03 Lab Manual

UTAR

FHSC1134 Inorganic Chemistry

Trimester 3

Lab manual version 4.1

Foundation in Science

1

Experiment 1

________________________________________________________________________

Title: Investigating the Properties of Period 3 Oxides

Aim:

To examine the oxides of Period 3 elements and describe their structure and bonding.

Introduction:

Generally, there are oxides of metals and non-metals. Metals burn in oxygen to form

basic oxides while non-metals form acidic oxides. Structurally, they are covalent or ionic

compounds. You are to do some simple observations and tests, to find out the differences

between the types of oxides provided and to account for these differences.

Apparatus:

Test tubes Test tube rack

Measuring cylinders Thermometer

Wooden splinter Glass rod

Materials:

Sodium peroxide Magnesium oxide

Silicon (IV) oxide Phosphorus pentoxide

Universal indicator solution Litmus paper

Safety measurements: Safety spectacle

**Warning:

Phosphorus (V) oxide is corrosive and irritates eyes, skin and lungs.

Sodium peroxide is also corrosive and a powerful oxidant.

Procedure:

Part A: Appearance:

Examine your oxide samples, and in Table 1, note for the physical states at room

temperature:

(a) whether it is solid, liquid or gas,

(b) its color (if any)

Part B: On mixing with water:

1. Set up 4 test tubes, side by side.

2. Into each test tube pour about 5 cm3

of distilled water.

3. In the test tube, place a thermometer.

a. Note the temperature.

b. Add half a spatula-tip of sodium peroxide and stir carefully with the glass rod.

c. Note after one minute, (i) the temperature, (ii) whether the solid has dissolved and

(iii) anything else you see. For example, is gas evolved at any time? If so, is the

gas acidic? Can you identify it using a simple test?

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d. Add 2 -4 drops of universal indicator solution, compare the color with the chart

provided, and note the pH indicated or use a piece of pH paper.

4. Repeat the above steps 3 (a) – (d), using, in turn, magnesium oxide, silicon (IV) oxide

and phosphorus (V) oxide.

5. Measure the pH of the water in the fifth test tube by adding 2-4 drops of universal

indicator solution for comparison with the above.

Results:

Table 1

Na2O2 MgO SiO2 P4O10

Appearance

Initial temperature/ºC

Final temperature/ºC

Solubility

pH of solution

Other observation(s)

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Questions:

1. Use your experimental results and your test books (if necessary) to complete a larger

copy of Table 2.

Table 2

Formula of oxide

Na2O2 MgO Al2O3 SiO2 P4O10 Cl2O

Melting-point/ o

C

Boiling-point/ o

C

State at s.t.p.

Action of water

pH of aqueous

solution

Acid/base nature

Conductivity of

liquid

Solubility in

hexane

Structure

Bonding

2. Write equations for any reactions which took place when you add the oxides to water.

3. Comment on the change in structure and bonding in the oxides of the elements in the

period between sodium and chlorine.

4. How does the acid-base nature of the oxides of the elements in Period 3 change with

increasing atomic number?

5. Can you relate this change in structure and bonding that take place along the period?

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Experiment 2

________________________________________________________________________

Title: The Solubility of Some Salts of Group II Elements

Aim:

To study the solubility of the sulphates, sulphites and hydroxides of Group II elements.

Introduction:

In this experiment, you add each of the anion solutions to 1 cm3 of each cation solution

provided, drop by drop, until the first sign of a precipitate appears. For each salt, the

solubility is proportional to the number of drops of anion added.

Apparatus:

Test tubes Test tubes racks

Graduated pipette Pasteur pipette

Beaker (50ml)

Materials:

0.1 M solution of the following cations: Mg2+

, Ca2+

, Ba2+

1.0 M solution of OH-

0.5 M solution of SO42-

0.5 M solution of SO32-

Procedure:

1. Set up three rows of three test tubes each.

2. For each row, label the 1st test tube Mg

2+, the 2

nd test tube Ca

2+, and 3

rd test tube Ba

2+.

3. Add 1 cm3

of the appropriate cation solution to each test tube by using a 1 cm3

graduated pipette.

4. To each test tube in the first row, add OH- solution drop by drop and shake until the

first sign of precipitate appears.

5. Record the number of drops of OH- solution used in Table 1.

6. If a precipitate is produced, classify the precipitate as slight (s) or heavy (h).

7. If no precipitate is produced after forty drops, then write ‘40+’ and assign the salt as

soluble.

8. Repeat steps 4 to 7 by replacing OH- with SO4

2- and SO3

2- to the 2

nd and 3

rd rows of

test tubes respectively.

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Results:

Table 1

Cation

Solution

Number of drops of anion solution added to form a precipitate

OH- SO3

2- SO4

2-

Mg2+

Ca2+

Ba2+

Questions:

1. For Group II, what are the solubility trends of the salts listed below:

(a) Hydroxides (b) Sulphites (c) Sulphates

2. Explain your answer for (1).

3. Use Table 2 to answer the following questions.

Table 2: Solubility of Group II compound in water at 298 K

Singly-charged anions Doubly-charged anions

Compound

Solubility

/mol per 100 g

of water

Compound

Solubility

/mol per 100g

of water

MgCl2 5.6 x 10-1

MgCO3 1.8 x 10-4

CaCl2 5.4 x 10-1

CaCO3 0.13 x 10-4

SrCl2 3.5 x 10-1

SrCO3 0.07 x 10-4

BaCl2 1.5 x 10-1

BaCO3 0.09 x 10-4

Mg (NO3)2 4.9 x 10-1

MgSO4 3600 x 10-4

Ca (NO3)2 6.2 x 10-1

CaSO4 11 x 10-4

Sr (NO3)2 1.6 x 10-1

SrSO4 0.62 x 10-4

Ba (NO3)2 0.4 x 10-1

BaSO4 0.009 x 10-4

Mg (OH)2 0.020 x 10-3

MgCrO4 8500 x 10-4

Ca(OH)2 1.5 x 10-3

CaCrO4 870 x 10-4

Sr (OH)2 3.4 x 10-3

SrCrO4 5.9 x 10-4

Ba (OH)2 15 x 10-3

BaCrO4 0.011 x 10-4

(a) Identify the group trends in solubility for each type of salt listed in Table 2.

(b) Does the solubility listed above for carbonates, sulphates and hydroxides match

with your findings in this experiment. If no, why?

(c) Which anions give more soluble compounds, singly-charged or doubly-charged

anions?

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Experiment 3

________________________________________________________________________

Title: Investigating the Properties of Aluminium and Its Compounds

Aim:

To investigate the properties of aluminium and its compounds.

Introduction:

Aluminium is the most abundant metal in the earth’s surface (7.5% by mass). The

abundance of Al, coupled with its attractive combination of physical and chemical

properties, accounts for the fact that it is one of the principal industrial raw materials used

by industrialized societies.

Aluminium has more metallic character compared to boron, which is in the same group in

the Periodic Table. It forms trivalent compounds due to the 3 electrons in its valence shell.

Generally, aluminium compounds are ionic in nature but many compounds of aluminium

show both ionic and covalent character. In this experiment, we will investigate the

properties of aluminium and some of its salts.

Apparatus: Test tube Glass dropper

Beaker Watch glass

Bunsen burner Glass rod

Materials: Aluminium strips Filter paper

Aluminium foil Litmus paper

Sand paper Diluted sodium hydroxide

Diluted hydrochloric acid 20% aluminium sulphate

Diluted nitric acid Diluted sulphuric acid

Diluted ammonium hydroxide Distilled water

Concentrated sodium hydroxide

Safety Precaution Steps:

Wear safety goggles. Wear gloves when using concentrated acid and base.

Procedure:

Part A: Reaction with air 1. Heat a small piece of aluminium foil using a bunsen burner. State the observation.

Part B: Reaction with acids 1. Put three small pieces of aluminium (clean with sand paper) into three separate test

tubes, each containing 5 cm3 of diluted hydrochloric acid, diluted nitric acid and

diluted sulphuric acid respectively.

2. Heat the solutions when necessary. Which acid reacts faster with aluminium?

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Part C: Reactions with alkali 1. Add 5 cm

3 of concentrated sodium hydroxide solution into a test tube containing a

piece of aluminium.

2. Identify the gas released.

Part D: Aluminium hydroxide and aluminium sulphate 1. Fill in about 10 cm

3 of 20% aluminium sulphate solution into a small beaker, and then

add 10 cm3 of diluted ammonium hydroxide solution.

2. Boil the mixture and filter it. Wash the aluminium hydroxide collected on the filter

paper with water.

3. Transfer a portion of the aluminium hydroxide prepared in step (2) on to a watch

glass. Using a glass dropper, add 6 drops of diluted hydrochloric acid.

4. Put a small amount of aluminium hydroxide from step (2) into a test tube. Add four

drops of sodium hydroxide solution, one drop at a time. Shake the test tube after

adding each drop and observe carefully for the appearance and disappearance of any

precipitate.

5. Test the pH of a small amount of aluminium sulphate solution using litmus paper and

record your observations.

6. Pour 5 cm3 of aluminium sulphate solution into a small beaker and add drop wise

sodium hydroxide solution slowly with stirring. Observe the reaction when excess

sodium hydroxide is added to the mixture.

Results:

Experiment Observation

A. Reaction with air

B. Reaction with acids

C. Reaction with alkali

D. Aluminium hydroxide and

aluminium sulphate

Question:

1. Write down all the chemical equations involved in this experiment.

2. Why must we use clean aluminium foil for this experiment?

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Experiment 4

________________________________________________________________________

Title: The Synthesis of Potassium Aluminium Sulfate (Alum)

Aim(s):

To prepare common aluminium complexes, alum, KAl(SO4)2.12H2O.

Introduction: Aluminium occurs naturally as mineral bauxite (primarily a mixture of Al2O3.3H2O,

Fe2O3 and SiO2), and it is purified in the following process:

Step 1: Purification of raw materials

Bauxite is firstly mined, crushed and then washed to remove water soluble impurities.

The remaining material is then dissolved in NaOH and heated up. Al2O3 is selectively

dissolved due to its properties of amphoteric oxide. The reaction is shown as below.

Al2O3 + 6NaOH + 3 H2O → 2Na3Al(OH)6

However, some of the crystalline forms of SiO2 can also be dissolved under the same

reaction. The equation is given as below.

SiO2 + 4NaOH → Na4SiO4 +2H2O

Both of these species are soluble, but Fe2O3 is a basic oxide and hence it is insoluble in

NaOH solution and can be filtered out. Over time the Na3Al(OH)6 decomposes to

Al(OH)3 (an insoluble species), which can also be filtered out.

Na3Al(OH)6 + 2H2O → 3NaOH + Al(OH)3 .2H2O

This is then decomposed by heating to temperatures above 1000 °C to give alumina,

Al2O3.

Al(OH)3 .2H2O → Al2O3 + 9 H2O

Step 2: Reduction of the alumina

The resultant alumina (Al2O3 ) is dissolved in molten cryolite (Na3AlF6), forming an ionic

and electrically conductive solution. This is decomposed by electrolysis later by using a

consumable carbon as anode with two concurrent reactions proceeding according to the

following equations:

Al2O3 + 3C → 2Al + 3CO

2 Al2O3 + 3C → 4Al + 3CO2

The use of consumable carbon anode lowers the required voltage by 1.0 V at the

operating temperature of 950-980 °C.

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Figure 1: Electrolysis process on extracting aluminium.

The environmental impact on extracting aluminium from its natural source (bauxite) is

quite high. Yet, the need for aluminium continues to grow due to its attractive properties.

Recycling aluminium products is one of the alternatives on saving environmental cost

due to its highly cost from the extraction process. Aluminium recycled from beverage

cans only requires 5 percent of energy if compared to the production of aluminium from

ores. In the year 2000, it was estimated that nearly 63 billion aluminium cans were

recycled in the United States. The commercial recycling of aluminium is a melting or

purification process that keeps the aluminium as aluminium metal. However, it needs to

dissolve aluminium in hot KOH in this experiment and use the dissolved aluminium to

make crystals of potassium aluminium sulphate dodecahydrate, KAl(SO4)2.12H2O. This

compound belongs to a family of metal sulphates called alums.

In this experiment, you will be recycling aluminium scrap in a very unusual way, and you

will produce two products which are potentially very useful: hydrogen gas (H2) and very

pure potassium aluminium sulphate (KAl(SO4)2.12H2O or alum). Hydrogen gas has great

potential use as fuel, if some of its dangerous properties can be controlled (mixtures of H2

and air are highly explosive). Hot aqueous hydroxide solution used in this experiment

will quickly remove the oxide layer on the aluminium surface and then dissolve the

aluminium metal by converting aluminium atoms to Al3+

. There are a number of Al3+

species formed in an aqueous environment, depending on the abundances of OH- and H

+

ions (the pH of the solution). Adding the H2SO4 to the mixture can lead to the desired

KAl(SO4)2.12H2O alum product.

Apparatus and Equipments:

Beaker, 250ml Measuring cylinder

Butchner funnel Pasteur pipette

Filtration flask Sandpaper

Filter paper Glass rod

Ice bath

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Materials: Aluminium beverage can

Boiling chips

Distilled water

Ethanol

2 M Potassium hydroxide, KOH

9 M Sulphuric acid, H2SO4

Safety measurements: Protective glove

Safety spectacle

Procedure:

Part A: Dissolution of the aluminium 1. Cut about 2 x 5 cm

2 of aluminium strip from aluminium can provided. Scrap off

thoroughly any paint and/or plastic coating from both sides by using a piece of

sandpaper.

2. Weigh out approximately 1.0 g of the scrap aluminium strip into a 250 cm3 beaker.

Record the mass.

3. Cut the weighted aluminium strip into small pieces and place the pieces in a clean

beaker, followed by the addition of 50 cm3 of 2 M KOH into the beaker. Perform this

operation in the fume hood.

4. Heat the mixture very gently on a hot plate if the reaction is proceeding too slowly.

Hydrogen gas produced is very explosive.

Therefore, you MUST:

ENSURE THAT YOU ONLY HEAT THE BOTTOM OF THE BEAKER AND

NOT LETTING THE FLAME GET NEAR THE TOP.

Concentrated sulphuric acid is very corrosive and reacts vigorously with water.


WEAR GLOVES AND SAFETY SPECTACLES.

DISPOSE UNWANTED ACID BY COOLING AND POURING SLOWLY

INTO AN EXCESS OF WATER.

Ethanol is very flammable.


KEEP THE BOTTLE CLOSED ALWAYS.

KEEP THE BOTTLE AWAY FROM FLAMES.

WEAR SAFETY SPECTACLES.

Potassium hydroxide is very harmful if swallowed and causes severe burns.


WEAR APPROPIATE PROTECTIVE CLOTHING.

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5. Bubbles should form from the reaction between aluminium and aqueous potassium

hydroxide. Add distilled water to maintain the volume at approximately 25 cm3 (if the

liquid level in the beaker drops to less than half of its original volume). The reaction

is complete when the gas evolution ceases and without any trace of aluminium pieces.

6. Vacuum filter the solution (if the solution contains any remaining solid). Use a small

amount of water (not more than 200 cm3) to rinse the beaker and funnel.

7. Allow the filtrate to cool down to room temperature.

Part B: Formation of Al(OH)3 and removal of OH-

1. Transfer 20 cm3 of 9 M H2SO4 to the filtrate.

2. Add 20 cm3 of 9 M H2SO4 drop wise to the filtrate slowly and with stirring.

3. White precipitate will form upon the addition of sulphuric acid and will dissolve

when excess of sulphuric acid is added into the beaker.

(Note: There may be a small amount of undissolved white precipitate. Add 2 or 3

boiling chips (if necessary) and heat the solution gently until the solution becomes

clear.)

Part C: Precipitation of alum crystals 1. Prepare a half-full ice bath and add water until the bowl is three quarters full.

2. Cool the solution in Part B to room temperature and then place the beaker in the ice

bath. Crystals of alum should form. Allow it to cool for 15 minutes.

3. Vacuum filter the product and wash with 10– 20 cm3 of ethanol/water solution (1:1).

4. Allow the product to dry for a few minutes, remove the boiling chips and record the

weigh.

5. Calculate the percentage yield of alum obtained.

Results:

Part A: Observation(s)

Table 1: Observation(s) for each part of experiment

Observation(s)

(A) Dissolving the aluminium

(B) Formation of Al(OH)3

(C) Precipitating alum crystals

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Part B: Calculation

Table 2: Mass of aluminium

Mass of empty beaker (g)

Mass of beaker + aluminium strips (g)

Mass of aluminium strips used (g)

Table3: Mass of alum

Mass of filter paper (g)

Mass of filter paper + alum (g)

Mass of alum (g)

Questions: 1. Identify the gas released during the reaction of aluminium beverage can and

potassium hydroxide.

2. What is the white precipitate formed during the addition of concentrated sulphuric

acid? Write down the balanced equation that involved.

3. Why the white precipitate disappears when excess of sulphuric acid is added? Write

down the balanced equation that involved.

4. Write down the complete balanced equation of aluminium hydroxide, Al(OH)3, to

form alum, KAl(SO4)2.12H2O.

5. What is the theoretical yield of alum, in grams, obtained from the experiment?

6. What is the percentage yield of alum obtained from the experiment?

%100)(

)(×=

gyieldltheoretica

gyieldactualyieldPercentage

7. List out several common usages of aluminium.

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Experiment 5

________________________________________________________________________

Title: Reaction of Tin and its Common Ions

Aim:

To shows the reactions of tin with acids and to carry out some of the common reactions

of Sn2+

(aq).

Introduction:

In this experiment you will find out if the metallic character of tin is evident from its

reactions with acids. You treat the elements with two acids: an oxidizing agent (nitric

acid) and a non-oxidizing acid (hydrochloric acid). In each case you attempt to identify

any gases evolved.

The reactions of the divalent ions are included here to demonstrate the relative stability of

the +2 state in tin.

Apparatus:


Bunsen burner Test tube holder/ clamp

Asbestos sheet Beaker

Graduated pipette wooden splinter

Materials:

Diluted hydrochloric acid (2 M) Concentrated hydrochloric acid

Concentrated nitric acid 2 M NaOH solution

2.0 M ammonia solution 0.1 M K2CrO4

0.02 M Na2S solution 0.1 M KI solution

Tin (small pieces/granule) 0.1 M solution of Sn2+

ions (in diluted HCl acid)

0.01 M KMnO4 (in diluted ethanoic acid)

Safety measurement:

Safety spectacles

**Warning:

Sodium sulphide is toxic and corrosive and evolves highly poisonous hydrogen

sulphide gas on contact with acids.

Concentrated hydrochloric acid is very corrosive.

Concentrated nitric acid is very corrosive and a powerful oxidant.

WEAR PROTECTIVE GLOVES AND SAFETY SPECTACLES. USE SODIUM SULPHIDE IN THE FUME-CUPBOARD.

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Procedure:

1. Place a small piece of tin in each of three test tubes.

2. To one test tube add about 2 cm3 of diluted hydrochloric acid and heat gently.

Can you see or detect a gas? If not, repeat the experiment carefully using about 2 cm3

of concentrated hydrochloric acid.

3. In the remaining test tube add about 2 cm3 of concentrated nitric acid to the tin and

heat gently.

4. Record your results in a larger copy of Table 1.

5. Add approximately 2 cm3 of the Sn

2+ solution to each of seven test tubes.

6. Add each of the following reagents to Sn2+

:

(a) Diluted sodium hydroxide solution, initially drop-by-drop, and then to excess;

(b) Ammonia solution, initially drop-by-drop, and then to excess;

(c) About 2 cm3 of diluted hydrochloric acid, heat the mixtures and then cool them

under running cold water;

(d) About 2 cm3 of acidified potassium manganate (VII) solution;

(e) About 1 cm3 of potassium chromate (VI) solution;

(f) 5 drops of sodium sulphide solution (do this in the fume-cupboard and dispose of

the mixture by pouring into the fume-cupboard sink);

(g) About 2 cm3 of aqueous potassium iodide.

7. Record your observations in a larger copy of results Table 2.

Result:

Table 1: Reaction of Tin granules with acids:

Acid Observations

Diluted hydrochloric acid

Concentrated hydrochloric acid

Concentrated nitric acid

Table 2: Reaction of acidified Sn2+

(aq) with other reagents:

Reagent Observations

a) Sodium hydroxide solution

b) Ammonia solution

c) Diluted hydrochloric acid

d) Acidified potassium manganate (VII)

solution

e) Potassium chromate (VI)

f) Sodium sulphide solution

g) Potassium iodide solution

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Questions:

1. Complete the following equations:

Sn(s) + HCl(aq.) →

2. Are the reactions between the elements and hydrochloric acid typical of metals?

Explain your answer.

3. Reactions with nitric acid tend to be complex and equations are not generally required.

The questions illustrate some general points.

(a) Which gas did you detect when nitric acid reacted with tin?

(b) Do other metals behave in similar way with nitric acid? Give one example.

(c) Why does nitric acid behave differently from hydrochloric acid?

4. Use your text-book(s) to complete and balance the equations in Table 3.

In the comments column you should describe the type of chemical reaction occurring

and any other important feature(s).

Table 3:

Equations Comments

Sn2+

(aq) + OH-(aq) →

Sn (OH)2 (s) + OH-(aq) →

The precipitate dissolves in excess NaOH

to form a stannate (II) ion2 , Sn (OH)6

4-

2MnO4-(aq) + 16H

+ (aq) + 5Sn

2+ (aq) →

Cr2O72-

(aq)** + H+ (aq) + Sn

2+(aq) →

Sn2+

(aq) + S2-

(aq) →

* Various formulae have been proposed for the stannate (II) ion ranging from SnO22-

for the

anhydrous form to Sn(OH)42-

and Sn(OH)64-

for the hydrated forms. Sn(OH)64-

seems the most

probable.

** Chromate (VI) (CrO42-

) changes to dichromate (VI) (Cr2O72-

) when acidified

[2CrO42-

(aq) + 2H+ (aq) → Cr2O7

2-(aq) + H2O (1)]

5. In the experiment, potassium manganate (VII) solution was acidified with diluted

ethanoic acid and hot, as is usual, diluted hydrochloric acid or diluted sulphuric acid.

With the aid of your text-book (s), explain why you think this change was made. Give

any relevant equations in your answer and state what you would observe if the MnO4-

solution were acidified with HCl (aq) or H2SO4 (aq).

6. Predict the reaction between aqueous tin (II) ions and a solution of mercury (II)

chloride. What do you think you would observe in this reaction?

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Experiment 6

________________________________________________________________________

Title: Investigation of thermal stability of nitrates

Aim(s):

To identify the products of thermal decomposition of s-block nitrates.

To predict the order of thermal stability of these compounds.

Introduction:

This experiment attempts to test the effect of heat on the solid nitrates of Groups I and II.

You are to heat small samples of the nitrate of each element, using the same size flame.

Take note the time taken to detect the product of decomposition. Refer to standard

textbook for the common tests for nitrogen dioxide (brown gas), produced from nitrates.

Apparatus:

Boiling tubes L-shaped delivery tube with bung (boiling tube)


Stop-watch Wooden splinter

Bunsen burner Boiling tube holder/ clamp

Materials:

Solid nitrates (anhydrous) of Groups I and II

Lime water (saturated calcium hydroxide solution)

2 M hydrochloric acid (diluted)

**Warning:

Nitrogen dioxide gas is poisonous. Heat nitrates in a fume cupboard. Nitrates are

strong oxidizing agents.

DO NOT ALLOW PIECES OF GLOWING SPLINT TO DROP ON TO

HOT NITRATES.

Procedure:

Effect of Heat on Nitrates:

1. Set up two rows of boiling tubes. First row with 2 boiling tubes and second row with

4 boiling tubes. Add a spatula-measured of the nitrates as stated in Table 1.

2. Start the stop-watch at the moment you begin heating the first nitrate by holding the

end of the boling tube just above the blue cone of a roaring Bunsen flame.

3. Test for oxygen at short regular intervals (place the glowing splinter in the same part

of the boling tube) and also look for the first sign of a brown gas. (A white

background is helpful)

4. In Table 1, record the time for the first appearance of brown fumes or when oxygen is

detected. (Whichever method of detection you choose for one member of a group you

must also apply to the other members of the same group) Continue heating for

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another minute after gas is detected.

5. Repeat steps 2, 3 and 4 for the other nitrates in turn.

6. To the cold solid residue remaining after each nitrate is heated, add a few drops of

diluted hydrochloric acid and warm.

Table 1: Effect of Heat on s-Block Nitrates

Nitrate Time to detect

O2/NO2 Observations

Effect of adding diluted

HCl to cold residue

NaNO3

KNO3

Mg(NO3)2

Ca(NO3)2

Sr(NO3)2

Ba(NO3)2

Questions:

1. Explain why many of these nitrates rapidly turn into colorless liquids on first heating.

On further heating, they become white solids again, before they decompose.

2. Describe the trend in thermal stability of the nitrates when going down a group in the

periodic table.

3. How does Group I nitrates compare with their Group II counterparts in terms of

thermal stability?

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Experiment 7

________________________________________________________________________

Title: Halogen-Halide Reactions in Aqueous Solution

Aim:

To investigate the order of oxidizing ability of the halogens Cl2, Br2 and I2 in aqueous

solution.

Introduction:

You mix each of the aqueous solutions with halide ion solutions, C1- (aq), Br

- (aq), and I

-

(aq) in turn, and see whether a reaction takes place. The addition of hexane to the

halogen-halide mixture enables you to recognize the halogen molecules present. The

halogen which oxidizes most of the other halide ions will clearly be the strongest

oxidizing agent.

Apparatus:

Test tubes fitted with cork Test tube rack

Graduated pipettes

Materials:

Bromine water, Br2 (aq) --------------- Chlorine water, C12 (aq) --------------------------------

Iodine solution, I2 in KI (aq) Potassium bromide solution, KBr

Potassium chloride solution, KC1 Potassium iodide solution, KI

Hexane, C6H14 ----------------------------------------------------------------------------------------------------------------

Safety measurement:

Safety spectacle

Hazard warning

KEEP HEXANE WELL STOPPERED AND AWAY FROM FLAMES

Halogen and organic vapors must not be inhaled.

THE LABORATORY MUST BE WELL VENTILATED AND REAGENT

BOTTLES AND TEST TUBES STOPPERED AS MUCH AS POSSIBLE.

Procedure:

1. Reaction (if any) of iodide with chlorine and bromine:

(a) To each of two test tubes add about 1 cm3 of potassium iodide solution.

(b) To one of those tubes, add about the same volume of chlorine water, and to the

other add the same volume of bromine water.

(c) Cork and shake the tubes and note the colour change (if any).

(d) To each tube add about 1 cm3

of hexane. Cork and shake, allow it to settle. Note

the colour of each layer.

(e) Decide which reactions have taken place, and complete Table 1.

2. Reaction (if any) of bromide with chlorine and iodine.

Repeat the above steps, 1 (a) – (e), using potassium bromide with chlorine water and

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iodine solution.

3. Reaction (if any) of chloride with bromine and iodine

Repeat steps 1 (a) – (e) using potassium chloride with bromine water and iodine

solution.

Results:

Table 1

Chlorine

water

Bromine

water

Iodine

solution

Initial colour

1.

Colour after shaking with KI solution

Colour of each layer

after shaking with

hexane

Upper

Lower

Conclusion

2.

Colour after shaking with KBr solution


after shaking with

hexane

Upper

Lower

Conclusion

3.

Colour after shaking with KCl solution


after shaking with

hexane

Upper

Lower

Conclusion

Questions:

1. From the results:

(a) Does I2 (aq KI) oxidize Cl – (aq) and Br

– (aq)?

(b) Does Br2 (aq) oxidize Cl –

(aq) and I- (aq)?

(c) Does Cl2 (aq) oxidize Br –

(aq) and I –

(aq)?

2. Explain your answer for question (1).

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3. State the purpose of using hexane in this experiment.

4. Write ionic equations for the reactions that have taken place.

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Experiment 8 ________________________________________________________________________

Title: Reaction of Halides in Solution

Aim:

To find out whether the ions, Cl-, Br

- and I

-, react in solution with certain reagents and

identify the products formed in the reactions.

Introduction:

In this experiment, you will add various reagents to separate samples of solutions

containing the Cl-, Br

- and I

- ions. In many of the reactions, precipitates are formed.

Where you are asked to add another reagent to excess, you should look carefully to see if

any of the precipitate dissolves.

Apparatus:

15 test tubes with corks Test tube racks

Pasteur pipettes

Materials:

Potassium bromide solution, KBr Potassium chloride solution, KCl

Potassium iodine solution, KI Silver nitrate solution, AgNO3

Diluted nitric acid, HNO3 Ammonia solution, NH3

Hydrogen peroxide solution, H2O2 Diluted sulphuric acid, H2SO4

Procedure:

Add the following reagents to 1 cm3 of the chloride, bromide and iodide solutions in turn,

and record your observations in Table 1.

1. Add approximately 1 cm3 of silver nitrate solution and shake gently. State the

observation. Move the three test tubes to a dark cupboard and leave them there till the

end of the lesson and note their appearance again.

2. Add silver nitrate solution as in (1). Leave these test tubes in their racks until the end

of the lesson, noting their appearance every 10-15 minutes.

3. Add approximately 1 cm3 of silver nitrate solution followed by excess (e.g. 5 cm

3)

diluted nitric acid. Cork the test–tubes and shake vigorously.

4. Add approximately 1 cm3 of silver nitrate solution followed by excess (e.g. 5cm

3)

ammonia solution. Cork the test tubes and shake.

5. Add approximately 1 cm3 of hydrogen peroxide solution followed by approximately 1

cm3 of diluted sulphuric acid. Cork these test tubes and allow them to stand. Add any

further reagent (s) which you think will help you to decide what has happened.

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Results:

Table 1:

Test Chloride Bromide Iodide

Action of AgNO3 (aq)

Effect of standing in

(a) dark

(b) light

Action of AgNO3(aq)

followed by diluted HNO3

Action of AgNO3 (aq)

followed by NH3 (aq)

Action of H2O2 (aq) and

diluted H2SO4 (aq)

Question:

1. Write ionic equations for the reactions between each of the three halide solutions and

silver nitrate solution.

2. Suggest chemical tests that can be used to distinguish between:

(a) Cl-(aq) and Br

-(aq),

(b) Br-(aq) and I

-(aq).

3. Write an ionic equation for the reaction between aqueous iodine and acidified

hydrogen peroxide.

4. Why do you think there is no reaction occurs between acidified hydrogen peroxide

and the other halide ions?

5. Suggest a reason for the darkening effect of light on the silver chloride and silver

bromide precipitates.

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Experiment 9 ________________________________________________________________________

Title: Complex Formation and Precipitation

Aim:

To investigate complex formation and precipitation reactions.

Introduction:

Complex formation and precipitation reactions are the basis of qualitative inorganic

analysis. The reactions involve the manipulation of chemical equilibria by either adding

specific ions or adjusting their concentrations, in order to cause a precipitate to form or to

dissolve.

Substances that show conductivities proportional to their concentrations in aqueous

solution are classified as strong electrolytes. Strong electrolytes will dissociate

completely into ions in aqueous solution. Examples of salts are Li, Na, K, all nitrates,

many metallic chloride salts, all hydroxides, HCl, HNO3, and HClO4. In the present study

the ions Na+, K

+, and NO3

- can be regarded as spectator ions, not taking part in the

reactions.

Complex formation

Addition of ammonia or sodium hydroxide to aqueous solutions of metal ions will often

cause precipitation of the metal hydroxide, eg:

Sn2+

(aq) + 2OH-(aq) Sn(OH)2(s)

However, on addition of excess ammonia, or hydroxide ion, some metal hydroxides

dissolve by forming complex ions.

Zn(OH)2(s) + 4NH3(aq) [Zn(NH3)4]2+

(aq) + 2OH-(aq)

Sn(OH)2(s) + 3OH-(aq) + 3H2O(l) [Sn(H2O)3(OH)3

-](aq)

The metal ions Zn2+

, Cu2+

and Pb2+

form complexes in solution with a coordination

number of 4. This means that four ligands are joined to the central metal ion to form

complex ion such as [Cu(NH3)4]2+

. Zn2+

and Pb2+

can also form six-coordinate complexes.

Precipitation Reaction

Almost all salts of the alkali metals are soluble in water and strong electrolytes, but

anions such as phosphate, carbonate or oxalate form insoluble salts with most other

cations. Some such precipitates can be made to dissolve by changing the pH or by adding

suitable ligands, soluble complexes can be formed whose large formation constant

overwhelms the small solubility product of the original salt.

Silver halides are insoluble in water. They can be distinguished by the differing actions of

dilute and concentrated ammonia solutions, thereby furnishing tests for the presence of

halide ions.

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Apparatus:

Test tubes Droppers

Red litmus paper Pipette

Materials:

1 M hydrochloric acid Potassium chloride

4 M ammonia solution Manganese (II) chloride

2 M sodium hydroxide Magnesium chloride

Zinc nitrate Calcium chloride

Copper (II) sulphate Trisodium phosphate

Caution:

Cleanliness of test tubes is essential.

Procedures:

Part A: Complex formation

1. Label two test tubes and add 4 M ammonia solution into respective test tubes of 0.5

cm3 of dilute zinc nitrate and copper (II) sulphate solutions drop wise until each is

just alkaline (test with red litmus paper).

2. Add more ammonia solution, and observe if there is any precipitate dissolves.

3. Repeat step 1 and 2 by using 2 M sodium hydroxide in place of ammonia solution.

Part B: Phosphate precipitation reaction

1. Label five test tubes and add 1 cm3 of each solution (potassium chloride, manganese

(II) chloride, magnesium chloride, calcium chloride and copper (II) sulphate) into

respective test tubes.

2. Add 1 cm3 of trisodium phosphate solution to each of the test tubes and shake.

Observe for any precipitation and record in the table.

3. Add a few drops of 1 M HCl to any test tube with precipitate formed and shake.

Note: The precipitates are not necessarily phosphates. Some precipitates are due to

hydrolysis of the anion in solution.

Results:

Part A: Complex formation

Observations

Limited NH3 Excess NH3

Zn(NO3)2

CuSO4

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Observations

Limited NaOH Excess NaOH

Zn(NO3)2

CuSO4

Part B: Phosphate precipitation reaction

Observations

KCl

MnCl2

MgCl2

CaCl2

CuSO4

Questions:

1. Explain briefly each of the reactions involved.

2. Write down appropriate equation(s) to describe your observations.

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Experiment 10 ________________________________________________________________________

Title: Identification of Unknown Metal Ions

Aim:

To identify the presence of unknown metal ions in series of solutions.

Introduction:

Common metallic cations can be separated from one another and identified by using

qualitative analysis. During analysis, inferences can be made on the colour, solubility, the

effect of heat on the salt, the type of gas given out, reaction with other reagents and

confirmatory test. When the unknown cation has been identified, equations should be

added to support the conclusions.

Colours of common cation solutions:

Colourless Blue Yellow Pink Green

Ag+, Pb

2+, Ca

2+, Ba

2+,

Zn2+

, Mg2+

, Cd2+

Cu2+

Fe3+

, Cr3+

Co2+

, Mn2+

(very pale)

Ni2+

, Fe2+

(pale)

The cations supplied belong to four main groups:

Group 1: Cations whose chlorides are relatively insoluble

Group 2: Cations whose sulfates are relatively insoluble

Group 3: Cations which form soluble amine complexes in excess ammonia (at a pH of

approximately 10)

Group 4: Cations which form insoluble metal hydroxides in a solution of excess ammonia

at pH10

Apparatus:

Test tubes Dropper

Boiling tubes Centrifuge

Water bath Centrifuge tubes

Materials:

6 M ammonia solution Silver nitrate, AgNO3

6 M hydrochloric acid Barium nitrate, Ba(NO3)2

6 M sulphuric acid Iron (III) nitrate, Fe(NO3)3

1 M hydrogen peroxide Chromium (III) nitrate, Cr(NO3)3

Universal indicator paper Copper (II) nitrate, Cu(NO3)2

Boiling chips Ethanol

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Safety Precautions:

• Safety glasses must be worn at all times.

• Most of the metal ions used are poisonous. Do not throw residues containing metal

ions or sulfides down the sink. Containers are provided in the fume cupboards for

such residues.

• When heating any acid always heat cautiously and do not allow any vigorous reaction

to occur.

• Do not take the test tubes away from fume hood until the reaction has ceased

completely.

• All test tubes used for tests should be cleaned thoroughly.

Procedures:

1. In a set of 5 clean test tubes, add each of them with 10 drops of one unknown

solutions. Label the test tubes A - E.

2. Add 4 drops of 6 M HCl to each test tube. Mix the solutions thoroughly. Record your

observations about precipitation (if any) and colors.

3. Follow the flow chart illustrated in the next page for further investigations.

Notes for Tests:

• Transferring the solution to boiling tubes is important to prevent loss of solution due

to over rapid boiling off of ethanol. Removal of ethanol is complete when vigorous

bubbling ceases.

• The 6 M NH3 should be added drop wise and the pH carefully monitored by using the

universal indicator paper provided. The simplest way is to dip a clean stirring rod into

the solution and touch a drop to a small section of universal indicator paper which is

resting on a watch glass.

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NO YES

NO

NO

YES

YES

Some or all precipitate dissolves to

give a yellow supernatant solution.

Cr present as H2CrO4

Precipitate remains, no colour change.

Fe present

Precipitate

Insoluble hydroxide

Subgroup: Fe3+,

Cr3+

- Centrifuge and discard supernatant.

- Transfer as much solid as possible to a

larger test tube.

- Add 6 drops – 1 cm3 of 1 M H2O2. Stir

thoroughly and place in a boiling water

bath for 5-10 minutes

Add excess of 6 M

NH3 (fume hood)

to the original

solution.

Solution - Transfer to boiling tube, add 3 boiling chips. Boil

to expel ethanol (water bath).

- Return the solution to a small test tube and add 6 M

NH3, dropwise, until has pH 8-10 (test with

universal indicator paper)

Subgroup: Fe3+,

Cr3+

, Cu2+

Add 4 drops of 6 M HCl

and mix thoroughly

Precipitate Insoluble sulphate

Group: Ba2+

Solution Add 1cm

3 of ethanol, 2 drops of 6 M

H2SO4 and stand for 5 minutes.

Group: Ba2+

, Fe3+,

Cr3+

, Cu2+

Precipitate Insoluble Chloride

Group: Ag+

Unknown metal ions

(Ag+, Ba

2+, Fe

3+, Cr

3+, Cu

2+)

Solution Soluble amine

Subgroup: Cu2+

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Results:

Qualitative Tests for the Individual Ion Solutions – Observations

Ion Chloride test

(HCl)

Ethanol test NH3 test H2O2 test Possible ion

A

B

C

D

E

Questions:

1. Write down all the balanced equations involved.

2. Describe briefly the reactions involved.

Documents

03 Lab Manual