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Titrations

Titrations

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If you've ever needed to figure out how much alkali you need to neutralize acid, you've probably done a titration. But have you ever heard of redox titrations? They help us figure out how much oxidizing agent we need to react with a reducing agent. What's cool about redox titrations with transition metals is that they have different oxidation states, which means they change colors as they react. So, you don't even need an indicator to tell when the reaction is done! Let's take potassium manganate(VII) as an example to see how it works.

In this article, we'll explore what titration is and how it works. We'll walk you through the redox titration of manganate(VII) with iron and the redox titration of manganate(VII) with ethanedioate ions. Plus, we'll give you tips on how to calculate the results for these titrations. So, get ready to dive into the exciting world of titrations! And if you're looking to impress your chemistry teacher with your knowledge of redox titrations, this article is for you!

Titration meaning

Titration is a useful method for finding out the concentration of an unknown chemical by using a substance with a known concentration. The process involves slowly adding a standard solution of the titrant from a burette to the analyte in a conical flask. To know when the reaction has reached its endpoint, we can use a colour indicator.

If you want to understand how titration works, performing one is a great way to do so. Now, let's take a closer look at the steps involved in a titration. The chemical with the known concentration is called the titrant, while the unknown substance is the analyte. The solution with the exact concentration that we use as the titrant is called a standard solution.

Titration method

The method of performing a redox titration is similar to the method for acid-base titrations. You may read about it in pH Curves and Titrations. We will use the redox titrations between iron(II) and ethanedioate ions with manganate(VII) as examples.

Redox titration of iron(II) with manganate(VII)

If you have low iron concentration in your blood, also known as anaemia, your doctor or pharmacist may prescribe you iron tablets. These are usually made with anhydrous iron(II) sulphate because it is both cheap and soluble.

To estimate the amount of iron(II) sulphate in each tablet, you can perform a titration with a standard solution of potassium manganate(VII). However, before you can do this, you will need to dissolve each tablet in diluted sulfuric acid.

In a laboratory, you can carry out the following experiment with iron tablets from the pharmacist:

  1. Weigh out 8 iron tablets and dissolve them in a beaker containing about 100 cm3 of 2 M sulfuric acid. Note that some tablets may have an outer coating that doesn't dissolve, so you may need to filter the solution.
  2. Pour the filtrate into a 250cm3 volumetric flask, making sure to include the washings from the conical flask and filter paper. Use distilled water to fill up to the mark.
  3. Pipette 25 cm3 of the solution into a conical flask and add 25 cm3 of dilute sulfuric acid. Titrate against 0.02 M potassium manganate(VII) until the solution changes from colourless to pale pink.
  4. Repeat the experiment until you get a concordance of ±0.10 cm3.

You will need specific equipment for this experiment, which is shown in the diagram below.

Diagram of titration equipment and method
Diagram of titration equipment and method

In this reaction, Fe2+ gets oxidised to Fe3+ while Mn7+ gets reduced to Mn2+. You write the half equations for the process as follows:

Redox reactions between manganate(VII) and iron(II)
Redox reactions between manganate(VII) and iron(II)

When performing the titration, it is important to use diluted sulfuric acid as potassium permanganate functions best as an oxidizer in acidic conditions. This is because transition metal ions require strongly acidic conditions when transitioning from higher to lower oxidation states. However, not just any acid can be used.

Using weak acids like ethanoic acid will not provide enough H+ ions, and using concentrated sulphuric acid or nitric acid may oxidize the analyte. Therefore, it is crucial to use diluted sulfuric acid to ensure that the reaction proceeds as expected.

There is no need to use an indicator with the titration because potassium manganate(VII) serves as the indicator. As the reaction progresses, the purple manganate(VII) reduces to manganate(II) (a colourless solution). One drop of excess manganate(VII) gives the solution a permanent pale pink colour. Now, let's consider the reaction between manganate ions and ethanedioate ions.

Titrations with ethanedioate ions

The reaction between manganate and ethanedioate ions (C2O42-) is quite fascinating because it is an autocatalytic reaction. Autocatalysts are substances that catalyze their own reaction. In this case, the reaction between manganate and ethanedioate ions speeds up as it progresses, due to the autocatalytic nature of the reaction.

Chemists often use ethanedioic acid, also known as oxalic acid, to standardize or determine the strength of permanganate solutions. Oxalic acid can be found in plants such as spinach and rhubarb, and its salts contain the ethanedioate ion (C2O42-). By titrating against potassium permanganate, we can determine the concentration of free oxalate ions in a solution. This reaction is commonly used to analyze the ethanedioate content of spinach leaves, for example.

The redox reaction between manganate(VII) and ethanedioate ions takes place as follows:

MnO4- is reduced to Mn2+ and C2O42- is oxidized to CO2.

During the reaction, the manganate(VII) ions are reduced to manganese(II) ions, while the ethanedioate ions are oxidized to carbon dioxide. The Mn2+ ions produced in the reaction act as autocatalysts, which means they speed up the reaction. Overall, the reaction between manganate and ethanedioate ions is an important one in analytical chemistry, and is frequently used to determine the concentration of oxalate ions in various substances.

Redox reaction between manganate (VII) and ethanedioate ions
Redox reaction between manganate (VII) and ethanedioate ions

Yes, I have got that. Let's move on to the calculations!

To calculate the concentration of the ethanedioic acid solution, we need to use the following formula:

M1V1 = M2V2

Where:
M1 = concentration of potassium permanganate solution
V1 = volume of potassium permanganate solution used
M2 = concentration of ethanedioic acid solution
V2 = volume of ethanedioic acid solution used

Let's assume that the volume of potassium permanganate solution used is 25.0 cm3, and the concentration of potassium permanganate solution is 0.1000 mol dm-3. The volume of ethanedioic acid solution used is 10.00 cm3.

Using the formula above, we can calculate the concentration of the ethanedioic acid solution as follows:

M2 = (M1V1) / V2
M2 = (0.1000 mol dm-3 x 25.0 cm3) / 10.00 cm3
M2 = 0.250 mol dm-3

Therefore, the concentration of the ethanedioic acid solution is 0.250 mol dm-3.

Note that we need to repeat the experiment several times to obtain a concordance of ∓0.10cm3. This helps to ensure that our results are reliable and accurate.

Titration calculations

After you have completed a titration, you will need to do some calculations to determine the concentration of the analyte. Let us try a few together!

24.55cm3 of 0.020M aqueous potassium manganate(VII) reacted with 25.0cm3 of acidified iron(II) sulfate solution. Find the concentration of Fe2+ ions in the solution.  Step 1: Write out the balanced equation5Fe2+ (aq) + MnO4- (aq) + 8H+ (aq) ➔ 5Fe3+ (aq) + Mn2+ (aq) + 4H2O (l)Step 2: Work out the number of moles of MnO4- ions added to the flask. Moles of MnO4- = We divide by 1000 to convert the volume from cm3 to dm3.Moles of MnO4- = Step 3: The equation tells you that 1 mole of MnO4- reacts with 5 moles of Fe2+. 5Fe2+ (aq) + MnO4- (aq) + 8H+ (aq) ➔ 5Fe3+ (aq) + Mn2+ (aq) + 4H2O (l)Step 4: Multiply moles of MnO4- by 5.0.000419 x 5 = 0.002455 moles of Fe2+ Step 5: Work out the concentration of Fe2+ moles of Fe2+ = Rearrange so that concentration = Concentration = Concentration = 0.0982 mol dm-3

Titration calculations generally follow the same principles as you will see in the next example.

Sammy checked the concentration of a solution of potassium permanganate against an ethanedioic acid solution of concentration 0.04 mol dm-3.

He placed 25cm3 of the ethanedioic acid solution in a flask with excess dilute sulphuric acid. After warming the solution, he carried out a titration. He needed 25cm3 of potassium permanganate solution to reach the endpoint.

Calculate the actual concentration of the permanganate solution.

Step 1: Write the balanced equation for the reaction

2MnO4– + 16H+ + 5C2O42- → 2Mn2+ + 10CO2 + 8H2O

Step 2: Find the number of moles of ethanedioic acid

No. of moles of 5C2O42- =

Step 3: Find the number of moles of potassium manganate(VII).

The balanced equation tells us that we need  as many moles of permanganate ions as ethanedioate ions.

2MnO4- (aq) + 16H+ (aq) + 5C2O42- (aq) ➔ 2Mn2+ (aq) + 10CO2 (aq) + 8H2O (l)

Step 4: Multiply the no. of moles of ethanedioate ions by

0.001 x  in 25cm3 = 0.0004 moles of manganate (VIII)

Step 5: Find the concentration by rearranging the formula

Rearrange the formula so that

Concentration = 0.0004 x

Concentration = 0.016 mol dm-3

It takes practice to get the hang of titration calculations. Try the examples in the exercises section to improve your skills!

Titrations - Key takeaways

To estimate the amount of iron(II) sulphate in each tablet, we can perform a redox titration using a standard solution of potassium manganate(VII). In this reaction, Fe2+ is oxidized to Fe3+ while Mn7+ is reduced to Mn2+.

It's important to note that redox titrations of iron(II) and ethanedioate ions with manganate(VII) must be acidified with dilute sulfuric acid. This is because hydrochloric acid is oxidized to chlorine, and ethanoic acid is too weak and doesn't provide enough H+ ions. Nitric acid and concentrated sulfuric acid also oxidize the analyte.

When performing the redox titration between permanganate and ethanedioic acid, we heat the ethanedioic acid solution to about 60-70ºC to speed up the reaction with potassium permanganate. Adding sulfuric acid to the analyte in titrations with permanganate prevents manganese from oxidizing to manganese dioxide.

The redox reaction between manganate(VII) and ethanedioate ions takes place as follows:

2MnO4– (aq) + 16H+ (aq) + 5C2O42- (aq) ➔ 2Mn2+ (aq) + 10CO2 (aq) + 8H2O (l)

During the reaction, the purple MnO4- ions are reduced to colorless Mn2+ ions. One drop of excess MnO4- ions presents a pale pink color, indicating that the reaction is complete.

Using the same formula as before (M1V1 = M2V2), we can calculate the amount of iron(II) sulphate in each tablet. To do this, we need to know the concentration of the potassium manganate(VII) solution, the volume of the potassium manganate(VII) solution used, and the volume of the iron(II) sulphate solution used.

It's important to repeat the experiment several times to obtain a concordance of ∓0.10cm3 to ensure that our results are reliable and accurate.

Titrations

What is titration?

Titration is a way of analysing chemicals to find an unknown concentration by using a substance with known concentration.

Why is universal indicator not used in titrations?

Universal indicator gives a different colour for different pH ranges. That makes it hard to titrate to a specific pH value. On the other hand, specialised indicators like phenolphthalein change from colourless to deep red at pH above 9.0.

How do you perform a titration?

Using a clean pipette, measure a set volume of a solution of unknown concentration into a clean conical flask.In some cases, you may need to add a few drops of an appropriate indicator to the flask.Place a white tile under the flask so you can easily see any colour changes. Rinse and fill the burette with the solution of known concentration. Record the starting volume in the burette. Slowly add the solution in the burette to the solution in the conical flask while gently swirling the flask. Stop the titration when you reach the endpoint. The endpoint is when one drop of excess solution from the burette changes the colour of the solution in the flask. Record the final volume from the burette. Subtract the initial burette reading from the final burette reading to obtain the titre. Repeat the titration until you get concordant titre values of ∓0.10 cm3.

How do you carry out titration calculations?

Let us try an example together. 23.9cm3 of 0.040 mol dm-3 of aqueous potassium permanganate reacted with 25cm3 of acidified iron(II) sulphate solution.  What was the concentration of Fe2+ ions in the solution? First things first, write down the equation for the reaction. The redox process between manganate(VII) and iron(II) takes place as follows:5Fe2+ (aq) + MnO4- (aq) + 8H+ (aq) ➔ 5Fe3+ (aq) + Mn2+ (aq) + 4H2O (l)Next, use the values provided to find the number of moles of MnO4- ions added to the flask. Use the formula: no. of moles = concentration x volume / 10000.04 x 23.9/1000 = 0.000956 or 9.56x10-4 moles of MnO4-Now we can figure out the number of moles of Fe2+ in the flask!Use the reaction equation to find the proportion of the reaction between the titrant and analyte. From the equation, we can see that 1 mole of manganate(VII) reacts with 5 moles of iron(II).So multiply 9.56x10-4 by 5.9.56x10-4 x 5 = 0.00478 or 4.78x10-3 of Fe2+ ionsUse the previous formula to calculate the concentration of Fe2+ ions.0.00478 = Conc. x 25 / 1000Conc. = 0.00478 x 1000 / 25Conc. = 0.1912 mol dm-3Congratulations, you have completed a titration calculation! Getting the hang of titration calculations takes practice. Have a go at the examples in the exercises section.

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