Introduction to Titration
Quick Notes
- Titration is a method used to determine the concentration of a substance (the analyte) in solution.
- A solution of known concentration (the titrant) is added gradually until the reaction is complete.
- The equivalence point is the moment when the moles of added titrant react perfectly with moles of analyte, based on the reaction’s stoichiometry.
- A visible change (usually a color change) signals the endpoint, which should closely match the equivalence point.
- Titrations are typically used in acid–base and redox reactions.
Full Notes
What Is a Titration?
A titration is an analytical technique used to determine the concentration of an unknown solution (the analyte) by reacting it with a solution of known concentration (the titrant). Titrations are particularly useful when a reaction occurs completely and predictably between the two substances.
The Equivalence Point
- The amount of titrant added exactly reacts with all of the analyte.
- The moles of titrant and analyte are present in the correct stoichiometric ratio.
For a simple 1:1 acid–base reaction:
HA + OH- → A- + H2O
Equivalence point is reached when:
moles of OH- added = moles of HA in solution
For more complex stoichiometries (e.g., 1:2 for NaOH reacting with H2SO4), adjust calculations accordingly.
The Endpoint
The endpoint is the observable signal that the equivalence point has been reached (e.g., a color change in an indicator). A good indicator changes color very close to the actual equivalence point.
Example:Phenolphthalein turns pink in basic (alkaline) solution. In a titration of a strong acid with a strong base, the endpoint is when a faint pink color remains.
Acid–Base Titration Typical Equipment and Setup
Prepare the solution for titration:
- Wash out a 250 cm³ volumetric flask with distilled water.
- Use a pipette to transfer exactly 25.0 cm³ of the unknown HCl solution into the flask.
- Fill the flask to the mark with deionised water – this dilutes the acid to a known volume.
Set up titration equipment:
- Fill a burette with standard sodium hydroxide solution.
- Place 25.0 cm³ of the diluted hydrochloric acid in a conical flask using a clean pipette.
Add indicator: Add 2 drops of phenolphthalein – the solution should remain colourless (acidic).
Titrate:
- Slowly add NaOH from the burette while swirling.
- The endpoint is reached when the solution turns pale pink and stays pink for at least 5 seconds.
- Record the initial and final burette readings to the nearest 0.05 cm³.
- Repeat until you obtain two concordant titres (within 0.10 cm³). Record results in a table.
Question: What volume of 0.100 mol/L NaOH is needed to neutralize 25.0 mL of 0.100 mol/L HCl?
- Balanced equation: NaOH + HCl → NaCl + H2O
Mole ratio = 1:1 - Find moles of HCl:
n = 0.100 mol/L × 0.0250 L = 0.00250 mol - Find volume of NaOH needed:
Moles of NaOH = 0.00250 mol
Volume = 0.00250 ÷ 0.100 = 0.0250 L = 25.0 mL
Answer: 25.0 mL of NaOH is needed to reach the equivalence point.
Always write the balanced chemical equation first to identify the correct mole ratio between titrant and analyte. Keep units consistent (especially volume in liters) and be precise when reading from the burette. Remember: the endpoint should be just after the last drop that causes a permanent color change.
Summary
- Titration is a key experimental technique used to determine the concentration of a solution by reacting it with a measured amount of another solution of known concentration.
- The equivalence point occurs when the amounts of titrant and analyte are in the correct stoichiometric ratio.
- The endpoint provides a visible cue that this point has been reached, often through a color change from an indicator.
- Mastering titration requires both accurate measurement and a clear understanding of the chemical reaction involved.