This document describes an experiment to prepare and standardize a 0.1 M sodium hydroxide (NaOH) solution through titration. The student measured and diluted a 3 M NaOH solution to make an approximate 0.1 M NaOH solution. This solution was then standardized through titration against a 0.104 M sulfamic acid standard solution. The student determined the concentration of the NaOH solution to be 0.092 M on average, with a percent error of -9% due to imprecise measurement of the initial 3 M NaOH solution. The student concluded the procedure was successful in standardizing the NaOH solution and learning the titration technique.

1. CALIFORNIA STATE UNIVERSITY, LOS ANGELES
Department of Chemistry
Chemistry 103 / Section 03- 94356
Prepare and standardize a 0.1 M NaOH solutions
Prepared by: Rodney Pujada
Performance Date: Tuesday, September 29, 2015
Submission Due: Tuesday, October 6, 2015
Professor: Dr. Xin Wen
Tuesday and Thursday: 1 pm. – 3:50 p.m.
September 2015

2. Experiment No 1: Prepare and standardize a 0.1 M NaOH solutions
I. PURPOSE
Performing this lab successfully will allow us to learn an important lab technique called titration.
Titration is the technique to find the concentration of unknown solution by the formula Ma x Va =
Mb x Vb to determine the exact concentration of sodium hydroxide solution.
II. PRINCIPLES
Titration is a laboratory technique that can be used to determine the concentration of
certain solutions by chemical reaction. A standard solution of known concentration is “titrated
against” (reacted with) a solution of unknown concentration. An indicator can signal the
completion of the reaction (by color change) and the concentration of the unknown solution can be
determined. Any chemicals that react in solution can be titrated with each other. Since acids and
bases are usually found in solution, they are commonly involved in titrations. Titrations involving a
strong acid or a strong base involve the neutralization reaction between hydrogen ions (or
hydronium) and hydroxide ions. These ions combine to form the neutral water molecule:
H+ + OH- ---> H2O or H3O + + OH- ---> 2 H2O
An indicator is any substance in solution that changes its color as it reacts with either an
acid or a base. Selecting the proper indicator is important because each indicator changes its color
over a particular range of pH values.
This experiment use titrimetric analysis to determine the concentration of unknown
sample. It involves preparation of a solution that has the approximate concentration desired
(NaOH), determination of the concentration by direct titration against a primary standard, and a
test of the accuracy of your determined concentration by comparison with a known standard.
After the NaOH standarzided, we can calculate the concentration of the sample by the
principle of relation of moles that react with the unknown acid, and calculate its molarity, using:
Moles solute before dilution = moles solute after dilution
……………………………. Formula No1
III. EXPERIMENTAL PROCEDURES
3.1 Material
20-mL graduated cylinder,50 mL Buret, three 250 mL Erlenmeyer
flask, buret clamp, wash bottle, 1L plastic bottle, funnel, 50 mL and
500 mL beaker, phenolphthalein NaOH (known concentration),
distilled water, 3M NaOH concentrated, and sulfanic acid standard
solution (known concentration).

3. 3.2 Procedure
3.2.1 Preparation of 1 M Sodium Hydroxide Solution solution.
a. Measure 16.7 mL of 3M NaOH with 20mL graduated cylinder.
b. Add small amount of water and pour it into a volumetric flask of
250 mL.
c. Fill the volumetric flask until the mark of 250 mL.
d. Pour this volume into a 1 liter plastic bottle.
e. Fill the volumetric flask with distilled water to get 500 mL into the plastic bottle.
3.2.2 Calculate the Molarity of Sodium Hydroxide Solution by titration
a) Clean and rinse the buret with distilled water.
b) Precondition the buret by rinsing it with the NaOH solution two or three times, and
fill the buret with a funnel.
c) Fill the buret with NaOH to exact measurement.
d) Remove all gas bubbles in the liquid column of the buret and attach a buret clamp.
e) Read the initial volume of NaOH by recording the position of mark that line up with
the meniscus of the liquid column.
f) Prepare three Erlenmeyers.
g) Take 25 mL of sulfamic acid standard with known
concentration into three erlemeyers.
h) Add four drops of indicator (phenolphthalein) into the
Erlenmeyer flask of sulfamic acid standard.
i) Place the Erlenmeyer flask under the buret; open the
stopcock slowly to add the NaOH solution. The end-point is
near when the solution changes the color pink. At this
point, open the stopcock slowly, adding drop by drop.
j) Record the final volume Vf NaOH when the solution turn to
pink.
IV Data and Calculation
4.1 Calculate the amount of 0.1 M NaOH by dilution of 3M NaOH
Calculate the amount of 0.1 M NaOH by dilute 3 M of NaOH, using the following the
formula…..
Titration Data Run 1 Run 2 Run 3
Molaridad of sulfamic acid standard (M) 0.104 +0.47% 0.104 +0.47% 0.104 +0.47%
Initial buret reading Vi NaOH (mL) 0.0 0.0 0.0
Final buret reading Vf NaOH(mL) 28.6 28.0 28.6
Volume of NaOH used (mL) 28.6 28.0 28.6

4. ……………………………. Formula No1
Data:
M1 = 3 M of NaOH known.
V1 = ??
M2 = 0.1 M of NaOH that we expected.
V2 = 500 mL of NaOH that we expected.
Using the formula No1 to calculate V1 =?
Evalute V1:
V1 = ( M2* V2) / M1
Replace the data
V1 = ( 0.1 M *500 mL) / (3M)
V1 = volume of NaOH diluted= 16.7 mL
4.2 Standardization of the 0.1M NaOH solution.
……………………………. Formula No1
DATA for Run 1
M1 = 0.104 M of sulfamic acid standard
V1 = 25 mL of sulfamic acid standard
M2 = ??? of NaOH that we expected in the plastic bottle.
V2 = 500 mL of NaOH that we expected in the plastic bottle.
Using the formula No 1 to evaluate M1 =?
M2 = ( M1* V1) / V2
Replace the data
M2 = ( 0.104 M * 25mL) / (28.6)
M2 = 0.091 M of NaOH in the plastic bottle
Titration Data Run 1 Run 2 Run 3
Molaridad of sulfamic acid standard (M) 0.104 +0.47% 0.104 +0.47% 0.104 +0.47%
Initial buret reading Vi NaOH (mL) 0.0 0.0 0.0
Final buret reading Vf NaOH(mL) 28.6 28.0 28.6
Volume of NaOH used (mL) 28.6 28.0 28.6
Molarity NaOH bottle (M) 0.091 0.093 0.091

5. 4.3 Calculate the percent error %
Percent error = ( Vpractical — Vtheoric) x 100 % ………………. Formula No 2
V theoric
Data:
Molarity NaOH (approx) = V prac = 0.1 M
Molarity NaOH (standard) = V teoric = 0.087 M
Using the formula No 2 to evaluate percent of error.
Percent error = (0.091 - 0.1) x 100 % = -9 %
0.1
Percent error = -9 %
V Results and Discussion
 In this lab we calculate the concentration of our NaOH diluted solution is 0.092 M .
 By definition one equivalent (or equivalent weight) of a substance is the amount of that
substance which supplies or consumes one mol of reactive species. In acid-base chemistry
the reactive species is the hydrogen ion (H1+) while in oxidation-reduction chemistry the
reactive species is the electron.
 We recommended follow the correct procedure and use
Erlenmeyer flask for the titration that are appropriate
material. Also, we must follow the correct technique of
titration because these factor alter the end point of the
titration.
 Using the titration technique correctly is important. A right
handed person should titrate with the left hand, swirling
the flask with the right hand and agitate strongly to avoid to
past the end point of the titration as occurring in our run1
and run3.
Titration Data of 0.1 M NaOH
Run 1 Run 2 Run 3
Molarity NaOH bottle (M) 0.091 0.093 0.091
Molarity NaOH bottle Average (M) 0.092

6. VI Conclusions
 We conclude that the concentration of sodium hydroxide is 0.092 M.
 Our percent of error is consequence to measuring 16.7 mL of 3M NaOH with 20mL
graduated cylinder instead to use a volumetric pipet or buret with 3M NaOH solution.
 If we believe that run #2 should be omitted because we know that we accidentally
overshot the equivalence point and our titration flask was little bright pink; we can report
0.093 M NaOH with less percent of error.
VII References
 D.C. Harris, Quantitative Chemical Analysis (7th ed., W. H. Freeman, NY, 2007) pp. 121-124,
221–218,
 Skoog, D. A.; West, D. M. Fundamentals of Analytical Chemistry; Holt, Rinehart and Winston:
New York, 1963; pp 341-351.
 Sweeder, R. D.; Jeffery, K. A.; A comprehensive general chemistry demonstration. J. Chem. Ed.,
2013, 90, 96-98. doi:10.1021/ed300367y