SUBDOMAIN 204.3
TITRATION OF STRONG AND WEAK ACIDS
SCA4/5 Task Late Nite Lab
By Nienke Adamse
12/26/10
ABSTRACT
This experiment used the titration technique to determine the
unknown concentration of a strong acid (HCl (aq)) and the weak
(acetic)acid ( C2H4O2 (aq)).
The base (NaOH (aq)) was used in the quantitative neutralization
reaction.

The titration
technique showed that the unknown concentrations
of the strong acid as well as the weak acid could be determined.
The concentration of HCl (aq) was 0.6 mol/L
and that of C2H4O2
(aq) was .16 mol/L.
The analytical technique of titration, using a quantitative
 of a weak
neutralization reaction to determine the concentration
and a strong acid was fairly accurate; comparing the accepted Ka
value with the experimental Ka value, showed a 3.5% error.
EXPERIMENT
For this experiment the following tools, equipment and chemicals
were used:
Lab tools:
pH Meter
Lab glassware :
Burette, Erlenmeyer flask
Lab chemicals :
H2O,C2H4O2,HCl,NaOH,C20H14O4
4
Lab background :

Titration is one of the most useful techniques for determining the
concentration of an unknown solution, and it is often used with
acidic and basic solutions. The equivalence point is reached
when a stoichiometric number of moles are added to the
unknown solution to neutralize the acid or base.
The determination of the equivalence point will be detected by
the following two methods:
* an indicator - a change in the color an indicator added to the
solution.
* pH meter - a potentiometric titration, with a sudden change in
the pH of the solution as it transitions from being acidic to basic
or vice versa.
Procedure1 for Late Nite (Laboratory Simulations)
Lab:
1. Take a clean Erlenmeyer flask from the Glassware shelf tab
and place it on the workbench.
2. Open the Chemical Shelf and add 25 mL HCl (of unknown
concentration) to the flask.
3. Add 2 drops of phenolphthalein to the flask.
4. Take a burette from the Glassware shelf and place it on the
workbench.
5. Fill the burette with 50 mL of 1M NaOH solution. Record this
initial volume.
6. Drag the flask to the lower half of the burette so that the two
will be connected. (see figure 1)
7. Open the Data window and click on the flask. Click the
Pushpin icon in the blue bar of the Data window to lock its
display to the contents of the flask.
8. Take a pH meter from the Tools shelf and drop it on the flask.
Record the initial pH of the solution.
9. Open the Properties window and click on the burette. Enter "1"
in the amount to add, and click the flow button to drip 1 mL of
NaOH into the flask. Record the pH of the solution in the flask.
10. Continue to add NaOH in 1 mL increments. Record the pH for
each milliliter added.
11. The pink color will appear in flask all at once when the
endpoint is either reached or crossed. Record the burette volume
and pH at which this occurs.
12. Continue to add 5 more increments of 1 mL and record the
pH at each point.
13. Detach the flask from the burette and drag it to the recycling
chute.
14. Take a new flask from the Lab ware Shelf, under the
Glassware Tab, and place it on the workbench.
15. Add 25 mL of HCl and 2 drops of phenolphthalein to the
flask.
16. Refill the burette to 50 mL NaOH.
17. Based on the results of the previous titration, add enough
NaOH solution - all at once - to get to 1 mL BEFORE the
endpoint.
18. Using the NaOH from the Chemical Shelf, add .0001 mL to
the solution until you reach the endpoint. Record the pH and
volume until the endpoint is reached and several drops after it is
reached.
Procedure 2
Repeat all steps but use C2H4O2 (aq) instead of HCl(aq)
It is important follow the exact procedure because we need to find

out what the accurate
amount of NaOH added to the acids is, in
order to find the equivalence points and thus the concentrations
of the unknown acids. The first titration is a rough titration to
find out in which area we can make the second titration with
the smaller increments of added base. If we start with the smaller
increments it might take a long time before we see the change of
color. Using two titrations narrows down the area of volume
added base and thus makes it possible to add smaller increments
of NaOH.
Figure 1
OBSERVATION
First titration HCl (1 ml increments of NaOH)
ml added NaOH
pH
Color solution
2
0.32
no
3
0.37
no
4
0.42
no
5
0.48
no
6
0.54
no
7
0.60
no
8
0.67
no
9
0.75
no
10
0.84
no
11
0.95
no
12
1.09
no
13
1.28
no
14
1.59
no
15
7.00
no
16
12.39
pink
17
12.68
pink
18
12.84
pink
19
12.96
pink
20
13.05
pink
21
13.12
pink
When adding 16 ml of NaOH, the solution turned pink and the
pH jumped to 7. In order to find a more accurate point at which
the solution becomes pink, a second titration is necessary with
smaller increments of NaOH added.
End point: 16 ml added NaOH , pH of 12.39
Second titration HCl (0.0001ml increments of NaOH)
ml NaOH
pH
color
15.0001
8.40
no
15.0002
8.70
no
15.0003
8.88
pink
End point: 15.0003 ml added NaOH, pH of 8.88
Half equivalence point: 7.5 ml, pH of about 0.635
First titration C2H4O2
(1ml increments of NaOH)
ml NaOH

0
pH
color
2.78
no
1
4.28
no
2
4.76
no
3
5.23
no
4
8.95
pink
5
12.52
no
6
12.81
No
7
12.97
No
8
13.08
No
9
13.17
No
End point: 4 ml added NaOH, pH of 8.95
Second titration C2H4O2 (.05 ml increments of NaOH) starting 1ml
before the end point.
ml NaOH 
pH
color
3.05
5.26
No
3.10
5.29
No
3.15
5.33
No
3.20
5.36
No
3.25
5.39
No
3.30
5.43
No
3.35
5.47
No
3.40
5.51
No
3.45
5.55
No
3.50
5.60
No
3.55
5.65
No
3.60
5.71
No
3.65
5.78
No
3.70
5.85
No
3.75
5.93
No
3.80
6.04
No
3.85
6.17
No
3.90
6.35
No
3.95
6.65
No
4.00
6.95
pink
4.05
11.23
Pink
4.10
11.53
Pink
4.15
11.71
pink
End point: 4.00 ml added NaOH, pH of 6.95
Half equivalence point: 2.05 ml NaOH added, pH of about 4.76
(did not measure exact pH)
ANALYSIS
The equivalence point from this graph = 15 ml added NaOH with a pH of 7
The equivalence point of this graph= 4.5 ml NaOH , pH 9.5
The difference between end-point and equivalence point is, if
one is trying to get precise, that the equivalence point is the exact
point where the acid and the base in the reaction have canceled
each other out and when acid and base have the same number of
moles (it is almost impossible to know when this point has been
reached with the use of a color indicator). The end-point is the
perceptual event at which the indicator actually does change,
and you decide you are "done", and you actually read the
quantity of measuring reagent you have added. Such as when the
solution finally changes color and does not change back right
away. The end-point may be slightly different from the true
equivalence point, due to the indicator you use and the slowness
of pH changing with reagent addition
In the Acid-Base titration of NaOH and HCl, using a
phenolphthalein indicato,r there is an equivalent amount (same
number of moles) of HCl and NaOH together in the flask when the
pH is near pH=7. Nothing visible happens right then, because the
phenolphthalein indicator changes color closer to pH=10. If
adding base from the burette to acid in the flask, you might need
to add an extra drop of NaOH solution to make the indicator
change color. Then the volume you have added would be "wrong"
by about 1 drop in 100ml. This means an error of about 1%.
When titrating weak acids and bases, it matters a lot more.
Because the pH changes more slowly, the error can be bigger than
20%.
When titrating a monoprotic acid with NaOH, the molar ratio for
the reaction is 1:1 and all of the following are true at the
equivalence point:
(a) Moles of acid in flask = moles of base added from the burette
(b) (molarity of acid) x (volume of acid) = (molarity of the base) x
(volume of added base)
(c) Molarity of the acid = (molarity of the base) x (added volume
of the base) / (volume of the acid)
A weak acid, such as CH3COOH (acetic acid), exists in
equilibrium with water, and only some of the acid donates the
protons to form hydronium ions:
CH3COOH + H2O <-> CH3COO- + H3O+
This also explains why the amount of NaOH added to the weak
acid to reach equilibrium was about half of the volume of NaOH
that needed to be added to the strong acid to reach equilibrium.
The equilibrium expression for acetic acid is:
Ka= [H3O+][CH3COO-] / [CH3COOH]
While titrating the acid with NaOH, the reaction is:
NaOH(aq) + CH3COOH(aq) -> H2O(l) + Na+(aq) + CH3COO-(aq)
During the titration there is a point where EXACTLY HALF of the
acid has reacted with NaOH. The concentration of the acetate ion
is exactly the concentration of the unreacted acid at this point
and it is at the HALF EQUIVALENCE point. From the equilibrium
expression above, we have:
Ka = [H3O+] (at the half equivalence point, where [CH3COO-] =
[CH3COOH])
Taking the log of both sides at this point, we have:
pKa = pH, with pKa defined as pKa = -log(Ka).
The acid's dissociation constant, Ka, then, can be found from the
pH of the solution at the half equivalence point. On the titration
curve, this is the point where the number of moles of added
NaOH is equal to half of the number of moles added to reach the
equivalence point.
The experimental Ka value of HCl = 1/-logpH (at the half
equivalence point) therefore is 1/-log0.64= 0.23
The experimental Ka value of CH3COOH = 1/-logpH (at the half
equivalence point)= 1/-log4.76 = 0.0000174
Errors
When calculating the percent error of the experimental Ka values
of HCl and [CH3COOH I looked up the accepted values in the
chemistry handbook. The percent error between the experimental
and accepted values are, according to %error= experimental Ka –
acceptedKa/acceptedKa x100%
For the HCl Ka value the percent error therefore becomes: 0.231300000/1300000= 100%
The Ka value for a strong acid is a really big number. It does not
matter in an aqueous solution, since everything will be converted
into hydronium and chloride ions:
HCl (aq) ==> H+ + ClThe acid is 100% dissociated, so all of HCl will be converted into
H+ and Cl-. If you are dealing with aqueous solutions, talking
about Ka values for strong acids is irrelevant.
Any strong acid when dissolved in water is converted into H+,
since H+ is the strongest acid that can exist in water. That is
called the leveling effect.
In essence, using this approach, you can say that the Ka of any
strong acid approaches infinity, any number divided by zero is
in essence infinity.
For the [CH3COOH Ka value the percent error becomes:
0.0000174-0.000018/0.000018= 0.035x100% = 3.5% error
The use of the interactive “Nite Lab” simulation equipment
eliminates any major human errors, especially when the
experiment can be done again and again following the exact
same procedures.
DISCUSSION AND CONCLUSION
The titration method to find the unknown concentration of a
substance is a fairly accurate method. A conjugated acid-base
pair consists out of two substances related to each other by the
donating and accepting of a single proton. The equilibrium
expression for Ka (the acid dissociation constant) in simple form
of a conjugated acid-base pair is:
(When H represents the acid and A the base:)
HA (aq) -----------------> H+ (aq) + A- (aq)
When titrating a monoprotic acid with the base NaOH, the precise
equivalence point (where the equivalence point of a chemical
reaction occurs when the amount of tyrant (NaOH) added is
stoichiometrically equal to the amount of analyte (acid) present
in the sample: the smallest amount of titrant that is sufficient to
fully neutralize or react with the analyte) can be determined by
a titration curve, a graph where the change of pH of the solution
is plotted versus the number of moles (or milliliters of) base
added. When working with a color indicator, the equivalence
point is the point where the neutralization is indicated by a
change of color.
At the equivalence point the moles of acid = the moles of base
added. With equation M1V1= M2V2, the unknown concentration
of the acid can be calculated.
The molarity of the HCl in the first experiment therefore was
calculated as 1M (given molarity of NaOH) x volume added NaOH
(15 ml)/ volume HCl (25 ml)=
0.6 mol/L or 0.6 M
The molarity of the CH3COOH in the second experiment was
calculated as 1Mx 4.5 ml/25 ml= 0.18 mol/L or 0.18 M
The difference in acid strength does not influence the found
concentrations with the titration method, however the differences
in found Ka values show that a weak acid has a strong conjugate
base which means that not all acid is dissociated in the aqueous
solution, which explains that only a small amount of titrant
(base) is needed to reach the equivalence point.
REFERENCES
Late Nite Simulation Labs
http://66.246.168.132/csp/lab/ViewLabReport
pH Titration curves
http://www.chemguide.co.uk/physical/acidbaseeqia/phcurves.htm
l
Zumdahl, S.S and Zumdahl, S.A (2009) Chemistry. Seventh
Edition
Boston, MA: Houghton Mifflin Company