Titration | Turkish Chemistry
May 11

Vitamin C Determination by Iodine Titration
(ascorbic ) is an antioxidant that is essential for human nutrition. deficiency can lead to a disease called scurvy, which is characterized by abnormalities in the bones and teeth. Many fruits and vegetables contain , but cooking destroys the , so raw citrus fruits and their juices are the main source of ascorbic for most people.

One way to determine the amount of in food is to use a redox titration. The redox reaction is better than an - titration since there are additional acids in a juice, but few of them interfere with the oxidation of ascorbic by iodine.

Iodine is relatively insoluble, but this can be improved by complexing the iodine with iodide to form triiodide:

I2 + I- <–> I3-

Triiodide oxidizes to form dehydroascorbic :

6H8O6 + I3- + H2O –> 6H6O6 + 3I- + 2H+

As long as is present in the solution, the triiodide is converted to the iodide ion very quickly. Howevever, when the all the is oxidized, iodine and triiodide will be present, which react with starch to form a blue-black complex. The blue-black color is the of the titration.

This titration procedure is appropriate for testing the amount of in tablets, juices, and fresh, frozen, or packaged fruits and vegetables. The titration can be performed using just iodine solution and not iodate, but the iodate solution is more stable and gives a more accurate result.

Purpose

The goal of this laboratory exercise is to determine the amount of in samples, such as fruit juice.

Procedure

The first step is to prepare the solutions. I’ve listed examples of quantities, but they aren’t important. What matters is that you know the of the solutions and the volumes that you use.

Preparing Solutions

1% Starch Indicator Solution

 

  1. Add 0.50 g soluble starch to 50 near-boiling distilled .
  2. Mix well and allow to cool before use. (doesn’t have to be 1%; 0.5% is fine)

Iodine Solution

 

  1. Dissolve 5.00 g potassium iodide (KI) and 0.268 g potassium iodate (KIO3) in 200 ml of distilled .
  2. Add 30 ml of 3 M sulfuric .
  3. Pour this solution into a 500 ml graduted cylinder and dilute it to a final volume of 500 ml with distilled .
  4. Mix the solution.
  5. Transfer the solution to a 600 ml beaker. Label the beaker as your iodine solution.

Standard Solution

 

  1. Dissolve 0.250 g (ascorbic ) in 100 ml distilled .
  2. Dilute to 250 ml with distilled in a volumetric flask. Label the flask as your standard solution.

Standardizing Solutions

 

  1. Add 25.00 ml of standard solution to a 125 ml Erlenmeyer flask.
  2. Add 10 drops of 1% starch solution.
  3. Rinse your buret with a small volume of the iodine solution and then fill it. Record the initial volume.
  4. Titrate the solution until the is reached. This will be when you see the first sign of blue color that persists after 20 seconds of swirling the solution.
  5. Record the final volume of iodine solution. The volume that was required is the starting volume minus the final volume.
  6. Repeat the titration at least twice more. The results should agree within 0.1 ml.
You titrate samples exactly the same as you did your standard. Record the initial and final volume of iodine solution required to produce the color change at the .

Titrating Juice Samples

 

  1. Add 25.00 ml of juice sample to a 125 ml Erlenmeyer flask.
  2. Titrate until the is reached. (Add iodine solution until you get a color that persists longer than 20 seconds.)
  3. Repeat the titration until you have at least three measurement that agree to within 0.1 ml.

Titrating Real Lemon

Real Lemon is nice to use because the maker lists , so you can compare your value with the packaged value.

  1. Add 10.00 ml of Real Lemon into a 125 ml Erlenmeyer flask.
  2. Titrate until you have at least three measurements that agree within 0.1 ml of iodine solution.

Other Samples

 

  • Tablet – Dissolve the tablet in ~100 ml distilled . Add distilled to make 200 ml of solution in a volumetric flask. 
  • Fresh Fruit Juice – Strain the juice through a coffee filter or cheese cloth to remove pulp and seeds, since they could get stuck in the glassware. 
  • Packaged Fruit Juice – This also may require straining. 
  • Fruits & Vegetables – Blend a 100 g sample with ~50 ml of distilled . Strain the mixture. Wash the filter with a few milliliters of distilled . Add distilled to make a final solution of 100 ml in a volumetric flask.

Titrate these samples in the same way as the juice sample described above.

Titration Calculations

 

  1. Calculate the ml of titrant used for each flask. Take the measurements you obtained and average them.average volume = total volume / number of trials

     

  2. Determine how much titrant was required for your standard.If you needed an average of 10.00 ml of iodine solution to react 0.250 grams of , then you can determine how much was in a sample. For example, if you needed 6.00 ml to react your juice (a made-up value – don’t worry if you get something totally different):

    10.00 ml iodine solution / 0.250 g Vit = 6.00 ml iodine solution / X ml Vit

    40.00 X = 6.00

    X = 0.15 g Vit in that sample

     

  3. Keep in mind the volume of your sample, so you can make other calculations, such as grams per liter. For a 25 ml juice sample, for example:0.15 g / 25 ml = 0.15 g / 0.025 L = 6.00 g/L of in that sample

Oca 31

Acid-Base Titration
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An acid-base titration is the determination of the concentration of an acid or base by exactly neutralizing the acid/base with an acid or base of known concentration. This allows for quantitative analysis of the concentration of an unknown acid or base solution. It makes use of the neutralization reaction that occurs between and the knowledge of how will react if their formulas are known.

Acid-Base titrations can also be used to find percent purity of chemicals.


Before starting the titration a suitable pH indicator must be chosen. The endpoint of the reaction, the point at which all the reactants have reacted, will have a pH dependent on the relative strengths of the acid and base used. The pH of the endpoint can be estimated using the following rules:

A strong acid will react with a strong base to form a neutral (pH=7) solution.
A strong acid will react with a weak base to form an acidic (pH<7) solution.
A weak acid will react with a strong base to form a basic (pH>7) solution.
When a weak acid reacts with a weak base, the endpoint solution will be basic if the base is stronger and acidic if the acid is stronger. If both are of equal strength, then the endpoint pH will be neutral. However weak acids are not often titrated against weak bases because the color change shown with the indicator is often quick, and therefore very difficult for the observer to see the change of color.

A suitable indicator should be chosen, preferably one that will experience a change in color close to the end point of the reaction.

First, the should be rinsed with the standard solution, the with the unknown solution, and the with distilled water.

Secondly, a known volume of the unknown concentration solution should be taken with the and placed into the , along with a small amount of the indicator chosen. The should always be filled to the top of its scale with the known solution for ease of reading.

The known solution should then be allowed out of the , into the . At this stage we want a rough estimate of the amount of this solution it took to neutralize the unknown solution. Let the solution out of the until the indicator changes color and then record the value on the buret. This is the first titre and should be discluded from any calculations.

Perform three more titrations, this time more accurately, taking into account we know roughly where the end point will occur. Take note of each of the readings on the at the end point, and average these at the end. Endpoint is reached when the indicator just changes color permanently. This is best achieved by washing a hanging drop from the tip of the into the flask right at the end of the titration to achieve a drop that is smaller in volume than what can usually be achieved by just dripping titre off the .

Acid-base titration is performed with a phenolphthalein indicator, when it is a weak acid – strong base titration, a bromthymol blue indicator in strong acid- strong base reactions, and a methyl orange indicator for strong acid – weak base reactions. If the base is off the scale, i.e. a pH of >13.5, and the acid has a pH >5.5, then an Alizarin yellow indicator may be used.On the other hand, if the acid is off the scale, i.e. a pH of <0.5, and the base has a pH <8.5, then an Thymol Blue indicator may be used.

When titrating a weak acid with a strong base, pH can be calculated by the following formula: [1]

pH = pK_s + log( \frac{[HO^-]_{added}}{[HA]_{total}-[HO^-]_{added}} )

where:

  • pKs is the acid dissociation constant of the weak acid.
  • [HO-]added is the concentration of added strong base in the final solution (not in original standard solution)
  • [HA]total is the summed concentration of both the weak acid and its conjugate base in the final solution.

Thus, at an addition of strong base that is half the amount of weak acid in the solution ([HO-]added = 0.5[HA]total), pH becomes equal to pKs.

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