Eka Surya Online

Acid-Base Titration

Posted on: 2 Juli 2008


I. Objective

Students can acquire knowledge and ability in determining concentration of the solution by divalen acid-base titration (H2SO4) by using phenoltalein indicator (PP).

II. Theory

According to Bronsted and Lowry, an acid is a species which donates a proton, H+, and a base is a species which accepts a proton in a proton-transfer reaction.

For example, in the following acid-base reaction:

H2SO4 (aq) + 2NaOH (aq) 2H2O (l) + 2Na+ (aq) + SO42 (aq)

H2SO4 (aq) is an acid which donates a proton, H+, to the base NaOH (aq). This reaction is often called a neutralization reaction.

An acid-base titration is a procedure for determining the concentration of an acid (or a base) in a solution by measuring the volume of base (or acid) of a known concentration that completely reacts with it.

The solution of accurately known concentration is called the standard solution (titrant), it contains a definite number of gram-equivalents per liter. Standard solution is usually added from a graduated vessel called a burette. The process of adding titrant until the reaction just complete is termed a titration and the substance to be determined is titrated (analyte). The point at which the reaction is complete is called the equivalence point or the theoretical (or stoichiometric) end point. This point must be detectable by some change unmistakable to the eye and this can be done by adding an auxiliary reagent, known as an indicator which should give a clear visual change (color change) in the liquid being titrated.

In order to perform a titration procedure a reaction must fulfill the following condition.

1. It must be simple reaction, which can be expressed by a chemical reaction. The substance to be determined should react completely with the titrant in stoichiometric or equivalent proportions.

2. The reaction should be practically instantaneous or proceed with very great speed.

3. There must be a marked change in some physical or chemical property (as color change) of the solution at the equivalence point.

4. An indicator should be available which, by a change in physical properties (Color), should sharply define the end point of the reaction.

5. If no visible indicator is available for the detection of the equivalence point, the latter can often be determined by other method as potentiometer or conduct metric or spectrophotometer titration.

Titration can be used for many types of reactions:

a. Neutralization (reaction of acid with base.).

b. Precipitation reaction.

c. Oxidation – reduction reactions

d. Complex formation reactions.

III. Materials and equipments

v Materials

NaOH solution 0,1 N

H2SO4 solution x N

Phenolphthalein Indicator (PP)

v Equipments

Burette 50 mL 1 pcs

Erlenmeyer flask 250 mL 2 pcs

Drooping Pipette 2 pcs

Graduated Glass Cylinder 50 mL 2 pcs

Stirring Stick 1 pcs

Funnel 1 pcs

Volumetric Pipette 10 mL 1 pcs

Stand, retort and klem (stative) 1 set

IV. Procedure

  1. Equipments are set for titration process
  2. Burette is washed up with aquades and the tap’s leakage is checked accuracy.
  3. Burette is washed up with some titrant (NaOH 0,1 N).
  4. Titrant (NaOH 0,1 N) are entered into burette using funnel to zero numeral.
  5. 10 mL titrated (H2SO4) are entered into Erlenmeyer flask.
  6. Titrated are added phenoltalein indicator (PP) around 1-2 crop.
  7. Titration is done by drooping titrant into titrated bit by bit from burette.
  8. Titration are stopped until the titrated’s color changed.
  9. Titrant volumes are written down.
  10. 5-9 procedure are repeated 3 times.
  11. Average titrant’s volume that used are calculated.
  12. Unknown concentration of solution are calculated.

V. Observe result

Observe table of H2SO4 solution.



Titrate Volume (H2SO4)

Titrant Volume (NaOH 0,1 N)


10 mL

9,6 mL


10 mL

9,1 mL


10 mL

9,4 mL

Average Volume

10 mL

9,4 mL

VI. Discussion

The experiment that we did is name alkalimetri titration because it is done to determine concentration of acid solution (H2SO4) by using base standard solution (NaOH 0,1 M). The titrated in this experiment is H2SO4 x M around 10 mL. PP indicator is added to differentiate acid solution with base solution so that by adding 2 drops PP indicator in H2SO4 x M make the H2SO4 x M solution still uncolored. Then in the titrated is titrationed by NaOH 0,1 M solution. This procedure is repeated 3 times. End point indication of titration is the changed of titrated color from uncolored into red when the volume of NaOH 0,1 M solution is used around 9,6 mL;9,1 mL; and 9,4 mL each other.

Next, will be done a calculation of concentration based on the theory and the experiment’s result.

For H2SO4 Tirtation with NaOH as titrant :

Reaction between H2SO4 and NaOH is:

Reaction: H2SO4 (aq) + 2NaOH(aq) Na2SO4 (aq) + 2H2O(l)

From reaction above we can know that the ratio of mol between H2SO4 and NaOH is 1:2. So, to calculate concentration of H2SO4 solution based on the experiment’s result, we can use the equation below:

mmol ion H+ = mmol ion OH

Ma . x . Va = Mb . n . Vb

Information: Ma = acid molarity HXA

Mb = strong base molarity L(OH)n

Va = acid volume

Vb = base volume

X = acid valence

n = base valence

According to The Theory :

To determine NaOH solution’s volume that needed:

Titrant (NaOH) that used: 0,1 N same as 0,1 M (because it is monovalen).

Titrated (H2SO4) that used: 0,1 N same as 0,05 M (because it is divalen). This concentration is acquired based on the calculation H2SO4 concentration that made by group 2 and laboratory assistant.

Ma . x . va = Mb . n . vb

0.05 M . 2 .10 mL = 0,1 M . 1 . vb

1 mL = 0,1 . vb

Vb = 1 mL / 0,1

= 10 mL

So, to calculate H2SO4 0,05 M need 10 mL NaOH 0,1 N (based on the theory)

Based on the experiment’s result :

Experiment I (needed NaOH around 0,96 mL)

Ma . x . va = Mb . n . vb

Ma . 2 . 10 ml = 0,1 M . 1 . 9,6 ml

Ma . 20 ml = 0,96 mmol

Ma = 0,96 mmol/20 ml

= 0,048 M

Experiment II (needed NaOH around 0,91 mL)

Ma . x . va = Mb . n . vb

Ma . 2 . 10 ml = 0,1 M . 1 . 9,1 ml

Ma . 20 ml = 0,91 mmol

Ma = 0,91 mmol/20 ml

= 0,0455 M

Experiment III (needed NaOH around 0,94 mL)

Ma . x . va = Mb . n . vb

Ma . 2 . 10 ml = 0,1 M . 1 . 9,4 ml

Ma . 20 ml = 0,94 mmol

Ma = 0,94 mmol/20 ml

= 0,047 M

Calculate using average volume:

Ma . x . va = Mb . n . vb

Ma . 2 . 10 ml = 0,1 M . 1 . 9,4 ml

Ma . 20 ml = 0,94 mmol

Ma = 0,94 mmol/20 ml

= 0,047 M

Calculate using average molarity:

Average morality of H2SO4 = (0,048 M + 0,0455 M + 0,047 M)/3

= 0,0468 M = 0,047 M

So from the whole of experiment we get concentration of H2SO4 solution is 0,047 M

The difference of molarity:

The difference = M – M experiment

= 0,05 M – 0,047 M

= 0,003 M

Based on the calculation from the experiment’s result and the real concentration of H2SO4, we can know that the relative mistake from the titration is:

The relative mistake : (0,05 M – 0,047 M) x 100 % = 6 %

0,05 M

From the difference above, there are view mistakes of concentration between concentration H2SO4 that used and concentration that acquired from the result of calculation NaOH volume that used. It is caused by:

1. Accurate less of titration process.

2. There is a leakage on titration tool.

3. The titration tool is not working well, we can see from the accurateness number.

4. Accurate less of solution’s making process.

5. Accurate less in noting the color change of indicator.

VII. Conclusion

1. Titration is used to determine solution’s concentration with reacting some solution’s volume into some other solution which its concentration has known before.

2. The end titration’s point is marked by the color changed of the titrated.

3. Adding PP indicator is used to differentiate whether it is acid or base solution.

4. Thus, from the titration’s result, we acquire H2SO4 concentration is 0,047 M, on the other hand from the liquiding’s result we acquire H2SO4 concentration is 0,05 M. It happens may be because of accurate less of liquiding H2SO4. Beside that it is also caused by accurate less in titration process..

VIII. Booklist

1. Sunardi.2007. Buku Kimia untuk SMA/MA kelas XI. Yrama Widya : Bandung

2. Brady, JE.1999. Kimia Universitas Azas dan Struktur. Jilid 1 edisi ke-5, alih bahasa Sukmariah Maun, Kamanti Anas, Tilda S. Sally. Jakarta : Binarupa Aksara

3. Simamora, Maruli, dkk. 2004. Buku Ajar Kimia Dasar 2. Singaraja : IKIP Negeri Singaraja.

4. Subagia, I Wayan dan Suhemi Sya’ban. 2005. Materi Praktikum Kimia Dasar 2. Singaraja : IKIP Negeri Singaraja.

5. Syukri. 1999. Kimia Dasar II. Bandung : ITB


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