Lab
PHC # 213
TA:Halah AlMutairi
2
Lab # 1
Standardization of 0.1N sodium hydroxide (NaOH)
 PURPOSE:
To standardize a solution of sodium hydroxide (secondary standard solution) by
titration with a primary standard, potassium hydrogen phthalate( KHP)
 Type of chemical reaction :
Acid-base reaction in aqueous solution
 PRINCIPLE:
-Sodium hydroxide solution need standardization , why ??
A solution of NaOH is secondary standard solution because its tends to absorb
atmospheric carbon dioxide which is weakly acidic. The reaction between the CO2
and NaOH partially neutralizes the NaOH solution, which means that the 0.100 M
solution of NaOH you think you're preparing today will NOT be 0.100 M
tomorrow (concentration will change). For this reason it is necessary to standardize
NaOH solutions to be used in titrations,
-Standardization is typically carried out by titrating the NaOH solution with a
primary standard, acid such as potassium hydrogen phthalate (KHP)
KHP + NaOH ------------------- >> KNaP + H2O
-The indicator you will use phenolphthalein, an acid-base indicator, which is
colorless in acid solution and pink in basic solution
 safety precautions in handling sodium hydroxide:
Sodium hydroxide Corrosive ,, in high concentration Causes eye damage and
severe skin burns
- If contact the eye,, Immediately flush eyes with water ,Remove any contact
lenses and continue to flush eyes with water
3
-If contact the skin,, Flush skin with water , then wash thoroughly with soap and
water (( Contact a physician if need ))
 Procedure:
1-Pipette 10 ml of 0.1N KHP into conical flask
2- Add 10 drops of phenolphthalein indicator
3- Titrate with 0.1N NaOH
The end point: when the solution change from colorless to faint pink color
 Calculation:
N.V.f = N''.V''.f ''
4
Lab. 2
Acid – Base Titration
1- Determination of Ammonium Chloride (Formol Titration)
Principle
When formaldehyde is added to ammonium salt solution, hexamethylene
tetramine is produced together with an amount of acid equivalent to the amount of
ammonium salt used.
The liberated acid is titrated with standard alkali using phenolphethalein as
indicator (ph.ph.).
4NH4Cl + 6HCHO ⎯⎯→ (CH2)6N4 + 4 HCl + 6 H2O
|
| # 4 NaOH
↓
4 NaCl + 4 H2O
Procedure
1. Pipette 10 ml of ammonium chloride solution (NH4Cl) into a 250 ml conical flask.
2. Add 5 ml of formalin solution (neutral) by measuring cylinder (caution).
3. Leave to stand for 5 min.
4. Add 8 drops of ph.ph indicator.
5. Titrate with 0.1 N NaOH.
The end point: faint pink color in the solution.
Calculation:
Conc. Of NH4Cl (g/L)= (ml of titrant) (F) (f) (1000) / ml of sample
5
2- Determination of Borax sample:
Principle:
Borax, also known as sodium borate, sodium tetraborate, or disodium tetraborate, is a salt of
boric acid. It is a direct acid-base titration in which borax react with HCl to give a neutral salt
i.e sodium chloride in addition to a weak acid-boric acid.
Boric acid has a very low ionization constant and has no effect on indicators.
Procedure
1. Pipette 10 ml of borax into conical flask.
2. Add 1-2 drops of M.O indicator.
3. Titrate with 0.1 N HCL.
The end point: orange colored solution.
Calculation:
Conc. Of Borax (g/L)= (ml of titrant) (F) (f) (1000) / ml of sample
6
3- Determination of Boric Acid (indirect titration)
Principle
Boric acid is too weak as acid to be titrated quantitatively in aqueous solution
with sodium hydroxide solution using visual indicators. However, in presence of a
polyhydric alcohol such as glycerol, the boric acid is esterified forming a stronger
monobasic acid (glyceryl boric acid) which is strong enough to give a satisfactory end
point when titrated with standard alkali using ph.ph. as indicator.
CH2OH
CH2OH
CH2OH
+ H3BO3
O
HC
B
Boric acid
Glycerol
CH2OH
HC
O
O
BO
B
H2C
H2O
Glyceryl boric acid
CH3OH
HC
+
O
H2C
CH2OH
OH
OH + NaOH
O
H2C
O
+
H2O
Procedure
1. Add 10 ml of neutral glycerol to the final solution in experimental No.2
2. Add 2 drops of phenolphthalein indicator.
3. Titrate with 0.1 N NaOH.
The end point: formation of onion red color solution.
Calculation:
Conc. Of Boric acid (g/L)= (ml of titrant) (F) (f) (1000) / ml of sample
7
Na
lab.3
Precipitation titration
Titrations with precipitating agents are useful for determining certain analytes containing halides e.g. Clcan be determined when titrated with AgNO3.
Detection of end point:
–
–
–
Formation of Precipitation - Mohr’s method
Formation of colored adsorped on the ppt. – Fajan’s method
Formation of colored solution –Volhard method
1- Mohr’s method for determining chloride:
Mohr’s method is used for determination of chloride and bromide ions (but it is not
used for iodide and cyanide ions why?).
principle:
Mohr’s method is a direct titration using standard solution of silver nitrate in a
neutral medium (why?).
Chloride is titrated with AgNO3 solution. A soluble chromate salt (K2CrO4) is added as the indicator. This
produces a yellow color solution. When the precipitation of the chloride is complete.
NaCl
+
AgNo3 ⇆ AgCl + NaNo3
At the end point: The first excess of Ag+ reacts with the indicator to precipitate red
silver chromate as a second precipitate after precipitation of all chlorides as silver
chloride.
2 Ag+(aq) + CrO42–(aq) → Ag2CrO4(s)
Yellow
red ppt
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The Mohr’s method must be performed at a pH about 8. This method is useful for determining Cl- in
neutral or unbuffered solutions such as drinking water.
Procedure:
1. Pipette 10 ml of sodium chloride solution (NaCL) into stoppered conical flask.
3. Add 1 ml of 2% neutral potassium chromate indicator.
4. Titrate with 0.05 N AgNO3 solution, swirling the liquid constantly, until the red
colour-formed by addition of each drop of AgNO3 solution begins to disappear
more slowly; this is an indication that most of chloride has been precipitated and
that the end point is near.
5. Continue the titration, dropwise, until a faint, but distinct, brick red color is formed
and does not disappear on vigorous shaking.
End point: brick red color
Calculation:
Conc. Of NaCl (g/L)= (ml of titrant) (F) (f) (1000) / ml of sample
9
2- Volhard method for determining of bromide:
Principle: This is an indirect titration procedure for the determination of anions that precipitate with
silver like CL-, Br-, I-, SCN-, and it is preferred in acid (HNO3) solution. A measured excess of AgNO3 is added
to ppt the anion, and the excess of Ag+ is determined by back titration with standard NH4SCN solution:
Ag+(aq) + Br–(aq) → AgBr(s) + excess Ag+
excess Ag+(aq) + SCN–(aq) → AgSCN(s)
The end point is detected by adding iron III (Fe3+) as ferric ammonium sulfate which forms a soluble red
complex with the first excess of titrant.
Fe3+(aq) + SCN–(aq) → [FeSCN]2+(aq)
These indicators must not form a compound with the titrant that is more stable than the precipitate or
the color reaction would occur on addition of the first drop of titrant.
Procedure
1. Pipette 10 ml of potassium bromide solution (KBr) into a glass stoppered conical flask.
2. Add 2 ml conc. HNO3.
3. Add accurately measured 20 ml of 0.05N AgNO3 and shake vigorously for
one minute.
4. Add 1 ml of ferric alum indicator.
5. Titrate with o.1 N NH4SCN.
End point : orange solution is formed
Calculation:
Conc. Of KBr (g/L)= (20 - ml of titrant) (F) (f) (1000) / ml of sample
10
3- Determination of iodide by Fajan’s method
The indicator reaction takes place on the surface of the precipitate. The indicator, which is a dye, exists in
solution as the ionized form, usually an anion.
Principle:
Iodide can be determined by direct titration with standard silver nitrate using an
adsorption indicator such as eosin, fluorescein or diiodofluorescein.
kI + AgNO3 → AgI + kNO3
The titration of I- with Ag+:
Before the equivalent point, I- is in excess and the primary layer is I- (go back to precipitation process in
gravimetry). This repulses the indicator anions; and the more loosely held the secondary (counter) layer
of adsorbed ions is cations
Beyond the equivalent point (end point as well), Ag+ is in excess and the surface of the precipitate becomes
positively charged, with the 1 layer being Ag+. This will now attract the indicator anion and adsorb it in
the 2 (counter) layer:
AgI : Ag+ : : indicatorThe color of the adsorbed indicator is different from that of the un-adsorbed indicator, and this difference
signals the completion of the titration.
Procedure
1. Pipette 10 ml of KI solution in a conical flask, dilute with 100 ml distilled water.
2. Add 10 drops of eosin indicator.
3. Titrate with 0.05N AgNO3 solution with constant rotation until the first separation of precipitate from
colloid state and assuming a pronounced pink colour.
Calculation:
Conc. of KI (g/L)= (ml of titrant) (F) (f) (1000) / ml of sample
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Lab # 4
Complexometric titration
• Complexometric titration is a type of titration useful for determination of metal
ions based on stable complex formation between the analyte (metal) and titrant
(complexing agent=chelating agent )
• An indicator with a marked color change is usually used to detect the end-point of
the titration
• The most useful complexing agent is EDTA
• (EDTA) Ethylenediamminetetraacetic acid has four carboxyl groups and two
amine groups that can act as electron pair donors ,The ability of EDTA to
potentially donate its six lone pairs of electrons for the formation of coordinate
covalent bonds to metal cations makes EDTA a hexadentate ligand (metal cationEDTA complex )
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•The titration of EDTA depends on PH stability, Performed at high (basic) PH, i.e.
PH 10 for Ca+2 or Mg+2
•EDTA is typically used as the disodium salt to increase solubility
•Application of EDTA:
- Medical application: Detoxification in heavy metal poisoning
cases (e.g.: hypercalcemia)
-Analytical application: As masking agent – form complexes with
unwanted/interfering metal ions
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1- Determination of Water Hardness
• Hard water is due to metal ions (minerals) that are dissolved in the ground
water, These minerals include Ca2+, Mg2+, Fe3+ and SO42- . Hard water does cause
soap scum
• Hardness (how much Ca and Mg?) is usually determined by titrating it with a
standard solution (EDTA) to capture the metal ions , by complexometric titration.
A- Determination of Ca2+ :
• Procedure:
1- Pipette 50 ml of tap water into conical flask
2- Add 5 ml of sodium hydroxide buffer (PH 12) (why added ?)
3- Add few crystal of murexide indicator
4- Titrate with 0.01M EDTA
End point: when the solution turns purple
• Calculation:
Conc. (g/L)= (ml of titrant) (F) (f) (1000) / ml of sample
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B- Determination of Mg2+:
• Procedure:
1- Pipette 50 ml of tap water into conical flask
2- Add 5 ml of ammonia buffer (PH 10)
3- Add few crystal of EBT indicator( Eriochrome Black T )
4- Titrate with 0.01 M EDTA
End point : when the solution turns blue
• Calculation:
The ml of EDAT in step A represents Ca, While that in step B represents both Ca
and Mg therefore:
Ml of EDTA equivalent to Mg = step B – step A
Conc. (g/L) = ( E.P Step B - E.P Step A ) (F) (f) (1000) / ml of sample
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2-Determination of ZN+2 concentration:
• Procedure:
1- Pipette 10ml of Zinc sulfate (ZnSO4) into conical flask
2- Add 5ml of ammonia buffer (PH 10)
3-Add few crystal of EBT indicator ( Eriochrome Black T )
4-Titrate with 0.01 M EDTA
End point: when the solution turns blue
• Calculation:
Conc. (g/L)= (ml of titrant) (F) (f) (1000) / ml of sample
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Lab # 5
Redox titration
Redox reactions, or oxidation-reduction reactions, involve the transfer of electrons
between two chemical species. The compound that loses an electron is said to be
oxidized, the one that gains an electron is said to be reduced.
Oxidation is
The half reaction in which there is loss of an electrons by a species (or increase
of oxidation number of an atom)
Oxidizing agent:
It is a substance which has the ability to gain electrons.
Reduction is
The half reaction in which there is a gain of an electrons by a species (or
decrease of oxidation number of an atom)
Reducing agent:
It is a substance which has ability to loss electrons.
-Classification of substances as an oxidizing or reducing agents depends on their
oxidation potential, which can be obtained from Nernest equation:
E = EO + 0.0591/n log[oxd]/[red]
-Substance with higher EO can act as an oxidizing agent for substance with lower
EO value.
Redox titration: is a type of titration based on a redox reaction between the
analyte and titrant.
17
Standardization of potassium permanganate ( KMnO4 )
Procedure:
1- Pipette 10 ml of 0.1N Na2C2O4 into a conical flask.
2- Add 20ml dil. H2 SO4.
3-Heat the flask on the hotplate to 70OC.
4- Titrate with 0.1N KMnO4.
End point: when faint pink color appear
Safety Note:
-Oxalic acid is poisonous and must be handled with care , If contact does occur
immediately wash the affected area with water
-KMnO4 (potassium permanganate) stains skin and clothes
The following points about this reaction should be noted:
1- The reaction occurs only in acid solution. Acidification is achieved by the
addition of sulfuric acid.
2 -The reaction rate is normally slow but in this experiment is increased a by
warming the contents of the conical flask.
3- No indicator is necessary in the titration because KMnO4 is a self-indicator.
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Principle:
The reaction that occurs here is oxidation and reduction reaction
MnO4- + C2O4-2 + H+ →
Mn+2 + CO2 + H2O
Oxidation half reaction:
C2O42- - 2 e- → 2 CO2
Reduction half reaction :
MnO4- + 5 e- →
Mn2+
Balanced net ionic equation
2 MnO4- + 5 C2O42- →
2 Mn2+ + 10 CO2
Calculation:
Conc. Of KMnO4 (g/L) = (ml of titrant) (f) (F) (1000) / ml of sample
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Determination of ferrous sulphate sample (FeSO4)
Procedure:
1-Pipette 10 ml of FeSO4 into a C.F.
2- Add 10 ml of a mixture of equal volumes of sulfuric and phosphoric acids (5 ml
H2SO4 + 5 ml H3PO4).
3- Add 3 drops of diphenylamine indicator.
4- Titrate with 0.1 N K2Cr2O7 until the formation of bluish-violet solution.
Principle :
Cr2O72 - + Fe2+
→
2 Cr3+ +
Fe3+
Calculation:
Conc. Of Fe2+ (g/L) = (ml of titrant) (f) (F) (1000) / ml of sample
20
Lab# 6
Derermination of KI (Andrew's Reaction )
Procedure




Pipette 10 ml of a sample into G.S.C.F
Add 5 ml of Conc HCL
Add 5 ml of CHCl3
Titrate with 0.05 M KIO3
End.Point : when chloroformic layer become colorless
Principle :
Calculation:
21
Determination of Phenol :
Procedure:









Pipette 10ml of phenol into G.S.C.F
Dilute with 25ml distilled water
Add 25ml of standard Br2 solution (using the pipette)
Moisten the stopper KI
Add 5ml conc HCl
Leave in dark place for 15min with occasional shaking
Add 10ml of 10% KI solution
Add 5ml of chloroform
Titrate the liberated iodine with 0.1N Na2S2O3
End.point : when chloroform became colorless
N.P: carry out the blank experiment by same procedure but
add 10ml water instead of the sample
Calculation :
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