the moles of benzoic acid, which is then converted to grams, and can be compared against the actual mass of benzoic acid produced to determine the purity of the product.


During reflux in the first trial the solution changed colour several times. First, it changed from dark purple to orange, then to light yellow, then to dark yellow, then burnt orange, and finally a brownish orange. 0.012 g of dullish light yellow crystals was retrieved, which was very unfortunate because 0.1 g was needed for determining the melting point and ~0.15 g to determine the purity through three titrations. Much of the loss of crystals was during reflux when huge amounts evaporated because the solution was too hot. During the second trial the experiment was done more carefully and only a few vapours were released when refluxation first started. Ten minutes into the procedure, the solution turned a pinkish colour, then light yellow, then darker-medium yellow at 15 minutes, and at 30 minutes the solution was bright yellow with a thin orange layer floating on the top. After vacuum filtration there was 1.509 g of light brown crystals. Two samples of the filtered crystals were put into beakers, one pre-warmed and one cooled in an ice bath. The latter crystals formed faster, cleaner, and smaller. Whereas the others, left at room temperature, formed slower, but were bigger. When the crystals on filter paper were examined, the cold crystals were off-white and very fine with a paper-like texture, and the warm crystals were coarse, off-white/pearly, and very bumpy. The average melting point for the first trial was 94 degrees Celsius, and 115 degrees Celsius for the second trial. The normal melting point of benzoic acid is 122.4 degrees. The purity of the crystals was measured using titration with sodium hydroxide as the titrant.

Calculations for Percent Yield

KMnO4 (Molar mass = 158.03 g/mol)
C6H5CH2OH (Molar mass: 108.14 g/mol ; Density: 1.04 g/ml)
(1.04 g/ml C6H5CH2OH * 4.00 ml) / 108.14 g/mol = 0.0385 mol C6H5CH2OH
Experiment 1: 9.00g KMnO4 / 158.03 g/mol = 0.570 mol KMnO4
Experiment 2: 9.02g KMnO4 / 158.03 g/mol = 0.571 mol KMnO4
C6H5CH2OH is limiting reagent so 0.0385 moles of C6H5CH2OH results in 0.0385 moles of C6H5COOH C6H5COOH (Molar mass: 122.12 g/mol)
0.0385 moles* 122.12 g/mol = 4.70 g C6H5COOH
Theoretical yield is 4.70 g C6H5COOH.
Grams of actual yield / grams of theoretical yield = Percent Yield
Experiment 1: 0.12 g / 4.70 g * 100% = 2.55%
Experiment 2: 1.509g / 4.70g * 100% = 32.1%

Titration Calculations

Titration #1

6.42 ml NaOH x 1L/1000 mL x 0.1357mol/1L = 8.71 x 10-4 moles
8.71 x 10-4 moles NaOH = 8.71 x 10-4 moles benzoic acid
8.71 x 10-4 moles benzoic acid x 122.12 g/mol = 0.106 g of benzoic acid
% purity = 0.106g/0.216g x 100% = 49.1% purity

Titration #2

27.9mL NaOH x 1L/1000mL x 0.1357mol/1L = 0.00379 mol NaOH
0.00379 mol NaOH = 0.00379 mol benzoic acid
0.00379 mol benzoic acid x 122.12 g/1mol = 0.463g of benzoic acid
% purity = 0.463g/0.173g = 268% purity


In the first trial run of this experiment, 0.12 g of crystals was obtained. These crude benzoic acid crystals were a dull yellow colour and they didn’t look very pure. The low yield, 2.57%, was a result of several errors. First of all, the 150 mL of H2SO4 was measured with a 400 mL beaker which is not very precise. Also, when the original solution was refluxed a large amount of fumes escaped through the top of the reflux apparatus without condensing. This was because the heat was turned on too high, and so some of the crystals evaporated, thereby reducing the yield. Also once the solution was finished refluxing, the solution was a brownish-orange colour and therefore the crystals were a dullish yellow colour versus being pure and white. Because the crystals were not very pure, the melting point was only found to be 94°C, whereas benzoic acid is supposed to have a melting point of 122°C. After recrystallizing and purifying the crude benzoic acid crystals, the cold crystal solution sat in the fridge and the room temperature crystal solution sat in the lab for a week. Then, because no crystals had formed over the week, the experiment was discontinued and therefore neither the titrations nor the percent purity calculations were completed. The lack of crystals was due to the fact that there was a low percent yield of 2.57%, and the hot gravity filtration apparatus was not completely hot when the crystals were filtered through.

The second trial run was more successful because 1.509 g of crude benzoic acid crystals was collected. The crude crystals were light brown with some clear spots. There was less error in this experiment because the 150 mL of H2SO4 was measured with a 10 mL graduated cylinder and care was taken to not release vapors during refluxation. Although some vapors did escape, it was considerably less than in the first trial run. The crude product resulted in a 32.3% yield. We obtained 1.509 g of crystals compared to the theoretical amount of 4.67 g. The melting point of these crude crystals was 115°C which is very close to the actual melting point of benzoic acid: 122°C. Once the crude crystals were recrystallized, purified, and separated into two beakers of different temperature, it was discovered that the cold crystals formed faster, yet were smaller, than the ones that formed at room temperature. The 0.216 g of cold crystals were very fine and off-white compared to the 0.173 g of lukewarm crystals which were coarse and had an off-white, pearly appearance. In the experiment, the melting point of the pure crystals was not determined, so it cannot be compared with the melting point of the crude crystals. Because not a lot of pure crystals were obtained, the titration was only completed twice although it should have been done three times. The first titration of the room temperature crystals determined that the crystals had 49.1% purity. The second titration, which was of the cold pure crystals, determined a purity of 268%. Obviously this percentage is inaccurate and that is due to a faulty titration. Phenolphthalein was not added to the solution and therefore the titration was off. When the phenolphthalein was added, the solution was already bright pink and the titration was well beyond the equivalence point.


In this experiment benzoic acid was synthesized through the oxidation of benzyl alcohol by acidified potassium permanganate.

The alteration of our initial solution into crude crystals was observed through using a reflux apparatus, collecting and cleaning the crystals with a vacuum filter, recrystallizing again through reflux, and finally purifying with a hot gravity filter. The melting point was examined, the percent yield calculated and the purification of the resulting crystals was discovered through titration with phenolphthalein. Although the first trial was not successful, the second trial resulted in a percent yield of crude crystals of 32.1%, a melting point of 115°C, and a percent purity of 49.1%.


[W2009] 2009. Benzoic Acid. Wikipedia, The Free Encyclopedia. (retrieved 2009-3-18).
[JC2006] Jim Clark. 2006. Oxidation of Alcohols. (retrieved 2009-3-18).
[ND2008] Nigel Dance et al.. 2008. Chemistry 114 Lab Manual. pp. 91–97.
[VB1986] Arthur Vogel and John Bassett. 1986. Textbook of Inorganic Analysis. p. 42. (retrieved 2009-3-18).