Introduction:

The degree of acidity is often responsible for the chemical behavior of substances present in solutions.  For example, too much coffee or other food or drink sometimes causes gastric distress because of an acid imbalance in the stomach.  A number of commercial antacid preparations are available to relieve this condition.  In a number of instances it is necessary to be able to accurately determine the concentration of an acid (or a base) present in a solution.  The most common units for expressing solution concentration are molarity (M, moles of solute/liter of solution) and normality (N, equivalents of solute/liter of solution).

Titration is a volumetric method of chemical analysis which involves taking an accurately measured volume of an acid and adding base until the solution becomes neutral (has the same number of H3O+ ions as OH- ions).  The point at which the amounts of acid and base become equivalent is called the "equivalence point" and is usually signaled by a color change caused by some acid sensitive dye (called an indicator) which has been added to the solution.  The two solutions (one of acid, the other of base) are delivered using volumetric glassware (in this experiment the glassware is a burrett) so that volumes will be accurately known.  In addition, the concentration of either the acid or base must be accurately known (must be a standard solution).  For a more in-depth introduction to titration, click on this hyper link:  An Overview of Titrations

To be successful, a titration must involve a chemical reaction between the two solutions which are being mixed.  This chemical reaction should be simple, rapid and complete.  Although a titration may involve an acid base reaction, a precipitation reaction, a complexation reaction, or an oxidation reduction reaction, only the acid base reaction will be studied in this experiment. 

In the Arrhenius (classical) sense, an acid is any substance which furnishes protons (H+) and a base is any substance which furnishes hydroxide ions (OH-) in water.  In this experiment, an aqueous solution of HCl is the acid and an aqueous solution of NaOH is the base.  The overall balance (formula unit) equation for this reaction may be written as:

HCl(aq)  +  NaOH(aq)    H2O(l)  +  NaCl(aq)

A more meaningful way to look at this reaction is to consider the total ionic equation:

H+(aq)  +  Cl-(aq)  +  Na+(aq) +  OH-(aq)    H2O(l)  +  Na+(aq)  +  Cl-(aq)

Because the Na+ and Cl- ions undergo no chemical change in the process, they are commonly called spectator ions and are often removed to give us the net ionic equation.  Therefore, the net chemical reaction is:

H+(aq)  +  OH-(aq)     H2O(l)

This reaction between an acid and a base is called neutralization .  The base is added to the acid until the solution contains an equivalent amounts of each.  At this point, the acid is said to be "neutralized."  If the proper chemical indicator has been added to the solution, a color change occurs.  Using the measured volumes of the acid and base and the concentration of the "standard solution" (either acid or base), the concentration of the other reactant may be readily calculated from the equation:

(Volume of acid)(Normality of acid)  =  (Volume of base)(Normality of base)

VaN =  VbNb

When using solutions of acids and bases it is convenient to express concentration in terms of NORMALITY, the number of equivalents of solute in a liter of solution. 


Lab Tips:

Although they seem simple, many people initially have trouble with titrations.  There is a good deal of eye-hand coordination involved and a lot of small errors than can creep  in to ruin your experiment. The following are some tips that should help you be successful:

  1. Make sure you have studied the video, before coming to lab: An Overview of Titrations
  2. Add your base to the burette over the sink.  If you try to add it in the burette clamp, some might spill into your acid sample and you would have to start over.
  3. Don't forget to put the indicator in your sample.
  4. Make sure you setup your burette so the tip is below the top of the beaker.
  5. Set your magnetic stirrer as fast as possible, BUT no splashing.  If any of the acid splashes on the side, you need to wash it down with distilled water.
  6. Make sure you have no bubbles in the tip of your burette.  This probably causes 75% of the problems students have with getting titrations to be reproducible.
  7. When you see that the pink color start to persist, slow the addition of base to a drop at a time.
  8. To get the best equivalence points, you will need to 'cut' drops.  Barely open your stopcock and let less than a drop form on the tip.  Then use your distilled water bottle to squirt it into the beaker.  This will assure that you obtain lightest pink color that you can:

  1. Your burette is a direct read delivery burette.  This means, 'What You See Is What You Get'.  If you start at 0.00mL and stop at 23.56mL, you used 23.56mL.
  2. Very important:  read all burette readings to 0.01mL!  Remember, you always read to one place past what is marked on the measuring device.  Also make sure you look directly at the burette at eye level, do not look down or up to read the meniscus, this will cause parallax errors.  The following figure shows an initial volume of 9.62mL and a final volume of 24.16mL:

Procedure:

Preparation of NaOH Solution:

  1. Dilute approximately 10 mL of 6 N NaOH with approximately 450 mL of distilled water in a 600 mL beaker.  Rinse the graduated cylinder with distilled water and add to the beaker to make sure you get all of the NaOH.
  2. Stir the solution thoroughly.  Insufficient mixing is a common source of error in titrations.  This sodium hydroxide solution will be standardized in the next step by titrating it versus the Standard  HCl.
Titration of NaOH Solution:
  1. Obtain a 50 mL burrett, close the stopcock and fill it to the top with distilled water.  Open the stopcock and allow all of the water to drain.
  2. Close the stopcock and fill your burrett with 50 mL of your NaOH solution (from above) so that the solution comes in contact with the entire inner surface of the burrett.
  3. Open the stopcock and allow all of the NaOH to drain through the tip.
  4. Fill the burrett to the top with the NaOH.  Open the stopcock all the way to flush all bubbles out of the tip. When all bubbles have been flushed out (it may take several tries), close the stopcock and refill the burrett.
  5. Read the bottom of the meniscus and record the initial reading to the nearest 0.01 mL. The Teflon stopcock should turn smoothly with a little resistance.  If the stopcock is too loose, tighten it a little, otherwise the solution will leak around the stopcock and the titration will be for naught.  A leaking burrett is a major cause of error in titrations.
  6. For more information about the proper use of a burrett and titration procedures, click on this hyper link:  "How to Read A Burrett" or "Everything you wanted to know about titrations but were afraid to ask".
  7. Use a repipetter to deliver 10.00 mL of the standard HCl into a clean 150 mL beaker.  Wash down the sides of the beaker with the wash bottle.  The addition of water at this stage has no effect on the total amount of acid already present in the beaker.  Be sure to record the concentration of the standard HCl!
  8. Add 2 drops of phenolphthalein indicator to the acid solution.  The solution should remain colorless.
  9. Add a magnetic stir bar to the beaker and place on the heating stir plate.  Adjust the stirring rate to obtain a vortex without any of the solution splattering on the sides of the beaker.  Avoid spilling any of the beaker contents.  Any loss of sample would render the titration worthless.
  10. Rinse down the inside of the beaker occasionally and continue, slowly adding NaOH until the first permanent, faint pink color persists for at least 30 seconds.  At this point the titration is complete (the endpoint).
  11. Read the final volume of NaOH and record to the nearest 0.01 mL.  Remember these are 'direct read' burretts, what you see is what you get!  Some of you may have used other burretts where they start at 50 mL and go to zero, so you subtract your reading from 50.  DO NOT do that with these burretts!
  12. Calculate the normality of the NaOH solution.
  13. Repeat Steps #7-12 with 20.00 mL of standard HCl.  Fill the burrett and read it before each titration if there is a  possibility of running out before a titration is complete.
  14. Repeat Steps #7-12 with 30.00 mL of standard HCl.
  15. The results from these three titrations should agree within 0.005 N, otherwise titrate a fourth sample and throw out the value which does not agree.  Average the results and use this value as the concentration of your STANDARDIZED BASE.  Do not discard the base which you have just standardized or clean the burrett.   This standardized NaOH solution will be used to determine the concentration of an unknown acid solution in the next section.
Titration of Unknown Acid:
  1. Use a repipetter to deliver 5.00 mL of the unknown acid into a clean 150 mL beaker.  Wash down the sides of the beaker with the wash bottle.  The addition of water at this stage has no effect on the total amount of acid already present in the beaker.  Make sure you record the number of the unknown in your notebook.
  2. Add 2 drops of phenolphthalein indicator to the acid solution.  The solution should remain colorless.
  3. Add a magnetic stir bar to the beaker and place on the heating stir plate.  Adjust the stirring rate to obtain a vortex without any of the solution splattering on the sides of the beak.  Avoid spilling any of the beaker contents.  Any loss of sample would render the titration worthless.
  4. Rinse down the inside of the beaker occasionally and continue, slowly adding NaOH until the first permanent faint pink color persists for at least 30 seconds.  At this point the titration is complete (the endpoint).
  5. Read the final volume of NaOH and record to the nearest 0.01 mL.
  6. Calculate the normality of the unknown acid solution. Use the equation:  VaNa  =  VbNb to solve for Na.
  7. Repeat Steps #1-6 with 10.00 mL of your unknown acid.  Fill the burrett and read it before each titration if there is a possibility of running out before a titration is complete.
  8. Repeat Steps #1-6 with 15.00 mL of your unknown acid..
  9. The results from these three titrations should agree within 0.005 N, otherwise titrate a fourth sample and throw out the value which does not agree.  Average the results and use this value as the concentration of your UNKNOWN ACID.
  10. When your calculations are complete and you are satisfied with your results, rinse the burrett with burrett rinse (dilute HCl) and then distilled water.  Leave the stopcock open when you return the burrett to the back bench.

(Updated 9/13/13 by C.R. Snelling)