Performing Volumetric Analysis Revision Notes for HSC SSCE Chemistry

Performing Volumetric Analysis

Volumetric analysis is a quantitative technique used to determine the concentration of an unknown solution. The most common application in chemistry laboratories is acid-base titration, where precise measurements and careful technique are essential for accurate results.

Understanding key terms

Before performing a titration, it's important to understand two fundamental terms:

infoNote

Titrant: This is the solution placed in the burette. It has a known, accurately determined concentration and is added gradually during the experiment.

Analyte: This is the solution being analysed - the one with the unknown concentration. It is typically placed in a conical flask and receives the titrant during the procedure.

In most school laboratories, the acid is placed in the burette (as the titrant) because bases can gradually etch glassware, potentially changing the burette's calibrated volume over time.

Basic technique for performing a titration

To conduct an accurate volumetric analysis, follow these systematic steps:

Preparation phase: Start with one solution that has an accurately known concentration. This is crucial because all your calculations will depend on this value. Place your chosen solution in the burette, ensuring the acid is typically the titrant to prevent glassware damage from prolonged base exposure.

Setting up the flask: Use a calibrated pipette to transfer a precise, known volume of the other solution into a conical flask. The flask should first be rinsed with distilled water only - never with the solution you're about to add, as this would alter its concentration. Add exactly two drops of an appropriate acid-base indicator to the solution in the flask.

Performing the titration: Slowly add the titrant from the burette to the flask, swirling constantly to ensure thorough mixing. Continue this addition until the indicator shows its first permanent colour change - this signals that you've reached the end point of the titration. Record the volume of solution added from the burette by reading the initial and final burette readings.

chatImportant

Ensuring accuracy: Repeat the entire experiment several times. Your results are considered reliable when consecutive titration volumes agree to within approximately one-third of a division on the burette scale (typically $0.03$ mL). During subsequent titrations, add the titrant rapidly at first, then slow to a dropwise rate, and finally to parts of drops as you approach the expected end point. Calculate the average titre from your most concordant (closely agreeing) results.

Calculations: Write a balanced chemical equation for the reaction occurring between your acid and base. Use stoichiometric mole calculations to determine the concentration of the unknown solution.

Worked example: standardising sodium hydroxide

Let's examine a practical calculation involving potassium hydrogen phthalate, an excellent primary standard for standardising alkali solutions.

lightbulbExample

Worked Example: Standardising Sodium Hydroxide

The problem: Potassium hydrogen phthalate ( $\text{KH}(\text{C}_8\text{H}_4\text{O}_4)$ ) contains one acidic hydrogen per formula unit. A sample weighing $0.917$ g was dissolved in water and titrated with approximately $0.2$ mol L $^{-1}$ sodium hydroxide solution. The titration required $27.2$ mL of the sodium hydroxide to reach the end point. What is the accurate molarity of the hydroxide solution?

Solution approach:

Step 1: Write the balanced equation for the neutralisation reaction:

$\text{KH}(\text{C}_8\text{H}_4\text{O}_4)(\text{aq}) + \text{NaOH}(\text{aq}) \rightarrow \text{Na}^+(\text{aq}) + \text{K}(\text{C}_8\text{H}_4\text{O}_4)^-(\text{aq}) + \text{H}_2\text{O}(\text{l})$

Step 2: Calculate the moles of potassium hydrogen phthalate using its mass and molar mass:

$n_{\text{KH}(\text{C}_8\text{H}_4\text{O}_4)} = \frac{m}{M} = \frac{0.917}{204.22} = 0.00449 \text{ mol}$

Step 3: From the balanced equation, the mole ratio shows that one mole of potassium hydrogen phthalate reacts with exactly one mole of sodium hydroxide. Therefore:

$n_{\text{NaOH}} = n_{\text{KH}(\text{C}_8\text{H}_4\text{O}_4)} = 0.00449 \text{ mol}$

Step 4: Calculate the molarity of the sodium hydroxide solution:

$C_{\text{NaOH}} = \frac{n_{\text{NaOH}}}{V_{\text{NaOH}}} = \frac{0.00449}{0.0272} = 0.165 \text{ mol L}^{-1}$

Answer: The accurate concentration of the sodium hydroxide solution is $0.165$ mol L $^{-1}$ .

Worked example: analysing wine acidity

This example demonstrates how to handle diprotic acids, which release two hydrogen ions per molecule.

lightbulbExample

Worked Example: Analysing Wine Acidity

The problem: The acidity of a white wine was determined by titrating $25.0$ mL of wine with $0.0511$ mol L $^{-1}$ sodium hydroxide solution. The titration required $8.70$ mL of sodium hydroxide. Calculate the molarity of hydrogen ions in the wine. Assuming all hydrogen ions come from diprotic tartaric acid ( $\text{H}_2\text{C}_4\text{H}_4\text{O}_6$ ), express the concentration of tartaric acid in grams per $100$ mL.

Solution approach:

Step 1: Begin with the balanced chemical equation:

$\text{H}_2\text{C}_4\text{H}_4\text{O}_6(\text{aq}) + 2\text{NaOH}(\text{aq}) \rightarrow 2\text{Na}^+(\text{aq}) + \text{C}_4\text{H}_4\text{O}_6^{2-}(\text{aq}) + 2\text{H}_2\text{O}(\text{l})$

Step 2: Calculate the moles of sodium hydroxide used:

$n_{\text{NaOH}} = C_{\text{NaOH}} \times V_{\text{NaOH}} = 0.0511 \times 0.00870 = 0.000445 \text{ mol}$

Step 3: Since the equation shows that each mole of sodium hydroxide neutralises one mole of hydrogen ions:

$n_{\text{H}^+} = n_{\text{NaOH}} = 0.000445 \text{ mol}$

Step 4: Calculate the molarity of hydrogen ions:

$C_{\text{H}^+} = \frac{n_{\text{H}^+}}{V_{\text{H}_2\text{C}_4\text{H}_4\text{O}_6}} = \frac{0.000445}{0.0250} = 0.0178 \text{ mol L}^{-1}$

Step 5: Remember that tartaric acid is diprotic - each molecule releases two hydrogen ions. Therefore, the moles of tartaric acid are half the moles of hydrogen ions:

$n_{\text{H}_2\text{C}_4\text{H}_4\text{O}_6} = \frac{1}{2}n_{\text{NaOH}} = 0.000222 \text{ mol}$

Step 6: Convert moles to mass using the molar mass of tartaric acid ( $150.088$ g mol $^{-1}$ ):

$m_{\text{H}_2\text{C}_4\text{H}_4\text{O}_6} = n_{\text{H}_2\text{C}_4\text{H}_4\text{O}_6} \times M_{\text{H}_2\text{C}_4\text{H}_4\text{O}_6} = 0.000222 \times 150.088 = 0.0334 \text{ g}$

Step 7: Express this as concentration in g per $100$ mL:

$[\text{H}_2\text{C}_4\text{H}_4\text{O}_6] = \frac{0.0334}{0.0250} = 1.33 \text{ g L}^{-1} = 0.133 \text{ g per } 100 \text{ mL}$

Answer: The wine contains $0.133$ g of tartaric acid per $100$ mL.

Choosing the correct indicator

Selecting an appropriate indicator is critical for accurate titrations. The indicator must change colour at a pH close to the equivalence point of your specific acid-base reaction.

Understanding acid and base strength

Strong acids (like hydrochloric acid, sulfuric acid, and nitric acid) ionise almost completely in solution, releasing nearly all their hydrogen ions. Weak acids (like acetic acid) only partially ionise, with most molecules remaining intact.

Similarly, strong bases (like sodium hydroxide and other hydroxide salts) ionise completely, while weak bases (like carbonates and ammonia) ionise only partially.

When an acid neutralises a base, the fundamental reaction is:

$\text{H}^+(\text{aq}) + \text{OH}^-(\text{aq}) \rightarrow \text{H}_2\text{O}(\text{l})$

infoNote

However, the salt formed during this neutralisation affects the final pH. Not all salts are neutral - some produce acidic solutions, others basic solutions, depending on whether they're formed from strong or weak acids and bases.

Indicator selection criteria

To choose the correct indicator, follow these steps:

Identify the salt formed during the neutralisation
Determine if either ion in the salt is a weak acid or weak base
Predict whether the solution at equivalence will be acidic (pH < 7), neutral (pH = 7), or basic (pH > 7)
Select an indicator that changes colour in the appropriate pH range

Type of titration	Example	Predicted equivalence point	Indicator
Strong acid–strong base	Hydrochloric acid and sodium hydroxide	Neutral	Bromothymol blue
Strong acid–weak base	Hydrochloric acid and sodium carbonate	Acidic region	Methyl orange
Weak acid–strong base	Acetic acid and sodium hydroxide	Basic region	Phenolphthalein
Weak acid–weak base	Acetic acid and sodium carbonate	Generally do not use direct titration	Not applicable

chatImportant

Exam tip: For weak acid–weak base combinations, direct titration is typically not used because the pH change at the equivalence point is too gradual for accurate indicator-based detection.

Investigation: performing a practical titration

This investigation brings together all the theoretical knowledge into practical application. You'll use a primary standard solution to determine the unknown concentration of hydrochloric acid.

Safety considerations

Before beginning any practical work, it's essential to understand and manage the risks involved.

chatImportant

Chemical hazards: Dilute solutions of hydrochloric acid, sodium carbonate, and methyl orange indicator may splash onto your skin or into your eyes. Always wear safety glasses throughout the investigation and wash your hands thoroughly at the end of the practical session.

chatImportant

Glassware safety: Laboratory glassware can break and cause cuts. Keep all glassware away from the edge of the bench to prevent accidental falls. When the pipette is not in use, leave the pipette filler attached - this prevents the pipette from rolling off the bench. If any glassware breaks, inform your teacher immediately and never attempt to clean it up yourself.

Materials required

You'll need the following equipment and solutions:

$250$ mL of sodium carbonate solution with accurately known concentration (from a previous preparation)
$200$ mL hydrochloric acid of unknown concentration
$50$ mL burette with retort stand and clamp
$25$ mL pipette with pipette filler
Two $150$ mL beakers for holding solutions
Three $250$ mL conical flasks
Dropper bottle containing methyl orange indicator
Small labels for the beakers
Wash bottle with distilled water
Filter funnel
Safety glasses

Detailed procedure

Step 1 - Prepare the acid: Rinse one $150$ mL beaker with a small amount of the hydrochloric acid solution (to remove any water and prevent dilution), discard this rinse, then label the beaker clearly and fill it with approximately $100$ mL of hydrochloric acid.

Step 2 - Fill the burette: Prepare the burette by rinsing it with small portions of the hydrochloric acid solution, then fill it completely with the acid. Ensure no air bubbles remain in the tip.

Step 3 - Prepare the base: Rinse the second $150$ mL beaker with sodium carbonate solution, label it, and fill with approximately $100$ mL of the sodium carbonate solution.

Step 4 - Prepare the flask: Rinse a conical flask with distilled water only (not with the solution you'll be adding).

Step 5 - Add the analyte: Rinse the pipette with the sodium carbonate solution, then use it to transfer exactly $25.00$ mL of sodium carbonate solution to the conical flask.

Step 6 - Add indicator: Add precisely two drops of methyl orange indicator to the flask and swirl gently to mix thoroughly.

Step 7 - Perform the titration: Place the conical flask under the burette tip. Record the initial burette reading. Begin adding the hydrochloric acid slowly while swirling the flask constantly. As you approach the end point (when the colour begins to change), slow down to adding drops one at a time. The end point is reached when the indicator shows its first permanent colour change that persists after swirling.

Step 8 - Record results: Note the final burette reading and calculate the volume of acid used.

Step 9 - Repeat for accuracy: Perform several more titrations until your titre values are consistent within $0.03$ mL. This usually requires at least three concordant results.

Analysing your results

After completing your practical work, you'll need to process the data:

Calculate the average volume of hydrochloric acid from your concordant titres
Write a balanced chemical equation: $\text{Na}_2\text{CO}_3 + 2\text{HCl} \rightarrow 2\text{NaCl} + \text{H}_2\text{O} + \text{CO}_2$
Calculate moles of sodium carbonate in the $25.00$ mL sample using $n = CV$
Use the $1:2$ mole ratio from the equation to determine moles of hydrochloric acid
Calculate the molarity of hydrochloric acid using $C = \frac{n}{V}$

infoNote

Compare your calculated value with those obtained by your peers. Discuss any variations and consider how experimental technique could be improved to increase accuracy.

chatImportant

The conical flask is rinsed only with distilled water because the exact volume of analyte is measured using the pipette. If you rinsed the flask with the analyte solution, you'd alter the amount present, making your volumetric measurement meaningless.

bookmarkSummary

Key Points to Remember:

The titrant is the solution of known concentration in the burette; the analyte is the solution of unknown concentration being analysed.
Always perform multiple titrations and use the average of concordant results (those within $0.03$ mL) for calculations.
Select indicators based on the expected pH at the equivalence point: bromothymol blue for strong acid–strong base, methyl orange for strong acid–weak base, and phenolphthalein for weak acid–strong base.
The balanced chemical equation is essential for determining the correct mole ratio in your calculations.
Accuracy in volumetric analysis depends on careful technique: proper rinsing of equipment, precise volume measurements, and patient addition of titrant near the end point.

Performing Volumetric Analysis (HSC SSCE Chemistry): Revision Notes