Chem experiments

UTAR
FHSC1134 Inorganic Chemistry
Trimester 3

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
Measuring cylinders
Wooden splinter

Test tube rack
Thermometer
Glass rod

Materials:
Sodium peroxide
Silicon (IV) oxide
Universal indicator solution

Magnesium oxide
Phosphorus pentoxide
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/ oC
Boiling-point/ oC
State at s.t.p.
Action of water pH of solution aqueous

Acid/base nature
Conductivity
liquid
Solubility
hexane
Structure

of in 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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FHSC1134 Inorganic Chemistry
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Experiment 2
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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
Graduated pipette
Beaker (50ml)

Test tubes racks
Pasteur pipette

Materials:
0.1 M solution of the following cations: Mg2+, Ca2+, Ba2+
1.0 M solution of OH0.5 M solution of SO420.5 M solution of SO32Procedure:
1. Set up three rows of three test tubes each.
2. For each row, label the 1st test tube Mg2+, the 2nd test tube Ca2+, and 3rd test tube Ba2+.
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 SO42- and SO32- to the 2nd and 3rd rows of test tubes respectively.

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Results:
Table 1
Number of drops of anion solution added to form a precipitate
Cation
Solution
OHSO32SO42Mg2+
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
Solubility
Solubility
Compound
/mol per 100 g
Compound
/mol per 100g of water of water
MgCl2
5.6 x 10-1
MgCO3
1.8 x 10-4
-1
CaCl2
5.4 x 10
CaCO3
0.13 x 10-4
SrCl2
3.5 x 10-1
SrCO3
0.07 x 10-4
-1
BaCl2
1.5 x 10
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
-1
Sr (NO3)2
1.6 x 10
SrSO4
0.62 x 10-4
Ba (NO3)2
0.4 x 10-1
BaSO4
0.009 x 10-4
-3
Mg (OH)2
0.020 x 10
MgCrO4
8500 x 10-4
Ca(OH)2
1.5 x 10-3
CaCrO4
870 x 10-4
-3
Sr (OH)2
3.4 x 10
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? Lab manual version 4.1
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Experiment 3
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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
Beaker
Bunsen burner
Materials:
Aluminium strips
Aluminium foil
Sand paper
Diluted hydrochloric acid
Diluted nitric acid
Diluted ammonium hydroxide
Concentrated sodium hydroxide

Glass dropper
Watch glass
Glass rod

Filter paper
Litmus paper
Diluted sodium hydroxide
20% aluminium sulphate
Diluted sulphuric acid
Distilled water

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 cm3 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 cm3 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
A. Reaction with air

Observation

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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FHSC1134 Inorganic Chemistry
Trimester 3

Experiment 4
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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
Butchner funnel
Filtration flask
Filter paper
Ice bath

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Measuring cylinder
Pasteur pipette
Sandpaper
Glass rod

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FHSC1134 Inorganic Chemistry
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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
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.
Therefore, you MUST:
WEAR GLOVES AND SAFETY SPECTACLES.
DISPOSE UNWANTED ACID BY COOLING AND POURING SLOWLY
INTO AN EXCESS OF WATER.
Ethanol is very flammable.
Therefore, you MUST:
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.
Therefore, you MUST:
WEAR APPROPIATE PROTECTIVE CLOTHING.
Procedure:
Part A: Dissolution of the aluminium
1. Cut about 2 x 5 cm2 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.

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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 OH1. 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?

Percentage yield =

actual yield ( g )
× 100 % theoretical yield ( g )

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:
Test tubes
Bunsen burner
Asbestos sheet
Graduated pipette

Test tube rack
Test tube holder/ clamp
Beaker
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
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 Sn2+ 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
Sn2+ (aq) + OH-(aq) →
Sn (OH)2 (s) + OH-(aq) →

Comments
The precipitate dissolves in excess NaOH to form a stannate (II) ion2 , Sn (OH)64-

2MnO4-(aq) + 16H+ (aq) + 5Sn2+ (aq) →
Cr2O72-(aq)** + H+ (aq) + Sn2+(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)
→
Cr2O72-(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 MnO4solution 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
Test tubes
Stop-watch
Bunsen burner

L-shaped delivery tube with bung (boiling tube)
Test tube rack
Wooden splinter
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
Observations
O2/NO2

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
Graduated pipettes

Test tube rack

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
Colour after shaking with KI solution
Colour of each layer
1. after shaking with hexane Upper
Lower

Conclusion
Colour after shaking with KBr solution
Colour of each layer Upper
2. after shaking with hexane Lower
Conclusion
Colour after shaking with KCl solution
Colour of each layer
3. 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
Pasteur pipettes
Materials:
Potassium bromide solution, KBr
Potassium iodine solution, KI
Diluted nitric acid, HNO3
Hydrogen peroxide solution, H2O2

Test tube racks

Potassium chloride solution, KCl
Silver nitrate solution, AgNO3
Ammonia solution, NH3
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 cm3) diluted nitric acid. Cork the test–tubes and shake vigorously.
4. Add approximately 1 cm3 of silver nitrate solution followed by excess (e.g. 5cm3) 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
Red litmus paper

Droppers
Pipette

Materials:
1 M hydrochloric acid
4 M ammonia solution
2 M sodium hydroxide
Zinc nitrate
Copper (II) sulphate

Potassium chloride
Manganese (II) chloride
Magnesium chloride
Calcium chloride
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
+
2+
2+,
2+
Ag , Pb , Ca Ba , Cu2+
Zn2+, Mg2+, Cd2+

Yellow
Fe3+, Cr3+

Pink
Green
2+
2+
Co ,
Mn
Ni2+, Fe2+
(very pale)
(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
Boiling tubes
Water bath

Dropper
Centrifuge
Centrifuge tubes

Materials:
6 M ammonia solution
6 M hydrochloric acid
6 M sulphuric acid
1 M hydrogen peroxide
Universal indicator paper
Boiling chips

Silver nitrate, AgNO3
Barium nitrate, Ba(NO3)2
Iron (III) nitrate, Fe(NO3)3
Chromium (III) nitrate, Cr(NO3)3
Copper (II) nitrate, Cu(NO3)2
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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Unknown metal ions
(Ag+, Ba2+, Fe3+, Cr3+, Cu2+)
Add 4 drops of 6 M HCl and mix thoroughly
YES

NO

Precipitate
Insoluble Chloride

Solution
Add 1cm3 of ethanol, 2 drops of 6 M
H2SO4 and stand for 5 minutes.

Group: Ag+

Group: Ba2+, Fe3+, Cr3+, Cu2+
YES

Precipitate
Insoluble sulphate
Group: Ba2+

NO
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+
YES

NO

Precipitate
Insoluble hydroxide

Solution
Soluble amine

Subgroup: Fe3+, Cr3+

Subgroup: Cu2+

- 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

Some or all precipitate dissolves to give a yellow supernatant solution.
Cr present as H2CrO4
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Add excess of 6 M
NH3 (fume hood) to the original solution. Precipitate remains, no colour change.
Fe present

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Results:
Qualitative Tests for the Individual Ion Solutions – Observations
Ion Chloride test
(HCl)
A

Ethanol test

NH3 test

H2O2 test

Possible ion

B

C

D

E

Questions:
1. Write down all the balanced equations involved.
2. Describe briefly the reactions involved.

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