10 – Determining the Concentration of Ethanoic Acid in Vinegar (LC 2027) (Leaving Cert Chemistry): Revision Notes
10 – Determining the Concentration of Ethanoic Acid in Vinegar
Introduction and purpose
This experiment helps us find out exactly how much ethanoic acid (also called acetic acid) is present in household vinegar. Vinegar is a dilute solution of ethanoic acid in water, and we can determine its concentration using a technique called titration. This is a classic example of an acid-base neutralisation reaction where we use a strong base to neutralise the weak acid in vinegar.
Titration is one of the most important quantitative analysis techniques in chemistry, allowing us to determine unknown concentrations with high precision. This experiment demonstrates both acid-base chemistry principles and practical analytical skills used in laboratories worldwide.
The method involves carefully measuring how much sodium hydroxide solution is needed to completely neutralise a known amount of vinegar. From this information, we can calculate the original concentration of ethanoic acid.
Chemical reaction involved
When ethanoic acid reacts with sodium hydroxide, a neutralisation reaction occurs:
- CH₃COOH = ethanoic acid (the acid in vinegar)
- NaOH = sodium hydroxide (strong base)
- CH₃COONa = sodium ethanoate (salt formed)
- H₂O = water
This is a 1:1 molar ratio reaction, meaning one molecule of ethanoic acid reacts with exactly one molecule of sodium hydroxide. This stoichiometric relationship is fundamental to all our calculations in this experiment.
Experimental setup
The experimental setup includes standard analytical equipment: a volumetric flask for dilution, beakers for holding solutions, a burette for precise measurement of sodium hydroxide, and phenolphthalein indicator to detect the end point.
Each piece of equipment serves a specific purpose in ensuring accurate results. The volumetric flask provides precise dilution, the burette allows controlled addition of titrant, and the conical flask facilitates easy swirling and observation of colour changes.
Key experimental procedure
Step 1: Dilution preparation
First, we need to dilute the vinegar because it's too concentrated for accurate titration. We take a small sample of vinegar (typically 25 cm³) and dilute it in a 250 cm³ volumetric flask. This creates a dilution factor of 10 (from 25 cm³ to 250 cm³), making the solution easier to titrate accurately.
Worked Example: Calculating Dilution Factor
Step 1: Identify the volumes
- Initial volume = 25 cm³
- Final volume = 250 cm³
Step 2: Calculate dilution factor Dilution factor = Final volume ÷ Initial volume Dilution factor = 250 ÷ 25 = 10
This means the original vinegar is diluted 10 times.
The dilution is important because undiluted vinegar would require such a small volume of sodium hydroxide that measurement errors would be significant.
Step 2: Equipment preparation
- Clean and rinse the burette with the sodium hydroxide solution
- Transfer some diluted vinegar to a conical flask
- Add 2-3 drops of phenolphthalein indicator to the vinegar sample
Step 3: Titration process
- Fill the burette with standard sodium hydroxide solution
- Record the initial burette reading
- Slowly add sodium hydroxide to the vinegar while swirling the flask
- Watch carefully for the end point - this is when the indicator changes colour
Near the end point, add the sodium hydroxide drop by drop while continuously swirling. The end point can be reached very quickly, and overshooting will ruin your titration results.
Observations and explanations
Indicator behaviour
Phenolphthalein is the perfect indicator for this experiment because:
- It remains colourless in acidic solutions (like our vinegar sample)
- It turns pink in basic (alkaline) solutions
Phenolphthalein has a transition range of pH 8.2-10.0, making it ideal for weak acid-strong base titrations. The sharp colour change occurs very close to the equivalence point of this reaction.
At the end point
When we reach the end point, the phenolphthalein changes from colourless to a permanent pale pink colour. This tells us that all the ethanoic acid has been neutralised and there's now a tiny excess of sodium hydroxide, making the solution slightly basic.
The colour change should persist for at least 30 seconds - if it fades quickly, you haven't reached the true end point yet.
Calculations
Working out the concentration
From the titration results, we can calculate the concentration of ethanoic acid using these steps:
Worked Example: Complete Concentration Calculation
Given data:
- Volume of NaOH used = 23.5 cm³ = 0.0235 dm³
- Concentration of NaOH = 0.100 mol/dm³
- Volume of diluted vinegar = 25.0 cm³ = 0.0250 dm³
- Dilution factor = 10
Step 1: Calculate moles of NaOH used Moles of NaOH = Volume × Concentration Moles of NaOH = 0.0235 × 0.100 = 0.00235 mol
Step 2: Find moles of ethanoic acid Since reaction ratio is 1:1, moles of CH₃COOH = 0.00235 mol
Step 3: Calculate concentration in diluted sample Concentration = moles ÷ volume Concentration = 0.00235 ÷ 0.0250 = 0.094 mol/dm³
Step 4: Find original concentration Original concentration = 0.094 × 10 = 0.94 mol/dm³
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Calculate moles of NaOH used: Volume of NaOH (in dm³) × Concentration of NaOH (in mol/dm³)
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Find moles of ethanoic acid: Since the reaction ratio is 1:1, moles of CH₃COOH = moles of NaOH
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Calculate concentration in the diluted sample: Concentration = moles ÷ volume of diluted sample (in dm³)
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Find original concentration: Multiply by the dilution factor to get the concentration in the original vinegar
Important calculation considerations
Remember to account for the dilution factor when calculating the final concentration. The concentration you calculate from the titration is for the diluted vinegar, so you must multiply by the dilution factor to find the concentration in the original vinegar sample.
Important experimental notes
Why we dilute the vinegar
Diluting the vinegar serves several purposes:
- Makes titration volumes more manageable and accurate
- Reduces the concentration to a level suitable for the indicator
- Minimises measurement errors that would occur with very small volumes
Without dilution, the volume of NaOH required would be so small (around 2-3 cm³) that even tiny measurement errors would significantly affect the accuracy of results. The dilution increases this to a more manageable 20-25 cm³.
Accuracy tips
- Rinse all glassware with distilled water
- Ensure the burette tap doesn't leak
- Add sodium hydroxide drop by drop near the end point
- Repeat the titration several times for reliability
- Take the average of concordant results (within 0.1 cm³ of each other)
Common Mistakes to Avoid:
- Forgetting to account for the dilution factor in final calculations
- Reading the burette incorrectly (always read at eye level to the bottom of the meniscus)
- Adding NaOH too quickly near the end point
- Using results that aren't concordant (differ by more than 0.1 cm³)
Safety considerations
- Sodium hydroxide is corrosive - handle with care
- Wear safety glasses and protective clothing
- Wash hands thoroughly after the experiment
Key Points to Remember:
- Ethanoic acid reacts with sodium hydroxide in a 1:1 ratio - this is crucial for calculations
- Phenolphthalein is colourless in acid and pink in base - the end point is a permanent pale pink colour
- Always account for dilution factors when calculating the final concentration of the original vinegar
- Accurate burette readings and multiple titrations are essential for reliable results
- The experiment demonstrates acid-base neutralisation and quantitative analysis techniques used in analytical chemistry