0.1 Stoichiometry: laws to moles to molarity  (Page 2/3)

Your teaching assistant will check the accuracy of the solution that you have made by titration, which is a method of quantitatively determining the concentration of a solution. A standard solution (known concentration) is slowly added from a burette to a solution of the analyte (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 acids and bases. 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, since there is 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. Since:

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 using:

$\begin{array}{}\text{Balance Base}\text{Bb}\times \text{Moles of Acid}=\text{Moles of Base}\times \text{Balance Acid}\text{Ba}\\ {{\rm B}}_b\times C_a\times V_a=\text{Βa}\times C_b\times V_b\end{array}$

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.

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

Example:

Calculate the concentration of a nitric acid solution ${\text{HNO}}_3$ if a 20 ml sample of the acid required an average volume of 55 ml of a 0.047 mol/l solution of $\text{Ba}{\text{OH}}_2$ to reach the endpoint of the titration.

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

Step 2:Given informationVolume Acid = 20 mlVolume Base (average) = 55 ml Concentration of Base = 0.047 mol/l

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

Experimental

Materials list

sodium bicarbonate ${\text{NaHCO}}_3$

3M hydrochloric acid (HCl) solution

Procedure

Part 1

• Weigh an empty 150-mL beaker on the electronic balance. Record this value in your data table.
• Remove the beaker from the balance and add one spoonful of sodium bicarbonate (approximately 5 g). Re-weigh and record this value.
• Pour approximately 20 mL of 3M hydrochloric acid into a 100-mL beaker. Rest a Pasteur pipette in the beaker.
• Add 3 drops of acid to the ${\text{NaHCO}}_3$ beaker, moving the pipette so that no drops land on each other. The key point is to spread out the adding of acid so as to hold all splatter within the walls of the beaker.
• 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.
• 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.
• Weigh the beaker and contents, record.
• Using a microwave oven, dry to constant weigh, initially for 1 min, when there is plenty of solution, and then 10 second intervals thereafter. Measure weight to the nearest milligram.