0.1 Stoichiometry: laws to moles to molarity  (Page 2/2)

1. The formula for potassium chromate is K ₂ CrO ₄ .
2. The elements present are potassium, chromium, and oxygen with atomic masses of 39.10, 52.00 and 16.00 respectively. Adding up these numbers in the correct ratios dictated by the formula yields the following: 2 x 39.10 + 1 x 52.00 + 4 x 16.00 = 194.20 g/mol.
3. For one liter of solution use a 1000 mL volumetric flask. So a 1M solution would require 194.2g of solid K ₂ CrO ₄ in 1 L, 0.1M 19.42g of solid K ₂ CrO ₄ and so on.

Your teaching assistant will check the accuracy of the solution that you have made by titration, which is a method of quantitative ly determining the concentration of a solution . A standard solution (a solution of known concentration ) is slowly added from a buret te to a solution of the analyte (a solution of unknown concentration – your solution) until the reaction between them is judged to be complete ( equivalence point ). In colorimetric titration , some indicator must be used to locate the equivalence point . One example is the addition of acid to base using phenolphthalein ( indicator ) to turn a pink solution colorless in order to determine the concentration of unknown acid s and base s. Record your TAs value of the molarity of your solution on your report form along with your percent error.

Figure 1: Reading the Burette

When an acid is neutralized by a base, there are stoichiometrically equal amounts of acid and base and the pH = 7. It is possible to accurately determine the concentration of either the acid or base solution.

Moles of a substance = Concentration of solution (moles/L) x Volume (L)

We can calculate the concentration of the acid or base in the solution by using the following equation where balance base and balance acid refer to the stoichiometric ratio of the base and acid to each other.

Titration calculations:

Step 1: Balance the neutralization equation. Determine Balance of Acid and Base.

Step 2: Determine what information is given.

Step 3: Determine what information is required.

Step 4: Solve using the equation below.

${{\rm B}}_b\times C_a\times V_a={{\rm B}}_a\times C_b\times V_b$

Example:

Calculate the concentration of a nitric acid solution HNO ₃ if a 20 mL sample of the acid required an average volume of 55 mL of a 0.047 mol/L solution of Ba(OH) ₂ to reach the endpoint of the titration.

Step 1: ${\text{2HNO}}_3+\text{Ba}{\left(\text{OH}\right)}_2\to \text{Ba}{\left({\text{NO}}_3\right)}_2+{\text{2H}}_{2}O$ Balance Base = 1Balance Acid = 2

Step 2: Given information Volume Acid = 20 mLVolume Base (average) = 55 mL Concentration of Base = 0.047 mol/L

Step 3: Required information Concentration of Acid Step 4: Solve using the equation. ${{\rm B}}_b\times C_a\times V_a=\text{Βa}\times C_b\times V_b$ $1\times \text{Ca}\times \text{20mL}=2\times 0\text{.}\text{047mol/L}\times \text{55mL}$ Ca = 0.2585 mol/L (considering significant figures 0.26 mol/L)

Procedure

Materials list

sodium bicarbonate NaHCO ₃

3M hydrochloric acid (HCl) solution

Part 1

1. Weigh an empty 150-mL beaker on the electronic balance. Record this value in your data table.
2. Remove the beaker from the balance and add one spoonful of sodium bicarbonate. Re-weigh and record this value. Use a small spoonful.
3. Subtract the mass of the beaker and sodium bicarbonate from the mass of the beaker to get the mass of the sodium bicarbonate added.
4. Pour approximately 20 mL of 3M hydrochloric acid into a 100-mL beaker. Rest a Pasteur pipette in the beaker.
5. Add 3 drops of acid to the NaHCO ₃ beaker, moving the pipette so that no drops land on each other. The key point is to spread out the addition of acid so as to hold all splatter within the walls of the beaker.
6. Continue to add acid slowly drop by drop. As liquid begins to build up, gently swirl the beaker. This is done to make sure any unreacted acid reaches any unreacted sodium bicarbonate. Do not add acid while swirling.
7. Stop adding the hydrochloric acid when all bubbling has ceased so that the minimum amount of HCl has reacted with all of the sodium bicarbonate. Check when all the bubbling has ceased, by swirling the beaker and to ensure that there is no more bubbling. When all the bubbling has ceased add one drop more of acid and swirl.
8. Using a microwave oven, dry to constant weight. Initially place the beaker in the microwave for 1 min when there is plenty of solution present. Then in 5 second intervals thereafter. Drying to a constant weight means that when you think your sample is dry, you weigh the beaker, dry for 5 more seconds, weigh the beaker again, dry for 5 more seconds, and weigh the beaker one last time. The three masses should be within one milligram. If they are not, continue to dry your sample until they are.
9. Calculate the percent yield of NaCl. Percent yield can be calculated as follows:

Materials list

100 mL volumetric flask

3M hydrochloric acid (HCl) solution

sodium bicarbonate (NaCHO ₃ )

methyl red indicator

Part 2

1. Ask you TA for your assigned molarity – it will range from 0.1M to 1M.
2. From the formula of the solute, find the molecular weight of the solute, in g/mol.
3. Determine how many grams of sodium bicarbonate you will need to make a 100mL solution for the assigned molarity.
4. Using your 100mL volumetric flask, add the correct amount of sodium bicarbonate. Make sure to use a funnel to pour the solute into the flask. Rinse the funnel with deionised water to make sure all of the solute makes it into the volumetric flask. Fill the volumetric flask half way with water, cap it, and invert repeatedly until all of the solute is dissolved. Hold your thumb over the cap to make sure that the flask does not leak. Once all of the solute is dissolved, remove the cap, fill the flask to the mark with water, recap the flask, and invert several more times.
5. Take your solution to your TA to check the molarity by titration, record value on your report form and your percent error. Percent error can be calculated as follows: