Question

You wish to prepare 1 L of a 0.02-M potassium iodate solution. You require that the final concentration be within 1% of 0.02 M and that the concentration must be known accurately to the fourth decimal place. How would you prepare this solution? Specify the glassware you would use, the accuracy needed for the balance, and the ranges of acceptable masses of \(\mathrm{KIO}_{3}\) that can be used

Short answer

To prepare a 1 L 0.02 M potassium iodate (KIO3) solution within 1% of the target concentration and with an accuracy of four decimal places, you should use a 1 L Class A volumetric flask and a high-precision balance with a readability of 0.0001 g (0.1 mg). The mass of KIO3 required is between 3.29 g and 3.35 g. Weigh the KIO3 with the high-precision balance, dissolve it in distilled water, transfer the solution to the volumetric flask, and fill the flask up to the 1 L mark to obtain the desired concentration and accuracy.

Step-by-step solution

Calculate the target mass of KIO3 required in the solution

Firstly, we need to find out the target mass of KIO3 required to make a 1 L solution with a concentration of 0.02 M. We will use the following formula: \[mass = (m p ) á (V) á (M)\] where: \(mass\) = mass of KIO3 in grams \(MW\) = molecular weight of KIO3 \(V\) = volume of the solution in liters \(M\) = molarity of the solution in moles/liter The molecular weight of KIO3 is 166 g/mol (39 g/mol from potassium, 127 g/mol from iodine, and 16x3 g/mol from oxygen). The volume and molarity of the solution are 1 L and 0.02 M, respectively. Plugging in the values into the formula, we find: \(mass = (166 g \cdot mol^{-1}) á (1 L) á (0.02 mol \cdot L^{-1}) = 3.32 g\)

Calculate the acceptable mass range of KIO3 for 1% variation

Next, we will determine the acceptable range of KIO3 masses that will result in a concentration within 1% of the desired 0.02 M. To find the upper and lower bounds of the acceptable mass range, we will multiply the target mass (3.32 g) by 1.01 and 0.99, respectively: Upper bound mass = \(3.32 g × 1.01 = 3.35 g\) Lower bound mass = \(3.32 g × 0.99 = 3.29 g\) Hence, the acceptable mass range of KIO3 needed for this solution is between 3.29 g and 3.35 g.

Determine the appropriate glassware and balance needed

To ensure the high level of accuracy needed for the final concentration, we should use high-quality glassware and devices. For the glassware, we can use a volumetric flask that is rated to hold 1 L of solution and has a high uncertainty tolerance to ensure precise volumes. A 1 L Class A volumetric flask would be suitable for this purpose, as it has an uncertainty of ± 0.1 mL in volume, which is sufficient for a four decimal place accuracy in the final concentration. As for the balance, we will need to choose one with a high level of precision and accuracy. Since we need to measure an acceptable mass range of 3.29 g to 3.35 g, a balance with a readability of 0.0001 g (0.1 mg) would be suitable. This level of precision will allow us to obtain the required four decimal place accuracy in the final concentration.

Preparing the solution with the appropriate glassware and balance

Finally, we will weigh out the KIO3 and prepare the solution using the recommended glassware and balance. 1. Use the high-precision balance (0.0001 g readability) to weigh out a mass of KIO3 between the acceptable range of 3.29 g to 3.35 g. 2. Dissolve the KIO3 in a small volume of distilled water in a beaker. 3. Transfer the dissolved KIO3 solution into the 1 L Class A volumetric flask. 4. Rinse the beaker with distilled water and add the rinsate to the volumetric flask to ensure complete transfer. 5. Fill the volumetric flask with distilled water up to the 1 L mark. 6. Mix the solution well by gently inverting the flask several times to distribute the KIO3 evenly throughout the solution. Now you have prepared a 1 L solution with a potassium iodate concentration within 1% of 0.02 M, known accurately to the fourth decimal place. Keep in mind that the range is calculated to assess the accuracy needs and may be different from the analytical needs of the experiment.

Key concepts

Molarity

Molarity is a key term in chemistry that measures the concentration of a solute in a solution. Specifically, it is the number of moles of a solute per liter of solution, expressed as moles per liter (mol/L). This makes it an important unit when preparing chemical solutions for experiments or industrial processes.

To calculate molarity, the formula is:

• M = moles of solute / liters of solution

In the exercise, we are asked to prepare a 1-liter solution with a molarity of 0.02 M for potassium iodate (KIO extsubscript{3}). This ensures that every liter of the solution contains 0.02 moles of potassium iodate, providing a defined quantity necessary for consistency in experimental results.

Understanding molarity is crucial for various applications such as titrations, where precise concentrations determine the outcome of the chemical reaction or process.

Volumetric Flask

A volumetric flask is a type of laboratory glassware used to prepare solutions with a specific volume and concentration. It is distinguished by its flat bottom, pear shape, and a long neck with a single graduation mark that indicates a precise volume.

Volumetric flasks are essential for creating accurate solutions because their design minimizes volume errors. In this exercise, a 1 L volumetric flask is recommended to ensure the solution's volume reaches exactly 1 liter, aligning with the molarity desired. It's crucial to use a Class A volumetric flask for this preparation because they offer a high degree of precision with a typical uncertainty of ±0.1 mL.

These flasks are particularly useful in analytical chemistry where precise volumes are necessary for accurate results. The narrow neck allows for a meniscus to be formed, and solutions can be mixed thoroughly by inverting the flask. This mixing is vital for ensuring the even distribution of the solute throughout the solvent.

Analytical Balance

An analytical balance is a high-precision instrument used to measure masses, providing readings to the nearest 0.0001 grams (0.1 mg). This high degree of precision is necessary when preparing solutions that require exact concentrations.

In this particular task, we aim to measure out potassium iodate (KIO extsubscript{3}) within a narrow mass range of 3.29 g to 3.35 g. An analytical balance with a 0.0001 g readability is recommended to achieve this precision and ensure that the solution's molarity falls within the acceptable limits.

Using an analytical balance correctly involves letting the device warm up, calibrating it if necessary, and using appropriate weighing protocols. This includes taring the scale after placing weighing paper or a container on it, which ensures only the mass of the chemical itself is measured. Such accuracy is critical for experiments requiring reliable and consistent outcomes.

Potassium Iodate

Potassium iodate (KIO extsubscript{3}) is an inorganic compound commonly used as a source of iodine in various chemical and biological applications. It appears as a white crystalline powder and is often used in analytical chemistry and iodine supplement preparations.

In this exercise, potassium iodate is the solute in preparing a 0.02 M solution. The molecular weight of KIO extsubscript{3} is calculated to be 166 g/mol. This value includes the atomic weights of its constituent elements: potassium (K), iodine (I), and oxygen (O).

• Potassium: 39 g/mol
• Iodine: 127 g/mol
• Oxygen (3): 16 g/mol each

This structure and precise molecular weight are essential for calculating the exact mass needed to achieve the desired molarity.

Handling potassium iodate requires careful attention to purity and weighing precision to ensure accurate solution preparation, underscoring its importance in scientific studies and iodine fortification processes.