CHAPTER 8 : SALTS. Meaning and uses of Salts A salt is an ionic compound formed when the hydrogen ion, from an acid is replaced by a metal ion or an ammonium.

CHAPTER 8 : SALTS

Meaning and uses of Salts

A salt is an ionic compound formed when the hydrogen ion, from an acid is replaced by a metal ion or an ammonium ion

Example of salts :

(i) sodium chloride

(ii) potassium carbonate

(iii) copper(II) sulphate

Examples of salts :

Acid Formula of acid

Salt Formula Cation Anion

Hydrochloric acid

HCl Sodium chloride

NaCl Na+ Cl-

Carbonic acid

H2CO3 Potassium carbonate

K2CO3 K+ SO42-

Sulphuric acid

H2SO4 Copper(II) sulphate

CuSO4 Cu2+ SO42-

Examples of salts :

Acid Formula of acid

Salt Formula Cation Anion

Nitric acid

HNO3 Ammonium nitrate

NH4NO3 NH4+ NO3

Ethanoic acid

CH3COOH Sodium ethanoate

CH3COONa Na+ CH3COO-

Nitric acid

HNO3Magnesium nitrate

Mg(NO3)2 Mg2+ NO3-

Salts with their uses.

Salt Uses

Barium sulphate BaSO4 X-ray ‘meals’ in hospital

Calsium sulphate CaSO4

Plaster of Paris for broken bone

Iron sulphate FeSO4 Iron tablets for anaemia patient

Ammonium nitrate NH4NO3

Nitrogenous fertilizer

Salt Uses

Copper(II) sulphate CuSO4

Fungicide

Sodium chloride NaCl A flavouring agent

sodium hydrogen carbonate

Baking powder

Sodium nitrite NaNO2 For preserving food/ food preservative

Salt Uses

Sodium hypochlorite NaOCl

Bleaching agent

Tin(II) fluoride SnF2 Toothpaste

Lead(II) chromate PbCrO4

Paint for yellow line on road

Identify soluble and insoluble salt.

NaNa++

KK++ CO CO332-2-

NHNH44++

Water Water

NONO33--

PbPb2+2+

BaBa2+2+ SO SO442-2-

CaCa2+2+

PbPb2+2+

AgAg++ Cl Cl--

HgHg++

All sodium, potassium and ammonium salts are soluble in water.

All nitrate salts are soluble in water. All sulphate salts are soluble in water except lead(II)

sulphate, barium sulphate and calcium sulphate. All chloride salts are soluble in water except lead(II)

chloride, silver chloride and mercury chloride. All carbonate salts are insoluble in water except

sodium carbonate, potassium carbonate and ammonium carbonate

State whether each of the following salt is soluble or insoluble in water

Formula of salt

Solubility

( / or x)

Formula of salt

Solubility

( / or x)

PbCO3 x NaCl /

CaSO4 x AgNO3 /

K2CO3 / FeCl3 /

Na2SO4 / NH4NO3 /

CuSO4 / PbCl2 x

Chemical and ionic equations for reactions used in the preparation of soluble salts

General equation for preparing soluble salts.

a. metal + acid salt + hydrogen

b. metal oxide (or metal hydroxide) + acid salt + water

c. alkali + acid salt + water

d. metal carbonate + acid salt + water + carbon dioxide

Complete the following chemical equation.

Mg + H2SO4 MgSO4 + H2

CuO + HCl CuCl2 + H2O

Zn(OH)2 + HNO3 Zn(NO3)2 + H2O

NaOH + HCl NaCl + H2O

MgCO3 + H2SO4 MgSO4 + CO2 + H2O

Preparation of soluble salt Method of Preparation

Reactants Salt Formed

Other Product

metal + acid Magnesium + hydrogen chloride

Magnesium chloride

Hydrogen

metal oxide + acid

Copper(II) oxide + sulphuric acid

Copper(II) sulphate

Metal carbonate + acid

Zinc carbonate + sulphuric acid

Zinc sulphate

Water + carbon dioxide

Preparation of soluble salt Method of Preparation

Reactants Salt Formed

Other Product

metal hydroxide + acid

Potassium hydroxide + nitric acid

Potassium nitrate

e) alkali + acid

Sodium hydroxide + hydrochloric acid

Sodium chloride

The reactants which are needed to prepare the following soluble salts:

Copper(II) sulphate : Copper(II) oxide / hydroxide / carbonate + sulphuric acid

Zinc chloride : Zinc / (zinc oxide / hydroxide / carbonate) + hydrochloric acid

Potassium nitrate: potassium hydroxide + nitric acid

Ammonium sulphate : aqueous ammonia + sulphuric acid

Magnesium nitrate : Magnesium / (magnesium oxide / hydroxide / carbonate) + nitric acid

Chemical equation can be simplified into an ionic equation.

Example : Chemical equation : Zn(s) + H2SO4(aq)

ZnSO4(aq) + H2(g)

Zn(s) + 2H+(aq) + SO42-(aq) Zn2+(aq)

+ SO42-(aq) + H2(g)

Ionic equation : Zn(s) + 2H+ (aq) Zn2+ (aq) + H2(g)

Chemical equation : Mg(s) + 2HCl(aq) MgCl2(aq) + H2 (g)

Mg(s) + 2H+(aq) + 2Cl-(aq) Mg2+(aq) + 2Cl-(aq) + H2(g)

Ionic equation : Mg(s) + 2H+ (aq) Mg2+ (aq) + H2 (g)

Chemical equation :MgO(s) + 2HCl(aq) MgCl2(aq) + H2O (l)

MgO(s) + 2H+(aq) + 2Cl-(aq) Mg2+

(aq) + 2Cl-(aq) + H2O(l)

Ionic equation : MgO(s) + 2H+ (aq) Mg2+ (aq) + H2O (l)

Chemical equation :NaOH (aq) + HNO3 (aq) NaNO3 (aq) + H2O (l)

Na+(aq) + OH-(aq) + H+(aq) + NO3-(aq)

Na+(aq) + NO3-(aq) + H2O(l)

Ionic equation :OH- (aq) + H+ (aq) H2O (l)

Chemical equation :CuCO3 (s) + H2SO4 (aq) CuSO4 (aq) + CO2 (g) + H2O (l)

CuCO3(s) + 2H+(aq) + SO42-(aq)

Cu2+(aq) + SO42-(aq) + CO2(g) + H2O(l)

Ionic equation : CuCO3 (s) + 2H+ (aq) Cu2+ (aq) + CO2 (g) + H2O (l)

The procedure for the preparation of soluble salts of sodium, potassium and ammonium

Soluble salt Sodium Chloride, NaCl

Name two chemical substances to prepare the salt

sodium hydroxide

hydrochloric acid

Chemical equation NaOH + HCl NaCl + H2O

Procedure: (Diagram) Description

1. A pipette is used to transfer 25.0of sodium hydroxide solution to a conical flask. 2 to 3 drops of phenolphthalein is added.

2. A burette is filled with hydrochloric acid and record the initial burette reading.

3. Titration is carried out carefully by slowly adding the acid into the conical flask and the flask is shaken well.

4. The acid is added continuously until the indicator turns from pink to colourless. The final burette

reading is recorded.

5. The volume of acid used to neutralize 25.0of the alkali is determined. (let the volume be V)

6. 25.0of the same sodium hydroxide solution is pipetted into a conical flask. No indicator is addeded.

7. From the burette, exactly Vof hydrochloric acid is added to the alkali and is shaken well.

8. The contents of the conical flask is poured into an evaporating dish.

9. The solution is heated gently to evaporate most of the water to produce a saturated solution.

10. The hot saturated salt solution is cooled for crystallization to occur.

11.The sodium chloride crystals is filtered, and the salt is rinsed with a little distilled water.

12. The crystals are dried by pressing them between filter papers

Describe the physical characteristics of the crystals that you obtained

Salt crystal characteristic

1. Flat surfaces, straight edges and sharp angles

2. Fixed geometrical shape

3. Fixed angles between two neighbouring surfaces

4. Crystals of some substance have same hapes but maybe in different sizes

Preparation of soluble salts (not sodium, potassium or ammonium salt)

Soluble salt Copper(II) sulphate, CuSO4

Name two chemical substances to prepare the salt

copper(II) oxide and sulphuric acid

Chemical equation CuO + H2SO4 CuSO4 + H2O

1. 50of sulphuric acid 1 mol dm-3 is poured into a beaker. The acid is warmed.

2. By using a spatula, copper(II) oxide powder is added bit by bit into the acid. The mixture is stirred well.

3. Copper(II) oxide is added continuously until some of it no longer dissolves.

3. The unreacted copper(II) oxide is removed by filtration

4. The filtrate is filtered into an evaporating dish. The solution is heated gently to produce a saturated salt solution.

5. The saturated solution is cooled until crystals are formed

6. The copper(II) nitrate crystals are filtered, and are then rinsed with a little distilled water.

Describe the purification process of the crystals

Purification process – Recrystallisation

1.The copper(II) sulphate crystals are placed in a beaker.

2.Enough distilled water is added to cover the crystals. The solution is gently heated and stirred with a glass rod. Water is added bit by bit until all the crystals are dissolved.

3. Impurities is removed by filtration and filtrate is poured into an evaporating dish.

4.The solution is heated gently to evaporate most of the water to produce a saturated solution.

5.The hot saturated salt solution is cooled for crystallization to occur.

6.The copper(II) nitrate crystals are filtered, and the salt is rinsed with a little distilled water.

Chemical and ionic equations for reactions used in the preparation of insoluble salts

Insoluble salts can be prepared by precipitation method through double decomposition reaction. In this reaction, two different aqueous solution mutually exchange their ions , to form precipitate.

Soluble salt solution + Soluble salt solution Insoluble salt MX

containing cation M+ containing anion X-

Chemical equation : AgNO3 (aq) + NaCl (aq) AgCl (s) + NaNO3(aq)

Ionic equation : Ag+ (aq) + Cl- (aq) AgCl (s)

Preparation of insoluble salts

Example 1: Barium sulphate, Solution 1: Barium chloride/nitrate Solution 2 :sodium/potassium sulphate Chemical equation :BaCl2 + Na2SO4

BaSO4 + 2NaCl Ionic Equation : Ba2+ + SO4

2- BaSO4

Observation : White precipitate formed

Example 2 :Copper(II) carbonate, CuCO3 ,

Solution 1:copper(II) nitrate/sulphate/chloride Solution 2 :Sodium/potassium carbonate Chemical equation :Cu(NO3)2 + Na2CO3

CuCO3 + 2NaNO3

Ionic Equation : Cu2+ + CO32- CuCO3

Observation : Green precipitate formed

Example 3 : Lead(II) chromate(VI),PbCrO4 ,

Solution 1:lead(II) nitrate Solution 2 :Sodium/potassium chromate(VI) Chemical equation :Pb(NO3)2 (aq) +

K2CrO4(aq) PbCrO4(s) + 2KNO3(aq)

Ionic Equation : Pb2+ + CrO42- PbCrO4

Observation : Yellow precipitate formed

The preparation of insoluble salts

Insoluble salt Name Lead(II) iodide, PbI2

Two chemical substances to prepare the salt

(i) Lead(II) nitrate

(ii) Sodium/potassium iodide

Chemical equation Pb(NO3)2 (aq) + 2KI (aq) PbI2 (s) + 2KNO3 (aq)

Ionic equation Pb2+(aq) + 2I-(aq)

PbI2(s)

1. 50 cm3 of 0.5 mol dm-3 lead(II) nitrate solution is poured into 50 cm3 of 1.0 mol dm-3 potassium iodide in a beaker

2. The mixture is stirred with a glass rod

3. A yellow precipitate of lead(II) iodide is formed immediately

4. The resulting mixture is then filtered

5. The yellow precipitate is rinsed with distilled water to remove impurities

6. The yellow solid, lead(II) iodide is then pressed between a few pieces of filter papers to be dried

Solve problems involving calculation of quantities of reactants or

products in stoichiometric reactions

Example 1 : A student prepares copper (II) nitrate by reacting copper (II) oxide with 100 cm3 1.5 mol dm-3 nitric acid. Calculate the mass of copper (II) oxide needed to react completely with the acid. [Relative atomic mass: Cu, 64 ; O, 16]Solution : Chemical equation : CuO + 2HNO3 Cu(NO3)2 + H2O

Mole ratio : 1 mole 2 mole 1 mole 1 mole Number of moles of HNO3 = 1.5 x 100 = 0.15 mol 1000

Mole ratio of CuO : HNO3 = 1 : 2 Number of mole of CuO = 1 x 0.15 = 0.075 mole 2 Mass of CuO = 0.075 x (64 + 16) = 6 g

Question :

1.Excess zinc powder is added to react completely with 50 cm3 of 2.0 mol dm-3 hydrochloric acid.

(a) Write an ionic equation for the reaction between zinc and hydrochloric acid.

Zn + 2HCl ZnCl2 + H2

Question :

(b) Calculate the number of moles of hydrochloric acid used.

No of mole of HCl mol

= 2.0 x 50 = 0.1 mol

Question :

(c) Calculate the volume of hydrogen gas liberated at room conditions.

[Molar volume: 24 dm3 mol-1]

Mole ratio HCl : H2 = 2 : 1

No of mole of H2 = 1/2 x 0.1 = 0.05 mol

Volume of H2 = 0.05 x 24 dm3

= 1.2 dm3

Question :

2 Excess of magnesium carbonate powder, MgCO3, is reacted with 100 cm3 of a 1 mol dm-3 sulphuric acid H2SO4 , What is the mass of magnesium sulphate formed?

[Relative atomic mass : Mg =24, O=16, S = 32 ]

Question :

No of mole of H2SO4 = 1.0 x 100 = 0.1 mol

Mole ratio H2SO4 : MgSO4 = 1 : 1

No of mole of MgSO4 = 0.1 mol

Mass of MgSO4 = 0.1 x (24 + 32 + 4x16) g = 12.0 g

Question :

3. 0.12 g of magnesium reacts with excess hydrochloric acid to produce hydrogen gas. Given that the relative molecular mass of H=1, Mg = 24, CI =35.5 and 1 mol of gas occupies 24 dm3 at room temperature and pressure.

Find the

(a) mass of salt formed

Question :

No of mole of Mg = 0.12 = 0.005

No of mole of MgCl2 = 0.005

Mass of MgCl2

= 0.005 x (24 + 2 x 35.5)

= 0.475 g

Question :

(b) volume of gas produced

No of mole of H2 = 0.005 mole

Volume of H2 = 0.005 x 24 dm3

= 0.12 dm3 or 120 cm3

Question :

Example 2 : A sample of insoluble lead (II) sulphate is prepared by mixing 50 cm3 of 1.0 mol dm-3 lead (II) nitrate solution and y of 1.5 mol dm-3 sulphuric acid.

[Relative atomic mass: O, 16 ; S, 32 ; Pb, 207]

(a) Calculate the volume, y, of the sulphuric acid needed to react completely with the lead (II) nitrate solution.

Question :

Solution :Chemical equation :

Pb(NO3)2 + H2SO4 PbSO4 + 2 HNO3

Mole ratio : 1 mole 1 mole 1 mole 2 mole Number of moles of Pb(NO3)2 = 1.0 x 50 = 0.05 mol 1000 Mole ratio of Pb(NO3)2 : H2SO4 = 1 : 1 Number of mole of H2SO4 reacted = 0.05 mol 1.5 x y = 0.05 mole 1000 y = 0.05 x 1000 = 33.33 cm3

Question :

(b) Calculate the mass of lead (II) sulphate obtained.Solution :

Number of mole of PbSO4 = Number of moles of Pb(NO3)2 = 0.05 mol

Mass of PbSO4 = 0.05 x (207 + 32 + 4 x 16) g

= 15.15 g

Question :

4. A sample of insoluble silver chloride is prepared by mixing 50 cm3 of 1.0 mol dm-3 silver nitrate solution and z cm3 of 0.5 mol dm-3 sodium chloride solution.

[Relative atomic mass: Ag 108; Cl 35.5]

(a) Write the chemical equation for the reaction between silver nitrate and sodium chloride.

AgNO3 + NaCl AgCl + NaNO3

Question :

(b) Calculate the volume, z, of the sodium chloride needed to react completely with the silver nitrate solution

Number of moles of AgNO3 = 1.0 x 50 = 0.05 mol 1000

Mole ratio of AgNO3 : NaCl = 1 : 1 Number of mole of NaCl = 0.5 x z = 0.05 mole 1000 Z = 0.05 x 1000 = 100 cm3

Question :

(c) Calculate the mass of silver chloride obtained.

Number of mole of AgCl = Number of moles of AgNO3 = 0.05 mol

Mass of AgCl = 0.05 x (108 + 35.5) g

= 7.175 g

Qualitative Analysis

Qualtitative analysis of a salt is a chemical technique used to identify the ions that are present in a salt by analysing its physical and chemical properties.

Colour (solid or solution)

Substance or cation or anion

Green powder , CuCO3

Blue powder Fe2+ Cu2+

Brown powder Fe3+

Black powder CuO, MnO2

Yellow powder when hot and white when cold

Brown powder when hot and yellow when cold

Blue solution Cu 2+

Pale green solution Fe 2+

Brown solution Fe3+

Solid : White

Solution : colourless

Cation : Ca2+ , Al3+ , Mg2+ , Pb2+ , Zn2+ NH4

Solid : White

Solution : colourless

Anion : Cl- , CO32- ,

SO42- , NO3

Salts Solubility in water Colour

Lead(II) chloride, silver chloride, barium sulphate, lead(II) sulphate and calcium sulphate

Insoluble white

Copper(II) carbonate Insoluble green

Iron(II) sulphate soluble green

Iron(III) salts except carbonate

Soluble Brown

Salts Solubility in water

Colour

Lead(II) sulphate

Insoluble white

Zinc chloride soluble white

Magnesium carbonate

Insoluble white

Salts Solubility in water

Colour

Ammonium carbonate

soluble white

Lead(II) iodide / chromate(VI)

Insoluble Yellow

Confirmatory Tests for gases

Gas Method Diagram Observation

Carbon dioxide

Bubble the gas produced into lime water

Lime water turn milky/chalky

Oxygen Insert a glowing splinter into the test tube

Glowing splinter will be relighted

Nitrogen dioxide

Observe the colour of gas produced. Bring a piece of moist blue litmus paper to the mouth of the test tube

Brown gas, blue litmus paper change to red

Chlorine Observe the colour of the gas.Bring a piece of moist blue litmus paper to the mouth of the test tube

Greenish yellow gas The colour of litmus paper change from blue red white

Ammonia Dip a glass rod into concentrated hydrochloric acid and bring a drop of acid to the mouth of the test tube/place moist red litmus paper at the mouth of the test tube

White fume formed Red litmus paper change to blue.

Hydrogen Bring a lighted splinter to the mouth of the test tube. Mg + HCl release hydrogen gas

‘Pop' sound is heard

Hydrogen chloride

Dip a glass rod into concentrated ammonia solution and bring a drop of ammonia to the mouth of test tube

White fume formed

Action of Heat On Carbonate Salts

Carbonate salts (except Na+ & K+ ) decompose on heating giving off carbon dioxide gas and residue metal oxide

Lime water turns chalky

Metal oxide colour

Copper(II) oxide

Zinc oxide Hot : Yellow

Cold : White

Lead(II) oxide

Hot : Brown

Cold : Yellow

Iron(III) oxide

Carbonate salt Action of heat

Potassium carbonate K2CO3 , Sodium carbonate Na2CO3

Not decompose by heat

Metal Carbonate metal oxide + carbon dioxide

Calcium carbonate

CaCO3 CaO + CO2 Observation : White solid formed. Gas liberated turn lime water chalky

Magnesium carbonate

MgCO3 MgO + CO2 Observation : White solid formed. Gas liberated turn lime water chalky

Aluminium carbonate

Al2(CO3)3 Al2O3 + 3CO2 Observation : White solid formed. Gas liberated turn lime water chalky

Carbonate salt

Action of heat

Zinc carbonate

ZnCO3 ZnO + CO2 Observation : The residue is yellow when hot and white when cold. Gas liberated turn lime water chalky

Lead(II) carbonate

PbCO3 PbO + CO2 Observation : The residue is brown when hot and yellow when cold. Gas liberated turn lime water chalky

Copper(II) carbonate

CuCO3 CuO + CO2 Observation : Black solid formed. Gas liberated turn lime water chalky

Action of Heat On Nitrate Salts

Nitrates Salts - Decompose on heating liberate nitrogen dioxide gas and oxygen gas except NaNO3 and KNO3 which liberate oxygen gas only.

Action of Heat On Nitrate Salts

Brown gas turn moist blue litmus to red(NO2)

Colourless gas relighted glowing splinter (O2 )

Action of Heat On Nitrates Salts

Nitrate salt

Action of heat

Metal Nitrate metal nitrite + oxygen

Potassium nitrate

2KNO3 2 KNO2 + O2 Observation : white solid formed, gas released relighted glowing splinter

Sodium nitrate

2NaNO3 2 NaNO2 + O2 Observation : white solid formed, gas released relighted glowing splinter

Nitrate salt

Action of heat

Metal Nitrate metal oxide + nitrogen dioxide + oxygen

Calcium nitrate

2Ca(NO3)2 2CaO + 4NO2 + O2 Observation : white solid formed, Brown gas which turns moist blue litmus red released. Another gas released relighted glowing splinter

Nitrate salt Action of heat

Magnesium nitrate

2Mg(NO3)2 2MgO + 4NO2 + O2 Observation : white solid formed, Brown gas which turns moist blue litmus red released. Another gas released relighted glowing splinter

Zinc nitrate 2Zn(NO3)2 2ZnO + 4NO2 + O2 Observation : The residue is yellow when hot and white when cold., Brown gas which turns moist blue litmus red released. Another gas released relighted glowing splinter

Nitrate salt Action of heat

Lead(II) nitrate

2Pb(NO3)2 2PbO + 4NO2 + O2 Observation : The residue is brown when hot and yellow when cold, Brown gas which turns moist blue litmus red released. Another gas released relighted glowing splinter

Copper(II) nitrate

2Cu(NO3)2 2CuO + 4NO2 + O2 Observation : black solid formed, Brown gas which turns moist blue litmus red released. Another gas released relighted glowing splinter

Confirmatory Tests for AnionsANIONS

Cl- SO42-

No Clue So Nothing

NO3- NO3

-Cl- SO42-

Observation : White precipitate

Observation :Brown ring

Confirmatory Tests for AnionsTEST FOR ANIONS

Anion test Anion test Anion test

Clue So Nothing

Cl- SO42- NO3

Reagent

Reagent Reagent

Clue So

NO3- Cl- SO4

Hati H+ NO3-

Agong Ag+ NO3-

Balik Ba 2+ Cl -

Haji H+ Cl -

Hendak H+ SO42-

Fetrah Fe 2+ SO42-

Harta H+ SO42-

Confirmatory Tests for Anions

Anions

CO3 2-

SO4 2-

+ HNO3

+ AgNO3

+ Dilute acid

+ HCl+ BaCl2

+ dilute H2SO4

+ FeSO4

+ concentrated H2SO4

Effervescence – CO2

Lime water turns milky

Ionic equation : 2H+ + CO32- H2O +

CO2 White precipitate

Ionic equation : Ag+ + Cl- AgCl

White precipitate

Ionic equation : Ba2+ + SO42- BaSO4

Brown ring

Salt K1

Add BaCl2 solution + HCl acid

Inference : sulphate ion

White precipitateformed

Inference : chloride ion

Add AgNO3 solution + HNO3 acid

Salt K2

Inference :Nitrate ion

Brown ring formed

Add FeSO4 solution+ concentrated sulphuric acid

Salt K3

Effervescence,Gas bubbles,Gas turn lime water chalky

Add sulphuric acid

Inference :Carbonate ion

Salt K4

CONFIRMATORY TEST FOR CATIONS

(SODIUM HYDROXIDE AS REAGENT)

CATION

Colourless / Unchanges

Add/put in NaOH SOLUTIONWhite precipitate

Precipitate

Coloured

Pb 2+ Zn 2+ Al 3+

Mg 2+ Ca 2+

Undissolved / not solubleIn excess NaOH solution

Dissolved / solubleIn excess NaOH solution

Add / Put in EXCESS NaOH solution

Pb 2+ Zn 2+ Al 3+ Ca 2+Mg 2+

MgCPZAL

Dissolved / solublein excess NaOH solution

Undissolved / not solublein excess NaOH solution

Al 3+Pb 2+ Zn 2+ Mg 2+

PZALDissolved Undissolved

White Unchanged

Ca 2+ Mg 2+

Calcium sulphate [ undissolved salt ]White Yellow Unchanged

Add KI solution

Add ion SO42- solution

Pb 2+ Zn 2+ Al 3+Lead (II) sulphate [ undissolved salt

Dissolved Unchanged

Zn 2+ Al 3+

Zn MAP

Add / put in NH3 solution

CONFIRMATORY TEST FOR CATIONS

(AMMONIA AS REAGENT)

CATION SOLUTION

Colourless / UnchangedPrecipitate

Add / put in NH3

NH4+ Ca 2+

Coloured White

Mg 2+ Al 3+ Pb 2+

Dissolved / solublein excess NH3 solution

Undissolved / not solublein excess NH3

Add / put in EXCESS NH3

Zn 2+ Mg 2+ Al 3+ Pb 2+

Zn MAPDissolved Undissolved

Zn MAP

Mg 2+ Al 3+ Pb 2+

Dissolved Undissolved

White Soluble / colourless

Add / put in SO42- solution

Lead (II) sulphate [ undissolved salt ]

Mg 2+ Al 3+

Colourless / unchangesWhite

Mg 2+ Al 3+

Add / put in EXCESS NaOH

Reaction of Cations with alkali solution

1. Positive ions are identified by their reactions with

a. sodium hydroxide NaOH solution b. Ammonia solution NH3

2. In these reactions, the cations (positive metal ions) produce different coloured precipitate which may or may not be soluble in excess alkali.

See if Precipitatedissolves

5 drops of alkali(NaOH or NH3)

Solution of cations

Look for Look for precipitateprecipitate

NaOH solution

A little In excess

Soluble ( , X )

Ca 2+ White precipitate X

Zn 2+ White precipitate Al 3+ White precipitate Pb 2+ White precipitate Mg 2+ White precipitate X

Cu 2+ Blue precipitate X

Fe 2+ Green precipitate X

Fe 3+ Brown precipitate X

Ammonia solution , NH3

A little In excess

Soluble ( , X )

Ca 2+ No change Zn 2+ White precipitate Al 3+ White precipitate X

Pb 2+ White precipitate X

Mg 2+ White precipitate X

Cu 2+ Blue precipitate X

Fe 2+ Green precipitate X

Fe 3+ Brown precipitate X

Salt K5

Add excess NaOH solution

Inference 2 : zinc, aluminium and lead(II) ionsWhite

precipitateDissolves in excess NaOH solution alkali

White precipitatedoes not dissolve in excess NaOH solution

Inference 3: magnesium or calcium ions

Inference 1 :

ammonium ion

Add 5 drops of NaOH solution

Salt K6

Add 5 drops of NH3 solution Add NH3 solution in

excess

Inference 5 : zinc ion

White precipitateDissolve in excess NH3 solution

White precipitatedoes not dissolve in excess NH3 solution

Inference 6: magneisum, aluminium, lead(II) ions

No White precipitateformed

Inference 4 : calcium ion

Confirmatory Tests for Fe 2+ , Fe 3+ , Pb 2+ and NH4

Cation Name of Reagent Observation

Pb 2+ Add a few drops of potassium iodide

to the test tube containing 2 cm3 of lead(II) nitrate solution (Pb 2+ ions)

Add 2 cm3 of distilled water and boil the mixture. Cool the contents using running water from the tap.

Yellow precipitate is formed

Which dissolve in the hot water

and is reappear on cooling

Fe 2+ Add a few drops of Potassium hexacyanoferrate(III) solution to the test tube containing 2 cm3 of iron(II) sulphate solution (Fe 2+ ions)

Dark blue precipitate is formed

Fe 3+ Add a few drops of potassium thiocyanate solution to the test tube containing 2 cm3 of iron(III) sulphate solution (Fe 3+ ions)

Blood red solution is formed

NH4 + Add a few drops of Nessler reagent to the test tube containing 2 cm3 of ammonium chloride solution (NH4 + ions)

Brown precipitate is formed

Chemical test of Fe2+ ions and Fe 3+ ions.

Solution contains Fe2+ ions or Fe3+ ions.

Light blue precipitate

Dark blue precipitate

Fe2+ ions

Fe3+ ions

K4Fe(CN)6

Potassium hexacyanoferrate(II)

Test I

(i) Pour 2 cm3 of iron(II) sulphate solution and 2 cm3 of iron(III) chloride solution into two test tubes respectively. Then add a few drops of potassium hexacyanoferrate(II) solution to two test tubes, Fe 2+ ions solution will form light blue precipitate whereas Fe 3+ ions solution will form dark blue precipitate

Test II

Solution contains Fe2+ ions or Fe3+ ions.

No change

Blood red solution

Fe2+ ions

Fe3+ ions

Potassium thiocyanate

(i) Pour 2 cm3 of iron(II) sulphate solution and 2 cm3 of iron(III) chloride solution into two test tube respectively. Then add a few drops of potassium thiocyanate solution to two test tubes, there is no change in Fe 2+ ions solution whereas Fe 3+ ions solution will form blood red solution.

Qualitative analysis to identify salts

Identify the salt S1

The following tests were carried out to identify salt S1. Based on the observations given for each test, state its inference. Finally, identify salt S1

Test Observation Inference

1. Heat S1 strongly in a test tube. Identify any gas liberated.Brown gas and a gas relights a glowing splinter are liberated.

Residue is brown when hot and yellow when cold

Nitrogen dioxide gas, oxygen gas released.

Residue is lead(II) oxide

2. Dissolve a spatulaful of S1 in distilled water. Divide into four portions and carry out the following tests:

Residue dissolve in acid to produce colourless solution

Fe 2+ , Fe 3+ or Cu 2+ ions may not be present

(a) add solution until excess.

White precipitate, dissolve in excess NaOH solution

Zn 2+ , Al 3+ or

Pb 2+ ions may be present

(b) add solution until excess

White precipitate, insoluble in excess ammonia solution

Pb 2+ or Al 3+ ion may be present

(c) add potassium iodide solution

Yellow precipitate formed

Confirm lead(II), Pb 2+ ions present

(d) add dilute , H2SO4 followed by FeSO4 solution. Carefully add about 1cm3 of concentrated H2SO4

Brown ring formed

Confirm nitrate , NO3

- ions present

Conclusion for salt S1 : Lead(II) nitrate

Identify the salt S2

1. Pour about 2 cm3 of S2 into a test tube. Add NaOH solution until excess

White precipitate, dissolve in excess NaOH solution

Zn 2+ , Al 3+ or Pb 2+ ions may be present

2. Pour about 2 cm3 of S2 into a test tube. Add NH3 solution until excess

White precipitate, dissolve in excess ammonia solution

Zn 2+ ions may be present

3. Pour about 2 cm3 of S2 into a test tube. Add dilute, HNO3 followed by silver nitrate, AgNO3 solution

No change Cl- ions not present

4. Pour about 2 cm3 of S2 into a test tube. Add dilute HCl solution, then add BaCl2 solution

White precipitate

SO42- ions

may be present

Conclusion for salt S2 : Zinc sulphate

SODIUM CARBONATE

SODIUM NITRATE

Test 1Add dilute (any acid)

Test 2add dilute H2SO4 followed by FeSO4 solution. Carefully add 1of concentrated H2SO4

Test 3Add dilute , followed by silver nitrate, solution

Test 4Add dilute HCl, followed by barium chloride, BaCl2 solution

S O D I U M C A R B O N A T E S O D I U M N I T R A T E

Result 1Effervescence

Result 2No change

Result 3No change

Result 4No change

Result 1No change

Result 2Brown ring

Result 3No change

Result 4No change

SODIUM CHLORIDE AND

SODIUM SULPHATE

Test 1Add dilute (any acid)

Test 2add dilute H2SO4 followed by FeSO4 solution. Carefully add 1of concentrated H2SO4

Test 3Add dilute , followed by silver nitrate, solution

Test 4Add dilute HCl, followed by barium chloride, BaCl2 solution

S O D I U M C H L O R I D E S O D I U M S U L P H A T E

Result 1No change

Result 2No change

Result 3White precipitate

Result 4No change

Result 1No change

Result 2No change

Result 3No change

Result 4White precipitate

Exercise :Page 144: Quick Review APage 147: Quick Review B

CHAPTER 8 : SALTS. Meaning and uses of Salts A salt is an ionic compound formed when the hydrogen ion, from an acid is replaced by a metal ion or an ammonium.

aq mg2 aq

aq h aq h2o

h aq mg2 aq h2o

aq h2g ionic equation

carbonate salts

following soluble salts

sulphate salts

chloride salts

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