AI English Translation, Pages 676-700
Page 2
c) Gases: Are those substances which have
variable shape and volume, their molecules are
far apart from each other, and have less
attraction among them than in the case of liquids.
8. Properties of Matter
a) Physical Properties:-
Are those properties that concern the outside
appearance of the substance and have no relation
to its composition.
b) Chemical properties:-
Are those properties that concern the actual
composition and behaviour of the substance.
9. Changes of matter:-
a) Physical changes
Are temporary changes that occur on the appearance
of matter such as a change of state without any
alteration of mass e.g. 1- Melting of Wax.
2- Heating a piece of iron until it becomes red,
then white. 3- Evaporation of water.
b) Chemical changes
Are permanent changes that occur on the actual
composition of the substance with the loss of
original properties e.g.
1- The burning of wood. 2- The Rusting of iron etc.
10. Constitution of Matter
All substances are composed of very small particles
invisible to the naked eye and the microscope which
are called "molecules". These molecules are in turn
composed of smaller particles called " atoms ".
11. A Molecule:
Is the smallest part of the substance that has all
the properties of that substance.
12. An Atom:
Is the smallest part of the substance that can
enter into a chemical reaction.
Page 3
13. Types of Matter:-
a) An Element: Is that substance which has <del>so far not been</del>
subdivided into simpler substances. The atoms of an
element are all similar. e.g. Silver, mercury, oxygen
and hydrogen. The molecules of an element are composed
of similar atoms.
b) A Compound:-
Is that substance formed by the union of two or more
elements in definite proportions by weight so that it
differs in its properties from its constituents. The
atoms of a compound are not similar. e.g. water is a
compound formed by the union of hydrogen and oxygen in
the ratio of 1:8 by weight.
c) A Mixture:-
Is that substance made by the addition of two or more
substances in any proportion, so that all the constit-
uents retain their original properties. e.g. Air is a
mixture of many gases.
14. Differences between Mixtures and Compounds
Compound | Mixture
1. The properties of a compound are entirely different from those of its constituents. | 1. A mixture possesses the common properties of its constituents.
2. The constituents of a com- pound are combined in definite proportions by weight. | 2. The proportions of the constituents of a mixture are variable.
3. The constituents of a com- pound cannot be separated, but by difficult chemical means. e.g. electrolysis, reduction, etc.. | 3. The constituents of a mixture can be separated easily by mechanical or physical processes e.g. by a magnet or by distillation, etc..
4. A compound is a homogeneous substance. | 4. A mixture is usually a heterogeneous substance.
Page 4
Kinds of Elements
15. There are ninety six elements in nature. They are divided into two
kinds according to their properties.
1. Metals
2. Non-Metals
16. Some important Metals and their Symbols
Name of Metal | Symbol | Name of Metal | Symbol
Potassium | K | Nickel | Ni
Sodium | Na | Tin | Sn
Barium | Ba | Lead | Pb
Magnesium | Mg | |
Aluminium | Al | Arsenic | As
Chromium | Cr | Copper | Cu
Manganese | Mn | Mercury | Hg
Zinc | Zn | Silver | Ag
Iron | Fe | Gold | Au
Cobalt | Co | Platinium | Pt
17. Some important Non-Metals and their Symbols
Oxygen | O2 | Phosphorus | P
Chlorine | Cl2 | Nitrogen | N2
Bromine | Br2 | Carbon | C
Iodine | I2 | Silicon | Si
Sulphur | S | Helium | He
| | Hydrogen | H2
Note: 1. All symbols begin with capital letters, the second letter
is always small.
2. All gases are made up of 2 atoms except the inert gases,
they are of one atom and ozone is of three atoms.
18. Comparison Between Metals & Non-Metals
Metals | Non-Metals
1. They are bright and reflect light when freshly cut. | 1. They are dull, and do not reflect light when freshly cut.
2. They are all solids except mercury. | 2. They are solids, liquids or gases e.g. S, Br2, N2
3. They are malleable & ductile. | 3. They are brittle.
4. They are good conductors of heat & electricity. | 4. They are bad conductors of heat & electricity.
5. They form basic oxides. | 5. They form acidic oxides.
6. Generally they dissolve in mineral acids and give off hydrogen. | 6. Usually they do not dissolve in mineral acids.
7. They form amalgams & alloys. | 7. They do not form amalgams & alloys.
8. Their Specific gravities are generally high. | 8. Their Specific gravities are usually low.
Page 5
19. A Solution: Is a homogeneous mixture formed by the disappearance
of the molecules of a substance called 'Solute' among the molecules
of another substance called "Solvent".
For example, the behaviour of some solids such as sugar or salt when
brought into contact with water is a familiar fact of experience,
The solid seems to disappear into the liquid, and we say that it has
"Dissolved", while the resulting homogeneous mixture is termed a
"Solution". The dissolved substance is called the "Solute" and the
liquid in which it is dissolved, the "Solvent". If we dissolve sugar
(solute) in water (solvent) the sugar will be divided into molecules.
These molecules will disappear in the spaces found among the molecules
of water and the resulting mixture is Solution of sugar in water.
20. Saturation: Is that limit beyond which apparently we cannot dissolve
any more of the solute in a certain quantity of the solvent at a
definite temperature.
21. Solubility: Is the weight (number of grams) of solute which is
sufficient to saturate 100 grams of the solvent at a certain
temperature. e.g., the solubility of sodium chloride (Na Cl) in
water at 0 ºC is 35.6 grams, i.e. 35.6 grams of sodium chloride will
saturate 100 gms. of water at 0 ºC.
22. The general bases of Solubility:
1. The solubility of solids in liquids increases with the increase
of temperature.
2. The solubility of solids in liquids does not increase with the
increase of agitation and stirring.
3. The Mass of gases dissolved in liquids increases with the
increase of pressure.
4. The solubility of gases in liquids decreases with the increase
of temperature.
5. The boiling point (B.P.) of a liquid rises and its freezing
point becomes lower when a solid substance dissolves in it.
For examples, A solution of sugar and water boils at a
temperature higher than 100 ºC and it freezes at a temperature
lower than 0 ºC.
23. Kinds of Solutions:
1. A true solution: Is that solution which is homogeneous and
transparent, in which the solid particles (solute) are very
small and pass through filter paper, and do not separate
under gravity. e.g. solution of sugar in water.
2. A Colloidal solution: Is that solution which is semi-homogeneous
and semi-transparent, in which the solid particles are comparatively
large in size, do not pass easily through filter paper, and do not
settle under gravity, e.g. milk, gum, blood.
CLASSIFICATION OF SOLUTIONS: Solutions are classified according
to the proportion of solute and solvent which they contain as
follows:-
(a) A Dilute Solution:- Is that solution which contains a small
amount of solute in proportion to the solvent.
which
(b) A Concentrated Solution:- Is that solution/contains a large
amount of solute in proportion to the solvent.
(c) An Unsaturated Solution:- Is that solution which contains a
quantity of solute less than that which is sufficicnt to
saturate it at a certain temperature and pressure. Such a
solution can dissolve more of the solute under the same
conditions of temperature and pressure. For example, if a
pinch of sugar is dissolved in a glass of water, the
<del>solution is unsaturated because it can readily dissolve</del>
more sugar.
(d) A Saturated Solution:- Is that solution which contains all
the solute it can normally dissolve at a given temperature
and pressure. For example, if several teaspoonfuls of sugar
are stirred in a glass of water, some of the sugar will
dissolve while the rest will drop to the bottom of the
glass. The clear solution is now said to be saturated at
the given temperature and pressure.
(e) A Supersaturated Solution:- Is that solution which contains
more of the solute than it can normally hold at the given
temperature and pressure. For example, if we dissolve 10 grams
of K Cl O3 in 100 grams of water we get a saturated solution
at 30°C. Upon raising the temperature of this solution to
50°C, we find that the 100 grams of water will now dissolve
20 grams of K Cl O3 then producing a saturated solution.
If the solution is carefully cooled back to 30°C, we observe
that the excess K Cl O3 does not crystallize out, but remains
dissolved, thus yielding a solution ⟦of⟧ K Cl O3 at 30°C which
has more solute in it than a saturated solution at that
temperature should normally have. Such a solution is said
to be "Supersaturated". This condition is unstable, for if
the solution is shaken, the 10 gms of K Cl O3 which are in
Page 7
excess will quickly crystallize out, leaving the original
saturated solution which contains 10 gms of K Cl O₃ dissolved
in 100 grams of water at 30°C.
25. DISTILLATION:- Is the process of converting a liquid into vapour by
heat, and then condensing the vapour back to liquid again, thus
leaving behind non-volatile impurities. The pure liquid thus
gotten is called the "Distillate". A suitable apparatus for this
purpose is shown in the diagram (Fig. 1).
⟦Diagram of distillation apparatus with labels and arrows⟧
26. FILTRATION:- Is the process of running a liquid through a porous
material, such as coarse paper, sand, etc... in order to remove
suspended solid matter. The clear liquid which runs through the
filter is called "Filtrate".
⟦Diagram of filtration apparatus with labels and arrows⟧
(Fig 2)
27. DECOMPOSITION:- Is the process of separating a compound into its
elements or into simpler substances. For example, heating
K Cl O₃ to yield K Cl and oxygen.
Δ
2 K Cl O₃ ⟦line⟧> 2 K Cl + 3 O₂
Potassium Chlorate Potassium Chloride + Oxygen
Page 8
28. Electrolysis: Is the process of decomposing a compound by means
of an electric current.
29. ANALYSIS:- Is the process of finding the composition of a substance
by decomposition into its elements or into simpler substances. For
example, water can be analysed by electrolysis yielding hydrogen
and oxygen.
Water Electric current Hydrogen + Oxygen.
⟦line⟧
Sulphuric acid
30. SYNTHESIS:- Is the combination of two or more substances, resulting
in the formation of a more complex substance. For example,
C + O₂ ⟦line⟧ C O₂
Carbon + Oxygen Carbon dioxide
31. REACTION: Any chemical change is called ⟦a⟧ reaction.
32. An Alloy: Is a substance formed by mixing together two or more
molten metals. For example, Brass is an alloy which is composed
of Copper and Zinc.
33. An AMALGAM: Is an alloy of mercury with some other metal. For
example, gold amalgam.
34. SUBLIMATION: Is the process of a solid vaporizing and then
condensing back to a solid without passing through the liquid
state. For example, sublimation of iodine.
35. PRECIPITATE: Is an insoluble substance resulting from a chemical
action between substances in solutions. For example,
Na Cl + Ag N O₃ ⟦line⟧ Ag Cl + Na N O₃
Sodium Chloride + Silver Nitrate ⟦line⟧ Silver Chloridge + Sodium
Nitrate.
x x x x x x x
36. CHEMICAL REACTIONS:
Substances often do not remain as they are, but they either
react with each other or decompose into simpler substances. Such
changes which alter the composition of a substance and form one
or more new substances, are called "Chemical Reactions". Chemistry
is a science which deals with such changes as distinguished from
other sciences.
Page 9
The Chemist has devised a highly practical system to represent elements and compounds by means of symbols and formulas indicating how substances are built up, and how they react with each other. By this means he can indicate what is present in a test tube, and what happens when something is added to it. To make this clear, the student has to be acquainted with the following:-
(a) Symbols: A chemical symbol consists of one or two letters of the English name or the latin name of an element denoting one atom of it. For example, the symbol for Oxygen is O, and that of Silver is Ag, (from Latin, Argentum).
(b) Valency or Valence: is the "Force" or "ability" of an element to combine with other elements to form a chemical compound. This force or ability is measured by the number of hydrogen atoms which will combine with one atom of the element. So the valency of an element is defined as the number of hydrogen atoms (or any monovalent element, e.g. Chlorine) with which one atom of the element can combine, or which it can displace. Accordingly, the valency of an element which does not combine with any other element is Zero, e.g. Helium. The valency of an element which combines with one atom of hydrogen is one. Such an element is said to be "Monovalent"; e.g. Chlorine, in (HCl), or hydrochloric acid. The valency of an element which combines with two atoms of hydrogen is 2. Such an element is "Divalent"; e.g. Oxygen in (H₂O) or water. The valency of an element which combines with three atoms of hydrogen is 3, and it is "Trivalent"; e.g. nitrogen in (NH₃) or ammonia and so on. Valency could reach 8 as a maximum. Some of the elements have constant valencies in all their compounds, while others have variable valencies according to the kind of compound they are present in, and to the condition of the reactions. The following table gives the symbols of the important elements and their valencies:-
Element | Symbol | Valence | Element | Symbol | Valence
Sodium | Na | 1 | Tin | Sn | 2,4
Potassium | K | 1 | Aluminium | Al | 3
Silver | Ag | 1 | Gold | Au | 3,1
Magnesium | Mg | 2 | Chlorine | Cl | 1
Calcium | Ca | 2 | Iodine | I | 1
Zinc | Zn | 2 | Oxygen | O | 2,4
Barium | Ba | 2 | Nitrogen | N | 2,3,5
Copper | Cu | 2,1 | Phosphorus | P | 3,5
Mercury | Hg | 2,1 | Carbon | C | 4,2
Iron | Fe | 2,3 | Silicon | Si | 4
Lead | Pb | 2,4 | Sulphur | S | 2,4,6
Page 22
4- Amphoteric oxides:- Are those oxides which possess the proper-
ties of both acidic and basic Oxides. They combine with both
acids and bases to form salts.
Al₂ O₃ + 6H Cl ⟦line⟧> 2 Al Cl₃ + 3H₂ O
Aluminium oxide hydrochloric Aluminium chloride
amphoteric oxide acid salt water
Al₂ O₃ + 2 Na OH ⟦line⟧> 2 Na Al O₂ + H₂ O
Aluminium oxide Sod. hydroxide Sod. Aluminate
base salt water
Examples:- Zinc oxide (Zn O), Aluminium oxide (Al₂O₃), Lead
Monoxide (Pb O), Stannous oxide (Sn O).
5- Peroxides:- These are usually the oxides of metals containing
a greater percentage of oxygen than their basic oxides. They give
off oxygen on heating.
2 Ba O₂ ———θ———> 2 Ba O + O₂ ↑
Barium peroxide Barium oxide oxygen
Examples:- Barium peroxide (Ba O₂), Sodium peroxide (Na₂ O₂),
Zinc peroxide (Zn O₂) & hydrogen peroxide (H₂ O₂).
A C I D S
59.
DEFINITION:- An acid is a hydrogen compound whose hydrogen may
be replaced by a metal and whose water solution changes the
colour of litmus from blue to red.
Many compounds such as sugar and alcohol, contain hydrogen, but
are not acids, because their hydrogen atoms cannot be replaced
by a metal. The hydrogen of the acid is always accompanied by
a non-metallic element such as Cl, S or by a radical such as (N O₃),
(SO₄) and (P O₄).
Page 23
60. PREPARATION OF ACIDS:- There are two common methods for the
preparation of acids:
1- By the action of acidic oxides (acid anhydride) with water e.g.
S O₃ + H₂ O ⟦line⟧> H₂ S O₄
C O₂ + H₂ O ⟦line⟧> H₂ C O₃
2- By the action of sulphuric acid with a salt of the required
acid e.g.
H₂ SO₄ + 2 Na Cl ⟦line⟧> 2 H Cl + Na₂ SO₄
61. PROPERTIES OF ACIDS:-
1- Acids taste sour.
2- They change litmus colour from blue to red.
3- All acids contain hydrogen which is replaceable by active
metals e.g.
Zn + 2H Cl ⟦line⟧> Zn Cl₂ + H₂
4- They react with marble (Ca CO₃) causing effervescence and
evolution of CO₂.
2H Cl + Ca CO₃ ⟦line⟧> Ca Cl₂ + C O₂ + H₂ O
5- Acids react with bases forming salt and water.
H Cl + Na O H ⟦line⟧> Na Cl + H₂ O
62. NAMING OF ACIDS:-
a- In case the acids do not contain oxygen i.e. when they are only
composed of hydrogen and a non-metal, they are named by using the
full name of the non-metal with the prefix hydro - and ending
with - ic. Examples hydrochloric acid ( H Cl), hydrosulphuric acid
H₂ S.
b- In case they contain enough oxygen, they are named by using the name
of non-metal ending with - ic. Examples, sulphuric acid ( H₂ So₄) ,
carbonic acid ( H₂CO₃) and nitric acid ( H N O₃).
c- When the acids contain one less atom of oxygen, they are named
by using the name of the non-metal ending with - ous. Examples,
sulphurous acid ( H₂ SO₃) and nitrous acid ( H N O₂)
⟦line⟧
Page 24
B A S E S
63. DEFINITION:- A base is a compound of a metallic element or a radical
with one or more hydroxyl ( O H) groups. The water solution of a base
change the colour of litmus from red to blue.
64. PREPARATION OF BASES:-
1- By the action of basic oxides with water:-
Ca O + H₂ O ⟦line⟧> Ca ( O H)₂
Mg O + H₂ O ⟦line⟧> Mg ( O H)₂
2- By the action between a base and a salt.
Na₂ CO₃ + Ca (O H)₂ ⟦line⟧> 2 Na O H + Ca CO₃ ↓
Fe Cl₃ + Na OH ⟦line⟧> 3 Na Cl + Fe (O H)₃ ↓
3- By the action of an active metal and water:
2 Na + H₂ O ⟦line⟧> 2 Na OH + H₂ ↑
65. PROPERTIES OF BASES.-
1- In general bases have a bitter taste.
2- Bases turn red litmus into blue.
3- Bases contain hydroxyl ( O H) group.
4- Bases do not react with marble.
5- Bases neutralize acids forming salt and water.
66. NAMING OF BASES:- Bases are named by placing the name of the metal
before the word " hydroxide ".
Examples: Sodium hydroxide ( Na OH), Calcium hydroxide (Ca (O H₂) ).
Strong soluble bases such as ( Na OH) and (K O H) are often called
alkalies.
Page 25
( S A L T S )
67 DEFINITION:- A "salt" is a compound consisting of a metal or a metallic
radical combined with a non-metal or an acid radical. In other words,
a salt is an acid the hydrogen of which is replaced by a metal or a
radical. The term salt, as commonly used, refers to sodium chloride,
but in chemistry salt is the general name of a class of compounds which
resemble Sod. Chloride.
68. PREPARATION OF SALTS:- Salts are prepared by a variety of methods.
1- Neutralization:- Neutralization is the reaction between an acid and
a base, forming salt and water. For example:
Na OH + H Cl ⟦line⟧> Na Cl + H_2 O
In making Na Cl by this method a small quantity of Na OH solution is
placed in a dish with a piece of blue litmus paper. Hydrochloric acid
is then added drop by drop until the litmus paper just turns red. If
this neutral solution is now evaporated to dryness, a white deposit
of (Na Cl) remains, and may be identified by its salty taste.
2- Action of an Acid on a Metal:-
Zn + H_2 SO_4 ⟦line⟧> Zn SO_4 + H_2
3- Action of an Acid on metallic oxide:
Ca O + 2 H Cl ⟦line⟧> Ca Cl_2 + H_2 O
4- Action of an Acid on the Salt of a More Volatile Acid
H_2 SO_4 + 2 Na Cl ⟦line⟧> Na_2 SO_4 + 2 H Cl
5- Double Replacement Resulting in the Formation of an Insoluble product:
Ba Cl_2 + Na_2 SO_4 ⟦line⟧> 2 Na Cl + Ba SO_4
6- Direct Union of the Elements:
Zn + S ⟦line⟧> Zn S
7- Union of a metallic oxide and a Non-Metallic Oxide
Ca O + Si O_2 ⟦line⟧> Ca Si O_3
Page 26
69. PROPERTIES OF SALTS: Salts vary widely in characteristics.
As a general rule, salts are white, crystalline solids with a salty
taste. They are usually soluble in water, and Neutral to litmus
paper. Among the many exceptions to the above, are copper sulphate,
which is blue; calcium carbonate (marble) which is insoluble in
water and sodium carbonate which in solution reacts basic to litmus,
owing to hydrolysis.
70. TYPES OF SALTS:
1- An "acidic salt" is one in which only part of the hydrogen of
the acid has been replaced by a metal. Thus Sodium acid carbonate,
or sodium bicarbonate (Na HCO₃), and sodium acid phosphate
(Na H₂ PO₄) are acid salts.
2- A "basic salt" is one which contains one or more hydroxyl (OH)
radicals. Thus bismuth subnitrate Bi (OH)₂ NO₃ is a basic salt used
in medicine.
3- A "Normal salt" is one which contains only a metal or a
metallic radical, combined with a non-metal or an acidic radical.
It contains neither hydrogen replaceable by a metal nor a hydroxyl
radical. Thus Ca CO₃ & NaCl are normal salts.
4- A double salt is one which contains two metals combined with one
acidic radical. Thus common alum is a double sulphate of potassium
and aluminum, having the formula K₂ SO₄ . Al₂ (SO₄)₃ . 24 H₂ O
or K Al (SO₄)₂ . 12 H₂ O .
71. NAMING SALTS: The name of a salt consists of two parts, the name
of the metal and the name derived from the acid. For example :
Acids of Chlorine | Formulas | Sodium Salts of the acids | Formulas
Hydrochloric acid | H Cl | Sod. Chloride | Na Cl
Chlorous acid | H Cl O₂ | Sod. Chlorite | Na Cl O₂
Chloric acid | H Cl O₃ | Sod. Chlorate | Na Cl O₃
Page 27
( HYDROGEN )
72. OCCURENCE Hydrogen is not as abundant as oxygen. It occurs both
free and in the combined state in nature. In the combined state it
occurs nearly <del>everywhere</del>, but in the free state, it occurs in natural
gas and in volcanic gases. Traces of hydrogen are found in the lower
strata of the earth's atmosphere and gradually its quantity increases
at high altitudes. Hydrogen is known to occur abundantly in the
sun's atmosphere.
In the combined state hydrogen constitutes about 1% of the earth crust
and one ninth of the weight of water. It occurs in all acids, in
animal and vegetable matter, such as butter, starch, sugar, in
hydrocarbons such as kerosene, gasoline, etc..
73. METHODS OF PREPARATION:--
1- LABORATORY METHODS.
a) by the action of certain metals on non oxidizing acids.
Metals such as zinc and iron react with certain dilute acids,
notably hydrochloric acid and sulphuric acid.
Zinc is put in a flask closed by a rubber stopper, through
which a thistle tube extends to the bottom of the flask, and a
delivery tube also extends to a trough containing water, over which
a large test tube filled with water is inverted. The dilute acid is
poured through the thistle tube until it covers its lower end.
Zinc replaces the hydrogen of the acid.
thistle tube
delivery tube
hydrogen gas
test tube
water
trough
HCl
Zn
Zn + 2H Cl ⟦line⟧> Zn Cl₂ + H₂
Fig.5
Page 28
and liberates this hydrogen as a gas, which passes through the
delivery tube and is collected by an upward displacement of water,
at the same time forming a new salt compound called zinc chloride.
Zn + 2 H Cl ⟦line⟧> Zn Cl₂ + H₂
Fe + H₂ SO₄ ⟦line⟧> Fe SO₄ + H₂
dil.
Ag + H Cl ⟦line⟧> no reaction
b- By the action of certain metals on water:-
Various active metals (above hydrogen) react with water under
certain conditions liberating hydrogen and forming either a hydro-
xide or an oxide of the metal as follows:-
1. With cold water: as potassium, sodium & calcium
2 Na + H₂ O ⟦line⟧> 2 Na OH + H₂ vigorously
sodium water sodium hydro-
xide
Ca + 2 H₂ O ⟦line⟧> Ca (OH)₂ + H₂ slowly
calcium + water calcium hydro-
xide
2. With hot water: as Magnesium, using Na OH as a catalyst.
Mg + 2 H₂ O ⟦line⟧> Mg (OH)₂ + H₂ slowly
Na OH
magnesium water Magnesium hydro-
xide
3. With steam of water: as red hot iron
3 Fe + 4 H₂ O ⟦line⟧> Fe₃ O₄ + 4 H2 Rapid
iron water magnetic oxide hydrogen
Metals below hydrogen do not react with water
Cu + H₂ O ⟦line⟧> no reaction
Note: The above reactions demonstrate the order of activity
of metals.
Page 29
( REPLACEMENT SERIES )
1 | Potassium | K ) | Highest
2 | Sodium | Na ) | activity
3 | Calcium | Ca ) | decreasing
4 | Magnesium | Mg ) | to
5 | Aluminum | Al ) | |
6 | Zinc | Zn ) | |
7 | Iron | Fe ) | |
8 | Nickle | Ni ) | |
9 | Tin | Sn ) | |
10 | Lead | Pb ) | V
11 | HYDROGEN | H ) | lowest
12 | Copper | Cu ) | activity
13 | Mercury | Mg ) |
14 | Silver | Ag ) |
15 | Platinum | Pt ) |
16 | Gold | Au ) |
Note: metals above hydrogen in the replacement series replace it from
dilute acids, while metals below hydrogen do not replace it
from acids.
II. COMMERCIAL METHODS :-
a- By electrolysis of water .- When an electric current is passed
through acidified water, hydrogen is liberated at the cathode.
b- By the action of Hot Coke on steam:-
When steam is passed over white - hot coke (carbon), a mixture of
hydrogen and carbon monoxide gas is formed which is called
" water gas ".
H2 O + C ⟦line⟧> CO + H2
When the mixture of both gases is cooled to a very low temperature,
the ⟦CO⟧ gas changes to a solid, leaving pure hydrogen. In other
cases, when the mixture of both hydrogen and carbon monoxide
gases is mixed with more steam and passed over a catalyst,
(divided Fe), CO will be changed into CO2
CO + H2 O + H2 ⟦line⟧> C O2 + 2 H2
Then C O2 + H2 are treated with water at a pressure of 30 atmos-
pheres when C O2 dissolves leaving H2
Page 30
74. PROPERTIES OF HYDROGEN:
a- Physical Properties:
1. It is a gas under ordinary conditions.
2. It is colourless, odourless & tasteless gas.
3. It is the lightest gas known, being 14.5 times lighter than air.
4. It is very slightly soluble in water.
5. It is absorbed in large volumes by the metals palladium
and platinum accompanied by the liberation of much heat.
b- Chemical Properties:
1. Hydrogen does not support combustion; it burns with a pale
flame. When it is mixed with air or oxygen it explodes at
kindling temperature.
2 H₂ + O₂ ⟦line⟧> 2 H₂ O
2. It is not active at ordinary temperature.
3. It combines directly with other elements at high temperatures
& sometimes it combines on mere exposure to light.
H₂ + Cl₂ ⟦line⟧> 2 H Cl
Hydrogen Chlorine hydrogen chloride
200 atm.
3 H₂ + N₂ ⟦line⟧> 2 N H₃
Fe + ⟦e⟧
Hydrogen nitrogen ammonia
H₂ + S ⟦line⟧> H₂ S
Hydrogen sulphur hydrogen sulphide
H₂ + Ca ⟦line⟧> Ca H₂
Hydrogen Calcium Calcium hydride
4. It is a powerful reducing agent. It removes oxygen from
so many of its compounds.
75. REDUCTION:- Is the process of removing oxygen from a
compound. A substance which can effect such a removal of
oxygen is called a " Reducing agent ".
Page 31
To prove that hydrogen is a reducing agent, we pass the dry
gas (H2) over heated copper oxide as in the figure below.
T
drying tube
containing CaCl2
H2
F
Cu O
drops of
water
HCl
Zn
Preparation of
Hydrogen
Fig. 6
The hydrogen combines with the oxygen present in the oxide,
forming water, the residue left behind in the tube is metallic
copper. Whenever reduction takes place, it is always accompa-
nied by oxidation, as can be observed by examining the equation
for this reaction:-
Cu O + H2 ⟦line⟧> Cu + H2 O
copper oxide hydrogen copper water
substance reduced reducing agent,
also oxidizing also substance
agent oxidized
76. TEST FOR HYDROGEN: Hydrogen may be identified by one chemical
test:- When hydrogen burns, the only product formed is water.
77. USES OF HYDROGEN:- Hydrogen is used:
1- To obtain a very high temperature, for welding metals.
2- For filling balloons and airships, but owing to its great
combustibility it is often substituted by helium (He).
3- For hydrogenation:- Many liquid oils and fats, such as cotton
seed oil & coconut oil have disagreeable odours & tastes & can
not in their simple state be used for cooking. They are con-
verted into palatable solid fats fit for cooking, by passing
hydrogen in the presence of powdered nickel as a catalyst.
The product is called VEGETALINE.
Page 32
4- For the preparation of ammonia by Habber process.
5- It is used as a reducing agent.
6- It is a valuable constituent of some fuel gases such as
water gas and producer gas.
⟦line⟧
- 33 -
W A T E R
78. OCCURRENCE. Water is present in nature, and it occupies about 3/4 of
the earth's surface. It is also present in many salts and animal and
vegetable bodies. For example, lean meat is about 60% water, while
tomatoes are about 95% water. The human body contains about 70%
water.
79. FORMATION. Being abundant in nature, water is not usually prepared
from other materials. However, it is a product of many chemical re-
actions some of which are:
1. Direct union of oxygen and hydrogen by an electric spark.
2 H₂ + O ⟦line⟧> 2 H₂O.
2. Oxidation of a compound of hydrogen.
CH₄ + 2O₂ ⟦line⟧> CO₂ + 2 H₂O
3. Reduction of an oxide by hydrogen.
CuO + H₂ ⟦line⟧> Cu + H₂O
4. Neutralization of acids and bases.
NaOH + HCl ⟦line⟧> Na Cl + H₂O
80. PROPERTIES.
a- Physical Properties:
1. Water is liquid at ordinary temperature. It is tasteless and
odourless. It has no colour when in thin layers, but it acquires
a bluish tinge when it is observed through great thickness.
2. Water freezes at 0°C (32°F), and boils at 100°C (212°F).
3. Its greatest density, namely 1, (i.e. 1cc weighs 1 gm), is reached
at 4 degrees centigrade.
4. It dissolves almost all substances to some extent, and may there-
fore be considered as a universal solvent. The pleasant taste of
drinking water is due to dissolved air and minerals.
5. Water when pure, does not conduct electricity.
b- Chemical Properties:
1. Water is extremely stable.
2. It reacts with various active metals such as Na, Ca, Fe, with
formation of hydrogen.
3. It is decomposed by an electric current into hydrogen and oxygen.
4. It combines with certain metallic oxides forming bases:
Ca O + H₂O ⟦line⟧> Ca (OH)₂
Calcium oxide + water Calcium hydroxide
- 34 -
5. It combines with certain non-metallic oxides forming acids.
C O₂ + H₂O ⟦line⟧> H₂C O₃
Carbon dioxide + water Carbonic acid
6. Water reacts with certain salts, forming both an acid and a
base. This process is called "hydrolysis".
Al₂(SO₄)₃ + 6 H₂O ⟦line⟧> 2 Al(OH)₃ + 3 H₂SO₄
Aluminium sulphate + water Aluminium + Sulphuric
hydroxide acid
7. Water combines with some compounds when they crystallize from
solution, forming hydrates, and this water is called water of
crystallization. For example:
Cu SO₄ + 5 H₂O ⟦line⟧> Cu SO₄.5H₂O
unhydrous copper sulphate Hydrated crystallized
plus water copper sulphate
8. Water acts as a catalyst. Many chemical reactions cannot
take place unless there is some moisture present, as for
example, rusting of iron.
81. WATER IS A COMPOUND. Water is a compound of oxygen and hydrogen.
in the ratio of 8:1 by weight and 1:2 by volume. One of the first
steps in the study of a chemical compound is to learn its composi-
tion, that is what elements it contains, and the exact amount of
each. There are two methods of obtaining this information.
1. Analysis. On passing an electric current in an acid water in
a voltametre (Hoffman apparatus), we get hydrogen at the Cathode
and oxygen at the anode. See pages 18 and 19.
2. Synthesis. On passing a spark through a mixture of definite
quantities of oxygen and hydrogen in an eudiometer, an explosion
occurs and drops of water form on the inner surface of the tube.
Spark wires connected to
gap induction coil
Mercury
Eudiometer
- 35 -
82. WATER CYCLE IN NATURE Water has the following cycle in nature:
1. Water evaporates from oceans and seas.
2. After evaporation it condenses and forms clouds.
3. Clouds change into rain.
4. Rain falls on the earth's surface dissolving some salts found
on its crust. Rain water then passes down through the soil
until it meets a solid layer where it collects. When sufficiently
large quantities of water are collected in this way, water is
again forced up to the earth's surface forming springs.
5. Waters from Springs and Brooks flow and meet together to form a
river. This river flows into the sea and thus the cycle is
repeated.
83. HARDNESS OF WATER. Hard water is that water which does not lather
with soap, because of the presence of SOLUBLE calcium or magnesium
salt in it which form an insoluble precipitate with soap.
Ca Cl₂ + 2 NaC₁₇H₃₅CO₂ ⟦line⟧→ Ca(C₁₇H₃₅CO₂)₂ + 2 NaCl
Sod. Stearate Calcium Stearate
soluble soap insoluble soap
Hardness is of two kinds:
a- Temporary Hardness: Is the hardness caused by the presence of
the bicarbonate of either Calcium or Magnesium (Ca(HCO₃)₂ or
Mg(HCO₃)₂). These are unstable salts, i.e. they easily
decompose on heating forming the insoluble carbonates. There-
fore this hardness can be got rid of by boiling the hard water,
as is represented by the following equations:
θ
Ca(HCO₃)₂ ⟦line⟧→ CaCO₃↓ + CO₂↑ + H₂O
Ca. bicarbonate Calcium Carbon Water
Carbonate dioxide
θ
Mg(HCO₃)₂ ⟦line⟧→ MgCO₃↓ + CO₂↑ + H₂O
Mg. bicarbonate Mg. carbonate Carbon Water
dioxide
We can also get rid of temporary hardness by adding lime water
that is calcium hydroxide (Ca(OH)₂), or washing soda, i.e., sodium
carbonate (Na₂CO₃) as follows:
Ca(HCO₃)₂ + Ca(OH)₂ ⟦line⟧→ 2 CaCO₃↓ + 2 H₂O
Ca. Bicarbonate Ca. Hydroxide Ca. Carbonate Water
Ca(HCO₃)₂ + Na₂CO₃ ⟦line⟧→ CaCO₃↓ + 2 NaHCO₃
Ca. Bicarbonate Sod. Carbonate Ca. Carbonate Sod. Bicarbonate
The precipitate of calcium carbonate or Magnesium carbonate can
be removed by filtration or decantation.
- 36 -
b- Permanent Hardness: Is that hardness due to the presence of
soluble salts of Mg and Ca in water. These salts do not decompose
on heating. They are: Calcium sulphate (CaSO₄), Calcium chloride
(CaCl₂), Magnesium sulphate (MgSO₄) and Magnesium Chloride (MgCl₂).
In other words, they are Sulphate and Chloride salts of Magnesium
and Calcium.
Permanent hardness is removed by adding washing soda (Na₂CO₃)
as follows:
CaSO₄ + Na₂CO₃ ⟦line⟧> Na₂SO₄ + CaCO₃
Calcium Sod. Sod. Cal.
sulphate Carbonate Sulphate Carbonate
MgCl₂ + Na₂CO₃ ⟦line⟧> 2 NaCl + Mg CO₃
Magnesium Sodium Sodium Magnesium
chloride Carbonate Chloride Carbonate
Sodium Chloride and Sodium sulphate remain soluble and have no
action on soap lather.
⟦line⟧
k.
- 37 -
THE HALOGENS
84 NOMEN⟦C⟧LATURE
The halogens refer to the four elements fluorine, chlorine, bromine,
and iodine. These four elements can combine with metals to form salts
such as common salt (NaCl) and hence the word halogen (meaning salt-
former) is given to this group of elements. These four elements are
grouped as one family because they resemble each other in properties
and in their chemical compounds. We shall choose the chief member of
this family, namely, Chlorine for our study as representative for the
others.
CHLORINE
85. OCCURRENCE: This element is not present in nature in the free state
due to its great chemical activity. Therefore it exists in nature in
the combined form. The common compound of it is sodium chloride (NaCl)
which is common salt. Other Chlorides exist in nature such as those of
K, Mg and other elements.
86. PREPARATION:
Laboratory Methods
By oxidation of hydrochloric acid with an oxidizing agent such
as manganes dioxide:
4HCl + MnO2 ⟦line⟧→ MnCl2 + 2H2O + Cl2
The apparatus commonly used for this preparation is illustrated
in the following diagram.
HCl ↘ A B
C D
MnO2
↓ Cl2
Water Sulphuric Acid
Fig. 7
- 38 -
Flask A contains MnO2 to which is added concentrated HCl through a
thistle tube.
The Chlorine gas which is librated in flask A is passed through the
connecting tube to bottle B. This bottle contains water to absorb any
gaseous (HCl) which might escape mixed with the librated chlorine.
Chlorine gas then passes through a 2nd connecting tube to bottle C which
contains concentrated H2SO4 (to absorb moisture). The dry chlorine gas is
lastly collected by the downward displacement of air. Chlorine gas, unlike
oxygen and hydrogen, cannot be collected over water because it is soluble
in it.
87. PROPERTIES:
(a) Physical properties:-
1- It is a greenish yellow gas.
2- It has a strong irritating odour and when inhaled it attacks
the delicate membranes of the nose; thorat & lungs.
3- It is extremely poisonous.
4- It is about 2 1/2 times as heavy as air.
5- It is somewhat soluble in water.
6- It is liquified easily by pr⟦ess⟧ure alone at room temperature.
(b) Chemical properties:-
1- It is extremely active.
2- It combines directly with many elements forming chlorides:
a) It combines with hydrogen:
Cl2 + H2 sun light -> 2HCl
b) It combines with metals such as heated Sodium:
Cl2 + 2Na -> 2Na Cl
c) It combines with phosphorous or sulphur
Cl2 + 2P ⟦line⟧ 2P Cl5 Phosphorus Penta chloride
Cl2 + 2S ⟦line⟧ S2 Cl4 Sulphur chloride
3- It reacts with hydrocarbons (compounds of carbon & hydrogen),
replacing part or all of the hydrogen
CH4 + 3Cl2 ⟦line⟧ Ch Cl3 + 3H Cl
Methane Chloroform
CH4 + 4Cl2 ⟦line⟧ C Cl4 + 4HCl
Carbon tetrachloride
4- It acts as a bleac⟦hing ag⟧ent in the presence of moisture.
This property is ex⟦plained b⟧y the following equation:
Cl2 + H2O ⟦line⟧ 2HCl + O
The nascent oxygen whi⟦ch is⟧ thus liberated oxidizes the colouring
materials & bleaches it.
Chlorine is also used according to the same principle for killing
Bacteria in the process of the purification of water.