Chemistry 141 - Chemical Principles I



Solution Stoichiometry – Quantitative Preparation of Aqueous Solutions

Introduction.

In this week’s lab, you and your lab partner will apply your prelab calculations to the preparation of aqueous solutions of known concentrations. The objectives of the lab include:

• calculation of solution concentrations based on the mass and molecular weight of solutes

• calculation of solution concentrations based on the dilution of stock solutions

• preparation of these solutions

• testing of your solution concentrations based on density and acid/base neutralization reactions

• observation of a rather unique example of “aqueous chemistry”

Part I. Preparation of Solutions.

Using the calculations completed in the pre-lab, prepare solutions B and C exactly as outlined below. Be sure to read the mass of reagents to three decimal places when using the electronic balances.

Solution A. 3.6 M H2O2 solution already prepared in hood.

(Caution!! Use gloves when handling!)

Solution B.

You will first make 100 mL solution of 0.08 M H2SO4 from the 1.00 M stock. You will then make 50 mL solution of 0.20 M potassium iodate (KIO3) using 50 ml of your 0.08 M H2SO4 solution just prepared, as the solvent. The remaining 50 ml of 0.08 M H2SO4 solution will be used for titration on Part II C.

Step 1 - 0.08 M H2SO4

Using a graduated cylinder, add _______ mL of 1.00 M H2SO4 into a 100 mL volumetric flask. Rinse the graduated cylinder once with about 15 mL of distilled water, and add this rinse to the same volumetric flask. Now, fill your volumetric flask to the 100 mL “mark” with distilled water. Swirl this solution.

Step 2 - 0.20 M KIO3 in 0.08 M H2SO4.

Add _______________ grams of potassium iodate (KIO3) to a 50 ml volumetric flask and then add 0.08 M H2SO4 , from step 1 above, up to the 50 ml mark. . You may need to swirl or stir this solution to dissolve the KIO3. Label this ‘Solution B’ and set aside.

Note: You still have 50 ml of 0.08 M H2SO4 left over to be used in titration!

Solution C.

You need to make 50 mL of a combined solution containing 0.15 M malonic acid (C3H4O4), and 0.02 M manganese sulfate monohydrate (MnSO4 ( H2O).

Add _______________grams of malonic acid (C3H4O4) and _______________ grams of manganese sulfate monohydrate (MnSO4 ( H2O) to a 50 ml volumetric flask and fill to the 50 ml mark with distilled water. Swirl to dissolve. Label this ‘Solution C’ and set aside.

NAME:___________________________ DATE: ______________

PARTNER:________________________ SECTION: ___________

Solutions

Data Sheet

Part II. Testing the concentration of your solutions.

You will be graded on the accuracy of your solution concentrations based on the results obtained below.

A) Density of solution C

Obtain a clean, dry 50 mL graduated cylinder, and determine its mass to three decimal places. Fill the graduated cylinder with approximately 40 mL of your Solution C, record the exact volume, and determine the mass of the filled graduated cylinder to three decimal places.

Data and Calculations (use correct units!).

Mass of empty graduated cylinder _______________

Volume of Solution C added to graduated cylinder _______________

Mass of graduated cylinder containing Solution C _______________

Mass of Solution C _______________

Calculated density of Solution C _______________

B) Stoichiometry of an acid/base neutralization reaction.

Complete the following equation by writing the Molecular, Ionic, and Net Ionic Equations:

(i) Balance the following complete molecular equation for the neutralization of sulfuric acid

(Molecular) H2SO4 + NaOH ( H2O + Na2SO4

(ii) Write the balanced ionic equation for this reaction and identify the spectator ions.

(Ionic equation)

(iii) Write net ionic equation for this reaction.

(Net Ionic)

C. Simple Titration of 0.08 M H2SO4.

Using an eye dropper add 20-30 drops of your 0.08 M H2SO4 (solution B, step 1) to a clean, dry large test tube and record the exact number of drops you added. Add one drop of phenolphthalein indicator to the test tube. The indicator should be colorless in the acid solution. Use a similar dropper to add 0.08 M NaOH solution to the same test tube, counting the drops as you go, and swirling the test tube to mix the solutions. Keep adding the NaOH dropwise until the solution is just barely pink. You now have a neutral solution. If the solution is deep pink or red, discard this sample and start over. Record the number of drops of NaOH solution added. Repeat this “titration” two more times. Show calculations.

Trial #

1 2 3

drops of 0.08 M H2SO4 used __________ __________ __________

drops of 0.08 M NaOH used __________ __________ __________

ml of 0.08 M H2SO4 used*** __________ __________ __________

ml of 0.08 M NaOH used*** __________ __________ __________

moles of H2SO4 used __________ __________ __________

moles of NaOH used __________ __________ __________

average moles of H2SO4 __________

average moles of NaOH __________

***Assume that one drop = 0.050 mL***

Determine the experimental NaOH : H2SO4 mole ratio and compare this to the NaOH : H2SO4 stoichiometry of your balanced molecular equation. Do they match or not? Explain.

D. An unusual example of “Solution Stoichiometry”.

Here’s the “big” pay-off. In a clean 250 mL Erlenmeyer flask mix 10 mL each (use a graduated cylinder) of Solutions B and C, and 5 drops of the Starch Solution provided to you. Now add 10 mL of Solution A (3.6M H2O2 solution made up for you), swirl to mix, then place it on a piece of white paper and record all observations. CAUTION: Hydrogen peroxide is a strong oxidizing agent that can cause burns. Should you get any on your hands, immediately rinse with water.

Count the total number of color “cycles” you observe over a 15 minute time period (colorless to blue): ____________

NAME:___________________________ DATE: ______________

SECTION: ___________

Solutions

Post-lab Questions

A student spilled 275 ml of a 0.350 M solution of H3PO4.There is available a solution of 0.500 M KOH to neutralize the acid. ( Hint: for now, we will assume that all H+ ions in H3PO4 will ionize)

1. Write and balance the chemical equation for this acid/base reaction.

2. Calculate the number of moles of H3PO4 spilled.

3. Based on the stoichiometry in your balanced molecular equation, determine the number of moles of KOH that will be needed to neutralize the acid.

4. Calculate the volume in ml of 0.500 M KOH that will be needed to neutralize the acid.

NAME: ___________________________ DATE: ______________

SECTION: __________

Solutions

Pre-lab Assignment

These answers will be used in the laboratory so you must be careful about your answers.

Part I. Definitions. Define each of the following terms in the space given below. Be Concise!

(a) solution

(b) percent concentration

(c) solubility

(d) precipitate

(e) molarity

Part II. Calculations.

The following calculations must be completed before you begin this experiment. You will use the results of these calculations to guide you through the preparation of solutions in the laboratory. SHOW ALL CALCULATION WORK! Be careful of your significant figures, and be sure to correctly label all units.

Solution A. This solution has already been made up for you, but you should work through the calculations for practice. You are required to make up 25.0 mL of a 3.6 M solution of hydrogen peroxide (H2O2). You have at your disposal a stock solution that is 30% by weight, H2O2, and has a density of 1.11 g/mL. Hint: A 30% by weight solution means that 100 g of the solution contains 30 g of H2O2. Use this information with the density to determine the Molar concentration of 30% H2O2, then apply your dilution equation to determine how you would make up 25.0 mL of 3.6 M H2O2.

1. 100 g of H2O2 solution contains 30 g of H2O2. Use the density of 1.11 g/mL to calculate the volume (mL) of 100 g of H2O2 solution.

Answer ________________________

2. Calculate the molecular mass of H2O2.

Answer_________________________

3. How many moles is 30.0 g of H2O2?

Answer_________________________

4. Based on your answers to 1 and 3, calculate the molar concentration of 30% H2O2.

Answer_________________________

5. Based on your answer to 4, determine the volume of 30% H2O2 needed to make 25.0 mL of 3.60 M H2O2.

Answer_________________________

Solution B. Make up a solution of 0.20 M potassium iodate (KIO3) in 0.08 M sulfuric acid (H2SO4). You have available a 1.00 M stock solution of H2SO4.

1. You will need to make up 100 mL of 0.08 M H2SO4 by dilution of a 1.00 M stock. Determine the volume (mL) of 1.00 M H2SO4 needed to complete the dilution.

Volume of 1.00 M H2SO4________________________

(enter this value on step 1 under Solution B on page 1)

2. What is the molecular mass of KIO3?

Answer ________________________

3. Calculate how many moles of KIO3 are needed to make 50 mL of 0.20 M solution.

Answer ________________________

4. Calculate how many grams of KIO3 are needed to make 50 mL of 0.20 M solution.

Answer_________________________

(enter this value on step 2 under Solution B on page 1)

Solution C. Make up a 50 mL solution containing both 0.15 M malonic acid (C3H4O4), and 0.02 M manganese sulfate (MnSO4 ( H2O, the ( means that the compound exists as a hydrate with a molecule of water trapped in the solid. You must include it in your calculation of molecular weight).

1. What is the molecular mass of malonic acid (C3H4O4)?

Answer_________________________

2. How many moles and grams of malonic acid are needed for 50mL of a 0.15M solution?

moles_________________________

grams_________________________

(enter this value on the appropriate line under Solution C on page 2)

3. What is the molecular mass of manganese sulfate (MnSO4 ( H2O)?

Answer_________________________

4. How many moles and grams of manganese sulfate are needed for 50mL of a 0.02M solution?

moles_________________________

grams_________________________

(enter this value on the appropriate line under Solution C on page 2)

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