1

Qualitative Analysis of

Anions

By

Ass.Prof.Dr. Nouruddin W. Ali

2

Introduction

- The purpose of chemical analysis is to establish the composition of naturally

occurring or artificially manufactured substances. This is usually done in two

distinct steps:

1. Qualitative analysis; which is used to identify the sample components

2. Quantitative analysis; in which the relative amounts of these components are

determined.

- In this course, the traditional methods of qualitative analysis are introduced.

These can be divided into two categories dry reactions which are carried out on

solid samples usually at elevated temperatures and wet reactions that uses

dissolved samples and reagent solutions. In both types, the chemical change that

appears or disappears is observed and used for the elucidation of sample

composition.

- It must be ensured that the study of classical qualitative inorganic analysis is

invaluable for any chemist as this where he first comes across and handles

materials related to chemistry science.

- The intelligent study of qualitative analysis requires a certain level of

theoretical background in general chemistry. Such a background involves chemical

symbols, formulae, equations, theory of electrolytes, equilibria in electrolyte

solutions, acid base theory, strength of acids, pH, buffer systems, hydrolysis,

…….etc.

3

Qualitative analysis of anions

- The methods available for detection of anions are not systematic i.e. no

really satisfactory scheme has yet been proposed which permits the separation of

the common anions into major groups and the subsequent separation of each

group into its independent constituents.

- The following scheme of classification has been found to work well in

practice; it is not a rigid one since some of the anions belong to more than one of

the subdivisions. Essentially the processes used may be divided into

1. Class A which involve the identification of volatile products obtained on

treatment with acids

This class is subdivided into

1- Gases evolved with dilute hydrochloric acid or dilute sulphuric acid

2- Gases evolved with concentrated sulphuric acid

2. Class B which involve the reactions done in solution

This class is subdivided into

1- Precipitation reactions

2- Complex formation reactions

3- Oxidation reduction reactions

- For convenience the reaction of certain groups of anions are given together

according to similarities in their reactions.

4

Carbonate and bicarbonate

CO32- HCO3

-

1- PARENT ACID: CARBONIC ACID ( H2CO3)

- WEAK ACID

- Existing only in solution

- Heating of a solution of H2CO3 CO2 evolves

H2CO3 CO2 + H2O

BICARBONATES decomposes on heating wit the liberation of CO2

bicarbonates are considered the first step of ionization of carbonic acid while

in the second step carbonates are formed

H2CO3 H+ + HCO3- H+ + CO3

2-

2- DRY REACTIONS:

A- ACTION OF DIL. HCl:

Decomposition with effervescence due to evolution of CO2 gas in case of CO32- and

HCO3-

CO32- + 2H+ CO2 + H2O

HCO3- + H+ CO2 + H2O

this is a type of displacement reaction in which stronger acid liberates the

very weak carbonic acid which spontaneously decomposes to CO2 & H2O

some natural carbonates e.g. magnesite (MgCO3) / Siderite ( FeCO3) /

dolomite ( Ca,Mg CO3) do not react appreciably in the cold so they must be finally

powdered and the reaction mixture warmed

if the acid used is very dilute no effervescence occurs with carbonate

because bicarbonates may be formed

Turbidity test for CO2 gas:

- This test is considered to be selective for CO2 and SO2 gases

- The solid substance is placed in a test tube, dil. HCl added the immediately

replaced gas which is evolved upon warming is passed into lime water [Ca(OH)2] or

baryta water [Ba(OH)2] contained in another test tube

the production of a turbidity indicates the presence of carbonates or

bicarbonates

5

CO2 + Ca(OH)2 CaCO3 + H2O

CO2 + Ba(OH)2 BaCO3 + H2O

with prolonged passage of CO2 , the turbidity formed due to the insoluble

carbonates slowly disappears as a result of the formation of a soluble bicarbonate

CaCO3 + CO2 + H2O Ca(HCO3)2

N.B. it should be remembered that lime water may be turned very

slightly cloudy by carbon dioxide from the air present in the test tube .

if Ba(OH)2 is used the turbidity can be quite appreciable

Therefore it is better to use Ca(OH)2 solution which must be tested first

B- ACTION OF DIL. H2SO4:

The same as with dil. HCl but the carbonates whose metal ions form insoluble

sulphates produce effervescence accompanied by precipitation

BaCO3 + H2SO4 BaSO4 + CO2 + H2O

SrSO4 + H2SO4 SrSO4 + CO2 + H2O

* ANY ACID WICH IS STRONGER THAN CARBONIC ACID WILL DISPLACE IT

ESPECIALLY ON WARMING THUS EVEN ACETIC ACID WILL DECOMPOSE

CARBONATES WHILE BORIC ACID AND HYDROCYANIC ACID WILL NOT

3- WET REACTIONS:

- all normal carbonates with the exception of those of the alkali metals and of

ammonium are insoluble in water

- all bicarbonates are soluble in water

a- reaction with AgNO3:

White precipitate of silver carbonate is formed immediately

CO32- + 2Ag+ Ag2CO3

* The precipitate is soluble in nitric acid and ammonia

Ag2CO3 + 2H+ 2Ag+ + CO2 + H2O

Ag2CO3 + 4NH3 2[Ag(NH3)2]+ +CO32-

* the precipitate becomes yellow or brown upon addition of excess reagent or

boiling of the mixture owing to the formation of silver oxide

Ag2CO3 Ag2O + CO2

* Ag2CO3 is insoluble in dil. HCl as HCl converts Ag2CO3 into AgCl (white ppt.)

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b- reaction with BaCl2,CaCl2, MgSO4:

White precipitate of BaCO3, CaCO3 & MgCO3 will be obtained upon the addition of

these reagents to samples of carbonate solutions.

BaCl2 + Na2CO3 BaCO3+ 2 NaCl

CaCl2 + Na2CO3 CaCO3+ 2 NaCl

MgSO4 + Na2CO3 MgCO3+ Na2SO4

N.B. BaCO3 dissolves in HCl, HNO3 & CH3COOH with evolution of CO2 the action

of H2SO4 converts barium carbonate into the even less soluble barium sulphate

FOR HCO3- no ppt. on cold since all bicarbonates are soluble in water but on

heating the mixture white ppt. of the produced carbonates are formed with the

above reagents

BaCl2 + 2 NaHCO3 Ba(HCO3)2 " SOLUBLE " + 2 NaCl

boiling

BaCO3 "WHITE PPT." + CO2 + H2O

c- reaction with HgCl2:

no precipitate is formed with bicarbonate while in a solution of carbonates a

reddish brown precipitate of basic mercuric carbonate is formed

CO32- + 4Hg2+ + 3H2O → Hg4O3CO3 + 6H+

The excess of carbonate acts as a buffer reacting with the hydrogen ions formed in

the reaction:

CO32- + 2H+ → CO2 + H2O

4- ANALYSIS OF MIXTURES:

" mixture of CO32- & HCO3

- "

PRINCIPLE OF SEPARATION both anions have similar reactions but CO32- form

ppt. immediately on cold upon the addition of CaCl2, BaCl2 or MgSO4 and other

metals while the bicarbonates of these metals are soluble

Separation:

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- add excess CaCl2 , BaCl2 or MgSO4 to a solution of the mixture CO32- & HCO3

-

a white ppt. indicates CO32-

Ca2+ + CO32- CaCO3

- on centrifuging or filtering HCO3- pass into the filtrate or the centrifugate

- on boiling or treatment of the filtrate or the centrifugate with ammonia

solution a white ppt. or cloudiness is obtained

2Ca2+ + 2HCO3- + 2NH3 2CaCO3 + 2NH4

+

5- BIOLOGICAL & ENVIRONMENTAL SIGNIFICANCE:

A) Carbonate:

Carbonate works as a buffer in the blood as follows: when pH is too low, the

concentration of hydrogen ions is too high, so you exhale CO2. This will cause the

equation to shift left, essentially decreasing the concentration of H+ ions, causing a

more basic pH.

When pH is too high, the concentration of hydrogen ions in the blood is too low,

so the kidneys excrete bicarbonate (HCO3−). This causes the equation to shift right,

essentially increasing the concentration of hydrogen ions, causing a more acidic

pH.

There are 3 important reversible reactions that control the above pH balance

1. H2CO3(aq) H+(aq) + HCO3

-(aq)

2. H2CO3(aq) CO2(aq) + H2O(l)

3. CO2(aq) CO2(g)

Exhaled CO2(g) depletes CO2(aq) which in turn consumes H2CO3 causing the

aforementioned shift left in the first reaction by L'Chatlier's principle. By the same

principle when the pH is too high, the kidneys excrete bicarbonate (HCO3-) into

urine as urea via the Urea Cycle (aka the Krebs-Henseleit Ornithine Cycle). By

removing the bicarbonate more H+ is generated from carbonic acid (H2CO3) which

has ultimately come from inhaled CO2(g)

B) Bicarbonate:

Bicarbonate is alkaline, and a vital component of the pH buffering system of the

body (maintaining acid-base homeostasis). 70 to 75 percent of CO2 in the body is

converted into carbonic acid (H2CO3), which can quickly turn into bicarbonate

(HCO3−).

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With carbonic acid as the central intermediate species, bicarbonate - in

conjunction with water, hydrogen ions, and carbon dioxide - forms this buffering

system, which is maintained at the volatile equilibrium required to provide prompt

resistance to drastic pH changes in both the acidic and basic directions. This is

especially important for protecting tissues of thecentral nervous system, where pH

changes too far outside of the normal range in either direction could prove

disastrous. (See acidosis, or alkalosis.)

Bicarbonate also acts to regulate pH in the small intestine. It is released from

the pancreas in response to the hormone secretin to neutralize the

acidic chyme entering the duodenum from the stomach.

In freshwater ecology, strong photosynthetic activity by freshwater plants in

daylight releases gaseous oxygen into the water and at the same time produces

bicarbonate ions. These shift the pH upward until in certain circumstances the

degree of alkalinity can become toxic to some organisms or can make other

chemical constituents such as ammonia toxic. In darkness, when no

photosynthesis occurs, respiration processes release carbon-dioxide, and no new

bicarbonate ions are produced, resulting in a rapid fall in pH Carbonate works as a

buffer in the blood as follows: when pH is too low, the concentration of hydrogen

ions is too high, so you exhale CO2. This will cause the equation to shift left,

essentially decreasing the concentration of H+ ions, causing a more basic pH.

The most common salt of the bicarbonate ion is sodium bicarbonate, NaHCO3,

which is used as baking soda. When exposed to an acid such as acetic

acid (vinegar), sodium bicarbonate releases carbon dioxide. This is used as

a leavening agent in baking.

The flow of bicarbonate ions from rocks weathered by the carbonic acid in

rainwater is an important part of the carbon cycle.

Bicarbonate also serves in the digestive system. It raises the internal pH of the

stomach, after highly acidic digestive juices have finished in their digestion of

food. Ammonium bicarbonate is used in digestive biscuit manufacture

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Sulphur containing anions

S2- / SO32- / S2O3

2- / SO42-

Sulphide / sulphite / thiosulphate / sulphates

* sulphur element has a great tendency to accept electrons to form sulphide ion

which has the ability to form covalent bonds with oxygen or with sulphur atoms to

form the other members of the sulphur group

* sulphides, sulphites & thiosulphates have general characters of being reducing

agents

therefore they can react with oxidizing agents like potassium dichromate (

K2Cr2O7) or potassium permanganate ( KMnO4) to give the corresponding oxidation

reduction products

1- PARENT ACIDS:

1- Hydrogen sulphide or hydrosulphuric acid (H2S)

- A gas with offensive rotten egg odor

- Poisonous

- In solution it gives a weak acid which ionizes in two steps

H2S H+ + HS- (HYDROSULPJIDE ION )

HS- H+ + S-2 (SULPHIDE ION )

Alkaline sulphide solutions dissolve sulphur to give polysulphide ion for

example ammonium polysulphide [ (NH4)2SX ]

2- Sulphurous acid (H2SO3)

- ONLY KNOWN IN SOLUTION

- It's a diprotic acid with moderate strong acidity

- Like H2CO3 in water it's present in equilibrium as follow

H2SO3 H+ + HSO3- H+ + SO3

-2

3- Thiosulphuric acid (H2S2O3)

- It's not known in the free form

- It decomposes to give H2O, SO2 & S

- It's more stronger acid than sulphurous acid in solutions

- It consists of SO32- & S solution which upon boiling gives S2O3

2-

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4- Sulphuric acid (H2SO4)

- It's a colorless oily liquid

- It's found as concentrated acid

General properties of H2SO4 acid :

A] acid properties:

It's a strong acid which is somewhat weaker than halide acids and nitric acid

it ionizes in dilute solutions in two steps

H2SO4 H+ + HSO4- ( HYDROGEN SULPHATE )

HSO4- H+ + SO4

2- (SULPHATE)

metals can liberate hydrogen from H2SO4 solution

H2SO4 + ZnO ZnSO4 + H2

being a strong acid it can replace weak acids like carbonic acid , boric acid ,

hydrocyanic acid and volatile acids or their decomposition products due to its high

B.P.

2NaCl + H2SO4 Na2SO4 + 2HCl

B] dehydrating properties:

conc. H2SO4 has a great tendency to combine with water to form stable hydrates

[H2SO4. x H2O ]

SO it is used as a dehydrating agent for certain substances and used mostly in

desiccators

H2SO4 causes charring for certain organic substances as sugars due to the

vigorous abstracting power of water from these substances

C] oxidizing properties:

It's considered to be as moderately strong oxidizing agent when heated with

most reducing agents where it is oxidized to SO2 while with more strong reducing

agents it may be reduced to S or H2S

H2SO4 H2O + SO2 + [O]

2- DRY REACTIONS:

A- ACTION OF DIL. HCl:

1- SULPHIDE:

evolution of H2S gas which is characterized by its rotten egg odor

S2- + 2H+ H2S

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it could be identified by

1- blackening of filter paper moistened with lead acetate solution

H2S + pb2+ pbS (black)

2- turns a filter paper moistened with cadmium acetate solution yellow

H2S + Cd2+ CdS (yellow)

3- reducing properties: H2S is a good reducing agent

it reacts with iodine solution (I2) bleaches the brown color of I2 solution

H2S + I2 2I- +2H+ + S

it reacts with acidic KMnO4 solution change its pink color into colorless

5H2S + 6H+ + 2 MnO4- 2Mn2+ + 8H2O +5S

it reacts with acidic K2Cr2O7 solution change its orange color into green

3H2S + 8H+ + Cr2O72- 2Cr3+ + 7H2O +3S

In each case sulphur is precipitated.

Small amounts of chlorine may be produced with KMnO4 and K2Cr2O7 if HCl

used is not dilute; this can be avoided by using dilute sulphuric acid.

polysulphide upon treatment with dilute HCl besides evolution of H2S they

give yellow ppt. of sulphur

(NH4)2SX +2H+ 2NH4+ + H2S + (x-1)S

2- SULPHITE:

decomposition more rapidly on warming with the evolution of SO2 gas

SO32- + 2H+ H2SO3 SO2 + H2O

SO2 could be identified by

1- characteristic suffocating odor of burnt sulphur

2- turbid lime water ( like CO2) due to the formation of the insoluble

CaSO3 which is soluble in water & upon prolonged passage of SO2

due

to the formation of soluble Ca(HSO3)2

SO2 + Ca(OH)2 CaSO3 + H2O

CaSO3 + SO2 + H2O Ca(HSO3)2

3- reducing properties: SO2 is a good reducing agent

it reacts with iodine solution (I2) bleaches the brown color of I2 solution

SO2 + I2 + H2O 2I- +2H+ + SO32-

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it reacts with acidic KMnO4 solution change its pink color into colorless

5 SO2+ 6H+ + 2 MnO4- 2Mn2+ + 8H2O +5 SO3

2-

it reacts with acidic K2Cr2O7 solution change its orange color into green

3 SO2 + 8H+ + Cr2O72- 2Cr3+ + 7H2O +3 SO3

2-

The reducing properties test can be done by holding a filter paper moistened

with acidified K2Cr2O7 solution over the mouth of the test tube where green

coloration due to chromium ions is observed.

Another method of identifying SO2 is to hold a filter paper moistened with

potassium iodate and starch solution in the vapor where a blue color due to the

formation of iodine is observed

5SO2 +2IO3- + 4H2O I2+ 5SO4

2- + 8H+

3- THIOSULPHATE:

no immediate change on cold but on warming with dilute HCl or

standing the solution becomes turbid due to the liberated yellow colloidal

sulphur with evolution of SO2 gas due to the decomposition of the liberated

unstable H2S2O3

S2O32- + 2H+ H2S2O3 SO2 + H2O + S (yellow colloidal ppt.)

4- SULPHATE:

no reaction with dilute HCl

3- WET REACTIONS:

- all sulphides are insoluble except Na+, K+, NH4+ , Ca2+, Ba2+& Sr2+ salts

- all sulphites are insoluble except Na+, K+ & NH4+ salts

- all thiosulphates are soluble except Ag+, Pb2+ , Hg2+ & Ba2+ salts

- all sulphates are soluble except Pb2+ ,Mg2+ , Ca2+, Ba2+& Sr2 salts

A- Reaction with AgNO3:

1- S2- :

Black ppt. of Ag2S is formed

2Ag+ + S2- Ag2S

soluble in hot dil. HNO3 insoluble in ammonia solution & KCN solution

2- SO32- :

White ppt. of Ag2SO3 is formed

2Ag+ + S2- Ag2SO3

13

Ag2SO3 on boiling with water undergoes self oxidation reduction reaction with the

production of gray ppt of metallic silver

2Ag2SO3 boiling 2Ago + Ag2SO4 + SO2

soluble in dil. HNO3 soluble in ammonia solution

soluble in excess sulphite solution to give a complex salt which on boiling gives a

grey ppt. of metallic silver

Ag2SO3 + SO32- 2 [Ag SO3 ]

-

2 (Ag SO3 )- boiling 2Ago + SO4

2-+ SO2

3- S2O32- :

white ppt. of Ag2S2O3

changes its color on standing to yellow brown and finally black

due to the formation of Ag2S

2Ag+ + S2O32- Ag2S2O3

Ag2S2O3 + H2O Ag2S + H2SO4

soluble in excess thiosulphate solution to give a complex salt

Ag2S2O3 + 3 S2O32- 2 [Ag(S2O3)2]3-

4- SO42- :

no ppt. with dil. solution but a ppt. may be formed in a very concentrated

solution

B- Reaction with BaCl2:

1- S2- :

no visible reaction

2- SO32- :

White ppt. of BaSO3 is formed

Ba2++ SO32- BaSO3

soluble in dil. HCl soluble in excess SO3 solution and reprecipitate on

boiling

N.B. on standing the precipitate is slowly oxidized to the sulphate and is then

insoluble in dilute mineral acids

this change is rapidly effected by warming with Br2 water or a little

concentrated nitric acid or with hydrogen peroxide

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2BaSO3 + O2 2 BaSO4

BaSO3 + Br2 + H2O BaSO4 + 2Br- + 2H+

3BaSO3 + 2HNO3 3BaSO4 + 2NO + H2O

BaSO3 + H2O2 BaSO4 + H2O

3- S2O32- :

no ppt. in dilute solution but a ppt. is formed from very concentrated solution

4- SO42- : white ppt. of BaSO4 is formed which is insoluble in dilute HCl even

upon boiling Ba2++ SO42- BaSO4

* this distinguishes BaSO4 from all other barium salts and this fact is used to test

for SO42-

* the test is complicated somewhat if the solution contains the S2O32- anion or a

mixture of SO32- and S2- ions as in both cases addition of acid produces a white ppt.

of sulphur which like BaSO4 is insoluble in acids

S2O32- + 2H+ S +SO2 +H2O

2 S2- + SO32- + 6H+ 3S + 3H2O

BaSO4 can be distinguished from sulphur as it forms mixed crystals of pink color

with KMnO4

C- Reaction with FeCl3:

1- S2- :

black ppt. of Fe2S3 is formed

2 Fe3+ + 3S2- Fe2S3

soluble in dil. acids

Methylene blue test: N,N-dimethyl-p-phenylenediamine is converted by ferric

chloride and hydrogen sulphide in strongly acid solution into the water soluble

dyestuff "methylene blue"

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(CH 3)

2N

NH 2

H 2S

H 2N

N(CH 3)

2

Fe3+

Fe2+

NH 4+ H

+

S

N

(CH 3)

2N N(CH

3)

2

n

+ + + 6

6 + + 4 +

++

this is a sensitive test for soluble sulphides and hydrogen sulphide

2- SO32- :

dark red color of ferric sulphite is produced on cold

2Fe3+ + 3 SO32- Fe2(SO3)3

upon boiling a reddish brown ppt. of the basic iron salt is formed

3- S2O32- :

a dark violet color of dithiosulphatoiron(III) complex is produced which

disappears on boiling as tetrathionate and Fe2+ are formed from the oxidation of

S2O32- with Fe3+ even on cold

Fe3+ + 2 S2O32- [ Fe (S2O3)2] –

2Fe3+ + 2 S2O32- 2 Fe2+ + S4O6

2-

the color fades and disappears after sometimes

4- SO42- :

do not react

D- Reaction with Pb( CH3COOH)2:

1- S2- :

black ppt. of PbS is formed

Pb2+ + S2- PbS

2- SO32- :

white ppt of lead sulphite

Pb2+ + SO32- PbSO3

soluble in cold HNO3

upon boiling it oxidizes to PbSO4 which is a white ppt.

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3- S2O32- :

white ppt of lead thiosulphate

Pb20 + S2O32- PbS2O3

soluble in cold HNO3

upon boiling it oxidizes to PbS which is a black ppt.

N:B this reaction can be used to distinguish sulphites and thiosulphates

4- SO42- :

white ppt. of lead sulphate which is insoluble in cold dilute mineral acids but

soluble in ammonium acetate and sodium hydroxide solutions

Pb2+ + SO42- PbSO4

PbSO4 + 4CH3COONH4 [Pb(CH3COO)4]2- + SO42-

PbSO4 + 3OH- HPbO2- + SO4

2- + H2O

N:B if acetic acid and potassium chromate are added to the precipitate, yellow lead

chromate is precipitated

E- Reducing action:

- Sulphides , poly sulphides , sulphites & thiosulphates are reducing agents

- They reduce solutions of I2 , KMnO4 & K2Cr2O7 with varying activities in

acidified solutions

a- with I2 solution:

S2- + I2 (ACIDIC MEDIUM) 2I- + S

SO32- + I2 + H2O 2I- +2H+ + SO4

2-

I2 + 2 S2O32- (ACIDIC MEDIUM) 2I- + S4O6

2-

b- with KMnO4 solution:

5S2- + 6H+ + 2 MnO4- 2Mn2+ + 3H2O +5 SO4

2-

5 SO32- + 6H+ + 2 MnO4

- 2Mn2+ + 3H2O +5 SO42-

5 S2O32-+ 14H+ + 8 MnO4

- 8Mn2+ + 7H2O +10 SO42-

c- with K2Cr2O7 solution:

3H2S + 8H+ + Cr2O72- 2Cr3+ + 7H2O +3S

3 SO32- + 8H+ + Cr2O7

2- 2Cr3+ + 4H2O +3 SO42-

3 S2O32-+ 26H+ +4 Cr2O7

2- 8 Cr3+ + 13 H2O +6 SO42-

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4- SPECIAL TESTS:

1- S2- :

Cadmium carbonate test:

a canary yellow ppt. of CdS is produced

S2-+ CdCO3 CdS (yellow) + CO32-

This test could be used for the identification and separation of sulphide when

present in a mixture with other sulphur containing anions or those anions which do

not react with CdCO3

if soluble cadmium salts were used CdSO3 would be precipitated together with

CdS

therefore solid cadmium carbonate should be used as the reagent

as cadmium carbonate is difficultly soluble in water it gives a very low

concentration of Cd2+ ions in solution which is quite sufficient to exceed the

solubility product of cadmium sulphide which is less soluble than cadmium

carbonate and is therefore completely ppted while on the other hand the solubility

product of CdSO3 is not reached and sulphite ion remains in solution

2- SO32- :

Zinc nitroprusside test:

to a cold saturated zinc sulphate solution is added an equal volume of

potassium ferrocyanide solution then a few drops of 1% sodium nitroprusside

solution is added to the above solution to obtain a solution of zinc nitroprusside

zinc nitroprusside solution is added to the sulphite solution

a red ppt. of Na5 [Fe (CN)5 SO3] is obtained

the test is not applicable in the presence of sulphides and thiosulphates which

can be removed by the addition of mercuric chloride that reacts forming the stable

mercuric sulphide

Hg2+ + S2- → HgS

Hg2+ + S2O32- + H2O → HgS + SO4

2- + 2H+

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3- S2O32- :

Formation of thiocyanate test:

by boiling with KCN solution in the presence of NaOH ,cooling acidifying then

adding FeCl3 solution blood red color of ferric thiocyanate complex is produced

S2O32- + CN- [OH- / BOIL] SCN- + SO3

2-

Fe3+ + SCN- COOL [Fe(SCN)]2+

4- SO42- :

Hepar test:

sulphate is reduced by carbon to sulphide by heating on a piece of charcoal in

the presence of sodium carbonate in the reducing zone of the flame

MSO4 + Na2CO3 (fusion) Na2SO4 + MCO3

Na2SO4 + C Na2S + 4CO

transfer the fusion product to a silver coin

moisten with a little water

a brownish black stain of Ag2S is obtained

S2- + 2H2O 2OH- + H2S

H2S + 2AgO Ag2S + H2

N:B The formed sulphide can also be tested by any other specific test for sulphides

5- ANALYSIS OF MIXTURES:

1- Mixture of S2- , SO32- , S2O3

2- & SO42- :

this mixture is liable to interference upon treatment with dil. HCl due to the

common reducing characters of H2S and SO2 gases

therefore separation technique must be done in order to identify each one of

them which is based on

(a) precipitation of S2- by the action of CdCO3 where SO32-

, S2O32- & SO4

2- remain in

solution ;

(b) precipitation of SO32- & SO4

2- by barium chloride which does not precipitate

S2O32- ; (c) the fact that BaSO4 is almost insoluble while BaSO3 is soluble in dilute

HCl

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S2- , SO32- , S2O32- & SO4

2-

+

CdCO3

CdS (YELLOW PPT) CENTRIFUGATE

S2- +

BaCl2

White ppt. centrifugate

BaSO3 & BaSO4 S2O32-

+ +

dilute HCl dilute HCl

heat

white ppt centrifugate SO2 + S

SO42- SO3

2- S2O32-

confirmed by oxidation

with Br2 OR H2O2 into SO42-

2- Mixture of CO32- , SO3

2- OR S2O32- :

this mixture is considered difficult due to the interference that occur upon the

addition of dilute HCl which liberates CO2 & SO2 gases which turbid lime water and

turbidity disappears on prolonged passage

SO2 can be detected by its reducing characters as discussed before but CO2 has

no reducing characters

therefore SO32- or S2O3

2- ions must be firstly oxidized into SO42- by an oxidizing

agent such as H2O2 , k2Cr2O7 or KMnO4 then dilute H2SO4 is added and the mixture

warmed

CO2 will only evolve which can be tested for using lime water test

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6- BIOLOGICAL & ENVIRONMENTAL SIGNIFICANCE:

A) Sulphide:

Dissolved free sulfides (H2S, HS− and S2−) are very aggressive species for the

corrosion of many metals such as, e.g., steel, stainless steel, and copper. Sulfides

present in aqueous solution are responsible for stress corrosion cracking (SCC) of

steel, and is also known as sulfide stress cracking. Corrosion is a major concern in

many industrial installations processing sulfides: sulfide ore mills, deep oil wells,

pipeline transporting soured oil, Kraft paper factories. Microbially-induced

corrosion (MIC) or biogenic sulfide corrosion are also caused by sulfate reducing

bacteria producing sulfide

B) Sulphite:

Sulfites are often used as preservatives in dried fruits, and dried potato products.

Although shrimp are sometimes treated with sulfites on fishing vessels, the

chemical may not appear on the label. In 1986, the Food and Drug

Administration in the United States banned the addition of sulfites to all fresh fruit

and vegetables that are eaten raw.

Sulfites are widely used to extend the shelf life of products. Because it is often

difficult to know whether a food contains sulfites, many people do not realize that

they may have a sensitivity to sulfite. Sulfites are also known to destroy vitamin

B1 (thiamin),[14] a vitamin essential for metabolism of carbohydrates and alcohol.

When pH is too high, the concentration of hydrogen ions in the blood is too low,

so the kidneys excrete bicarbonate (HCO3−). This causes the equation to shift right,

essentially increasing the concentration of hydrogen ions, causing a more acidic

pH.

C) Thiosulphate:

Thiosulphates occurs naturally in hot springs and geysers, and is produced by

certain biochemical processes. It rapidly dechlorinates water, and is notable for its

use to halt bleaching in the paper-making industry. Thiosulphates is also useful in

smelting silver ore, in producing leather goods, and to set dyes in textiles. Sodium

thiosulfate, commonly called hypo, was widely used in photography to fix black

and white negatives and prints after the developing stage; modern 'rapid' fixers

use ammonium thiosulfate as a fixing salt because it acts three to four times

faster.

21

D) Sulphate:

Sulfates are important in both the chemical industry and biological systems:

- The lead-acid battery typically uses sulfuric acid.

- Some anaerobic microorganisms, such as those living near deep sea thermal

vents use sulfates as an energy source for chemosynthesis.

- Copper sulfate is a common algaecide.

- Magnesium sulfate, commonly known as Epsom salts, is used in therapeutic

baths.

- Gypsum, the natural mineral form of hydrated calcium sulfate, is used to

produce plaster.

- The sulfate ion is used as counter ion for some cationic drugs

Sulfates occur as microscopic particles (aerosols) resulting from fossil

fuel and biomass combustion. They increase the acidity of the atmosphere and

form acid rain.

22

Halides

F- / Cl- / Br- / I-

Fluoride / chloride / bromide / iodide

* fluorides, chlorides, bromides and iodides are known as halides and are

belonging to group VIIA in the periodic table

* they are characterized by their higher electronegativity ( their tendency for

gaining electrons is very great)

so when these anions gain one electron each one attains the inert gas

structure of 8 electrons in the outermost shell

* they have similar properties and characters

all are monovalent

resemble each other in chemical reaction except fluoride

* regarding the electronegativity it increases according to the following order

I- Br- Cl- F-

as the ionic size increases the tendency to lose electrons increases

therefore I- ion is firstly and easily oxidized into free I2 by losing readily

an electron followed by Br- when present in a mixture

it's difficult to oxidize F- into F2 so F- ions are highly stable to held

strongly a proton therefore the order of stronger halogen acid is from HI HBr

HCl HF

1- PARENT ACIDS:

1- Hydrofluoric acid (HF)

- Colorless fuming highly corrosive and itching liquid

- Soluble in water producing the weakest acidic solution in the halogen acid

series

2- Hydrochloric acid (HCl)

- Colorless gas with irritating odor that fumes in moist air

- Extremely soluble in water to form acidic solution

* concentrated HCl contains 37% of HCl gas

an exactly standard solution is obtained by constant boiling mixture which

contains 20% of HCl gas that distills at 110oC without change in composition

"Constant boiling point HCl "

23

3- Hydrobromic acid (HBr)

- Colorless gas with irritating odor that fumes in moist air

- Extremely soluble in water forming very strongly acidic solution

- Like HF & HCl it forms constant boiling mixture containing 48% of HBr which

distills at 126oC

* On standing the solution becomes yellow due to the oxidation to bromine

4- Hydroiodic acid (HI)

- Colorless gas with irritating odor that fumes strongly in moist air

- Soluble in water forming the strongest acidic solution of the halo acid series

- A constant boiling mixture has a concentration of 57% at 127oC

* On standing the solution becomes brown due to the oxidation to iodine

2- DRY REACTIONS:

a. ACTION OF DIL. HCl:

HCl shows no reaction upon treatment of the solid sample with it even on heating

this reaction can differentiate carbonate and sulphur group from halides

b. ACTION OF CONC. H2SO4 :

1- FLOURIDE :

evolution of HF gas which is colorless and fumes with moist air

2F- + H2SO4 2HF + SO42-

it could be identified by its corrosive and itching action on the glass in

presence of water

the test tube or the glass rod subjected to the evolved HF gas acquire

oily appearance due to the formation of silicic acid and hydroflourosilicic acid

4 HF + SiO2 (GLASS) SiF4 + 2H2O

3SiF4 + 3H2O H2SiO3 + 2 H2SiF6

"silicic acid" "hydroflourosilicic acid"

2- CHLORIDE:

evolution of HCl gas that can be identified by

2Cl- + H2SO4 2HCl + SO42-

1- formation of white fumes with moist air

2- pungent irritating odor

3- changing a blue moistened litmus paper into red

4- formation of white fumes of NH4Cl when a glass rod moistened with

24

ammonium hydroxide solution is exposed to the evolved gas

NH4OH + HCl NH4Cl + H2O

3- BROMIDE :

a mixture of HBr &Br2 may be formed which have characteristic brown color

specially on warming

at the same time Sulphuric acid will be reduced into SO2 , H2S or S

2Br- + H2SO4 2HBr + SO42-

2HBr + H2SO4 Br2 + SO2 + 2H2O

4- IODIDE :

a mixture of HI & I2 may be formed but since HI is the most active reducing

agent it's readily oxidized to iodine which appears as violet fumes

2I- + H2SO4 2HI + SO42-

2HI + H2SO4 I2 + SO2 + 2H2O

6HI + H2SO4 3 I2 + S + 4H2O

8HI + H2SO4 4 I2 + H2S + 4H2O

I2 can be identified by exposing the evolved gas to paper moistened with

starch solution it changes into blue

c. ACTION OF CONC. H2SO4 AND MnO2:

if the solid halide is mixed with an equal quantity of precipitated manganese

dioxide and conc. H2SO4 is added then the mixture is gently warmed chlorine ,

bromine & iodine are evolved from Cl- ,Br- & I- but F- liberates HF since it has no

reducing properties

Cl- + 4H+ + MnO2 Mn2+ + 2H2O + Cl2

I- + 4H+ + MnO2 Mn2+ + 2H2O + I2

Br- + 4H+ + MnO2 Mn2+ + 2H2O + Br2

The free halogen could be identified by:

1- bleaching of a moistened colored litmus paper

2- suffocating and irritating odor

3- characteristic color of Br2 (brown), I2 (violet) and Cl2 gas (greenish

tint)

4- I2 changes starch paper into blue while Br2 turns it orange

5- Cl2 and Br2 change a starch /KI paper into blue due to the oxidation of

I- producing a blue adsorption complex of I2 with starch

25

d. CHROMYL CHLORIDE TEST:

"THIS IS A SPECIFIC TEST FOR CHLORIDE EVEN IN PRESENCE OF OTHER

HALIDES"

N:B it's classified as one of the dry reactions because it's carried out on the solid

sample

1- the solid chloride is mixed with three times its weight of powdered

potassium dichromate in a test tube

2- an equal bulk of conc. H2SO4 is added then the tube is attached to another

tube by a delivery tube dipped into a NaOH solution

4Cl- + Cr2O72- + 6 H+ 2 CrO2Cl2 + 3H2O

3- the deep red vapors of chromyl chloride CrO2Cl2 which are evolved are

passed into sodium hydroxide solution

4- the resulting yellow solution in the test tube contains sodium chromate

CrO2Cl2 + 4 OH- CrO42- + 2 Cl- + 2H2O

5- chromate is confirmed by:

a] perchromic acid test : it's carried out by acidifying with dilute H2SO4 ,adding 1-2

ml amyl alcohol followed by a little H2O2 solution the organic layer is colored blue

2 CrO42- + 2H+ Cr2O7

2- + H2O

Cr2O72- + 7H2O2 2 CrO8

3- + 5 H2O + 4H+

b] lead acetate: produce yellow ppt. of lead chromate

CrO42- + Pb2+ PbCrO4 (YELLOW)

N:B

some Cl2 may also be liberated owing to the reaction which can decrease the

sensitivity of the test

6Cl- + Cr2O72- + 14H+ 3Cl2 + 2Cr3+ + 7H2O

bromides and iodides give rise to the free halogens which yield colorless

solutions with NaOH due to the formation of hypobromide and hypoiodide

6Br- + Cr2O72- + 14H+ 3Br2 + 2Cr3+ + 7H2O

6I- + Cr2O72- + 14H+ 3I2 + 2Cr3+ + 7H2O

Br2 + 2OH- OBr- + Br- +H2O

I2 + 2 OH- OI- + I- + H2O

if the ratio of iodide to chloride exceeds 1:15 the chromyl chloride formation

is largely prevented and Cl2 is evolved

26

fluorides give rise to the volatile CHROMYLFLOURIDE (CrO2F2 ) which is

decomposed by water and hence it should be absent or removed

nitrites and nitrates interfere as nitrosyl chloride may be formed

5- WET REACTIONS:

- all fluorides are insoluble except Na+, K+, NH4+ & Ag+ salts

- all chlorides, bromides and iodides are soluble except Ag+, Hg22+ & Cu+ salts,

their lead salts are insoluble in cold water but soluble in hot water

A- Reaction with AgNO3:

1- F- :

no ppt. since AgF is soluble in water

2- Cl - :

White curdy ppt. of AgCl is formed

Ag+ + Cl- AgCl

insoluble in dil. HNO3 soluble in KCN & Na2S2O3

soluble in dilute ammonia solution to give the ammine complex

AgCl + 2 NH3 [Ag (NH3)2]Cl

AgCl is reprecipitated upon treatment of the ammine complex with acid

[Ag (NH3)2]Cl + 2H+ 2NH4+ + AgCl

3- Br- :

a curdy pale yellow ppt. of AgBr is formed

Ag+ + Br- AgBr

insoluble in dil. HNO3 soluble in KCN & Na2S2O3

sparingly soluble in dilute but readily soluble in conc. ammonia solution

AgBr + 2 NH3 [Ag (NH3)2]Br

4- I- :

a curdy yellow ppt. of AgI is formed

Ag+ + I- AgI

insoluble in dil. HNO3 soluble in KCN & Na2S2O3

insoluble in dilute ammonia but very slightly soluble in conc. ammonia

solution

27

N:B

there is a periodicity in character of the three silver halides since AgI is the

most insoluble one followed by AgBr and AgCl

therefore AgCl will dissolve in dilute ammonia followed by AgBr in conc.

Ammonia solution but AgI does not

this is attributed to that the conc. Of silver ions produced from the dissociation of

silver ammine complex according to its instability constant is insufficient to exceed

the high solubility product of AgCl (soluble) , approach that of AgBr ( partially

soluble) but exceeds that of AgI (insoluble)

therefore when Br- or I- solutions are added to AgCl yellow ppt. of AgBr or AgI

are formed

AgCl + Br- or I- AgBr or AgI + Cl- // AgBr + I- AgI + Br-

b. Reaction with BaCl2:

1- F- :

white gelatinous ppt. of BaF is formed

Ba2+ + F- BaF

partially soluble in dilute HCl or HNO3 insoluble in acetic acid

2- Cl - : no ppt

3- Br - : no ppt.

4- I - : no ppt.

c. Reaction with FeCl3:

1- F - :

white crystalline ppt. of the complex salt is formed

Fe3+ + 6F - [FeF6]3-

sparingly soluble in water

2- Cl - :

no reaction

3- Br - :

no reaction

4- I - :

28

reacts with FeCl3 due to its strong reducing action with the liberation of I2

(brown color)

2I- + 2Fe3+ I2 + 2Fe2+

d. Reaction with Pb( CH3COOH)2:

1- F - :

white ppt. of PbF2 is formed , sparingly soluble in cold water soluble in hot water

Pb2+ + F - PbF2

2- Cl - :

white ppt. of PbCl2 is formed , sparingly soluble in cold water soluble in hot

water

Pb2+ + Cl - PbCl2

3- Br - :

white ppt. of PbBr2 is formed , sparingly soluble in cold water soluble in hot

water

Pb2+ + Br - PbBr2

4- I - :

bright yellow ppt. of PbI2, soluble in hot ware and recrystallizes on cooling as

golden spangles

Pb2+ + I - PbI2

e. CHLORINE WATER TEST:

PRINCIPLE :

when chlorine water reagent is added drop wise to a solution of iodide or bromide

chlorine water oxidizes I- or Br- into I2 or Br2 which are liberated as brown

solutions

both I2 & Br2 liberated can be extracted with chloroform or carbon tetrachloride

and I2 appears as violet color while Br2 appears as brown color in the organic layer

2 I- + Cl2 I2 + 2Cl-

2 Br- + Cl2 Br2 + 2 Cl-

29

N:B

1- Chlorine water should be added drop wise why??

as excess chlorine water converts Br2 into yellow bromine monochloride

or into colorless hypobromus acid or bromic acid and the organic layer turns pale

yellow or colorless

Br2 + Cl2 2 BrCl (yellow)

Br2 + Cl2 (xss) +2H2O HOBr + 2 HCl

Br2 + 5 Cl2 + 6 H2O 2 HBrO3 + 10 HCl

as excess chlorine water converts I2 into colorless iodic acid

I2 + 5 Cl2 + 6 H2O 2 HIO3 + 10 HCl

2- Fluoride and chloride don't interfere with chlorine water test

4- SPECIAL TESTS:

1- F - :

Boron fluoride test:

when fluoride is mixed with borax and moistened with conc. H2SO4

the formed HF and boric acid react to produce boron fluoride gas which

when introduced into the flame gives green tinge to the flame

Na2B4O7 + H2SO4 + 5H2O 4 H3BO3 + Na2SO4

2NaF + H2SO4 2 HF + Na2SO4

H3BO3 + HF BF3 + 3H2O

2- I - :

a- oxidation test:

iodide is readily oxidized in acid medium (dil. H2SO4 ) with nitrite solution or

H2O2 into free I2

2I- + 2NO2- + 4H+ I2 + 2 NO + 2H2O

2 I- + H2O2 I2 + 2 H2O

b- Copper test:

I- reacts with Cu2+ forming a white ppt. of Cu2I2 and I- is oxidized to free I2

thus a white ppt. in brown solution is formed on treating I- with CuSO4 solution

2 Cu2+ + 4 I- Cu2I2 + I2

N:B * if CuSO4 solution is mixed with SO32- or FeSO4 all the I- is precipitated as

Cu2I2 and no free I2 is formed

31

c- mercuric chloride test:

I- reacts with HgCl2 to give mercuric iodide which is yellow scarlet red ppt. that

dissolves in excess iodide solution forming soluble colorless complex

the formed complex in alkaline medium is termed " NESSLER'S REAGENT " which

is used to test for ammonia

HgCl2 + 2 I- HgI2 +2 Cl-

HgI2 + 2I- [HgI4]2-

5- ANALYSIS OF MIXTURES:

1- Mixture of F - , Cl - , Br - & I- :

(a) Fluoride is separated by treating the mixture solution acidified with acetic acid

with barium nitrate then centrifuge

white ppt. centrifugate

BaF2 (confirm by conc. H2SO4test ) Cl- , Br- & I--

(b) for the centrifugate carry out chlorine water for both I- & Br-

(c) for Cl- carry out chromyl chloride test on a solid sample

6- BIOLOGICAL & ENVIRONMENTAL SIGNIFICANCE:

A) Chloride:

Chloride is a chemical the human body needs for metabolism (the process of

turning food into energy).[1] It also helps keep the body's acid-base balance. The

amount of chloride in the blood is carefully controlled by the kidneys

In the petroleum industry, the chlorides are a closely monitored constituent of

the mud system. The increase of the chlorides in the mud system could indicate the

possibility of drilling into a high-pressure saltwater formation. Its increase can also

indicate the poor quality of a target sand.

Chloride is also a useful and reliable chemical indicator of river / groundwater

fecal contamination, as chloride is a non-reactive solute and ubiquitous to sewage

& potable water

31

B) Bromide:

Bromide compounds, especially potassium bromide, were frequently used as

sedatives in the 19th and early 20th century. This gave the word "bromide" its

colloquial connotation of a boring cliché, a bit of conventional wisdom overused as

a sedative.

The bromide ion is antiepileptic, and bromide salts are still used as such,

particularly in veterinary medicine. The renal half-life of bromide in humans (12

days) is long compared with many pharmaceuticals, making dosing difficult to

adjust (a new dose may require several months to reach equilibrium). Bromide ion

concentrations in the cerebrospinal fluid are about 30% of those in blood, and are

strongly influenced by the body's chloride intake and metabolism.

Chronic toxicity from bromide can result in bromism, a syndrome with multiple

neurological symptoms. Bromide toxicity can also cause a type of skin eruption.

See potassium bromide.

Lithium bromide was used as a sedative beginning in the early 1900s, but it fell

into disfavor in the 1940s when some heart patients died after using it as a salt

substitute. Like lithium carbonate and lithium chloride it was used as treatment

for bipolar disorder

Bromide is needed by eosinophils (white blood cells of the granulocyte class,

specialised for dealing with multi-cellular parasites), which use it to generate

antiparasitic brominating compounds by the action of eosinophil peroxidase,

a haloperoxidase enzyme which is able to use chloride, but preferentially uses

bromide when available.[3] Despite this use by the body, bromide is not known to

be strictly necessary for life, as its functions may generally be replaced (though in

some cases not as well) by chloride.

Bromide salts are also sometimes used in hot tubs and spas as mild

germicidal agents, using the action of an added oxidizing agent to generate in

situ hypobromite, in a similar fashion to the peroxidase in eosinophils.

The average concentration of bromide in human blood is 5.3±1.4 mg/L and

varies with age and gender.[4] Much higher levels may indicate exposure to

brominated chemicals (e.g. methyl bromide). However, bromide occurs in relatively

high concentration in seawater and many types of seafood, and bromide

concentrations in the blood are heavily influenced by seafood contributions to the

diet

32

In some countries, bromide salts remain available in a liquid form at

pharmacies. Although since the 1950s they have been removed as over-the-

counter sedatives in most countries in the West.

It was rumored in particular by British troops during World War II that bromide

was regularly added to their tea to reduce incidence of erections for males

(see anaphrodisiac). Historically this actually had been bromide's initial

pharmacological use a century before. However, such an action is common to all

effective sedatives[citation needed] and not known to be especially particular to

bromide. In addition, stories of anaphrodesiacs being used for troops also were

told about a number of other chemical compounds, such as nitrates, and there has

not been good evidence produced for any of them

C) Iodide:

Iodide is a mild reducing agent, which is a chemical term for an antioxidant. Its

antioxidant properties can be expressed quantitatively as a redox potential:

I− ⇌ 1/2 I2 + e− (electrons) = - 0.54 Volt vs NHE

Because iodide is easily oxidized, some enzymes readily convert it

into electrophilic iodinating agents, as required for the biosynthesis of myriad

iodide-containing natural products. Iodide can function as an

antioxidant reducing species that can destroy reactive oxygen species such

as hydrogen peroxide:

2 I− + Peroxidase + H2O2 + tyrosine, histidine, lipid, etc. → iodo-Compounds + H2O

+ 2 e− (antioxidants)

In an evolutionary sense, iodides are probably one of the most ancient

antioxidants. Organoiodine compounds, called organic iodides, are similar

structurally to organochlorine and organobromine compounds, but the C-I bond is

weaker. Many organic iodides are known, but few are of major industrial

importance. Iodide compounds are mainly produced as nutritional supplements.

Biomedical use

The thyroxin hormones are essential for human health, hence the usefulness

of iodized salt.

6 mg of iodide a day can be used to treat patients with hyperthyroidism due to

its ability to inhibit the organification process in thyroid hormone synthesis, the so-

33

called Wolff-Chaikoff Effect. Prior to 1940, iodides were the predominant

antithyroid agents. In large doses, iodides inhibit proteolysis of thyroglobulin,

which permits TH to be synthesized and stored incolloid, but not released into the

bloodstream.

This treatment is seldom used today as a stand-alone therapy despite the rapid

improvement of patients immediately following administration. The major

disadvantage of iodide treatment lies in the fact that excessive stores of TH

accumulate, slowing the onset of action of thioamides (TH synthesis blockers).

Additionally, the functionality of iodides fade after the initial treatment period. An

"escape from block" is also a concern, as extra stored TH may spike following

discontinuation of treatment

34

Cyanogen group

CN- / SCN- / [Fe(CN)6]4- / [Fe(CN)6]3-

Cyanide / thiocyanate / ferrocyanide / ferricyanide

* All cyanide containing anions are highly poisonous

in all experiments in which the gas is likely to be evolved or those in which

cyanides are heated should be carried out cautiously in the fume cupboard.

* Cyanide ion has strong tendency to the formation of complexes which may be

"double cyanides" or" complex cyanides".

a- Argentocyanide complexes:

when a ppt. is formed upon reacting CN – with Ag+

at first a white turbidity is formed which is Ag CN

if CN- ions are present in excess a soluble complex is formed

Ag CN + CN- [Ag (CN)2]-

addition of excess metal ions disturb the complex formation with the

reprecipitation of insoluble cyanide

Ag+ + [Ag (CN)2]- 2 Ag CN

b- complex cyanides:

- Stable metallo-cyanogen complexes can be formed by reacting FeSO4 with CN-

in alkaline medium to give stable ferrocyanide complex. Similar complex is formed

with Fe3+ to give ferricyanide.

Fe2+ + 6CN- [Fe(CN)6]4-

Fe3+ + 6CN- [Fe(CN)6]3-

Therefore ferrocyanide and ferricyanide are considered to be stable

complexes from cyanide ions

- when cyanides are heated with polysulphides e.g. (NH4)2Sx or thiosulphate

(S2O32-) they give thiocyanate ion.

CN- + (NH4)2Sx (NH4)2Sx-1 + SCN-

CN- + S2O32- SO3

2- + SCN-

- Also Co2+ can form stable complexes with CN-.

35

1- PARENT ACIDS:

1- Hydrocyanic acid (HCN):

- VERY POISONOUS WEK MONOPROTIC ACID

- Colorless volatile liquid with and odor of bitter almond.

- Not stable in solution due to the formation of ammonium formate.

- Any dilute mineral acid can replace HCN in its solution.

* On passing CO2 in CN- solution HCN is produced with HCO3-

CN- + CO2 + H2O HCN + HCO3-

* Cyanide solutions have an alkaline reaction owing to hydrolysis.

KCN + H2O HCN + KOH

2- Thiocyanic acid (HSCN):

- Colorless toxic liquid with unpleasant odor.

- IT'S AS STRONG AS HCl but unstable

- Soluble in ether after the addition of HCl to an aqueous solution of

thiocyanate

On standing its aqueous solution decomposes to HCN and yellow solid polymer

3 HSCN HCN + H2N2C2S3

3- Ferrocyanic acid ( H4[Fe(CN)6] ):

- White crystalline solid.

- Its aqueous solution is strongly acidic [ the first two protons are nearly

completely ionized ]

4- Ferricyanic acid ( H3[Fe(CN)6] )

- Brownish crystalline solid.

- Its aqueous solution is strongly acidic [ the three protons are nearly

completely ionized ]

2- DRY REACTIONS:

a. ACTION OF DIL. HCl:

1- CYANIDE:

evolution of HCN gas which is colorless with characteristic bitter

almond odor

CN- + HCl HCN + Cl-

36

it could be identified by

1- converting HCN evolved into SCN- by exposing the evolved HCN gas to

a paper moistened with ammonium polysulphide

the resulted SCN- can be treated by adding a drop of FeCl3 solution

a blood red color is produced

2- by passing the evolved gas into AgNO3 solution white ppt. of Ag

CN

HCN + AgNO3 Ag CN + HNO3

insoluble in dil. HNO3 soluble in ammonia solution

3- Prussian blue test:

- the evolved HCN gas is passed into NaOH solution

- add drops of FeSO4 solution , heat to boiling

HCN is converted into ferrocyanide which can be tested by adding drops of

FeCl3 solution to produce a Prussian blue ppt.

2- Thiocyanate:

no reaction

3- Ferrocyanide:

no gases are evolved but hydroferrocyanic acid may precipitate

[Fe(CN)6]4- + 4H+ H4[Fe(CN)6]

4- Ferricyanide:

no gases are evolved but hydroferricyanic acid may precipitate

[Fe(CN)6]3- + 3H+ H3[Fe(CN)6]

b. ACTION OF CONC. H2SO4 :

1- CYANIDE:

all cyanides are decomposed on heating with evolution of carbon

monoxide that may be ignited and burns with a blue flame

M(CN)2 + 3H2SO4 +2H2O MSO4 +2NH4HSO4+ 2CO

2- Thiocyanate:

decomposition with evolution of carbonyl sulphide which burns with a

blue flame

SCN- + 4H++2SO42- + H2O NH4

++ 2HSO4- + 2COS

3- Ferrocyanide:

on heating , decomposition with evolution of carbon monoxide which

burns with a blue flame. SULPHUR DIOXIDE IS ALSO PRODUCED.

37

[Fe(CN)6]4- +6H2O + 22H+ + 10 SO42- Fe2+ + 6NH4++ 10 HSO4

- + 6CO

Fe2+ + 4H++SO42- SO2 + 2H2O + 2 Fe3+

4- Ferricyanide:

on heating , decomposition with evolution of carbon monoxide which

burns with a blue flame.

[Fe(CN)6]3- +6H2O + 22H+ + 10 SO42- Fe3+ + 6NH4++ 10 HSO4

- + 6CO

3- WET REACTIONS:

- all cyanides are insoluble except Na+, K+, NH4+ ,Ca2+, Ba2+, Sr2+& Hg2+

- all thiocyanates are soluble except Ag+, Hg22+ & Cu+ salts, their lead salts

are insoluble in cold water but soluble in hot water

- all ferrocyanides & ferricyanides are insoluble except Na+,K+,NH4+,Ca2+,

Ba2+& Sr2+

a. Reaction with AgNO3:

1- CN- :

white ppt. of silver cyanide is formed

Ag++ CN- Ag CN

insoluble in dil. HNO3 soluble ammonia solution

soluble in excess cyanide solution to give the dicyanoargentate complex

Ag CN+ CN- [Ag (CN)2]-

2- SCN - :

White curdy ppt. of silver thiocyanate is formed

Ag++ SCN- Ag SCN

insoluble in dil. HNO3 soluble in ammonia solution

3- [Fe(CN)6]4-:

white ppt. of silver ferrocyanide is formed

4 Ag+ + [Fe(CN)6]4- Ag4 [Fe(CN)6]

insoluble in dil. HNO3 insoluble in ammonia solution

soluble in KCN & Na2S2O3

4- [Fe(CN)6]3-:

orange red ppt. of silver ferricyanide is formed

3 Ag+ +[Fe(CN)6]3- Ag3[Fe(CN)6]

insoluble in dil. HNO3

soluble in dilute ammonia

38

N:B

The solubility of silver ferricyanide ppt. can be used for the separation of

ferrocyanide and ferricyanide when present in a mixture.

Oxidation of the white ppt. of silver ferrocyanide by warming with few drops of

conc. HNO3 leads to

formation of the orange red ppt. of silver ferricyanide which becomes soluble in

dilute ammonia solution

b- Reaction with BaCl2:

No observed reaction

c- Reaction with FeCl3: "differentiating reaction"

1- CN- :

ferric cyanide will be formed from dilute solution as a ppt. which dissolves in

excess cyanide forming ferricyanide.

Fe3++ 3 CN- Fe(CN)3 + 3CN- [Fe(CN)6]3-

2- SCN - :

" THIS TEST IS SPECIFIC FOR Fe3+ & SCN- IN ABSENCE OF OTHER INTERFERING

IONS"

a cold acidic solution of SCN- is treated with FeCl3 reagent a blood red color is

produced which is extractable with ether.

Fe3++ SCN- Fe(SCN)3 or [Fe(SCN)6]3- or [Fe(SCN)]2+

* In order to increase the sensitivity of the test the following precautions are

done

ensure the presence of iron in the Fe3+ state

acidification of the medium (dilute HCl is preferable)

cooling of the solution before testing

removal of interfering ions by precipitation or complexation

* Fluoride, phosphate, oxalate and tartrate must be absent as they bleach the color

as more stable complexes are formed

e.g. Fe(SCN)3 + 6F- [Fe F6]3- + 3 SCN-

Fe(SCN)3 + 3(COO)22- [Fe{(COO)2}3]3- + 3 SCN-

39

* mercuric ions must be absent as they bleach the color by forming the more

stable non-dissociated mercuric thiocyanate

2Fe(SCN)3 + 3Hg2+ 2Fe3+ + 3 Hg(SCN)2

* iodide also interferes by being oxidized by ferric ion to the brown colored iodine

2I- + 2Fe3+ I2 + 2Fe2+

the presence of nitrites should be avoided because in acidic solution they form

nitrosyl thiocyanate (NOSCN) which yields a red color disappearing on heating

3- [Fe(CN)6]4-:

a Prussian blue ppt. is formed from acidic solution of ferrocyanide

Fe3+ + 3 [Fe (CN)6]4- Fe4 [Fe(CN)6]3

insoluble in dil. HCl

soluble in alkali hydroxides [brown ferric hydroxide being formed]

4- [Fe(CN)6]3-:

brown color is formed of the non-ionized ferric ferricyanide

Fe3+ + [Fe(CN)6]3- Fe[Fe(CN)6]

N:B

this test can be used to differentiate between ferrocyanide and ferricyanide

when present in a mixture.

d- Reaction with FeSO4:

1- CN- :

ferrous cyanide will be formed from dilute solution as a yellow brown ppt. which

dissolves in excess cyanide forming ferrocyanide.

Fe2++ 2 CN- Fe(CN)2 + 4 CN- [Fe(CN)6]4-

2- SCN - :

NO REACTION

3- [Fe(CN)6]4-:

a white ppt. of ferrous ferrocyanide

2 K+ + Fe2+ + [Fe (CN)6]4- K2Fe [Fe(CN)6]

4- [Fe(CN)6]3-:

forms a similar blue ppt. " turnbull's blue" as that of Prussian blue but differ in

the distribution of iron [different oxidation state]

K+ + Fe2+ + [Fe (CN)6]3- KFe [Fe(CN)6]

41

e- Reaction with CuSO4:

1- CN- :

! IN ACIDIC MEDIUM cyanide like iodide reacts with Cu2+ which oxidizes cyanide

into Cyanogen (CN)2 [ HIGHLY POISONOUS GAS].

Cu2++ 2 CN- Cu (CN)2 (GREENISH YELLOW)

2 Cu(CN)2 Cu2(CN)2 (WHITE) + (CN)2

Cu2(CN)2 + 4CN- 2 [Cu(CN)3]2-

Cuprocyanide complex "soluble"

As a conclusion of this reaction cupric ions react with excess cyanide ions to

form the soluble Cuprocyanide complex and Cyanogen

2 Cu2++ 8 CN- 2 [Cu(CN)3]2- + (CN)2

! IN ALKALINE MEDIUM Cyanogen is converted to CN- and cyanate [CNO-]

(CN)2+ 2 OH- CN- + CNO- +H2O

2- SCN - :

at first green color which changes into a black ppt. of cupric thiocyanate

Cu2+ + SCN- Cu(SCN)2 " black"

with excess reagent Cu(SCN)2 decomposes gradually to white cuprous

thiocyanate [Cu2(SCN)2] and separation of thiocyanogen as a gummy mass

2 Cu(SCN)2 Cu2(SCN)2 "white" + (SCN)2 "gummy mass"

in the presence of a reducing agent e.g. sulphurous acid the addition of Cu2+

to SCN- would directly precipitate the white cuprous salt

2Cu2+ +H2SO3 + 2 SCN- + H2O Cu2(SCN)2 +4 H+ + SO42-

3- [Fe(CN)6]4-:

brown ppt. of cupric ferrocyanide insoluble in dilute acids

2 Cu2+ + [Fe (CN)6]4- Cu2 [Fe(CN)6]

4- [Fe(CN)6]3-:

green ppt. of cupric ferricyanide insoluble in dilute acids

3 Cu2+ + 2 [Fe (CN)6]3- Cu3 [Fe(CN)6]2

f- Reaction with CoNO3:

1- CN- :

a buff ppt. of cobaltous cyanide dehydrate is formed at first which dissolves in

excess cyanide forming soluble cobaltocyanid complex.

Co2++ 2 CN- + H2O Co(CN)2. 2 H2O + 4 CN- [Co(CN)6]4-

41

2- SCN - : " Vogel's reaction"

a characteristic blue color appears owing to the formation of

tetrathiocyanatocobaltate ion

Co2++ 4 SCN- [Co(SCN)4]2-

if amyl alcohol or diethyl ether is added the free acid H2[Co(SCN)4] is formed

and dissolved by the organic solvent [ the organic layer is colored blue]

N:B

other Cyanogen ions form ppts.

the test is rendered more sensitive if the solution is acidified with conc. HCl

where the equilibrium 2H+ + [Co(SCN)4] H2[Co(SCN)4]

is shifted towards the formation of the free acid which then can be extracted by

ether.

to prevent the interference of iron ammonium or sodium fluoride is added to the

test.

3- [Fe(CN)6]4-:

a grayish green ppt. of cobalt ferrocyanide

2 Co2+ + [Fe (CN)6]4- Co2 [Fe(CN)6]

4- [Fe(CN)6]3-:

red ppt. of cobalt ferricyanide

3 Co2+ +2 [Fe (CN)6]3- Co3 [Fe(CN)6]2

4- SPECIAL TESTS:

1- CN - :

a- Prussian blue test:

- cyanide solution is made strongly alkaline with NaOH solution

- add drops of freshly prepared FeSO4 solution , heat to boiling

CN- is converted into ferrocyanide which can be tested by adding drops of

FeCl3 solution to produce a Prussian blue ppt.

Fe2++ 2 CN- Fe(CN)2 + 4 CN- [Fe(CN)6]4-

Fe3+ + 3 [Fe (CN)6]4- Fe4 [Fe(CN)6]3

b- iron thiocyanate test:

- this test depends on the direct combination of alkali cyanides with sulphur

(ammonium polysulphide).

a blood red color is produced upon addition of FeCl3 .

42

CN- + (NH4)2Sx (NH4)2Sx-1 + SCN-

CN- + S2O32- SO3

2- + SCN-

Fe3++ SCN- Fe(SCN)3 or [Fe(SCN)6]3- or [Fe(SCN)]2+

this test is applicable to CN- in presence of S2- or SO32-

IF SCN- is originally present then cyanide must be isolated first by precipitation

e.g. as zinc cyanide

2- SCN - :

a- reduction test:

this reaction depends on the reduction of SCN- with metallic zinc and dilute

acid into H2S and HCN which can be tested for.

Zno + 2H+ 2(H) + Zn2+

2 SCN- + 4(H) 2 HCN + H2S + S2-

b- iron thiocyanate test:

this test can also be applied for SCN- but omitting the sulphide

3- [Fe(CN)6]4-:

a- as mild reducing agent:

ferrocyanide can be oxidized to ferricyanide by oxidizing agents e.g. MnO4- ,

NO3- , H2O2 & Cl2

2 [Fe (CN)6]4- + Cl2 2 [Fe (CN)6]3- + 2 Cl-

b- Thorium nitrate & calcium chloride test:

ferrocyanides give white ppt. with Th(NO3)4 & CaCl2 solutions

[Fe (CN)6]4- + Ca2+ + 2K+ K2Ca [Fe (CN)6]

[Fe (CN)6]4- + Th4+ Th [Fe (CN)6]

this test can be used to distinguish ferrocyanide ions from ferricyanide and

thiocyanate which don't react

4- [Fe(CN)6]3-:

a- as oxidizing agent:

ferricyanide can oxidize I- into a brown colored I2 which can be identified by

starch (blue color) or chloroform (violet color).

2 [Fe (CN)6]3- + 2I- 2 [Fe (CN)6]4- + I2

5- ANALYSIS OF MIXTURES:

1- Mixture of CN - , SCN - , [Fe (CN)6]4- & [Fe (CN)6]3- :

(a) cyanide must be tested first then removed from the mixture .

43

this is done depending on its strong affinity to protons , low ionization and

volatility of HCN this could be done by passing CO2 in the mixture solution using

acetic acid or NaHCO3 and heat until no more HCN is evolved which can be

confirmed by

passing in AgNO3 solution acidified passing in NaOH adding FeSO4 solution,

heat,

with dil. HNO3 WHITE PPT. add HCl then FeCl3 PRUSSIAN BLUE

PPT.

(b) to the remaining solution after removal of CN- acidify with dil. HCl , cool and

add FeCl3 solution and centrifuge

deep blue ppt. centrifugate

[Fe (CN)6]4-

blood red color extractable brown solution with ether + SCN- SnCl2

blue ppt.

[Fe (CN)6]3-

N:B

the analysis of a mixture containing the four cyanide containing anions can also

be applied for mixture containing two or three of them.

the application of knowledge of solubility difference of the precipitates with

various reagents can be also applied

44

2- Mixture of SCN - , Cl - , Br- & I- :

(a) thiocyanate is tested for by reaction with FeCl3 to give blood red color which is

extractable with ether and removed [ in presence of I- I2 is also formed which

can be extracted with CHCl3 ]

the blue complex formed with Co2+ can also be used to detect and remove

SCN- by extraction with ether or amyl alcohol.

(b) the halides are tested for in the usual way after the removal of SCN- as it

interferes with their precipitation

SCN- is removed by igniting the mixture till no more blackening or no odor of

burnt sulphur is observed.

the residue will contain only Cl-, Br- , I-

test for Cl- by chromyl chloride test

test for I- & Br- by chlorine water test

45

Arsenic and phosphorous containing anions

Arsine-containing anions

AsO4 3- / AsO3

3-

Arsenate / Arsenite

* Arsine - containing acids and salts are highly poisonous

1- PARENT ACIDS:

1- Orthoarsenic acid (H3AsO4):

- Crystalline solid

- Its aqueous solution is a moderately strong acid, slightly weaker than

phosphoric acid.

- It has the tendency for condensation and formation of pyroarsenic (H2As2O7)

and metaarsenic HAsO3 by gentle heating.

- H2O - H2O

2 H3AsO4 H4As2O7 2 HAsO3

(Orthoarsenic a') + H2O (pyroarsenic a') + H2O (metaarsenic a')

* arsenic acid and arsenate ion are mild oxidizing agents

* three series of salts of arsenates exist

1ry arsenates H2AsO4- 2ry arsenates HAsO4

2- 3ry arsenates AsO43-

2- arsenious acid (H3AsO3):

- it cannot be isolated as such because of thermal decomposition to the anhydride

As2O3 ( sometimes written as As4O6).

- The oxide is slightly soluble in water giving orthoarsenious acid and meta

arsenious acid.

+6 H2O

As4O6 4H3AsO3 4 HAsO2 + 4H2O

(orthoarsenious a') (metaarsenious a')

* arsenious acid and arsenite ion are mild reducing agents

* two series of salts of arsenites exist

orthoarsenites H2AsO3- metaarsenites AsO2

-

46

Reduction of As5+ & As3+:

Pentavalent arsenic salts or anions containing it can be reduced first to the

trivalent arsenious or the corresponding anion containing it and finally to the

metallic form

As5+ +2e- As3+ +3e- Aso

the reduction can be made using reducing agents with lower redox potential

e.g. saturated solution of stannous chloride [a powerful reducing agent in the

presence of conc. HCl].

As5+ +Sn2+ As3+ +Sn4+

2As3+ +3 Sn2+ 2 Aso + 3 Sn4+

Redox reaction with iodine /iodide:

- [I2 /I-] system has potential (EO=+0.54 V) while [AsO43- /AsO3

3-] system has a

more or less slightly higher potential (EO=+0.57 V) so arsenate ions oxidizes

iodide into iodine but the redox reaction is reversible due to the narrow difference

in EO values of the two redox systems.

Therefore the presence of acid converts the oxidized form (AsO43-) to the

reduced form while the presence of mild alkali e.g. NaHCO3 converts the reduced

form (AsO33-) to the oxidized form.

H+

AsO43- + 2H+ + 2 I- AsO3

3- + H2O + I2

OH-

- Arsenate oxidizes iodide into iodine in acid medium while arsenite reduces

iodine into iodide in alkaline medium.

Phosphorus-containing anions

PO4 3-

Phosphate

1- PARENT ACIDS:

Orthophosphoric acid (H3PO4):

- Crystalline solid

- Its aqueous solution is ACIDIC and ionizes into

H3PO4 H+ + H2PO4- [ DIHYDROGEN PHOSHATE]

H2PO4- H+ + HPO4

2- [ MONOHYDROGEN PHOSHATE]

HPO42- H+ + PO4

3- [ TRIBASIC PHOSHATE]

47

the intermolecular loss of water from two molecules of Orthophosphoric acid will

give pyrophosphoric acid and metaphosphoric acid

- H2O - H2O

2 H3PO4 H4P2O7 2 HPO3

(Orthophosphoric a') + H2O (pyrophosphoric a') + H2O (metaphosphoric a')

* Three series of salts of phosphates exist

1ry phosphates H2PO4- 2ry phosphates HPO4

2- 3ry phosphates PO43-

1ry phosphates aqueous solution is acidic

2ry phosphates aqueous solution is slightly alkaline

3ry phosphates aqueous solution is strongly alkaline

2- DRY REACTIONS:

a. ACTION OF DIL. HCl:

1- Arsenate:

no visible reaction since arsenic acid is not volatile

2- Arsenite:

no visible reaction since arsenious acid is not volatile

3- Phosphate:

no visible reaction since phosphorous acid is not volatile

b. ACTION OF CONC. HCl:

1- Arsenate:

on hot arsenate oxidizes HCl into free Cl2 while it will be reduced to

arsenite

2Cl- + AsO43- + 4H+ Cl2 + AsO2

- + 2H2O

2- Arsenite:

it will react and vapor of arsenious chloride is evolved

3Cl- + AsO2- + 4H+ AsCl3 + 2H2O

3- Phosphate:

no visible reaction since phosphorous acid is not volatile

c. ACTION OF CONC. H2SO4 :

1- Arsenate:

no visible reaction since arsenic acid is not volatile

2- Arsenite:

on heating some reduction to SO2 may occur

48

3- Phosphate:

no visible reaction since phosphorous acid is not volatile

3- WET REACTIONS:

- all arsenates, arsenites, phosphates are insoluble in water except those of

Na+,K+,NH4+, besides the alkali dihydrogen salts as Ba(H2AsO4)2

a. Reaction with AgNO3:

1- AsO43- :

chocolate brown ppt. of silver arsenate is formed

3Ag++ AsO43- Ag3AsO4

soluble in dil. HNO3 soluble in ammonia solution

2- AsO33- :

yellow ppt. of silver arsenite is formed

3Ag++ AsO33- Ag3AsO3

soluble in dil. HNO3 soluble in ammonia solution

3- PO43-:

yellow ppt. of silver phosphate is formed

3Ag++ PO43- Ag3PO4

soluble in dil. HNO3 soluble in ammonia solution

b- Reaction with BaCl2:

1- AsO43- :

from neutral medium white ppt. of BaHAsO4is formed

Ba2++ HAsO42- BaHAsO4

from ammoniacal or dilute alkaline medium white ppt. of Ba3(AsO4)2 is formed

3Ba2++2AsO43- Ba3(AsO4)2

the precipitates are soluble in dil. acids including acetic acid

2- AsO33- :

from neutral medium white ppt. of BaHAsO3is formed

Ba2++ HAsO32- BaHAsO3

from ammoniacal or dilute alkaline medium white ppt. of Ba3(AsO3)2 is formed

3Ba2++2AsO33- Ba3(AsO3)2

the precipitates are soluble in dil. acids including acetic acid

3- PO43-:

from neutral medium white ppt. of BaHPO4is formed

Ba2++ HPO42- BaHPO4

49

from ammoniacal or dilute alkaline medium white ppt. of Ba3(PO4)2 is formed

3Ba2++2PO43- Ba3(PO4)2

the precipitates are soluble in dil. acids including acetic acid

c- Reaction with magnesia mixture:

magnesia mixture reagent is formed of MgCl2 , NH4Cl & NH4OH

Mg2+ precipitating ions

NH4OH to render the medium ammoniacal

NH4Cl to reduce OH- concentration by common ion effect to be insufficient to

precipitate Mg(OH)2

1- AsO43- :

from neutral or ammoniacal medium white ppt. of magnesium ammonium

arsenate is formed

Mg2++ NH4+ + AsO4

3- Mg (NH4) AsO4

the precipitates are soluble in dil. acids including acetic acid

2- AsO33- :

NO PPT.

3- PO43-:

from neutral or ammoniacal medium white ppt. of magnesium ammonium

phosphate is formed

Mg2++ NH4+ + PO4

3- Mg (NH4) PO4

the precipitates are soluble in dil. acids including acetic acid

N:B

if the white precipitates are treated with AgNO3 (in weak acetic acid medium)

that of the phosphate will be transformed into yellow ppt. while that of arsenate

into chocolate brown ppt. due to the transformation to the less soluble Ag3PO4&

Ag3AsO4.

d- Reaction with ammonium molybdate:

In this test a large xss of the reagent in conc. Nitric acid is added to a small volume

of the test solution acidified with nitric acid and heating gradually

1- AsO43- :

canary yellow crystalline ppt. of ammonium arsenomolybdate on boiling is

formed

51

MoO42- + 2H+ H2 MoO4 MoO3 + H2O

12 MoO3+ 3 NH4+ + AsO4

3- (NH4)3 AsO4. 12 MoO3

OR

12 MoO42- + 3 NH4

+ +24 H+ +AsO43- (NH4)3 [As(Mo3O10)4] + 12H2O

the precipitate is insoluble in dil. HNO3

the precipitate is soluble in ammonia or alkali hydroxides solution

the precipitate is soluble in excess arsenate solution

the precipitate is soluble on boiling with ammonium acetate solution AND ON

COOLING IT YIELDS WHITE PPT.

2- AsO33- :

NO PPT.

3- PO43-:

canary yellow crystalline ppt. of ammonium phosphomolybdate on cold or

warming to 40OC is formed

MoO42- + 2H+ H2 MoO4 MoO3 + H2O

12 MoO3+ 3 NH4+ + PO4

3- (NH4)3 PO4. 12 MoO3

OR

12 MoO42- + 3 NH4

+ +23 H+ +HPO42- (NH4)3 [P(Mo3O10)4] + 12H2O

the precipitate is insoluble in dil. HNO3

the precipitate is soluble in ammonia or alkali hydroxides solution

the precipitate is soluble in excess phosphate solution

the precipitate is soluble on boiling with ammonium acetate solution AND ON

COOLING IT DOES NOT GIVE WHITE PPT.

N:B

(NH4)3 AsO4. 12 MoO3 & (NH4)3 PO4. 12 MoO3 were the former formulae for the

ppts.

the correct formulae are (NH4)3 [As(Mo3O10)4] & (NH4)3 [P(Mo3O10)4] in which

Mo3O10 group replace each oxygen atom in arsenate & phosphate

Chlorides and reducing agents such as S2-, SO32-, [Fe(CN)6]4- and tartrates

seriously affect the reaction and should be destroyed before carrying out the test.

e- Reaction with H2S:

1- AsO43- :

no immediate visible change but after prolonged passage of H2S yellow ppt. of

arsenious sulphide is formed

51

the first action of H2S is to reduce the arsenate into arsenite through the

formation of thioarsenate ion H2AsO3S- which decomposes slowly into arsenious

acid and sulphur

H2AsO4- + H2S H2AsO3S- +H2O

H2AsO3S- +H+ HAsO2 + H2O +S

2 HAsO2 + 3H2S As2S3 + 4H2O

N:B

if the acid concentration is high and the stream of H2S is rapid

no preliminary reduction to arsenite occurs and arsenic pentasulphide ppt. is

produced

2H2AsO4- +5H2S + 2H+ As2S5 + 8H2O

if the solution is heated under the same conditions

a mixture of As2S3 & As2S5 is produced

2- AsO33- :

immediate yellow ppt. of arsenious sulphide is formed

2 HAsO2 + 3H2S As2S3 + 4H2O

N:B

both As2S3 & As2S5 are soluble in nitric acid , alkali hydroxides, ammonium

sulphide and ammonium polysulphide but they are insoluble in hot conc. HCl.

in case of arsenite , the solution must be strongly acidic

as if there isn't enough acid present a yellow coloration is visible only

owing to the formation of colloidal As2S3

3- PO43-:

no ppt.

f- Reaction with CuSO4:

1- AsO43- :

bluish green ppt. of cupric arsenate.

on adding an xss of NaOH the ppt. assumes a pale blue color but does

not dissolve and on boiling no red ppt. is produced

HAsO42- + Cu2+ CuHAsO4

the ppt. is soluble in mineral acids and in ammonia

2- AsO33- :

yellowish green ppt. of copper arsenite from sample solution just alkaline with

NaOH

52

on adding an xss of NaOH the ppt. dissolves to give deep blue color of

CuO.HAsO2 and on boiling red ppt. is formed due to the reduction of cupric oxide to

cuprous oxide while arsenious acid is simultaneously partially oxidized to arsenic

acid

AsO2- + Cu2+ + OH- CuHAsO3 =[CuO.HAsO2 ]

2[CuO.HAsO2 ] + H2O Cu2O+ H3AsO4+ HAsO2

3- PO43-:

bluish green ppt. of cupric phosphate.

on adding an xss of NaOH the ppt. assumes a pale blue color but does

not dissolve and on boiling no red ppt. is produced

HPO42- + Cu2+ CuHPO4

the ppt. is soluble in mineral acids and in ammonia

g- Reaction with uranyl acetate:

1- AsO43- :

light yellow gelatinous ppt. of uranyl ammonium arsenate in presence of xss

ammonium acetate

AsO43- + UO2

2+ + NH4+ UO2(NH4)AsO4

the ppt. is soluble in mineral acids but insoluble in acetic acid

2- AsO33- :

no ppt.

3- PO43-:

light yellow gelatinous ppt. of uranyl ammonium phosphate in presence of xss

ammonium acetate

PO43- + UO2

2+ + NH4+ UO2(NH4)PO4

the ppt. is soluble in mineral acids but insoluble in acetic acid

4- SPECIAL TESTS:

1- AsO43- :

a- potassium iodide test:

- to the test solution add 1 ml chloroform , 3 ml potassium iodide solution and 5 ml

conc. HCl, shake vigorously and allow to settle

a violet color of free iodine appears in the organic layer

H+

AsO43- + 2H+ + 2 I- AsO3

3- + H2O + I2

OH-

53

this test can be used to detect arsenate in presence of phosphate and arsenite

this test can be used to detect arsenate in absence of other oxidizing agents

2- AsO33- :

a- iodine test:

- to the test solution add saturated NaHCO3, few drops iodine solution in KI

the brown color of iodine disappears immediately due to the reducing effect of

arsenite.

OH-

AsO33- + H2O + I2 AsO4

3- + 2H+ + 2 I-

H+

in absence of other reducing agents this test can be used to distinguish arsenite

from arsenate or phosphate

b- Bettendorf's test:

- a few drops of the test solution are added to 4 ml of conc. HCl and 1 ml of

saturated stannous chloride solution is added then the solution is gently warmed

it becomes dark brown and finally black ppt. of arsenic is formed

3Sn2+ + 8H+ + 2AsO2- HEAT 2As + 3 Sn4++ 4H2O

strong reducing agents as SnCl2 reduce arsenate or arsenite in presence of conc.

HCl to elemental arsenic [arsenate is first reduced to arsenite]

c- Marsh's test:

- in acidic solution arsenic (III) and (V) compounds are reduced by hydrogen to the

poisonous hydrogen arsenide gas [H3As] with garlic like odor

which when heated dissociates to elementary arsenic and hydrogen

3ZnO + 9H+ + AsO33- H3As + 3 Zn2++ 3H2O

2 H3As heat 2 Aso + 3 H2

3- PO43-:

Magnesium test:

depends on reduction of the stable phosphates into phosphide [P3-] by mixing

with magnesium powder and heating in an ignition tube

PO43- +4Mg heat 4MgO + P3-

the cold mass is moistened with water

54

Phosphine gas is produced which has unpleasant odor and is inflammable

P3- + 3 H2O PH3 + 3 OH-

5- ANALYSIS OF MIXTURES:

1- Mixture of arsenite and arsenate:

ammoniacal solution of the mixture + magnesia mixture and filter

white ppt. of Mg(NH4)AsO4 FILTRATE

Wash with dil. ammonia + AgNO3 1- acidify with dil. HCl & pass H2S

immediate

acidified with acetic acid yellow ppt. of As2S3 arsenite

or 2- add 5-7 ml of 30% H2O2 + magnesia mixture

chocolate brown ppt. of Ag3AsO4 drop wise with stirring white ppt. arsenite

ARSENATE or 3- add sat. sol. Of NaHCO3 + few drops I2

disappearance of I2 brown color arsenite

2- Mixture of arsenite & phosphate:

(a) ammoniacal solution of the mixture + magnesia mixture and filter

white ppt. of Mg(NH4)PO4 FILTRATE

Wash with dil. ammonia + AgNO3 1- acidify with dil. HCl & pass H2S

immediate

acidified with acetic acid yellow ppt. of As2S3 arsenite

or 2- add 5-7 ml of 30% H2O2 + magnesia mixture

YELLOW ppt. of Ag3PO4 drop wise with stirring white ppt. arsenite

Phosphate or 3- add sat. sol. Of NaHCO3 + few drops I2

disappearance of I2 brown color arsenite

55

(b) acidified solution of the mixture with dil. HCl + H2S and filter

yellow ppt. of As2S3 FILTRATE

arsenite evaporate to dryness & dissolve in conc. HNO3

add ammonium molybdate & warm

canary yellow ppt. Phosphate

3- Mixture of arsenate & phosphate:

acidified solution of the mixture with conc. HCl + pass H2S for 5 minutes and filter

yellow ppt. of As2S5 FILTRATE

arsenate evaporate to dryness & dissolve in conc. HNO3

add ammonium molybdate & warm

canary yellow ppt. Phosphate

4- Mixture of arsenate, arsenite & phosphate:

ammoniacal solution of the mixture + magnesia mixture and filter

white ppt. of Mg(NH4)PO4 FILTRATE

& Mg(NH4)AsO4 Wash with 1- acidify with dil. HCl & pass H2S

immediate

dil. ammonia + conc. HCl yellow ppt. of As2S3 arsenite

boil & pass H2S or 2- add 5-7 ml of 30% H2O2 + magnesia

mixture

proceed exactly as mixture (3) drop wise with stirring white ppt. arsenite

or 3- add sat. sol. Of NaHCO3 + few drops I2

disappearance of I2 brown color arsenite

6- BIOLOGICAL & ENVIRONMENTAL SIGNIFICANCE:

Phosphate:

In biological systems, phosphorus is found as a free phosphate ion in solution and

is called inorganic phosphate, to distinguish it from phosphates bound in various

56

phosphate esters. Inorganic phosphate is generally denoted Pi and at physiological

(neutral) pH primarily consists of a mixture of HPO2−4 and H2PO−4 ions.

Inorganic phosphate can be created by the hydrolysis of pyrophosphate, which is

denoted PPi:

P2O4−7 + H2O 2 HPO2−4

However, phosphates are most commonly found in the form of adenosine

phosphates, (AMP, ADP and ATP) and in DNA and RNA and can be released by the

hydrolysis of ATP or ADP. Similar reactions exist for the other nucleoside

diphosphates and triphosphates. Phosphoanhydride bonds in ADP and ATP, or

other nucleoside diphosphates and triphosphates, contain high amounts of energy

which give them their vital role in all living organisms. They are generally referred

to as high energy phosphate, as are the phosphagens in muscle tissue. Compounds

such as substituted phosphines have uses in organic chemistry but do not seem to

have any natural counterparts.

The addition and removal of phosphate from proteins in all cells is a pivotal

strategy in the regulation of metabolic processes.

Phosphate is useful in animal cells as a buffering agent. Phosphate salts that

are commonly used for preparing buffer solutions at cell pHs include Na2HPO4 ,

NaH2PO4 , and the corresponding potassium salts.

An important occurrence of phosphates in biological systems is as the

structural material of bone and teeth. These structures are made of

crystalline calcium phosphate in the form ofhydroxyapatite. The hard dense enamel

of mammalian teeth consists of fluoroapatite, an hydroxy calcium phosphate where

some of the hydroxyl groups have been replaced by fluorideions.

Insect exoskeleta are constructed of chitin containing crystalline calcium

phosphate as a strengthening material.

57

Nitrogen containing anions

NO3- / NO2

-

Nitrate / Nitrite

* the nitrate ion contains nitrogen in its highest oxidation state of +5 thus reacts

only as oxidizing agent while nitrite ion contains nitrogen with oxidation number of

+3 therefore it can react either as a reducing or as an oxidizing agent

1- PARENT ACIDS:

1- Nitric acid (HNO3):

- Colorless liquid

- Its aqueous solution is strongly acidic

- It decomposes on aging to nitrogen dioxide

4 HNO3 4 NO2 + O2 + 2 H2O

2- Nitrous acid (HNO2):

- the pure acid has never been isolated due to its thermal instability

2 HNO2 NO2 + NO + H2O

however the addition of a strong acid to a solid nitrite or its solution in the cold

yields a transient pale blue liquid ( due to the presence of free HNO2 acid or its

anhydride N2O3 ) and the evolution of brown fumes of NO2

2- DRY REACTIONS:

a- ACTION OF DIL. HCl:

1- NITRATE:

no reaction

2- NITRITE:

evolution of brown fumes of nitrogen dioxide and a transient pale blue

liquid

2 NO2- + 2 H+

2 HNO2 NO2 + NO + H2O

H+ ions from dilute acids including acetic a' displace nitrous a' from its salts

the acid spontaneously decomposes to colorless nitrogen monoxide &

brownish nitrogen dioxide gases

the brown fumes intensify when getting in contact with the atmosphere due to

combination of nitrogen monoxide with the oxygen of the air

2 NO + O2 2 NO2

58

b. ACTION OF CONC. H2SO4 :

1- NITRATE:

nitric acid is formed and some of it decomposes with evolution of brown

fumes of NO2 with characteristic odor

NO3- + H+ HNO3

4 HNO3 4 NO2 + O2 + 2 H2O

When copper turnings are added and the mixture heated to boiling

the brown fumes of NO2 increases due to the reduction of nitric acid by copper

metal which is oxidized to cupric ions that imparts a blue color to the solution

2 NO3- + 4 H+ + CuO 2 NO2 + Cu2+ + 2 H2O

2- NITRITE:

the reaction is the same as with dilute HCl but it occurs with considerable

violence

on adding copper turnings the same occurs as with nitrate

2- WET REACTIONS:

- all nitrates are soluble in water

- all nitrites are soluble in water except AgNO2 which is slightly soluble

a- Reaction with AgNO3:

1- NO3- :

NO PPT.

2- NO2- :

white crystalline ppt. of silver nitrite form concentrated solutions

Ag++ NO2- AgNO2

b- Reaction with BaCl2:

1- NO3- :

NO PPT.

2- NO2- :

NO PPT.

59

c- Reaction with potassium iodide:

the test solution is acidified wit dil. H2SO4 then KI solution is added with few drops

of starch solution

1- NO3- :

NO reaction

2- NO2- :

iodine is liberated imparting blue color to the starch

2NO2- + 2I- +4H+ 2 NO + I2 + 2H2O

d- Reaction with ferrous sulphate: "brown ring test"

the test solution is acidified wit dil. H2SO4 then freshly prepared FeSO4 solution is

added

1- NO3- :

no visible change in case of using only dilute H2SO4 but on adding conc. H2SO4

cautiously down the sides of the test tube

a brown ring is formed at the interface

2- NO2- :

brown color in the whole solution if ferrous sulphate solution is not cautiously

added or a brown ring at the junction of the two liquids if cautiously added

N:B

FeSO4 reduces nitrite or nitrate ions to nitrogen monoxide (NO)

nitrate ion is not reduced except in solutions containing high H+ ion

concentration [conc. H2SO4]

XSS Fe2+ ions then combines with nitrogen monoxide to form the unstable

brownish black complex ion [Fe(NO)]2+ which is readily decomposed by heat.

3Fe2+ + NO3- +4H+ 3 Fe3+ + NO + 2H2O

OR Fe2+ + NO2- +2H+ Fe3+ + NO + H2O

And Fe2+ + NO [Fe(NO)]2+

this test differentiates nitrate ion from nitrite ion as nitrite gives the brown ring

in presence of dil. H2SO4 or even acetic acid while nitrate ion does not form the ring

except in presence of conc. H2SO4

nitrite, iodide, bromide ions interfere with the brown ring test for nitrate

61

4- SPECIAL TESTS:

1- NO3- :

a- Ammonia test:

- if solution of nitrate is boiled with zinc or aluminum metals and NaOH solution

ammonia will be evolved which can be identified by its odor or with red litmus

paper

NO3- + 4Zno + 7OH- 3 NH3+4 [ZnO2]- +2H2O

(ZINCATE ION)

3 NO3- + 8Alo + 5OH- + 2H2O 3NH3 + 8 [AlO2]-

b- Reduction test:

- in acidic solution nitrate can be reduced with Zno metal to nitrite

4 NO3- + 2 Zno + 4H+ 3 NO2

-+2 [ZnO2]- + 2H2O

2- NO2-:

1- Permanganate test:

when a dilute potassium permanganate solution is added to an acid solution of

nitrite

its pink color is bleached

2 MnO43- +5NO2

- + 6H+ 2Mn2+ + 5NO3- + 3H2O

PINK COLORLESS

2- Urea test:

when a solution of nitrite is treated with urea and the mixture acidified with dil.

HCl

nitrite is decomposed with evolution of nitrogen and carbon dioxide gases

CO (NH2)2 + 2 HNO2 2 N2 + CO2+ 3H2O

3- Thiourea test:

when a dilute acetic acid solution of nitrite is treated with little Thiourea

nitrogen is evolved and Thiocyanic acid is formed which may be identified by

the red color produced with dil. HCl and FeCl3 solution

CS (NH2)2 + HNO2 N2 + CNS- + H++ 3H2O

N:B

Thiocyanates and iodides interfere and if present must be removed either with

Ag2SO4 (solid) or dil.AgNO3 solution

61

4- Ammonium chloride test:

by boiling a solution of nitrite with xss of solid ammonium chloride

nitrogen is evolved and nitrite is completely decomposed

NH4++ NO2

- N2 + 2 H2O

5- ANALYSIS OF MIXTURES:

1- Mixture of nitrate and nitrite:

Nitrite can be tested for in presence of nitrate by

treatment with dil. HCl, KI, KMnO4, FeSO4 in dil. H2SO4

the special tests for nitrite

Nitrate cannot be tested for in presence of nitrite since nitrite gives all the

reactions of nitrate ( conc.H2SO4, brown ring test and ammonia test)

Therefore nitrite must be removed before testing for nitrate by

1- decomposition of nitrite through its brown complex with ferrous sulphate

formed in dil. H2SO4 or acetic acid by heat and shaking

[Fe(NO)]2+ heat Fe2+ + NO

2- decomposition of nitrite through its reduction to nitrogen by boiling with NH4Cl

,warming with urea and few drops of dil.H2SO4 or warming with little sulphamic

acid

HO.SO2.NH2 + HNO2 N2 + H2SO4+ H2O

2- Mixture of nitrate & bromide and /or iodide:

(a) bromide and iodide can be detected in presence of nitrate by chlorine water

test

(b) nitrate can be detected in presence of bromide and iodide by ammonia test

N:B

the brown ring test for nitrate cannot be applied in the presence of bromide and

iodide since the liberation of free halogen with conc. H2SO4 will obscure the brown

ring due to nitrate

therefore bromide and iodide must be firstly removed by either

1 addition of saturated solution of silver sulphate where AgBr and AgI are

precipitated and then filtered off , xss Ag+ is precipitated with Na2CO3 OR

2 adding potassium persulphate and dil.H2SO4 and warming to about 80oC where

the halogen is removed by boiling or extraction with organic solvent

S2O82- + Br- Br2 + 2 SO4

2-

S2O82- + I- I2 + 2 SO4

2-

62

6- BIOLOGICAL & ENVIRONMENTAL SIGNIFICANCE:

A) Nitrate:

Nitrate toxicosis in humans occurs through enterohepatic metabolism of nitrate

to ammonia, with nitrite being an intermediate]. Nitrites oxidize the iron atoms

in hemoglobin from ferrous iron (2+) to ferric iron (3+), rendering it unable to

carry oxygen[4]. This process can lead to generalized lack of oxygen in organ tissue

and a dangerous condition calledmethemoglobinemia. Methemoglobinemia can be

treated with methylene blue, which reduces ferric iron (3+) in affected blood cells

back to ferrous iron (2+).

Infants in particular are especially vulnerable to methemoglobinemia due to

nitrate metabolizing triglycerides present at higher concentrations than at other

stages of development. Methemoglobinemia in infants is known as blue baby

syndrome. There are now significant scientific doubts as to whether there is a

causal link between nitrates in drinking water and 'blue baby syndrome. Blue baby

syndrome is now thought to be the product of a number of factors, which can

include any factor which causes gastric upset, such as diarrhoeal infection, protein

intolerance, heavy metal toxicity etc., with nitrates playing a minor role. Nitrates, if

a factor in a specific case, would most often be ingested by infants in high nitrate

drinking water. However, nitrate exposure may also occur if eating, for instance,

vegetables containing high levels of nitrate. Lettuce may contain elevated nitrate

under growth conditions such as reduced sunlight, undersupply of the essential

micronutrients molybdenum (Mo) and iron (Fe), or high concentrations of nitrate

due to reduced assimilation of nitrate in the plant. High levels of nitrate

fertilization also contribute to elevated levels of nitrate in the harvested plant.

Some adults can be more susceptible to the effects of nitrate than others.

The methemoglobin reductase enzyme may be under-produced or absent in certain

people that have an inherited mutation. Such individuals cannot break down

methemoglobin as rapidly as those that do have the enzyme, leading to increased

circulating levels of methemoglobin (the implication being that their blood is not as

oxygen-rich). Those with insufficient stomach acid (including some vegetarians

and vegans) may also be at risk. It is the increased consumption of green, leafy

vegetables that typically accompany these types of diets may lead to increased

nitrate intake. A wide variety of medical conditions, including food allergies,

63

asthma, hepatitis, and gallstones may be linked with low stomach acid; these

individuals may also be highly sensitive to the effects of nitrate

B) Nitrite:

Sodium nitrite is used for the curing of meat because it prevents bacterial

growth and, in a reaction with the meat's myoglobin, gives the product a desirable

dark red color. Because of the toxicity of nitrite (the lethal dose of nitrite for

humans is about 22 mg per kg body weight), the maximum allowed nitrite

concentration in meat products is 200 ppm. Under certain conditions, especially

during cooking, nitrites in meat can react with degradation products of amino

acids, forming nitrosamines, which are known carcinogens

64

Qualitative Analysis of

Cations

By

Dr. Ibrahim A. Naguib

65

B- Cations (basic radicles)

- Cations are divided into a number of groups according to the

difference in solubilities of their chlorides, sulphides, hydroxides

& carbonates into:

Distinguishing

features

Formula of

precipitate

Ions Group

reagent

Group

Chlorides are

insoluble in

cold dilute HCl

AgCl, PbCl2

Hg2Cl2

Ag+,Pb++,Hg2+

+

Dilute HCl I

(Silver

group)

Sulphides are

insoluble in

dilute HCl (ca.

0.3 N)

HgS, PbS

Bi2S3, CuS

CdS, SnS

As2S3,

Sb2S3,

SnS2

Hg++,pb++

Bi+++, Cu++

Cd++ , Sn++

As+++ , Sb+++

Sn++++

H2S in

presence

of dilute

HCl

II

(Copper

and

arsenic

groups)

Hydroxides

are

precipitated by

aqueous NH3 in

presence of

excess of

NH4Cl

Al(OH)3

Cr(OH)3

Fe(OH)3

Al+++

Cr+++

Fe+++

Aqueous

NH3 in

presence

of NH4Cl

III

(Iron

group)

Sulphides are

precipitated by

H2S in

presence of

aqueous NH3

and NH4Cl

NiS,

CoS,

MnS,

ZnS

Ni++

Co++

Mn++

Zn++

H2S in

presence

of aqueous

NH3 and

NH4Cl

IV

(Zinc

group)

Carbonates

are

precipitated by

(NH4)2CO3 in

presence of

NH4Cl.

BaCO3,

SrCO3,

CaCO3

Ba++

Sr++

Ca++

(NH4)2CO3

in

presence

of aqueous

NH3 and

NH4Cl

V

(Calcium

group)

66

Ions are not

precipitated in

previous

groups.

Different

precipitate

d, forms

Mg++ , Na+

K+ , NH4+

No

particular

reagent

VI

(Alkali

group)

.. Principle of their separation:

- Group reagents are added systematically where the group

reagent of any group is capable not only to precipitate the cations

of its own group, but also the cations of preceding groups.

( ... a test for complete precipitation must be carried out for each

group before starting to separate subsequent groups. )

ex. group IV reagent ( NH4Cl + NH4OH + H2S ) is capable of

precipitating the sulfides of group I , II , III and IV .

- After separating the precipitate of each group, it must be

washed well before anaylsis ??

→ to remove only adsorbed impurities.

but this washing must be carried out with an electrolyte

solution(not by water) ??

→ to avoid peptisation ??

→ peptisation of a precipitate means that water favors its

transformation into the colloidal state.

- Steps of analysis:

(1) Precipitation (2) Washing

(3) Separation (4) Confirmation

Group I cations : silver group ( Pb2+ , Ag+ , Hg22+ )

* precipitation:

Q. What is the group reagent (precipitant) ??

→ cold dil. HCl (or NH4Cl)

Q. Why HCl is preferred ??

67

→ as it prevents the formation of insoluble Oxy salts of bismuth

Bi3+, tin and antimony .

e.g. BiOCl + 2H+ Bi3+ + Cl- + H2O

- Account : a slight excess of HCl should be added ??

√ to obtain complete precipitation

√ to prevent peptisation of the insoluble Chlorides .

- Account : a very large excess of HCl should be avoided ?

→ not to increase the solubility of chlorides by complex formation

e.g.

PbCl2 + Cl- [PbCl3]- + Cl- [PbCl4]

2-

( so we don't use conc. HCl )

- Account : the diluteHCl used must be cold ??

→ because PbCl2 is solublein hot water, thus it may escape to

subsequent groups.

- Account : after addition of the cold diluteHCl , the solution must

be vigorously shaken for 2 mins. ??

→ because PbCl2 can form a supersaturated solution, hence escape

to subsequent groups , so we shake to ensure complete

precipitation

- PbCl2 Ksp = 1.6 x 10-5

AgCl Ksp = 1.1 x 10-10

Hg2Cl2 Ksp = 3 x 10-18

- after complete precipitation , the precipitate is centrifuged ,

then washed with cold diluteHCl ? why ?

→ to remove surface adsorbed cations of the next groups.

- Account : HCl is preferred to water in washing ?

→ (1) as it renders PbCl2 less soluble ( by common ion effect )

PbCl2 Pb2+ + 2Cl-

(2) to avoid peptisation of the precipitate.

68

H2

* Separation :

(1) The group precipitate is dissolved in boiling water (where PbCl2

dissolve in hot water, while AgCl and Hg2Cl2 don’t) , then

centrifuge.

hot centrifugate precipitate

(may be PbCl2) (may be AgCl, Hg2Cl2)

Confirm.

(2) The precipitate (AgCl, Hg2Cl2) is dissolved in diluteNH4OH

(where AgCl dissolves & forms the water soluble Silver ammine

complex [Ag(NH3)2]+ , while Hg2Cl2 is converted into insoluble black

mixture of aminomercuric chloride + finely divided mercury .

Hg2Cl2 + 2NH3 Hg(NH2)Cl ↓ + Hgo ↓ + NH4Cl

White precipitate black precipitate

centrifuge

precipitate centrifugate

(may be Hgo + (may be [Ag(NH3)2]-)

Hg(NH2)Cl) → confirm confirm

* Confirmation:

a) For lead ions:

the hot centrifugate is cooled → needle crystals of PbCl2 . The

precipitatere is dissolved by heating & divided into 3 parts :

(1) hot solution + dilute acetic acid + K2CrO4 solution→ PbCrO4 ↓

Yellow lead chromate precipitate

[notes :the precipitate is insoluble in acetic acid but soluble in

dilute mineral acids (e.g HNO3) , why ???

as CrO4-- will transform into Cr2O7

-- where lead dichromate is

water soluble.

69

OH

2CrO4-- 2HCrO4

- Cr2O7-- + H2O

Yellow solution orange solution

√ upon addition of dilute NaOH , the yellow precipitate

transforms into red, why??

2PbCrO4 ↓ + H2O Pb2CrO5 ↓ + H2CrO4

(= PbO.PbCrO4)

red precipitate of lead basic chromate

- addition of excess NaOH dissolves the precipitate ??

PbCrO4 + 2OH- → Pb(OH)2 ↓ + CrO4--

( Pb(OH)2 is amphoteric ... dissolves in excess NaOH )

Pb(OH)2 + OH- → [HPbO2]- + H2O

Colorless soluble Plumbite

(2) hot solution + KI solution→ PbI2 ↓

Yellow precipitate

√ the yellow precipitate is soluble in hot water

√ the yellow precipitate is soluble in excess I- solution :

PbI2 + 2I- → [PbI4]2- soluble Complex

(3) hot solution + dilute H2SO4 → PbSO4 ↓

white precipitate

√ the white precipitate is soluble in NaOH and ammonium acetate

solution ( …mentioned before in anions )

(b) For silver ions :

(1) ammoniacal centifugate (contains Ag+ as [Ag(NH3)2]+)

+ dilute HNO3 → white precipitate ???

[ Ag (NH3)2 ]+ Ag+ + 2NH3

AgCl ↓ Cl- 2H+ NH4+

white ppt in the solution from HNO3

71

.redn

.redn

(2) another portion + KI solution → yellow AgI precipitate

[ Ag(NH3)2]+ Ag+ + 2NH3

AgI ↓ I-

yellow precipitate

.. Account : addition of KI solution to [Ag(NH3)2]Cl gives a yellow

precipitate??

because ionization of the silver ammine chloride complex provide

amount of Ag+ which is sufficient to exceed the Ksp of AgI thus

favour its precipitation while not sufficient to exceed Ksp of AgCl.

(c) For mercurous :

- the black precipitate is dissolved in aquaregia (3 part conc. HCl +

1 part HNO3) :

Hgo + 6HCl + 2HNO3 3HgCl2 + 2NO + 4H2O

HgNH2Cl + HCl + HNO3 HgCl2 + NH4NO3

N2O + 2H2O

i.e. the black Hgo & the white HgNH2Cl both converted to HgCl2

which is water soluble.

- destroy excess aquaregia (by boiling) then divide the solution into

2 parts:

(1) part + SnCl2 solution → white Hg2Cl2 , which turns to gray then

to black (with excess SnCl2 )

2HgCl2 + Sn2+ Hg2Cl2 ↓ + Sn4+ + 2Cl-

Hg2Cl2 + Sn2+ Hgo ↓ + Sn4+ + 2Cl-

Excess black

Notes :

√ how is aquaregia destroyed ?

NO3- + 3Cl- + 4H+ → NOCl + Cl2 ↑ + 2H2O

nitrosyl chloride

71

√ if aquaregia was not destroyed well , it will oxidize Sn2+ to Sn4+

thus Sn2+ will not be available to react with HgCl2 .

√ why HgCl2 is written as such in equations (not Hg2+) ???

because HgCl2 is weakly ionised .

√ while destroying aquaregia, if evaporation is continued to

dryness , mercury will be completely lost ???

because HgCl2 is volatile .

(2) part + KI solution→ red precipitate of HgI2

HgCl2 + 2I- → HgI2 ↓ + 2Cl-

red precipitate

HgI2 + 2I- → [HgI4]--

excess Neesler reagent

(soluble Complex)

(3) neutral (or faintly ammoniacal solution) + excess KI + copper

ethylenediamine reagent → blue violet precipitate.

Q. Account: precipitation reactions involving Hg22+ are

complicated?

Hg22+ + H2S HgS ↓ + Hgo ↓ + 2H+

Hg22+ + 2OH- HgO ↓ + Hgo ↓ + H2O

because mercurous exists in a dimeric form where a covalent

bond exists between the 2Hg22+ atoms . so, it can be written as

[Hg:Hg]2+ or Hg.Hg2+ i.e. react as if it is Hgo + Hg2+ , so its

reactions are complicated by disproportionatn.

72

OHadd 2 HCladd

Group II cations √ group 2A (copper subgroup) : Pb2+, Cu2+, Bi3+, Cd2+, Hg2+

√ group 2B (arsenic subgroup) : (As3+, As5+), (Sb3+, Sb5+), (Sn2+, S4+)

* Precipitation :

.. Group II cations are precipitated as sulphides , how ?

by addition of excess H2S solution to acidic solution of the

group Cations .

.. acidity of the solution should correspond to 0.2 – 0.3 N HCl , why

???

H2S H+ + HS-

HS- H+ + S--

HCl H+ + Cl-

(1) because if the medium is strongly acidic , [S--] will decrease due

to the decrease in ionization of H2S , due to the common ion effect

achieved by the conc. acid [ thus no complete precipitation will

occur, & we will have to add too much excess H2S to achieve

complete precipitation of PbS, CdS, SnS]

(2) because if the medium is alkaline , this will enhance ionization

of H2S, thus [S--] will increase, thus group IV sulphides (which have

higher Ksp than those of group II sulphides ) will co-precipitate

with group II sulphides .

.. How can you adjust acidity to 0.2-0.3 N HCl ???

by using methyl violet (crystal violet) indicator, where

yellow bluish green blue

(conc.) 0.2-0.3 N HCl (dil)

i.e. adjust acidity till you obtain a bluish green color of crystal

violet where;

√ if the color is yellow (this means that acidity is high, so add

water to adjust it) .

73

√ if the color is blue (this means that acidity is too low, so add

dilute HCl to adjust it)

N.B. if any precipitate forms during adjusting the acidity, keep it

with the solution where it will further precipitate with H2S

(ex. of precipitates that may form :

PbCl2 , oxysalts of bismuth , tin , antimony)

* washing of the precipitate:

by 5% NH4NO3 , why ??

√ to remove the surface adsorbed cations of subsequent groups

√ to remove H2S (otherwise Hg2+ escape to group II B .. see later)

√ to prevent peptisation & to coagulate the precipitate.

N.B after precipitation, you must boil??

√ to coagulate the precipitate

√ to expel excess H2S (where S-- may be oxidized to So↓ or SO4

-- (

which precipitate group V )

* colors of sulphides :

Black : Pbs , HgS , CuS

brown : Bi2S3 , SnS

orange : Sb2S5 , Sb2S3

Dull yellow : SnS2

Yellow : As2S5 , As2S3

Canary yellow : CdS

* Separation of subgroups A & B :

group II precipitate+ excess NaOH solution (with or without a

drop of yellow (NH4)2Sx , shake , centrifuge .

precipitate centrifugate

(group II A) (group II B)

74

.. principle : group II B sulphides are amphoteric, so they dissolve

in excess NaOH :

e.g. 2As2S3 + 4OH- 3AsS2- + AsO2

- + 2H2O thioarsenite arsenite

3As2S5 + 6OH- 5AsS3- + AsO3

- + 3H2O thioarsenate arsenate

e.g. Sb2S3 + OH- SbS2- + SbO2

- + H2O thioantimonite antimonite

Sb2S5 + OH- SbS3- + SbO3

- + H2O thioantimonate antimonite

e.g. 2SnS + 4OH- 2SnS22- + SnO2

2- + 3H2O thiostamite stannite

3SnS2 + 6OH- 2SnS32- + SnO3

2- + 3H2O thiostannate stannate

.. Account : a drop of yellow (NH4)2Sx should be added to NaOH ?

to increase solubility of SnS & Sb2S3

( which dissolves a little in NaOH alone )

Sb2S3 + S-- 2SbS2-

thioantimonite (sol)

SbS2- + Sx-- SbS3

- + Sx-12-

Thioaintimonate (sol)

oxidn. [similarly do As2S3]

SnS + S-- SnS2--

thiostannite (sol)

SnS2--+ Sx

-- SnS32- + Sx-1 2-

thiostannate (sol)

oxidn.

i.e. yellow (NH4)2Sx allow oxidation of these sulphides to the

higher valence.

75

.. Account : only few drops of (NH4)2Sx should be used ???

to avoid precipitation of yellow collidal So when the solution is

subsequently acidified .

(NH4)2Sx + 2H+ 2NH4+ + H2S + (x-1)So ↓

to avoid (NH4)2Sx + NaOH Na2S which makes Hg2+ escape

From group IIA to group IIB .

* Separation of subgroup 2A :

(Pb2+, Cu2+ , Bi3+ , Cd2+ , Hg2+)

(1) precipitate is dissolved in boiling dilute HNO3 and dilute H2SO4

Principle : HgS don't dissolve , while Bi2S3, CdS, CuS, PbS all

dissolve then Pb2+ precipitate as SO4-- as follows :

3PbS + 8H+ + 2NO3

- 3Pb2+ + 2NO + 4H2O + 3So↓

3CuS + 8H+ + 2NO3

- 3Cu2+ + 2NO + 4H2O + 3So↓

3Bi2S3 + 24H+ + 6NO3

- 6Bi3+ + 6NO + 12H2O + 4So↓

3CdS + 8H+ + 2NO3

- 3Cd2+ + 2NO + 4H2O + 3So↓

Then :

Pb2+ + SO4-- PbSO4 ↓

Then cool , centrifuge :

Centrifugate precipitate

(Bi3+, Cu2+, Cd2+ (HgS + PbSO4)

as nitrates)

- precipitate (HgS + PbSO4) + boiling saturated Ammonium acetate,

shake, centrifuge .

Centrifugate precipitate (HgS)

(soluble lead acetate complex)

76

Cl2

3HNO

)(O

H

(2) centrifugate (Bi3+ , Cu2+ , Cd2+ soluble nitrates) + excess,

NH4OH , centrifuge .

White precipitate centrifugate

(Bi(OH)3) [Cu(NH3)4]2+ blue soluble complex

[Cd(NH3)4]2+ colorless soluble

complex

* identification & confirmation :

.. For lead :

Soluble lead acetate (saturated Ammonium acetate extract) +

dilute acetic acid + K2CrO4 solution yellow precipitate of lead

chromate.

.. For mercuric :

the black precipitate(HgS) + 5ml aqua regia (3HCl : 1HNO3) , boil

till dissolved:

3HgS + 6HCl + 2HNO3 3HgCl2 + 3So↓ + 2NO + 4H2O

[ HCl, HNO3 act as team work on HgS :

HgS Hg2+ HgCl2 (slightly ionised) + S-- So↓ + NO

Cool , then carry out the following tests :

(1)- SnCl2 test : salt solution + SnCl2 white precipitate which

turns to gray then to black with excess SnCl2

HgCl2 + Sn2+ HgCl2 ↓ + Sn4+ + 2Cl-

Hg2Cl2 ↓ + Sn2+ 2Hgo ↓ + Sn4+ + 2Cl-

white excess black

(2)- KI test : salt solution + KI red HgI2 precipitate which

dissolves in excess I- giving the water soluble Neesler reagent

[HgI4]2- .

(3)- put a clean Cuo foil in HgCl2:

Hg2+ + Cuo Cu2+ + Hgo ↓

gray film of mercury

deposit on Cuo foil .

77

.. For Bismuth:

dissolve the white Bi(OH)3 precipitate in dilute HCl & divide into 3

portions .

(1)- add to one portion a large volume of water

white precipitate of BiOCl ↓

Bi3+ + Cl- + H2O BiOCl ↓ + 2H+

(2)- add to another portion sodium stannite solution

[SnCl2 + NaOH till the precipitate formed redissolve]

black precipitate of Bio↓ metal

Bi3+ + OH- Bi(OH)3 ↓

2Bi(OH)3↓ + 3[HSnO2]- 2Bio↓ + 3[HSnO3]

- + H2O

White ppt stannite black precipitate stannate

N.B this test can be made directly on the white precipitate

blackening.

(3)- the third portion + dilute HCl, heat, + excess 10% pyragallol

solution* yellow precipitate↓ of Bi(C6H3O3)

.. for copper & cadmium :

[centrifugate containing copper & cadmium ammine complexes ]

* 3 methods to analyse Cu2+ & Cd2+ in presence of one another:

(A) divide into 2 portions:

(1) the smaller portion (to test for Cu2+) :

acidify with dilute acetic acid till the blue color discharge

[Cu(NH3)4] Cu2+ + 4NH3

CH3COONH4 CH3COOH

then divide the solution to 3 parts :

a- part + pot. Ferrocyanide Cu2[Fe(CN)6]↓

chocolate precipitate

of copper ferrocyanide

b- part + KCNS solution black precipitateof Cu(CNS)2↓

78

which changes to white precipitatein presence of SO3-- ?

Cu(CNS)2 + SO3-- + Cu2+ + H2O Cu2(CNS)2↓ + SO4

-- + H+

black precipitate white precipitate

c- part + KI solution white precipitate in a brown solution

2Cu2+ + 4I- Cu2I2↓ + I2

white precipitate brown solution

(2) the larger portion (to test for Cd2+) :

add KCN solution dropwise till the blue color of copper complex

disappears then add H2S solution canary yellow precipitate of

CdS .

[Cu(NH3)4]2+ + 2CN- Cu(CN)2↓ + 4NH3

Yellow ppt

2Cu(CN)2↓ Cu2(CN)2↓ + (CN)2↑

white precipitate cyanogens

Cu2(CN)2↓ + 4CN- 2[Cu(CN)3]2-

Cuprocyanide complex

water soluble & very stable

[Now we could mask Cu2+ by transforming it into cyanide complex

which is very stable & thus no Cu2+ ions will be available in the

medium to react with H2S ]

While

[Cd(NH3)4]2+ + 2CN- Cd(CN)2↓ + 4NH3

white precipitate

Cd(CN)2 + 2CN- [Cd(CN)4]2-

Cadmicyanide complex

water solublebut unstable

[Now we transformed Cd2+ into a cyanide complex which is

unstable, thus ionise & provide Cd2+ ions in the medium which react

with H2S & give CdS (canary yellow precipitate) ].

(B) Displacement procedure:

79

.Redn

.. Divide the solution into 2 portions:

(1) the smaller portion (to test for Cu2+)

as before

(2) the larger portion + dilute H2SO4 till acidic + Feo powder & stirr

:

Cu2+ + Feo Cuo↓ + Fe2+

then the precipitated Cuo & excess Feo are filtered off, then add

dilute NH3 till alkaline + H2S canary yellow precipitate of CdS.

[Here we removed Cu2+ from the medium by transforming it to Cuo

↓ by reduction then filtration ]

(c) High acidity procedure:

.. principle: CdS is soluble in sufficient dilute HCl while, CuS is

insoluble.

.. procedure: solution (of Cu2+, Cd2+ ammine complexes) + dilute HCl

till acidic + H2S + excess HCl, boil , filter :

√ the residue consists of CuS:

dissolve in boiling HNO3 + dilute NH3 till alkaline

blue ammine complex + dilute acetic acid till the blue color

discharge + ferrocyanide chocolate precipitate

√ the filtrate (containg soluble Cd2+) + excess water + H2S

canary yellow precipitate

81

* Group II B (arsenic subgroup)

.. precipitation:

the (alkali hydroxide – polysulphide) extract of group II

precipitate will contain ( the Thio & oxysalts of arsenic, tin,

antimony + some Hg(SNa)2 + excess OH- + polysulphide Sx2- )

[Account: before separation of subgroups. A,B sulphides should be

washed well to remove H2S ? why ?

Otherwise mercuric will escepe to group IIB as Hg(SNa)2 as

seen

H2S + 2NaOH Na2S + 2H2O

Na2S + HgS Hg(SNa)2

soluble comple ]

- Now to reprecipitate group II B, add 6-8 N HCl dropwise to the

extract (till acidic to litmus paper) + some H2S : where the thio

salt:

√ 2A2S3- + 2H+ As2S5 ↓ + H2S

thioarsenate yellow precipitate

(No thioarsenite exist ? why ?

because the polysulphide added transformed it to the higher

sulphide thioarsenate)

√ 2SbS3- + 2H+ Sb2S5 ↓ + H2S

thioantimonate orange precipitate

√ SnS32- + 2H+ SnS2 ↓ + H2S

Dull yellow precipitate

- while the oxysalts :

2AsO3- + 2H+ + 5H2S As2S5↓ + 6H2O

2SbO3- + 2H+ + 5H2S Sb2S5↓ + 6H2O

SnO32- + 2H+ + 2H2S SnS2↓ + 3H2O

- the excess polysulphide will precipitate yellow So↓ :

Sx2- + 2H+ H2S + (x – 1)So ↓

.. Account: 6-8 N HCl is preferred to conc HCl for acidificatn. ?

if conc. HCl used , some As2S3 dissolve .

81

* Separation of group II B members:

.. principle:

Precipitate ( arsenic, tin & antimony sulphides (+ some HgS) )

+ 6-8 N HCl, Δ , centrifuge

precipitate centrifugate

(As2S5 , HgS) (Sn4+ , Sb3+)

here As2S5 is insoluble in 6-8 N HCl while Sb2S5 , SnS2 are soluble

as follows :

Sb2S5 + 6H+ 2Sb3+ + 3H2S + 2So↓

(here a self oxidation Reduction reaction occurred)

SnS2 + 4H+ Sn4+ + 2H2S

√ the precipitate(containing As2S5 + some HgS) +

(NH4)2CO3 solution only As2S5 dissolve

3As2S5 + 3CO3-- 5AsS3

- + AsO3- + 3CO2↑

2As2S3 + 2CO3-- 3AsS2

- + AsO2- + 2CO2↑

Centrifuge

precipitate centrifugate

(contain HgS) (contain oxy & thio salts

dissolve in aqua regia of arsenic)

& confirm confirm

√ centrifugate (containing Sn4+ , Sb3+) :

the acidic solution is treated with iron wire , where

Fe + Sn4+ Fe2+ + Sn2+

3Fe + 2Sb3+ 3Fe2+ + 2Sbo↓

reducing precipitate black precipitate

centrifuge

precipitate(Sbo↓) centrifugate (Sn2+)

confirm confirm

82

* identification & confirmation :

.. mercuric :

HgS in dissolved in aquaregia & tested as before

.. Arsenic :

the ammonium carbonate extract (containing the thio & oxy salts

of arsenite or/& arsenate) :

√ acidify with acid yellow precipitate of As2S5 or As2S3

AsO2- + 3AsS2

- + 4H+ 2As2S3↓ + 2H2O

AsO3- + 5AsS3

- + 6H+ 3As2S5↓ + 3H2O

√ Filter, dissolve the yellow precipitate in conc. HNO3, evaporate

the acid , and then add

NH4OH + magnesia mixture Mg(NH4)AsO4↓

till alkaline white precipitate

.. Antimony :

(1) the black precipitate(Sbo↓) dissolved in dilute HNO3 ,

+ H2S Sb2S5 ↓ orange precipitate.

(2) the solution containing Sn4+ , Sb3+ (before use of Feo wire) +

NH4OH till alkaline + 5-6 gm oxalic acid + 50 ml water , boil until

dissolved + H2S (pass H2S gas for 3 min) Sb2S3 ↓

orange precipitate

[ N.B here oxalic acid forms a very stable complex with tin so,

when H2S passed only Sb2S3 precipitate ... oxalic acid acts as

masking agent here ]

.. Tin:

(1) The centrifugate containing Sn2+ (after reduction with Feo)

+ 2-3 drops HgCl2 Hg2Cl2 ↓

white precipitate which may turn into

gray then black (Hgo↓)

(2) after testing antimony using oxalic acid :

the solution containing tin-oxalic acid complex + conc HNO3 +

conc H2SO4 , evaporate (to destroy the complex) , cool , add

NH4OH till alkaline , then add dilute HCl till acidic

+ H2S SnS2 ↓ yellow precipitate

83

* Another method for separation of tin & antimony :

Solution (containing SbCl3 & SnCl4), boil , add KOH till the

precipitate formed redissolves, add bromine water till the liquid

remains yellow , add NH4Cl , boil for few mins. , filter

Filtrate Residue

(contains the soluble (Sn(OH)4↓ , dissolved in hot

antimenate, + dilute HCl conc. HCl , dilute with water

(add till acidic)

+ H2S add Feo wire SnCl2

(tested as

Sb2S5↓ orange precipitate before by HgCl2)

What happened ?? √ Br2 + 2OH- OBr- + Br- + H2O

hypobromite

hypobromite is an oxidizing agent stronger > Br2

√ Sb3+ + 4OH- SbO2- + 2H2O

Soluble antimonite (unstable)

SbO2- + OBr- SbO3

- + Br-

Soluble antimonite (stable)

√ Sn2+ + 4OH- SnO22- + 2H2O

Soluble stannite

SnO22- + OBr- SnO3

2- + Br-

Soluble stannate (unstable)

i.e. we have finally now a stable antimonite & unstable stannate :

stannate can be hydrolysed to Sn(OH)4 by boiling with NH4Cl :

SnO32- + 3H2O Sn(OH)4 ↓ + 2OH-

H2O + NH3↑ 2NH4OH 2NH4

+

84

i.e. NH4Cl helps to remove the excess alkali & thus push the

hydrolysis reaction forward.

while antimonite not (because it is very stable) .

... if NH4Cl is added before complete oxidation of antimonite to

antimonate Sb(OH)3 will precipitate together with Sn(OH)4

... Account : bromine water should be added till the solution

remains yellow ?

85

Group III cations : Iron group (Fe3+ , Al3+ , Cr3+)

* precipitation :

- group III cations precipitated as hydroxides by addition of

NH4Cl + NH4OH dropwise till alkaline.

- Q. why alkalinity must be controlled ??

(cations) (cations of later groups)

(Ksp of hydroxides) (Ksp of hydroxides)

Fe3+ 6 x 10-38

Al3+ 5 x 10-33 Mn2+ 2 x 10-13

Cr3+ 6.7 x 10-31 Zn2+ 4.5 x 10-17

Fe2+ 1.8 x 10-15 CO2+ 2.5 x 10-16

Ni2+ 1.6 x 10-14

Mg2+ 9 x 10-12

From this we notice that :

the Ksp of group III hydroxides is lower than that of the

remaining hydroxides.

... [OH-] must be just sufficient to exceed the Ksp of group III

hydroxides but not of the remaining hydroxides .

Q. How is [OH-] controlled ???

(1) use of NH4OH together with NH4Cl where

NH4OH NH4+ + OH-

NH4Cl NH4+ + Cl-

thus NH4Cl reduces ionization of NH4OH by the common ion effect

(i.e. control [OH-])

N.B. the ratio of NH4OH : NH4Cl should be 1 : 100

,where at this [OH-] :

the product [M3+] [OH-] 1.2 x 10-22 , thus hydroxides having Ksp

lower than this value will only precipitate (i.e. group III

hydroxides)

(2) addition of NH4OH dropwise till just alkaline to litmus paper

(then boil to remove excess NH3) .

N.B. NH4OH is preferred to other hydroxides (e.g. NaOH , KOH)

why ???

86

weak base .

volatile ... if excess added , can be removed by boiling .

form soluble ammine complexes with cations of later groups.

e.g. [Ni(NH3)6]2+ , [CO(NH3)6]

3+ , [Zn(NH3)4]2+ thus don't allow

them to precipitate with group III .

* important precautions :

(1) the filtrate from group II must be boiled well ?

to remove H2S ? why ?

√ otherwise sulphides of group IV precipitate upon rendering the

medium alkaline &

√ otherwise H2S will be oxidized to SO4-- ( upon addition of HNO3

to oxidize Fe2+ → Fe3+ ) , thus group V sulphates (Ca2+ , Ba2+ , Sr2+)

will precipitate with group III hydroxides.

(2) before adding group III reagent (NH4Cl + NH4OH) , we should

add few drops of conc. HNO3 , boil & cool ??

to convert any Fe2+ (produced from reduction of Fe3+ by H2S) to

Fe3+ , why ??

because Fe(OH)2 is not precipitated by group III reagent , why

?

because its Ksp is not attained , thus Fe(OH)2 will escape to the

subsequent groups.

(3) after additionof group III reagent , you should boil well ??

1 to drive off any excess NH3 ??

because Cr(OH)3 is slightly soluble in excess NH4OH forming a

violet soluble ammine complex .

... boil to decompose this complex

[ Cr(NH3)6 ] (OH)3 Cr(OH)3↓ + 6NH3↑

2 to coagulate the precipitate 3

to avoid dissolutn. of amphoteric

Al(OH)3 , Cr(OH)3

87

(4) after addn of solid NH4Cl , boil well ( to expel dissolved oxygen

) then add NH4OH , why ?

Mn(OH)2 though not precipitated with group III hydroxides ,

but it can be oxidized by atmospheric O2 into the insoluble

manganic hydrate MnO(OH)2↓ & thus coprecipitate with group III

... we boil to expel O2

(5) certain anions e.g oxalates , tartarates & citrates should be

absent when precipitating group III hydroxides ? why ?

as they prevent precipitation of group III hydroxides ( by

forming soluble complexes e.g. with Fe3+ )

- phosphates & arsenates also should be absent ?

as they interfere by precipitating cations of later groups .

(6) Colors of the precipitates :

Fe(OH)3 reddish brown

Al(OH)3 white gelatinous

Cr(OH)3 gray green

MnO(OH)2 reddish brown

(7) Washing :

the precipitate should be washed with 1 % hot NH4Cl ?

to remove the surface adsorbed cations of subsequent groups ?

why?

to prevent peptisation

* Role of NH4Cl in group III reagent :-

(1) to coagulate the precipitate( as it is electrolyte )

(2) to control [OH-] by the common ion effect it exerts on NH4OH,

thus doesn't allow hydroxides of subsequent groups. to precipitate

NH4Cl NH4+ + Cl-

NH4OH NH4+ + OH-

88

(3) Al(OH)3 & Cr(OH)3 are amphoteric & thus may dissolve in

excess [OH-] → AlO2- , CrO4

- , where NH4Cl prevent this rex

e.g. AlO2- + 2H2O Al(OH)3↓ + OH-

H2O + NH3↑ NH4OH + Cl- NH4Cl

(4) it acts as source of NH3 ( NH4Cl → NH3 + HCl ) , which makes

soluble ammine complexes with cations of group IV e.g. CO2+, Ni2+,

Zn2+ , thus doesn't allow their hydroxides to precipitate in group

III .

* Separation :

principle :

√ Al(OH)3 & Cr(OH)3 are amphoteric while Fe(OH)3 is not

√ Cr(OH)3 can be oxidized to the stable form (CrO4--) which is not

precipitated upon neutralization.

... precipitate + excess NaOH + H2O2 (or Na2O2) , boil :

√ Al(OH)3 + OH- [Al(OH)4]- AlO2

- + 2H2O

√ Cr(OH)3 + OH- CrO2- + 2H2O

(unstable)

aluminate is stable , while chromite is unstable

... chromite is further oxidized to the stable form chromate :

Na2O2 + 2H2O 2Na+ + 2OH- + H2O2

2CrO2- + 3H2O2 + 2OH- 2CrO4

-- + 4H2O (yellow , stable)

√ while Fe(OH)3 remains unchanged

... centrifuge

precipitate centrifugate

contain Fe(OH)3 (yellow)

& MnO(OH)2 contain AlO2- & CrO4

--

89

* identification & confirmation :

(A) the precipitate ( contains Fe(OH)3 + MnO(OH)2 ) + 40% HNO3 +

drops of H2O2 , boil till no evolution of O2 ↑ ( to destroy H2O2 )

Fe(OH)3 + 3 H+ Fe3+ + 3 H2O

MnO(OH)2 + 2H+ + H2O2 Mn2+ + 3H2O + O2 ↓ ↓ +4 +2

why SnCl2 is not used to reduce Mn4+ → Mn2+ ??

otherwise it will also reduce Fe3+ into Fe2+

... H2O2 is preferred .

(notice here that manganese is reduced to Mn2+ by H2O2) , then

divide the solution into 2 parts to test for iron & manganese :

(1) Test for iron :-

a- part + KCNS solution blood red color [Fe(CNS)6]3-

.. precautions :

√ the solution should be cold (as the color destroy by heat)

√ avoid presence of reducing agents e.g. SnCl2 ?

which reduce Fe3+ to Fe2+

√ avoid presence of F- or PO43- which form stable complex with

Fe3+ & prevent color formation .

√ avoid presence of Hg2+ which forms the stable complex

[Fe(CNS)6]3- + 3HgCl2 → Fe3+ + 3Hg(CNS)2 + 6Cl-

colorless

b- part + pot. Ferrocyanide solution → Prussian blue precipitate

FeK[Fe(CN)6]

(2) Test for manganese :

solution + 50% HNO3 + 1 gm red lead pb3O4 , boil for 2 min , allow

to settle → pink color of MnO4-

5 Pb3O4 + 2 Mn2+ + 24 H+ 2 MnO4

- + 15 Pb2+ + 12 H2O

5 PbO2 + 2 Mn2+ + 4 H+ 2 MnO4

- + 5 Pb2+ + 2 H2O

colorless pink color

91

Ag

N.B.

√ Account : HCl or chlorides should be absent ?

as they decompose the purple color of MnO4-

10Cl- + 2MnO4- + 16H+ Cl2 + 2Mn2+ + 8H2O

√ there are other methods of oxidation of Mn2+ e.g. :-

- boiling with persulfate & H2SO4 in presence of Ag+ catalyst .

or – boiling with bismuthate & HNO3

or – boiling with periodate & H2SO4

2Mn2+ + 5S2O8-- + 8H2O 2MnO4

- + 10SO4-- + 16H+

5BiO3- + 2Mn2+ + 14H+ 2MnO4

- + 5Bi3+ + 7H2O

5 IO4- + 2Mn2+ + 3H2O 2MnO4

- + 5IO3- + 6H+

periodate iodate

(B) the centrifugate ( contain AlO2- & CrO4

-- ) :-

divide into 2 porions :-

(3) Test for chromium :

if the centrifugate is yellow ... CrO4-- is present

a- part + dilute acetic acid till acidic + lead acetate slon → yellow

precipitate of lead chromate .

b- perchromic acid test :

yellow solution + dilute H2SO4 dropwise till orange + ether + drops.

H2O2, shake → blue color in etherial layer .

2CrO4-- + 2H+ 2HCrO4

- Cr2O7-- + H2O

yellow orange

Cr2O7-- + 7H2O2 2CrO8

3- + 5H2O + 4H+

Perchromic acid

( blue in ether )

N.B.

.. Account : exess peroxide present in the yellow CrO4-- solution

must be decomposed by boiling before the solution is acidified with

diluteH2SO4 ?

otherwise H2O2 will be oxidized with Cr2O7--

Cr2O7-- + 3H2O2 + 8H+ 2Cr3+ + 3O2 + 7H2O

91

.. Account : dilute H2SO4 should be added just till the yellow color

becomes orange ?

because excess acid will favour oxidation of H2O2 with Cr2O7-- .

(4) Test for Aluminium :

a- solution + 1 gm NH4Cl solid , boil until a point NH3 odour

persist → white gelatinous precipitate

Al(OH)3↓ + OH- [Al(OH)4]- AlO2

- + 2H2O

NH4

Cl H2O + Cl- + NH3↑

N.B.

[OH-] can also be lowered by carefull addition of HCl dropwise till

the white gelatinous precipitate of Al(OH)3 appear but if excess

HCl added → the white gel dissolve .

b- Al(OH)3 is dissolved in HCl , then add aluminon (or alizarin – s)

reagent , then render the medium slightly ammoniacal → red lake (red color or gelatinous precipitate)

92

Group IV cations (Zinc group) : Mn2+ , Zn2+ , Co2+ , Ni2+

* precipitation :

this group is precipitated as sulfides but from ammoniacal

solution

... the group reagent is ( NH4Cl + NH4OH + H2S )

Q. Account : group IV sulphides didn't precipitate with group II

sulphides ?

because the solubility products of group IV sulfides are higher

than those of group II sulfides .

(i.e. group IV sulfides need a higher [S--] to precipitate which was

not available before, due to the common ion effect induced by

HCl).

.. important notes :

(1) what is the importance of NH4Cl ??

√ NH4Cl NH4+ + Cl-

NH4OH NH4+ + OH-

i.e. NH4Cl controls [OH-] by the common ion effect it exerts, thus

doesn't allow Mg2+ to precipitate as Mg(OH)2↓ .

√ after complete precipitation , the solution should be boiled

because boiling in presence of NH4Cl (electrolyte) allows

coagulantion of the colloidal sulfides of Zn2+ & Mn2+ (otherwise the

colloidal precipitate will pass through the filter paper & escape to

the susbsequent groups) .

(2) boiling after precipitation is important ?

√ for coagulation of the colloidal precipitate.

√ to decompose any water soluble ammine complexes of group IV

cations (not to escape to following groups)

e.g. [Ni(NH3)6]S NiS↓ + 6NH3↑

soluble dark brown ammine complex

√ to increase insolubility of CoS, NiS in 1M HCl (by physical

change)

93

(3) if Fe2+ was not completely oxidized to Fe3+ before precipitation

of group III hydroxides , it will precipitate as FeS with group IV

sulfides .

[FeS is black , thus may be confused with the black CoS , NiS ] .

(4) colors of the precipitates :

CoS → black

NiS → black

ZnS → white

MnS → buff

(5) the precipitate is washed with 1% NH4Cl + little NH42S ?

√ to prevent peptisation

√ to decrease solubility of sulfides due to the common ion S--

√ to remove any surface adsorbed cations of following groups.

* Separation :-

group IV sulfides + cold very dilute HCl (1M) ,

shake for 2 min , centrifuge

centrifugate precipitate

(Zn2+ , Mn2+) (CoS↓ , NiS↓)

[ principle : MnS , ZnS are soluble in very dilute HCl , while CoS ,

NiS are not ???

because boiling after precipitation increases the degree of

aggregatn. of CoS , NiS molecules & thus increase their insolubility

in 1M HCl ]

precipitate of (CoS , NiS) is dissolved in aqua regia :

CoS + 2HNO3 + 6HCl 3CoCl2 + 2NO + 3S + 4H2O

blue solution

NiS + 2HNO3 + 6HCl 3NiCl2 + 2NO + 3S + 4H2O

Yellow solution

For the centrifugate (ZnCl2 , MnCl2) :

Principle : Zn2+ is amphoteric , while Mn2+ not .

94

... centrifugate + excess NaOH + Bromine water , ∆

Zn2+ + OH- Zn(OH)2↓ + 2OH- ZnO2-- + 2H2O

excess soluble zincate

while Mn(OH)2↓ is further oxidized to the more insolubleform

MnO2↓

Mn2+ + Br2 + 4OH- MnO2 + 2Br- + 2H2O

(( really : Br2 + 2OH- OBr- + Br- + H2O

then, Mn2+ + OBr- + 2OH- MnO2 + Br- + H2O ))

brown precipitate

centrifuge

centrifugate precipitate

(soluble ZnO2--) (MnO2)

* identification & confirmation :-

.. for manganese : as under group III

.. for zinc :

- add NH4OH & boil ?? to decompose OBr- ?? ( otherwise OBr- will

further oxidize H2S )

3OBr- + 2NH3 3Br- + N2 + 3H2O

- then divide into 2 parts :

a- part + H2S solution ZnS↓ white precipitate

ZnO2-- + S-- + 2H+ ZnS↓ + 2OH-

b- part + dilute acetic acid till acidic + pot. Ferrocyanide solution →

white precipitate

ZnO22- + 4CH3COOH Zn2+ + 4CH3COO- + 2H2O

2Zn2+ + [Fe(CN)6]4- Zn2[Fe(CN)6]↓

white precipitate

.. for centrifugate ( CoCl2 , NiCl2 ) :

divide into 2 portions :

.. Test for cobalt :

divide the 1st portion into 3 parts .

95

H

CN4

xss

xss

CN2

(a) part + NH4Cl + NH4OH + pot. Ferricyanide ,

warm CO3[Fe(CN)6]2 red precipitate

[ N.B Ni3[Fe(CN)6]2 is a yellow precipitate but is soluble in NH4OH

while the red precipitate not ]

(b) Vogel's rex : part + KF + NH4SCN solid + 1ml conc HCl + 1ml

ether (or amyl alcohol), shake

blue color in alcoholic or ethereal layer .

[ N.B Fe3+ may interfere with this test , so we add F- to form the

stable [FeF6]3- complex or we add SnCl2 to reduce Fe3+ to Fe2+]

(NH4)2[Co(CNS)4] H2[Co(CNS)4] Ammonium Cobahothiocyanate more soluble in amyl alc.

(c) part + dilute HCl + 1% - nitroso -naphthol, ∆

red precipitate of cobaltinitroso -naphthol

[C10H6(NO)O]3Co , ( while [C10H6(NO)O]2Ni is brown

precipitate but soluble in dilute HCl ) .

.. Test for Nickel :

divide the 2nd portion into 2 parts

(a) part + NH4Cl + NH4OH + 1% alcoholic solution of DMG (dimethyl

glyoxime) red precipitate

CH3 C = NOH CH3 C = NO

| + Ni2+ | Ni + 2H+

CH3 C = NOH CH3 C = NOH 2

red precipitate

why NH4Cl is added?? to control [OH-]

not to precipitate Ni(OH)2

(b) part + dilute KCN solution dropwise till the precipitate formed

redissolves , boil , add drops of NaOH + excess bromine water and warm

black precipitate of NiO2 Ni2+ + 2CN- Ni(CN)2 [Ni(CN)6]

2-

nickelocyanide complex

Co2+ + 2CN- Co(CN)2 [Co(CN)6]4-

cobaltocyanide complex

96

boil

nickelocyanide complex is not affected by boiling , while

cobaltocyanide complex is oxidized by boiling & O2 :

4[Co(CN)6]4- + O2 + 2H2O 4[Co(CN)6]

3- + 4OH-

2[Co(CN)6]4- + OBr- + H2O 2[Co(CN)6]

3- + 2OH- + Br-

cotaltocyanide cobalticyanide complex

(water soluble)

while

[Ni(CN)4]2- + OBr- + 2OH- NiO2 + Br- + 4CN- + H2O

black precipitate

.. Account : excess KCN should be avoided ??

otherwise it reacts with Br2

CN- + Br2 CNBr + Br-

cyanagen bromide

97

o60

OH2

Group V (alkaline earth group) Ca2+ , Ba2+ , Sr2+

* precipitation :

Group reagent : NH4Cl + NH4OH + (NH4)2CO3

... they are precipitated as white carbonates CaCO3 , BaCO3 ,

SrCO3

notes :-

(1) (NH4)2CO3 consists of an equimolar mixture of (NH4HCO3 &

ammonium carbamate) & NH3 should be add to convert

NH4HCO3 (NH4)2CO3 :

NH4HCO3 + NH3 (NH4)2CO3

NH4COONH2 + H2O (NH4)2CO3

Ammonium carbamate

(2) After complete precipitation, the solution should be heated at

60o C for several purposes : √ to convert ammoniumcarbamate (NH4)2CO3

√ to convert the soluble bicarbonates of group V into insoluble

carbonates :

e.g. Ca(HCO3)2 CaCO3 + CO2 + H2O

√ promote precipitation of carbonates (especially at low conc. of

ions). i.e. coagulation .

√ to allow precipitation of a crystalline precipitate (which is easily

filtered & washed)

(3) boiling should be avoided (especially in presence of much NH4+) ??

not to lose NH3 &

not to allow dissolution of carbonates as HCO3-

e.g. BaCO3 + NH4+ Ba(HCO3)2 + NH3

soluble

(4) avoid presence of much excess NH4+ salts ?

for the same previous reason .

98

H2

... excess NH4+ should be destroyed before precipitation of group V

, how ??

by evaporation of group IV centrifugate in presence of few

drops of HNO3 ?

to convert NH4Cl NH4NO3 , where NH4NO3 decompose by

heat at a lower temperature than NH4Cl

then evaporate to dryness & ignite till no more white fumes of

NH4+ salts are given off .

then the residue is extracted with (a little H2O + dilute HCl) ,

then add group V reagent .

.. Account : NH4OH is used in the group reagent ?

√ to convert NH4HCO3 (NH4)2CO3

√ to prevent transformation of CO3-- HCO3

-

where CO3-- + H+ HCO3

-

(NH4)2CO3 2NH4+ + CO3

--

NH4Cl NH4+ + Cl-

- why group V carbonates are dissolved in acetic acid not in HCl ?

because HCl converts CrO4-- Cr2O7

--

where Ca2+ , Ba2+ , Sr2+ dichromates are soluble

2CrO4-- 2HCrO4

- Cr2O7-- + H2O

.. Account : NH4OH should be added dropwise just till alkaline ??

if excess added , pH will be high & thus Mg(OH)2 may

coprecipitate with group V carbonates .

.. Account : NH4Cl is used in group V reagent ?

NH4OH NH4+ + OH-

NH4Cl NH4+ + Cl-

99

i.e. NH4Cl reduces [OH-] by the common ion effect it exerts on

NH4OH , thus doesn't allow Mg(OH)2 to precipitate .

(2)to coagulate the precipitate.

(3) prevent precipitation of MgCO3

N.B. the white precipitate is washed with hot water .

* Separation & identification :

the white precipitate of group V carbonates is dissolved in hot

dilute acetic acid clear solution

e.g. BaCO3 + 2CH3COOH Ba(CH3COO)2 + CO2 + H2O

then the solution + K2CrO4 solution , shake , centrifuge

centrifugate yellow precipitate

(Ca2+ , Sr2+) BaCrO4

... Ba2+ is present

( principle : BaCrO4 is insoluble in acetic acid , SrCrO4 is soluble in

acetic acid , while CaCrO4 is not precipitated at all neither in

acetic acid nor in neutral solution ??

because acetic acid lowers [CrO4--] as it converts it to HCrO4

-

thus [CrO4--] will not be sufficient to exceed Ksp of CaCrO4

√ then centrifugate +dilute NH3 till alkaline + excess conc.

(NH4)2SO4 solution, centrifuge

centrifugate white precipitate

( Ca2+ ) SrSO4

... Sr2+

( principle : BaSO4 is water insoluble > SrSO4 > CaSO4 but now, we

removed Ba2+ as BaCrO4 ... we can precipitate Sr2+ as SrSO4 while

Ca2+ is still soluble as (NH4)2[Ca(SO4)2] .

111

N.B. precipitation of Sr2+ as SO4-- can be done with H2SO4 but a

little of Ca2+ may also precipitate .

... (NH4)2SO4 is preferred where the water soluble

(NH4)2[Ca(SO4)2] can form .

√ Then the centrifugate (containing Ca2+) + dilute NH3 solution till

alkaline + ammonium oxalate solution ,

white precipitate of ca.oxalate (insoluble in acetic acid)

... Ca2+

... separation of Ba2+ , Sr2+ , Ca2+ depends on stepwise separation of

Ba2+ as BaCrO4 1st, then Sr2+ as SrSO4 2

nd then Ca2+ as Ca-oxalate

3rd .

N.B. Ca2+ can be removed 1st from a mixture of Ba2+ , Sr2+, Ca2+,

how ?

by additionof conc. NH4Cl solution + excess pot. ferrocyanide

solution:

Ca2+ + NH4+ + K+ + [Fe(CN)6]

4- CaNH4K[Fe(CN)6]

White creamy precipitate

(triple ferrocyanide salt)

where solubility of this triple salt is decreased by presence of

much NH4+, K+ ( due to the common ion effect ) & so we use them in

excess.

* Flame reactions :

we can confirm the presence of Ba2+ , Sr2+ and Ca2+ by flame

reactions

, how ??

the precipitates BaCrO4 , SrSO4 , Ca-oxalate each is hold on a

platinum wire , moistered with conc. HCl (to be converted to the

volatile form BaCl2 , SrCl2 , CaCl2) , then the flame test is run (on a

Bunsen flame) where

Ba2+ give green edge flame

Sr2+ give crimson flame

Ca2+ give brick red flame

111

N.B SrSO4 (before treating with conc. HCl) :

is heated first in the reducing flame ?

SrSO4 SrS (to react with HCl easily) then heated for 1/2

min in the oxidizing flame

to convert surface C CO2 , then treat with conc. HCl &

apply the test .

SrSO4 can be precipitated from Sr2+ solution with saturated

CaSO4 solution

112

Group VI (alkali group) Mg2+ , NH4+ , K+ , Na+

- There is no group reagent

... each cation will be identified separately .

- Before you test for any of group VI cations , you must first

remove any traces of alkaline earth metals ? why?

otherwise they interfere with Mg2+ test , where alkaline earth

metals give precipitates with phosphate .

Q. why group V escaped to group VI ?

group V cations may have dissolved in excess NH4+ added .

Q. How would you remove the traces of group V ?

Filtrate of group V + (NH4)2SO4 ( which precipitate BaSO4 &

SrSO4 ) + Ammonium oxalate ( which precipitate Ca. oxalate ) ,

heat , stand For 5 min , filter & reject the residue , then divide the

filtrate to 4 portions :

* Test for Magnesium :

(1) part + NH4Cl + NH4OH + Na2HPO4 white crystalline

precipitate

Mg2+ + NH4+ + PO4

3- MgNH4PO4

.. why NH4OH added ?

to convert HPO42- into PO4

3-

.. why NH4Cl added ?

to decrease ionization of NH4OH , so that [OH-] becomes

insufficient to precipitate Mg(OH)2 .

.. Large excess of NH4+ should be avoided ?

to avoid precipitation of Mg2+ as Mg(NH4)4(PO4)2 rather than

MgNH4PO4

(2) Lake formation :

113

part + dilute NaOH till alkaline + magneson 1

reagent blue precipitate

[ the test depends on adsorption of magneson 1 dye on Mg(OH)2

precipitate]

[N.B. NH4+ reduces sensitivity of the test ]

* Ammonium :

ammonium reactions may be :

(1) Displacement by a strong base :

NH4+ + OH- NH3 + H2O

(NaOH) ammonia odour

(2) precipitation reactions :

√ all NH4+ salts are water soluble, but NH4

+ can give precipitates

with :

a- cobaltinitrite (NH4)2Na[CO(NO2)6]

yellow precipitate

( like K2Na[CO(NO2)6 )

b- Hydrogen tartarate NH4HC4H4O6 white precipitate

(like pot. Hydrogen tartarate)

(i.e. kt interfere)

(3) Complex formation :

HgCl2 + 2NH3 HgNH2Cl + NH4+ + Cl-

amino mercuric chloride

( to prove that NH3 formed a complex with HgCl2 , add

NaOH No NH3 odour)

NH3 + 2[HgI4]-- + 3OH- NHg2I . H2O + I- + H2O

alkaline Neester's dimercuric ammoinium iodide

reagent monohydrate

(brown precipitateor color)

Now, how can you detect NH4+ ??

solid mix + NaOH (or Na2CO3) solution, warm

114

ammonia vapours , tested by

√ odour

√ turns red litmus paper blue

√ turns yellow turmeric paper brown

√ turns mercurous nitrate paper black ??

Hg22+ + NO3- + 4NH3 + H2O NHg2 . NO3 . H2O + Hgo

+ NH4+

black precipitate √ turns a filter paper moistened with

MnSO4 , H2O2 brown ??

2Mn2+ + H2O2 + 4NH3 + 4H2O Mn(OH)3 + 4NH4+

brown precipitate

[ N.B Hg2+ may interfere with NH4+ test

... to prevent its interference , add S-- ??

HgNH2Cl + HS- HgS + NH3 + Cl-

black precipitate]

Removal of NH4+ :

NH4+ interfere with Na+ , K+ tests

... to test for Na+ , K+ , we must remove NH4+ , how ??

1 by evaporating the solution to dryness , then ignition of the

NH4+ residue in the form of chlorides or nitrates :

NH4Cl NH3 + HCl the salt reforms

NH4NO3 NH3 + HNO3 on cooling

NH4NO3 N2O + H2O

Mg(NO3)2 MgO

N.B.

( .. Account : ignition as nitrate (by adding conc HNO3) is preferred

?

because decomposition of NH4NO3 is

√ less reversible

√ occur at lower temp. )

115

2 then dissolve the residue in water , filter , test for

Na+, K+ in the filtrate .

precipitate filtrate

(MgO) (Na+ , K+)

why Removal of NH4+ salts by ignition with HNO3 is better > HCl ,

H2SO4 ??

NH4+ + Cl- NH4Cl NH3 + HCl

NH4+ + SO4

-- (NH4)2SO4 NH3 + HSO4--

i.e. decomposition is reversible , while

NH4NO3 N2O + H2O

i.e. Forward

* Sodium :

(1) Salt solution + magnesium uranyl acetate solution, shake

Na+ + HMg(UO2)3 (CH3COO)4 NaMg(UO2)3 (CH3COO)4 + H+

Yellow Crystalline precipitate

√ insoluble in acetic acid √ soluble in strong acids

(2) Sodium antimonate test :

conc. Salt solution + KOH + KH2SbO4 pot. dihydrogen

antimonate NaH2SbO4 white precipitate.

KH2SbO4 + Na+ NaH2SbO4 + K+

.. Account : the solution should be neutral or weakly alkaline ??

because acids & NH4+ salts of strong acids interfere by giving a

white precipitate of antimonic acid .

KH2SbO4 + H+ H3SbO4 + K+

While precipitate

116

OH

(3) sodium pyroantimonate test :

Salt solution + pot. pyroantimonate white precipitate

K2H2Sb2O7 + 2Na+ Na2H2Sb2O7 + 2K+

white precipitate

(4) Flame test :

give intense golden yellow flame

N.B. Ca2+ , Sr2+ may interfere

N.B. organic matter ( which burn with a yellow flame ) interfere .

* Potassium :

(1) pot. cobaltinitrite test :

Salt solution + dilute acetic acid + sod. cobaltinitrite solution

yellow precipitate

Na3[Co(NO2)6] + 3K+ K3[Co(NO2)6] + 3Na+

Na3[Co(NO2)6] + 2K+ K2Na[Co(NO2)6] + 2Na+

[ N.B. if the medium is alkaline Co(OH)3

Na3Co(NO)6 Co(OH)3 black precipitate]

(2) Potassium hydrogen tartarate test :

Salt solution + acetic acid + sod. hydrogen tartarate (or tartaric

acid) solution white precipitate

HO CH COOH HO CH COOH

| |

HO CH COONa + k+ HO CH COOK + Na+

white precipitate of

pot. hydrogen tartarate

* ( soluble in strong acids ? due to formation of the little ionised

tartaric acid & soluble in strong bases ? due to the formation of

pot. tartarate )

(3) with perchlorate ( e.g. perchloric acid )

117

ClO4- + K+ KClO4

white precipitate

(4) Flame test :

give violet color to Bunsen flame .

( N.B. the violet color of K+ may be masked by yellow color of Na+

(if present)

So, to see the flame, view it through a thick cobalt glass (which

absorbs the yellow color of Na+ ) where you will see the K+ flame as

reddish brown against the blue background of the glass ) .

118

References

- G. Svehla; Vogel’s Qualitative Inorganic Analysis (6th Edition), Pearson Education ltd., 2008, Delhi, India.

- K. Whitten, R. Davis & M. Peck; General Chemistry with qualitative analysis (6

th edition), Saunders College

Publishing,2004, New York, U.S.A.