Claudia Ch25

7/28/2019 Claudia Ch25

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Chapter 25 Oxygen and its compoundsOccurrence:

Uncombined oxygen exists in the air, forming 21% by volume (or 23% by mass). Nearly half

of the mass of the Earths crust consists of oxygen in a combined state in the form of water,

silicates, many metallic and non-metallic oxides, and in the form of salts.Four methods to collect gas.

1) Collect over water (gas must not very soluble in water)

e.g. neutral oxide, oxygen, carbon dioxide.

2) Upward delivery (for gas with lower density than air)

3) Downward delivery (when gas has higher density than air)

4) Gas syringe (any gas)

can measure the rate of the reaction with volume of gas formed

Dry the gas with anhydrous Calcium Chloride

Experiment 84 (Pg. 295)

Laboratory preparation of oxygen from potassium chlorate

Since Oxygen has about the same density as air, it cannot be collected by displacement

of air. If required dry it can be dried by anhydrous Calcium Chloride and collected in

a syringe.

2KClO3(s) 2KCl(s) + 3O2(g)

2KClO3-(s)2Cl-(s) + 3O2(g)

Test for oxygen

It rekindles a glowing splint of wood. This distinguishes it from all gases exceptdinitrogen oxide, N2O. It is distinguished from this gas:


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a) by having no smell (dinitrogen oxide has a sweet, sickly smell)b) with nitrogen monoxide, oxygen produces brown fumes of nitrogen dioxide.

2NO(g) + O2(g) 2NO2(g) (nitrogen dioxide)


Laboratory preparation of oxygen from hydrogen peroxide

This is a convenient preparation because it requires no heat. Hydrogen peroxide

solution (20 vol.) is added, drop by drop, to manganese(IV) oxide, which catalyses

decomposition of the peroxide, Oxygen is collected over water.

2H2O2(aq) 2H2O(l) + O2(g)

An alternative preparation (in the same apparatus) is the drop by drop addition of

hydrogen peroxide solution (20 vol.) to potassium manganate(VII) in the presence of

excess of dilute sulphuric acid. Oxygen is liberated until all manganate(VII) is

decomposed, by which time the mixture is colourless.

5H2O2(aq) + 2KMnO4(aq) + 3H2SO4(aq)

K2SO4(aq) + 2MnSO4(aq) + 8H2O(l) +5O2(g)

INDUSTRIAL PREPARATION OF OXYGEN:

Since Oxygen exists to such a large extent in air, it is natural for attempts to bemade obtain it from this source. It is not easy to do this since nitrogen is an unreactive


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element and does not readily combine with anything to leave the oxygen pure. By far

the best process for obtaining oxygen industrially is from liquid air.

Liquid air. Air is first compressed to about 200 atmospheres pressure, cooled, and

allowed to escape from a small jet. Expansion cools the air further because heat

energy is used in separating the molecules. The cooled air is allowed to flow away bypassing round tubes containing the incoming compressed air. This cools the incoming

air and these successive coolings are finally sufficient to liquefy the air. On

evaporating of the liquid, nitrogen (b.p. 77K) is first evolved, leaving a liquid very

rich in oxygen (b.p. 90K at 760 mmHg). This is a fractional distillation of the liquid

air. Oxygen is sold either as liquid or as a gas compressed in strong steel cylinders at

about 150 atmospheres pressure.

USES:

(1) As an aid to breathing where the natural supply of oxygen is insufficient, for

example, in high-altitude flying or climbing, and also when anaesthetics are

administered to a patient.

(2) In the oxyacetylene (oxygen-ethyne) flame, which can be used for welding and

for cutting even very thick steel plate. The temperature of the flame reaches about

2200

(3) In the L-D process for making steel. A great and increasing oxygen tonnage is

now used in this way.

PROPERTIES OF OXYGEN:Oxygen is a colourless, ordourless, neutral gas, is only slightly soluble in water,

and has approximately the same density as that of air. It is an exceptionally active

element, combining vigorously with many metals and non-metals, forming basic and

acidic oxides respectively:

Metals + Oxygen metallic oxides

most of which are basic in character

non-metals + Oxygen non-metallic oxides

most of which are acidic in character

Action with metals

The manner in which oxygen reacts with metal is summarized in the list below:

K Na CaMg Al Zn Fe Pb Cu

When heated in air these metals oxidize

with a readiness indicated by the order

shown; that is, potassium most easily, copper

least readily

HgAg Au These metals do not oxidize easily; their

oxides yield oxygen on heating


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The following experiment illustrates the difference in reactivity between a metal high

in the list (magnesium) and a metal lower in the series (iron).


Action of oxygen with metalsMagnesium. Lower a piece of burning magnesium ribbon by means of tongs into

a gas-jar of oxygen. It burns with a more dazzling flame and forms a white ash,

magnesium oxide.

2Mg(s) + O2(g) 2MgO2(s)

Or 2Mg(s) + O2(g) 2(Mg2+O2-)(s)

Oxygen is acting as an oxidizing agent by accepting electrons (from the metal)

Iron. Attach a piece of iron wire to the end of a deflagrating spoon and dip the

end of the wire in sulphur (to start the action). Warm the wire in the Bunsen flame

until the sulphur begins to burn and then plunge it quickly into a gas-jar of oxygen

which contains a little water. The iron wire burns, giving off a shower of sparks, and

finally a molten bead of oxide drops into the water.

3Fe(s) + 2O2(g) Fe3O4(s) triiron tetraoxide


Action of phosphorus with oxygen

Piece a small piece of yellow phosphorus in a deflagrating spoon, warm it until it

begins to burn, and then plunge it into a gas-jar of oxygen into which you have

previously poured a little blue litmus solution. The phosphorus burns with a dazzling

flame, emitting white clouds of oxides of phosphorus which dissolve in the water toform acids of phosphorus, which turn the litmus red.

P4(s) + 5O2(g) P4O10(s) Phosphorus(V) oxide

On solution in water:

6H2O(l) + P4O10(s) 4H3PO4(aq) Phosphoric acid


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Action of sulphur with oxygen

Sulphur. In a similar manner, lower a piece of burning sulphur into a gas-jar ofoxygen containing blue litmus solution. Misty fumes of sulphur dioxide are given off

as the sulphur burns more brightly with its characteristic blue flame, and this gas

dissolves in the water to form sulphurous acid, which turns the litmus red.

S(s) + O2(g) SO2(g) Sulphur dioxide

SO2(g) + H2O(l) H2SO3(aq) sulphurous acid


Action of Carbon with oxygen

Perform the same experiment with wood charcoal (carbon). The charcoal burns

and emits a shower of sparks, combining vigorously with the oxygen to form a

colourless gas, carbon dioxide,, which dissolves in the water to form cabonic acid.

This is only a very weak acid and the litmus is turned claret-coloured, but not

definitely red.

C(s) + O2(g) CO2(g) Carbon dioxide

CO2(g) + H2O(l) H2CO3(aq) Cabonic acid

If the above experiment is performed with calcium hydroxide solution in the place of

litmus solution, the calcium hydroxide solution will become milky because of the

formation of a precipitate of calcium carbonate..Ca(OH)2(aq) + CO2(g) CaCO3(s) + H2O(l)

CLASSIFICATION OF OXIDES

A basic oxide is a metallic oxide which reacts with an acid to produce a salt and

water only; if solution in water, it forms an alkaline solution, e.g.:

CaO(s) + 2HCl(aq) CaCl2(aq) + H2O(l);

CaO(s) + H2O(l) Ca(OH)2(aq)

Other examples: Na2O, K2O (alkalis NaOH, KOH); CuO, MgO.

An acidic oxide is a non-metallic oxide which, when combined with the elements

of water, produces an acid, e.g.:

SO3(g) + H2O(l) H2SO4(aq): P4O10(s) + 6H2O(l) 4H3PO4(aq)

Other examples: CO2, SO2, SiO2 (acids H2CO3, H2SO3, H2SiO3).

An amphoteric oxide is a metallic oxide which can show both basic and acidic

properties, i.e. can react with both acid an alkali to produce a salt and water only.

Basic: ZnO(s) + H2SO4(aq) ZnSO4(aq) + H2O(l)

Al2O3(s) + 6HCl(aq) 2ALCl3(aq) + 3H2O(l)

Acidic: ZnO(s) + 2NaOH(aq) + H2O(l) Na2Zn(OH)4(aq)

Al2O3(s) + 2NaOH(aq) + 3H2O(l) 2NaAl(OH)4(aq)

The salts from the acidic reactions aresodium zincate andsodium aluminate.


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A neutral oxide is an oxide which shows neither basic nor acidic character,

e.g. dinitrogen oxide, carbon monoxide.

THE BASIC OXIDES

KNa Oxides are soluble in water forming alkalis

Ca Oxides of these metals are not reduced to

Mg metal by hydrogen

Al

Zn Oxides can be made from the metal

Fe by the action of nitric acid and then heat.

Pb

Cu

Hg

Ag Oxides of these metals decompose when heated

Au

Aluminium oxide Al2O3

This is a white solid. It is most conveniently prepared by first adding dilute

ammonia solution to a solution of an aluminium salt. This precipitates aluminium

hydroxide.

Al2(SO4)3(aq) + 6NH4OH(aq) 2Al(OH)3(s) + 3(NH4)2SO4(aq)The precipitate is then filtered, washed, dried, and heated.

2Al(OH)3(s)Al2O3(s) + 3H2O(l)

If prepared at the lowest temperature possible, it shows both basic and acidic

properties.

If strongly heated, it passes into a form which is insoluble in both acid and alkali.

Uses: The most important form of this oxide is bauxite Al2O3, 2H2O, from which the

metal is extracted. It also occurs in an impure form as emery and is used as an

abrasive.

Coloured by the presence of impurities, this oxide occurs as the gems, ruby(iron

and titanium),sapphire(chromium), and amethyst(manganese).

SUMMARY of preparation of the normal oxides of some common heavy metals

from the metals or their soluble salts.

Metal(Pb, Cu, Mg, Zn)

Dilute nitrate acid

Nitrate of metalIn solution


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Solid nitrate addNaOH solution add Na2CO3 solution

Heat hydroxide of metal carbonate of

Strongly as precipitate; metals as precipitate;Filter; wash; dry filter; wash; dry

Heat heat

Oxide of metal oxide of metal oxide of metal

(+ nitrogen (and water) (and carbon dioxide)

dioxide and

oxygen)

Calcium oxide (quicklime) CaO

Calcium oxide is made in industry by the action of strong heat upon limestones,

calcium carbonate, the latter being placed in a kiln.

CaCO3(s) CaO(s) + CO2(g)

Very large quantities of calcium oxide are made in this way. In the laboratory,

calcium oxide can be made by placing a piece of marble in a crucible and heating itstrongly in a gas-heated muffle furnace. A high temperature(1125)is required and

an hour or so is necessary to complete the action.


Preparation of calcium oxide

Make a loop in a stout iron wire large enough to hold a piece of marble in the

wire and so arrange it on a tripod that when a Bunsen burner is placed underneath it,

the marble is just above the inner cone of the roaring Bunsen flame. Leave in this

position for 5-10 minutes and then allow to cool until the solid can be comfortably

held in the fingers.

The original substance(calcium carbonate ) and the final product (calcium oxide)

are very similar in appearance, both being white solids. The difference between them

can be readily shown.

a) By the reaction with water

i) On calcium carbonate-no reaction

ii)On calcium oxide. Add water a drop at a time to the piece of calcium oxide in a

dish. Great heat is developed (there may be hissing as the water drops on the calcium

oxide), steam is formed and the oxide cracks and puffs up and finally crumbles to a

powder about three times as bulky. This is (calcium hydroxide) slaked lime.CaO(s) + H2O(l) Ca(OH)2(s)


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Allow more water to fall on to the calcium hydroxide until there is no further action.

If desired, at this stage, the mixture of calcium hydroxide and solution can be filtered

and the filtrate shown to be calcium hydroxide by expelling air from the

lungs(containing carbon dioxide) through a glass tube into the solution.

c) By the action of dilute hydrochloric acidi) On calcium carbonate. Effervescence is seen and the marble finally

disappears. Carbon dioxide is evolved, which, if passed into calcium

hydroxide solution, turns the latter milky.

CaCO3(s) + 2HCl(aq) CaCl2(aq) + H2O(l) + CO2(g)

ii) On calcium oxide. No evolution of carbon dioxide. The calcium oxide will

first give a similar action to a) ii) [slaking] but will give finally a colourless

solution of calcium chloride.

CaO(s) + H2O(l) Ca(OH)2(aq)

Ca(OH)2(aq) + 2HCl(aq) CaCl2(aq) + 2HCl(l)

Properties of calcium oxide:

1) white solid

2) very refractory

3) becomes incandescent and gives out a powerful light when heated to a very high

temperature

4) reacts vigorously with water to form slake lime

5) hygroscopic


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Zinc oxide ZnO

This compound is a white powder (yellow when hot) made in industry by distillating

zinc and burning the vapour at a jet.

2Zn(s) + O2(g) 2ZnO(s)

It is made in the laboratory from zinc by dissolving the metal in dilute nitric acid,

evaporating the zinc nitrate solution so formed to dryness, and heating the residue

strongly.

3Zn(s) + 8HNO3(aq) 3Zn(NO3)2(aq) + 4H2O(l) + 2NO(g)

2Zn(NO3)2(s) 2ZnO(s) + 4NO2(g) + O2(g)

1) amphoteric

2) cannot be converted into the metal by heating the oxides in a stream of hydrogen

3) used as a white pigment

Iron(III) oxide Fe2O3


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This compound is a red powder known as jewellers rouge. It is used for polishing

precious stones, and as a pigment. It is in the impure state as haematite. Iron(III)

oxide is made in the laboratory.

By heating iron(III) oxide is also the product formed if iron(II) hydroxide is heated

strongly in the air. All iron(II) compounds tend to become oxidized to iron(III)compounds by the oxygen of the atmosphere.

It has the usual properties of an oxide. It can be reduced to metallic iron by being

heated in a stream of hydrogen or carbon monoxide.

Fe2O3(s) + 3H2(g) 2Fe(s) + 3H2O(g)

Fe2O3(s) + 3CO(g) 2Fe(s) + 3CO2(g)

Triiron tetraoxide (magnetic oxide of iron) Fe3O4

This compound may be prepared by passing steam over red-hot iron or by

burning iron in oxygen.

It occurs naturally as magnetite and as such is a natural magnet or lodestone.

On heating it in a stream of hydrogen it is reduced to iron

Fe3O4(s) + 4H2(g) 3Fe(s) + 4H2O(g)

Lead(II) oxide PbO

The powder is orange when hot, yellow when cold. It can be made in the laboratory by

heating lead(II) nitrate, lead(II) carbonate, or lead(II) hydroxide.

Pb(NO3)2 PbO + NO2

PbCO3 PbO + CO2

Pb(OH)2 PbO + H2OAlthough lead(II) oxide can be considered a typical base, the only common acid

in which it will readily dissolve is nitric acid. The reason why it does not react

quantitatively with the others is a purely mechanical one.

Reaction of lead(II) oxide with dilute sulphuric or hydrochloric acids. Lead(II)

chloride and lead(II) sulphate are not formed quantitatively. These two substances are

almost insoluble in water and, as the action can proceed only on the outside of a

particle of oxide, the lead(II) chloride or sulphate forms as a layer on the outside. This

layer of chloride or sulphate is not permeable to the acids and hence reaction stops

before any appreciable amount of the salt has been formed(the inner core of PbO

cannot react with acid).

Lead(II) oxide is easily reduced to grey metallic lead by heating it in a stream of

hydrogen, town-gas, or carbon monoxide.

PbO(s) + H2(g) Pb(s) + H2O(g)

PbO(s) + CO(g) Pb(s) + CO2(g)

Lead(II) oxide is also an amphoteric oxide dissolving in caustic alkalis to form

plumbites.

NaOH(aq) + PbO(s) + H2O(l) NaPb(OH)3(aq) Sodium plumbite

Copper(II) oxide (black copper oxide) CuOCuO is hygroscopic, absorbing moisture from the air. It is a basic oxide and


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dissolves readily in warm dilute mineral acid, forming copper(II) salts, e.g.:

CuO(s) + 2HCl(aq) CuCl2(aq) + H2O(l)

Mercury(II) oxide HgO

This red oxide yields, when heated, a mirror of mercury on the cooler sides of thetest-tube, with oxygen evolved.

2HgO(s) 2Hg(s) + O2(g)

THE HIGHER OXIDES The higher oxides are oxides which contain more oxygen

per molecule than the corresponding basic oxide. E.g. H2O2, O3, HOCl, Na2O2, K2O2

Properties of ozone O3:

Ozone is a gas at ordinary temperature and pressure. It is extremely poisonous,

and rooms containing operating electrical machinery must be well ventilated. It is

obtained pure by liquefaction and is then a dark blue, explosive liquid, boiling at

about -112 under ordinary pressure.Ozone as an oxidizing agent:

Ozone is a vigorous oxidizer. It oxidizes lead(II) sulphide to lead(II) sulphate

PbS(s) + 4O3(g) PbSO4(s) + 4O2(g)

Hydrogen sulphide to sulphuric acid.

H2S(g) + 4O3(g) H2SO4(aq) + 4O2(g)

It also liberates iodine from potassium iodide in acidic solution2KI(aq) + H2SO4(aq) + O3(g) I2 + O2(g) + K2SO4(aq) + H2O(l)

Allotropy of oxygen and ozone

Oxygen and ozone are allotropes of the same element; the difference between

them is one of molecular complexity, oxygen having a diatomic molecule, O2, and

ozone a triatomic molecule, O3. Their chemical identity is shown by the fact that each

can be converted to the other without change of mass.

Oxygen O2 Ozone O3

Gas at s.t.p.

Density 16

Insoluble in turpentine

Heat has no action

No effect on mercury at room temperature

No effect on rubber

Has no effect on potassium iodide solution

Oxidizing agent

Gas at s.t.p.

Density 24

Absorbed by turpentine

Heat converts ozone into oxygen

2O3(g) 3O2(g)

Makes mercury wet glass

Attacks rubber

Liberates iodine from potassium iodide solution

Vigorous oxidizing agent

Hydrides of Oxygen


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Oxygen may be considered to form two principal hydrides. The commoner

hydride is water H2O. Oxygen also forms another hydride, hydrogen peroxide H2O2.


Oxidation of lead(II) sulphide by passing hydrogen sulphide into a solution oflead(II) nitrate in a boiling-tube. Allow the precipitate to settle, pour off the liquid,

and add to the black lead(II) sulphide some hydrogen peroxide solution. Leave it to

stand for some time, shaking occasionally. The precipitate gradually turns white,

because it is slowly converted to lead(II) sulphate.

PbS(s) + 4H2O2(aq) PbSO4(s) + 4H2O(l)

Experiment 96

Oxidation of acidified potassium iodide solution by hydrogen peroxide

Acidify a solution of potassium iodide with dilute sulphuric acid. Add hydrogen

peroxide. A brown coloration is caused by the production of free iodine.

2Kl(aq) + H2SO4(aq) + H2O2(aq) K2SO4(aq) + I2(aq) + 2H2O(l)

Hydrogen peroxide also oxidizes an iron(II) salt, e.g. FeSO4, in acidic solution to

the iron(III) state, the solution turning from green to yellow.

2Fe2+(aq) + H2O2(aq) + 2H+(aq) 2Fe3+(aq) + 2H2O(l)

Its powerful oxidizing action makes hydrogen peroxide useful as a bleaching

agent. It bleaches hair to a blonde colour, and is used commercially for the bleaching

of paper pulp, cotton, and other natural fibres.

Hydrogen peroxide as a reducing agent:H2O2(aq) 2H

+(aq) + O2(g) + 2e-

Silver oxide is reduced to metallic silver(a black precipitate)

2MnO4-(aq) + 5H2O(aq) + 6H

+(aq) 2Mn2+(aq) + 8H2O(l) + 5O2(g)

Decomposition of hydrogen peroxide

Hydrogen peroxide is decomposed catalytically by many substances, e.g.

Manganese(IV) oxide, finely powdered gold, and platinum.

2H2O2(aq) 2H2O(l) + O2(g)

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