東海大學九十二學年度第一學期
畜產品檢查學上課參考講義
Primary Standard Base
It has been shown that tris(hydroxymethyl)aminomethane, also known as THAM, can serve as
an excellent primary standard base. THAM has a molecular weight of 121.14 g/mol,
corresponding to the formula:
and can therefore be regarded as a derivative of ammonia, with a basic nitrogen atom. It has
been demonstrated that THAM is non-hygroscopic, does not pick up carbon dioxide from air,
is stable both as the solid and in aqueous solution, can be prepared in very pure form, and can
be dried at 100-103°C without decomposition. THAM reacts quickly and stoichiometrically
with hydronium ion.
THAM, which has been oven dried, will be provided in desiccators. Place 2.5 g of THAM into
a clean, dry, weighing bottle. By difference, accurately (use analytical balance) weigh four 0.5 g
samples of THAM into 4 labeled Erlenmeyer flasks. You must know exactly (±0.1 mg) how
much THAM is in each Erlenmeyer but they will not all be exactly the same. Use the exact
values for later calculation.
Standardization of HCl Solution - Provide a labeled clean 2 liter polyethylene storage
bottle to your TA. The bottle need not be completely dry since the TA will rinse it with reagent.
You will be given approximately 1 liter of HCl solution. Record the solution letters provided
by the TA, in your lab notebook, on the reagent bottle and on the report form. You must now
standardize this HCl solution. Do not waste solution or allow it to become contaminated as
this is your only supply of HCl reagent.
Add 50 mL of distilled water to each Erlenmeyer. Swirl to dissolve the THAM. Add 2-4 drops
of bromocresol green indicator. Titrate with your HCl standard solution. Bromocresol green
changes from blue, through green, to yellow. The proper stopping point is at the intermediate
green color -- this corresponds to a pH of 4.7. Repeat the titration with at least three more 50
mL aliquots or until results are reproducible. Always read the buret volume to the hundredth
place. Remember, the THAM samples are not the same size so the volume of the HCl will not
be identical, but the concentrations of HCl calculated should be reproducible.
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東海大學九十二學年度第一學期
畜產品檢查學上課參考講義
Primary Standard Acid
Titration of the Weak Acid Potassium Hydrogen Phthalate (KHP)
INTRODUCTION
Materials generally considered to possess acidic and/or basic properties are widely distributed in
nature and range from simple inorganic materials through organic and biological molecules of
great complexity. Since acid-base equilibrium is a general phenomenon, it is advantageous to use
it as an analytical tool. Acid-base titrations are conducted by adding a known amount of one
reagent (either acid or base) to a sample of the opposite nature (either base or acid) until all
available ionizable hydrogen ions of one solution have reacted with all available hydroxide ions
in the other to form water. This experiment is designed both to demonstrate the techniques of
titrimetric analysis and to explore the relationships involving stoichiometry, pH, and acid-base
equilibria.
The general reaction involved in neutralization titrations can be depicted as follows:
HA + MOH H2O + MA
(acid) (base)
In this experiment you will determine the amount of acid present by titration with the strong base
NaOH. Since it is hard to prepare a NaOH solution of accurately known concentration directly
from the solid, you will need to standardize your NaOH solution against a precisely weighed
amount of standard acid. The acid used is the weak monoprotic acid, potassium hydrogen. The
reaction of KHP with sodium hydroxide is shown below.
KHP Neutralization
REAGENTS AND APPARATUS
Potassium Hydrogen Phthalate (high purity)
Phenolphthalein indicator solution
50% NaOH solution
50.00 mL buret
Erlenmeyer flasks
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東海大學九十二學年度第一學期
畜產品檢查學上課參考講義
PROCEDURE
*** One week in advance*** You should dry the standard KHP and solid unknown one week in
advance. It may also be possible to boil the water in advance.
PART A -- PREPARATION OF 0.1 N SODIUM HYDROXIDE
1. Boil about 1.5 liter of distilled water for 5 minutes to remove carbon dioxide. This will allow you to
make enough NaOH solution to do several experiments. Allow the water to cool covered with a watch
glass; transfer while warm (40C) to your large plastic bottle.
2. Using a transfer pipet and a rubber bulb, transfer to this bottle the volume required of a clear saturated,
50%(w/w), solution of NaOH to make the desired concentration of 0.1 N. The density of the
50%(w/w) NaOH is 1.50 g/mL. Mix thoroughly and keep covered (How can atmospheric carbon
dioxide lower the concentration of OH-?). Keep this solution even after completion of this experiment,
as it will be required for other experiments.
PART B -- STANDARDIZATION OF NaOH SOLUTION
1. Dry 4 to 5 g of primary-standard KHP in a weighing bottle at 110C for 2 hours. At the same time dry
your unknown sample in a separate weighing bottle. Allow to cool 30 minutes in a desiccator before
weighing (why should the desiccator lid be left slightly "cracked" during the cool-down period?) Save
any unused KHP for use later in the semester.
2. Accurately weigh four samples of the pure, dry KHP standard. Each sample should weigh between 0.6
g - 0.7 g.
3. Dissolve each KHP sample in approximately 100 mL of distilled water in a 250 mL or 500 mL
Erlenmeyer flask (why is the exact volume of water unimportant?)
4. Add a few drops of phenolphthalein solution.
5. Titrate with the NaOH solution that you are standardizing. The endpoint is faint pink.
6. Calculate the formality of the NaOH for each titration and find the average formality.
The sample containing potassium hydrogen phthalate (potassium biphthalate or potassium
acid phthalate), is neutralized by NaOH solution, giving the corresponding phthalate salt.
Since phthalic acid is a moderately weak dibasic acid, the pH of the sample solution will be
approximately 4.2. The pH at the equivalence point, when only phthalate ion is present (no
biphthalate left), is about 9. The indicator phenolphthalein (see the section on indicators in
Skoog) at pH 10 is in the ionic form, which is deep red. At below pH 8 it is in the molecular
form, which is colorless. Since the pH rises fairly abruptly as the last traces of biphthalate are
neutralized, we can use the first appearance of the reddish color as an accurate indication of
the quantitative completion of the neutralization.
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東海大學九十二學年度第一學期
畜產品檢查學上課參考講義
Carbonate, bicarbonate, and CO2 interfere in several ways. Primarily, since carbon dioxide
reacts with water to form acidic solutions, the amount of acidic protons in the solution is
affected. Boiling water drives out dissolved gases, including carbon dioxide. This is the major
reason for using boiled, deionized water when preparing the NaOH an sample solutions.
NaOH solutions, once prepared, have to be protected from the CO2 in the air if they are to be
stored for more than a few days.
PART 1: PREPARATION OF STANDARD SODIUM HYDROXIDE
Reagent grade NaOH contains absorbed water and carbonate. This means that no matter
how careful you are, it is impossible to make up a solution of NaOH of known concentration
simply by weighing out some solid NaOH as accurately as possible and then dissolving it in
water to a known volume in a volumetric flask. Therefore, the molarity of an aqueous solution of
NaOH must be determined via a standardization titration. In a typical procedure, an accurately
weighed sample of potassium hydrogen phthalate (KHP, a primary standard) is titrated with a
solution of NaOH of unknown molarity. The reaction of KHP with NaOH is shown below.
It is important to keep the titration solutions and standard NaOH free of carbon dioxide,
which dissolves in water to form carbonic acid. CO2 can be virtually eliminated from the solution
by first boiling the water to be used in preparing the standard NaOH solution and allowing it to
cool to room temperature in a tightly sealed polyethylene bottle. When performing all titrations,
care should be taken to prevent CO2 from re-entering solutions by covering them with parafilm.
1. Instead of preparing the standard NaOH solution by dissolving some freshly weighed
sodium hydroxide pellets in water, we will dilute a 50% (wt/wt) aqueous NaOH solution
prepared in advance by the laboratory staff. Sodium carbonate is insoluble in this
solution and precipitates (this is the solid observed at the bottom of the bottle). The
solution is stored in a tightly sealed polyethylene bottle and handled gently to avoid
stirring the precipitate when supernate is taken. The density is close to 1.50 g/mL.
CAUTION: Do not spill any of this solution on your skin. If you do, immediately rinse the
affected area with plenty of water and then notify your instructor.
IF YOU SPILL THE CONCENTRATED NaOH SOLUTION (OR ANY OTHER CHEMICAL)
IN THE HOOD OR ANYWHERE ELSE IN THE LAB,
CLEAN IT UP IMMEDIATELY!!!
2. Primary standard-grade potassium hydrogen phthalate (KHP) should be dried for 1 hr at
110C and cooled and stored in a desiccator.
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東海大學九十二學年度第一學期
畜產品檢查學上課參考講義
3. Boil 1 L of water for 5 min to expel CO2. Pour the water into a polyethylene bottle, which
should be tightly capped whenever possible. Calculate the volume of aqueous 50%
NaOH needed to produce 1 L of ~0.1 M NaOH. Use a graduated cylinder to transfer this
much concentrated NaOH to the bottle of water. (Why is it OK to use a graduated
cylinder for this step?) Mix well and allow the solution to cool to room temperature
(preferably overnight).
4. Weigh four samples of solid potassium hydrogen phthalate and dissolve each in ~25 mL
of deionized water in a 125-mL flask. Each sample should contain enough solid to react
with ~25 mL of 0.1 M NaOH. Add 3 drops of phenolphthalein indicator (Table 12-4) to
each, and titrate one of them rapidly to find the approximate endpoint. The buret should
have a loosely fitted cap to minimize entry of CO2. Before you begin to titrate, be sure
that there are no air bubbles in the buret.
5. Calculate the volume of NaOH required for each of the other three samples and titrate
them carefully. During each titration, you should periodically tilt and rotate the flask to
wash all liquid from the walls into the bulk solution. When very near the end, you should
deliver less than one drop of titrant at a time. To do this, carefully suspend a fraction of
a drop from the buret tip, touch it to the inside wall of the flask, wash it into the bulk
solution by careful tilting, and swirl the solution. The end point is the first appearance of
faint pink color that persists for 15 s. (The color will slowly fade as CO2 from the air
dissolves in the solution.) If you have any questions about titration technique, be sure to
ask your instructor for help.
CALCULATIONS FOR PART 1
Calculate the average molarity, the standard deviation, and the relative standard deviation
(standard deviation / mean). The relative standard deviation should be <0.5%. If the relative
standard deviation is larger than this value you may be required to perform an additional
titration.
Accuracy and Precision
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東海大學九十二學年度第一學期
畜產品檢查學上課參考講義
Stoichiometry of Acid-Base Titrations
In an acid-base titration, we slowly add the titrant strong acid or strong base until the equivalence point
is reachedas indicated in the previous section. The equivalence point is that point at which the number of
moles of acid or base added as titrant is exactly equivalent to the number of moles of acid or base
present originally in the other solution in accordance with the stoichiometric reaction.
Example. The equivalence point for the titration of 50.00 mL of 0.100 molar HCl with 0.200 molar
NaOH could be calculated as follows: 50 mL x 0.1 mol/L = 5.0 mmol HCl. The titration uses the
stoichiometric reaction HCl + NaOH --> NaCl + H2O, which could just as accurately be written as H3O+
+ OH- --> 2H2O. Since the reaction is a 1:1 reaction, 5.00 mmol of HCl are equivalent to 5.00 mmol
NaOH. The volume of NaOH required can be calculated: 5.0 mmol NaOH = 0.2 mol/L x V mL, V =
5.00/0.200 = 25.00 mL NaOH
Example. In the titration of H2SO4, sulfuric acid, the reaction requires 2 moles of NaOH per mole of
H2SO4. A complete titration of 50.00 mL of 0.100 molar H2SO4 would therefore require 50.00 mL of
0.200 molar NaOH rather than the 25.00 mL needed for the monoprotic acid HCl in the preceding
example.
Detecting the Equivalence Point
In acid-base titrations, there is a sharp change in pH at the equivalence point. The titration of 2.5 mmol
HCl and then that of 2.5 mmol CH3COOH with 0.1 molar NaOH (dashed curve).
The stoichiometric chemical reactions are Na+ + OH- + H3O+ + Cl- --> H2O + Na+ + Cl- and Na+ + OH+ H3O+ + CH3COO- --> H2O + Na+ + CH3COO-. The reaction stoichiometry is 1:1 in both cases. The pH
change can be detected by a pH meter, an electrochemical device whose discussion we will defer to later
sections, or by a chemical indicator. Chemical indicators are acid-base conjugate pairs whose acid
form and base form are different in color. A table of useful chemical indicators is given below.
A chemical indicator is a compound which can change color to indicate that the endpoint of a titration
has been reached. The color of an indicator chages because it is affected by the concentrations of ions in
the solution. An acid-base indicator is an acid-base conjugate pair, a weak acid-weak base system in
which the two forms have different colors. This is added, in low concentration so that it exerts
essentially no control on the pH of the system. For an indicator, the acid ionization constant Ka =
[H3O+][A-]/[HA] is usually written as Ka = [H3O+][In-]/[HIn], where HIn is the acid form of the indicator
and In- is the base form of the indicator. Where the ratio [In-]/[HIn] is one, [H3O+] = Ka and pH = pKa.
This point is in the color-change region of the indicator, usually at its center. The color-change region of
the indicator is usually +/-1 pH unit around pKa. To select a proper indicator, then, requires
determination of the pH of the solution at the equivalence point of the titration and selection of an
indicator whose pKa is as close to that pH as possible.
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東海大學九十二學年度第一學期
畜產品檢查學上課參考講義
Indicators are usually organic structures of some complexity or natural products and, like dyeing
materials, are generally known by their trivial names as in the above table. A few indicators, of which
thymol blue is an example, are polyprotic acids which can change color more than once as pH is
continuously increased. Such indicators can be used at a pH equal to either of their pKa values.
Acid - Base Indicators
Acid - base indicators (also known as pH indicators) are substances which change color with pH. They
are usually weak acids or bases.
Consider an indicator which is a weak acid, with the formula HIn. At equilibrium, the following
chemical equation is established.
HIn(aq)
+
In-(aq)
H2O(l)
acid
base
color A
color B
+ H3O+(aq)
The acid and its conjugate base have different colors. At low pH, the concentration of H3O+ is high and
so the equilibrium position lies to the left. The equilibrium solution has the color A. At high pH, the
concentration of H3O+ is low and so the equilibrium position thus lies to the right and the equilibrium
solution has color B.
Phenolphthalein is an example of an indicator which establishes this type of equilibrium in aqueous
solution:
colorless (acid)
magenta (base)
Phenolphthalein is a colorless, weak acid which dissociates in water forming magenta anions. Under
acidic conditions, the equilibrium is to the left,and the concentration of the anions is too low for the
magenta color to be observed. However, under alkaline conditions, the equilibrium is to the right, and
the concentration of the anion becomes sufficient for the magenta color to be observed.
We can apply equilibrium law to indicator equilibria - in general for a weak acid indicator:
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東海大學九十二學年度第一學期
畜產品檢查學上課參考講義
Kln is known as the indicator dissociation constant. The color of the indicator turns from color A to
color B or vice versa at its turning point. At this point:
So from the equilibrium expression:
The pH of the solution at its turning point is pKln and is the pH at which half of the indicator is in its acid
form and the other half in the form of its conjugate base.
Indicator Range
At a low pH, a weak acid indicator is almost entirely in the HIn form, the color of which predominates.
As the pH increases, the intensity of the color of HIn decreases and the equilibrium is pushed to the
right. Therefore, the intensity of the color of In- increases. An indicator is most effective if the color
change is distinct and over a small pH range. For most indicators the range is within ±1 of the pKln
value.
Table: Properties of Aqueous Acid-Base Indicators at 25oC
Indicator
pH
Range
Quantity per 10 ml
Acid
Base
Thymol Blue
1.2-2.8
1-2 drops 0.1% soln. in aq.
red
Pentamethoxy red
1.2-2.3
1 drop 0.1% soln. in 70% alc.
red-violet colorless
Tropeolin OO
1.3-3.2
1 drop 1% aq. soln.
red
2,4-Dinitrophenol
2.4-4.0
1-2 drops 0.1% soln. in 50% alc.
colorless yellow
Methyl yellow
2.9-4.0
1 drop 0.1% soln. in 90% alc.
red
yellow
Methyl orange
3.1-4.4
1 drop 0.1% aq. soln.
red
orange
Bromphenol blue
3.0-4.6
1 drop 0.1% aq. soln.
yellow
blue-violet
Tetrabromphenol blue
3.0-4.6
1 drop 0.1% aq. soln.
yellow
blue
Alizarin sodium
sulfonate
3.7-5.2
1 drop 0.1% aq. soln.
yellow
violet
-Naphthyl red
3.7-5.0
1 drop 0.1% soln. in 70% alc.
red
yellow
p-Ethoxychrysoidine
3.5-5.5
1 drop 0.1% aq. soln.
red
yellow
Bromcresol green
4.0-5.6
1 drop 0.1% aq. soln.
yellow
blue
Methyl red
4.4-6.2
1 drop 0.1% aq. soln.
red
yellow
Bromcresol purple
5.2-6.8
1 drop 0.1% aq. soln.
yellow
purple
Chlorphenol red
5.4-6.8
1 drop 0.1% aq. soln.
yellow
red
Bromphenol blue
6.2-7.6
1 drop 0.1% aq. soln.
yellow
blue
p-Nitrophenol
5.0-7.0
1-5 drops 0.1% aq. soln.
colorless yellow
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yellow
yellow
東海大學九十二學年度第一學期
畜產品檢查學上課參考講義
Azolitmin
5.0-8.0
5 drops 0.5% aq. soln.
red
blue
Phenol red
6.4-8.0
1 drop 0.1% aq. soln.
yellow
red
Neutral red
6.8-8.0
1 drop 0.1% soln. in 70% alc.
red
yellow
Rosolic acid
6.8-8.0
1 drop 0.1% soln. in 90% alc.
yellow
red
Cresol red
7.2-8.8
1 drop 0.1% aq. soln.
yellow
red
7.3-8.7
1-5 drops 0.1% soln. in 70% alc.
rose
green
Tropeolin OOO
7.6-8.9
1 drop 0.1% aq. soln.
yellow
rose-red
Thymol blue
8.0-9.6
1-5 drops 0.1% aq. soln.
yellow
blue
Phenolphthalein
8.0-10.0
1-5 drops 0.1% soln. in 70% alc.
colorless red
9.0-11.0
1-5 drops 0.1% soln. in 90% alc.
yellow
Thymolphthalein
9.4-10.6
1 drop 0.1% soln. in 90% alc.
colorless blue
Nile blue
10.1-11.1 1 drop 0.1% aq. soln.
blue
red
Alizarin yellow
10.0-12.0 1 drop 0.1% aq. soln.
yellow
lilac
Salicyl yellow
10.0-12.0 1-5 drops 0.1% soln. in 90% alc.
yellow
orange-brown
Diazo violet
10.1-12.0 1 drop 0.1% aq. soln.
yellow
violet
Tropeolin O
11.0-13.0 1 drop 0.1% aq. soln.
yellow
orange-brown
Nitramine
11.0-13.0 1-2 drops 0.1% soln in 70% alc.
colorless orange-brown
Poirrier's blue
11.0-13.0 1 drop 0.1% aq. soln.
blue
Trinitrobenzoic acid
12.0-13.4 1 drop 0.1% aq. soln.
colorless orange-red
-Naphtholphthalein
-Naphtholbenzein
9
blue
violet-pink