Oxygen and Sulphur

The Properties of Oxygen Gas

Oxygen is one of the most abundant elements on this planet. Our atmosphere is 21% free elemental

oxygen. Oxygen is also extensively combined in compounds in the earths crust, such as water (89%)

and in mineral oxides. Even the human body is 65% oxygen by mass.

Free elemental oxygen occurs naturally as a gas in the form of diatomic molecules, O2 (g).

 Oxygen

exhibits many unique physical and chemical properties. For example, oxygen is a colorless and odorless

gas, with a density greater than that of air, and a very low solubility in water. In fact, the latter two

properties greatly facilitate the collection of oxygen in this lab. Among the unique chemical properties

of oxygen are its ability to support respiration in plants and animals, and its ability to support

combustion.

In this lab, oxygen will be generated as a product of the decomposition of hydrogen peroxide. A catalyst

is used to speed up the rate of the decomposition reaction, which would otherwise be too slow to use as

a source of oxygen. The catalyst does not get consumed by the reaction, and can be collected for re-use

once the reaction is complete. The particular catalyst used in this lab is manganese(IV) oxide.

Generating Oxygen Gas:

catalyst

2 H2O2 (aq) → 2 H2O (l) + O2 (g)

hydrogen peroxide water oxygen

The oxygen gas produced will be collected in bottles by a method known as the downward displacement

of water (see figure on page 3). Once collected, several tests will be performed in order to investigate

the role of oxygen in several combustion reactions.

A combustion reaction is commonly referred to as “burning”. During a combustion reaction, oxygen

reacts chemically with the substance being burned. Note that since our atmosphere is roughly 21%

oxygen, many substances readily burn in air. Both oxygen and the substance being burned (the

reactants) are consumed during the combustion reaction, while new substances (the products) and heat

energy are generated. Since heat is produced, this is an exothermic reaction.

Combustion Reactions:

Substance being burned + Oxygen → Products + Heat

The actual products of a combustion reaction depend on what substance is burned and how much

oxygen is present. In general, however, when a pure element burns in oxygen the product is called an

oxide. An oxide is a compound containing both the element and oxygen chemically combined together.

Some examples of element combustion are shown below. Several such reactions will be performed

using the oxygen gas collected in this lab.

Combustion of an Element:

Element + Oxygen → Oxide of Element + Heat

C (s) + O2 (g) → CO2 (g) + Heat

carbon oxygen carbon dioxide

2 Hg (l) + O2 (g) → 2 HgO (s) + Heat

mercury oxygen mercury(II) oxide

Procedure

Safety

First, be sure to exercise caution when using the hydrogen peroxide (H2O2) and the hydrochloric acid

(HCl) as they can cause chemical burns and skin irritation. If either of these chemicals comes into

contact with your skin, immediately rinse with water for a minimum of fifteen minutes and notify your

instructor. Second, do not look directly at the burning magnesium. In addition to being very bright, it

emits harmful UV radiation that could cause damage to the retina of your eyes.

Materials and Equipment

Materials: 9% Hydrogen peroxide solution, manganese(IV) oxide, wooden splints, candle, sulfur, steel

wool, magnesium ribbon, zinc metal and 6M hydrochloric acid

Equipment: 250-mL Erlenmeyer flask, five wide-mouth bottles, four glass ‘cover’ plates, pneumatic

trough, “stopper + thistle tube + tubing” apparatus*, utility clamp, stand, deflagration spoon, crucible

tongs, small beaker, medium beaker and a large test tube.

Part A: Generating and Collecting Oxygen Gas

1. Obtain the following equipment:

• A 250-mL Erlenmeyer flask (locker)

• The “two-hole stopper + thistle tube + glass tubing + rubber tubing” apparatus (stockroom)

• Five wide mouth ‘gas-collecting’ bottles (under sink)

• Four glass ‘cover’ plates (front desk)

• A pneumatic trough (under sink) filled with water to ½ inch above the metal shelf

2. Fill four of the five wide-mouth bottles to the brim with water (the fifth will be used later). Then

gently slide a glass plate over the mouth of each bottle. Make sure that there are no air bubbles at

the top of the glass plate.

3. While holding the glass plate with your fore and middle finger, gently invert a bottle and lower it

into the water in the pneumatic trough. Remove the glass plate when the mouth of the bottle is

below the water level in the pneumatic trough. Repeat this for all four bottles. Place the glass plates

aside on a paper towel, as they will be used later.

4. Place one gas-collecting bottle on the metal shelf. Make sure that the mouth of the bottle does not

come out of the water.

5. Now focus on your reaction vessel, the Erlenmeyer flask. Add a pea-sized amount of

manganese(IV) oxide (the catalyst) to the flask, followed by about 50-mL of tap water.

6. Finally, assemble all your equipment together as demonstrated by your instructor, or as shown in the

figure below. Make sure that

• the end of the thistle tube is completely covered with water at the bottom of the flask,

• the end of the glass tubing running from the Erlenmeyer flask is inserted under the opening in the

bottom of the metal shelf into the gas-collecting bottle (which is full of water),

• the Erlenmeyer flask is stabilized with a utility clamp.

7. Obtain about 30-mL of 9% aqueous hydrogen peroxide (H2O2) in your smallest beaker. Then

carefully add about 10-mL of this H2O2 through the thistle tube. The generation of oxygen gas

should begin immediately. If at any time the rate of the reaction in the Erlemeyer flask appears to

slow down, add another 10-mL portion of H2O2.

8. The oxygen produced will fill the inverted bottle by displacing the water in it. This is because

oxygen does not dissolve in water, due to its low solubility. When the first bottle is completely filled

with gas, place the second bottle on the metal rack in its place and allow it to fill in a like manner.

Repeat this for the third and fourth bottles.

9. As soon as each bottle is completely filled, remove it by placing a glass plate under the bottle’s

mouth while under water, then lifting the bottle and plate from the pneumatic trough. Place the

bottle on the lab bench mouth up and do not remove the glass plate. Since oxygen is denser than

air, it sinks to the bottom of the flask and will not readily leak out the top.

10. Using masking tape, label each bottle of gas in the order they are collected: Bottle #1, Bottle #2,

Bottle #3 and Bottle #4. Label the fifth unused empty bottle “Air Bottle”.

11. Once all four bottles are filled with oxygen, do not add any more H2O2 to the Erlenmeyer flask. Set

it aside and allow the reaction to go to completion. At the end of the lab, the chemicals remaining in

the reaction flask and any unused H2O2 must be disposed of in the labeled waste container in the

hood. In the meantime, proceed to Part B.

Part B: The Properties of Oxygen Gas

*Dispose of all chemicals used in these tests as indicated by your instructor*

Test 1: Combustion of wood

Light a wooden split, and then blow it out. While it is still glowing red, quickly insert the splint into

Bottle #1 (oxygen-filled). How many times can you repeat this? Record your observations. Now

re-light the same wooden split, and again blow it out. Place it in the empty bottle (air-filled) while it

is still glowing. Record your observations.

Test 2: Combustion of candle wax

Place a small candle on a glass plate and light it. Then uncover and carefully lower Bottle #2

(oxygen-filled) over the candle. Measure and record the number of seconds that the candle

continues to burn. Then re-light the candle lower the empty bottle (air-filled) over it. Again,

measure and record the number of seconds that the candle continues to burn. Also be sure to record

any other relevant observations.

Test 3: Combustion of sulfur

This test must be performed in the hood under instructor supervision. Take your Bottle #3

(oxygen-filled) and your empty bottle (air-filled) to the hood your instructor directs you to. Place a

small lump of sulfur in a deflagrating spoon (located in the hood). Light the Bunsen burner in the

hood, and heat the sulfur in the spoon. The sulfur will first melt, then burn with an almost invisible

blue flame. Insert the spoon with the burning sulfur in Bottle #3 and record your observations. Then

insert it in the empty bottle, and again record your observations. When finished, extinguish the

burning sulfur in the beaker of water provided in the hood.

Test 4: Combustion of iron

Pour about 20-mL of tap water into Bottle #4 (oxygen-filled) and replace the glass plate quickly.

Take a loose, frayed out 2-3 centimeter piece of steel wool and hold it in a Bunsen burner flame for a

very brief instant with your crucible tongs (it will glow red). Then immediately lower the steel wool

into Bottle #4. Record your observations. Repeat with the empty bottle (air-filled) and record your

observations.

Test 5: Combustion of hydrogen

This test must be performed in air only. (Note: The hydrogen burned in this test must be first

generated by a reaction between zinc and hydrochloric acid.) To a large test tube, add 1-2 pieces of

zinc metal followed by about 3-mL of hydrochloric acid. Rapid bubbling should begin immediately

as hydrogen gas is produced, and the bottom of the test tube will get quite hot. Place the test tube in

the medium beaker. After 60 seconds have elapsed, light a wooden splint. Do not blow it out. Hold

the burning splint to the mouth of test tube (where the hydrogen gas is being evolved) and record

your observations.

Test 6: Combustion of magnesium

This test is an instructor demonstration. It must be performed in air only. Hold a 1-inch piece of

magnesium metal in a Bunsen burner flame with your crucible tongs until it ignites (in air). Record

your observations, remembering not to look directly at the burning magnesium!

Liquid oxygen in an unsilvered flask. Oxygen in the liquid state is a pale blue substance.

The structure shown above is oxygen's cubic crystal element structure.

Physical Properties

          Oxygen exists in all three forms - liquid, solid, and gas. The liquid and solid forms are a pale blue colour. However, oxygen gas is colourless, odourless, and tasteless. The elemental structure is a cubic crystal shape.

          Oxygen changes from a gas to a liquid at a temperature of 182.96°C, and then can be solidified or frozen at a temperature of -218.4°C.

          Oxygen exists in all three allotropic forms. The three allotropic forms include normal oxygen, diatomic oxygen, or dioxygen; nascent, atomic, or monatomic oxygen; and ozone or triatomic oxygen. The three allotropes differ in several ways; such as, atoms and molecules. For example, the oxygen we're most familiar with in the atmosphere has two atoms in every molecule. Nascent oxygen only has one atom in every molecule, and the third allotrope (ozone) has three atoms in every molecule.

Sulphur – the element

Sulphur is a non-metallic element which has a very important role in the chemical industry. It is a yellow solid which is found in large quantities but in various forms throughout the world It is found in metal ores such as copper pyrites (CuFeS2) and zinc blende (ZnS) and in volcanic regions of the world. Natural gas and oil contain sulphur and its compounds, but the majority of this sulphur is removed as it would cause environmental problems. Sulphur obtained from these sources is known as 'recovered sulphur' and it is an important source of the element. It is also found as elemental sulphur in sulphur beds in Poland, Russia and the US (Louisiana). These sulphur beds are typically 200 m below the ground. Sulphur from these beds is extracted using the Frasch process, named after its inventor Hermann Frasch.

Uses of sulphur The vast majority of sulphur is used to produce perhaps the most important industrial chemical, sulphuric acid. Sulphur is also used to vulcanise rubber, a process which makes the rubber harder and increases its elasticity. Relatively small amounts are used in the manufacture of matches, fireworks and fungicides, as a sterilising agent and in medicines. Allotropes of sulphur Sulphur is one of the few non-metal elements which exist as allotropes

The main allotropes are called rhombic sulphur and monoclinic sulphur. Both of these solid forms of sulphur are made up of S8 molecules

The fact that there are two different allotropes of sulphur is due to the way in which these S8 molecules pack together. In rhombic sulphur the molecules are packed more closely than in the monoclinic form (Figure 16.4).

Although sulphur is insoluble in water, it will dissolve in an organic solvent such as methylbenzene. If a solution of sulphur in methylbenzene is heated and allowed to cool then crystals of monoclinic sulphur are produced. When the temperature of the solution falls below 96°C, rhombic sulphur crystals are produced. Rhombic sulphur is stable below 96°C and monoclinic sulphur is stable above 96°C. This temperature is called the transition temperature.

When solid sulphur is heated, it melts at 112°C and forms a runny (mobile) liquid. At this point the S8 molecules are moving freely around each other,

as the weak attractive forces between them have been overcome

However, if the sulphur is heated further the liquid becomes thicker (viscous). This is because the S8 rings have been broken by the energy given to the sulphur and they bond together, forming long chains of sulphur atoms which become tangled, making the liquid viscous (Figure 16.6). Continued heating, to 444°C, makes the liquid more mobile once again as the long chains are broken down into smaller ones which move around one another freely.

If this liquid is poured into a beaker of cold water, a substance called plastic sulphur is formed. This is an elastic, rubber-like substance. In plastic sulphur, the sulphur atoms remain bonded together in the form of chains, very similar to chains of carbon atoms in plastics such as polythene. After a few hours, however, the plastic sulphur loses its elasticity and once again becomes solid as the Ss molecular rings re-form.

Sulphur dioxide

Sulphur dioxide is a colourless gas produced when sulphur or substances containing sulphur, for example crude oil or natural gas, are burned in oxygen gas. It has a choking smell and is extremely posionous. The gas dissolves in water to produce an acidic solution of sulphurous acid.

sulphur dioxide + water  sulphurous acid

SO2(g) + H2O(l)  H2SO3(aq)

It is one of the major pollutant gases and is the gas principally responsible for acid rain. However, it does have some uses: as a bleaching agent, in fumigants and in the preservation of food by killing bacteria