Sir. La Rose Chemistry notes on Inorganic Chemistry. 647-5159 contact for private tutoring
Teacher:
Sir. LaRose
1
Sir. La Rose Chemistry notes on Inorganic Chemistry. 647-5159 contact for private tutoring
Table of Contents
Contents
1.
Describe the physical and chemical properties of metals .................................................................... 4
2.
Describe the reactions of metallic oxides, nitrates, carbonates and hydroxides ................................. 9
3.
Describe the extraction of Aluminium and Iron ................................................................................. 11
Aluminium ............................................................................................................................................... 11
Iron .......................................................................................................................................................... 13
4 Explain why metal alloys are often used in place of their metals ........................................................... 14
5. Relate the properties of the metals; aluminium, lead, iron, copper and their alloy to their uses. ........ 15
Aluminium ............................................................................................................................................... 15
Lead ......................................................................................................................................................... 15
Iron .......................................................................................................................................................... 16
Copper..................................................................................................................................................... 16
6. Explain the importance of metals and their compounds on living systems and the environment. ....... 17
7. Discuss the harmful effects of metals and their compounds on living systems and the environment.. 18
8. Describe the physical and chemical properties of non-metals ............................................................... 20
9. Describing the laboratory preparation of oxygen, carbon dioxide, ammonia. ...................................... 23
Oxygen .................................................................................................................................................... 23
Carbon Dioxide........................................................................................................................................ 24
Ammonia ................................................................................................................................................. 25
10. Explain the use of gases based on their properties .............................................................................. 26
Oxygen .................................................................................................................................................... 26
Hydrogen................................................................................................................................................. 26
Carbon Dioxide........................................................................................................................................ 26
Nitrogen .................................................................................................................................................. 27
Chlorine ................................................................................................................................................... 27
11. List the usage of non-metals, C, S, P, Si, N, Cl and their compounds.................................................... 28
Carbon ..................................................................................................................................................... 28
Sulphur .................................................................................................................................................... 28
Phosphorous ........................................................................................................................................... 28
Silicon ...................................................................................................................................................... 29
Nitrogen .................................................................................................................................................. 29
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Chlorine ................................................................................................................................................... 29
12. Describe the harmful effects of non-metals on living systems and the environment. ........................ 30
BIBLIOGRAPHY ............................................................................................................................................ 32
3
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1. Describe the physical and chemical properties of metals
Metal
Zinc (Zn)
Physical Properties
Chemical Properties
•
Bluish-white metal
•
Shiny
is formed. This reaction occurs slowly in
•
The boiling point is 420ºC
the cold but when heated, it occurs more
and the melting point is
rapidly
907ºC.
2Zn(s) + O2(g)
•
•
•
Reasonable conductor of
When zinc reacts with oxygen, zinc oxide
2Zn0(s)
When zinc reacts with water, it displaces
electricity
hydrogen from steam.
•
The hardness is 2.5
Zn(s) +H20(l)
•
The density of zinc is 7.140
•
g/mol
ZnO(s) +H2(g)
When zinc reacts dilute HCL, it produces
hydrogen gas and zinc chloride
Zn(s) + 2HCl(aq)
•
ZnCl2(aq) + H2(g)
Zinc metal dissolves slowly in dilute
sulphuric acid to form solutions containing
the ZnSO4 ion together with hydrogen gas.
Zn(s) + H2SO4(aq)
Iron (Fe)
•
Silver grey metal
•
The melting point is 1538°C
oxide is formed. Whereas, it rusts in cold
and the Boiling point is
temperatures if moisture is present.
2861°C.
3Fe(s) + 2O2(g)
•
•
ZnSO4(aq) + H2(g)
•
Good transmission of heat
When reacting with oxygen, if heated, and
Fe3O4(s)
When iron reacts with water, reacts
or electricity
reversibly with steam to produce hydrogen.
•
Has a shine or glow
3Fe(s) + 4H20(l)
•
The hardness is 4
•
The density of iron is
aqueous Hydrochloric acid, then Iron (II)
7.874g/cm3
chloride or ferrous chloride is formed.
•
When solid iron filings are added to dilute
Fe(s) + 2HCl(aq)
4
Fe3O4(s) + 4H2(g)
FeCl2(aq) + H2(g)
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•
When iron and dilute sulphuric acid reacts
with each other, hydrogen gets displaced
and iron sulphate is formed
Fe(s) + H2S04(aq)
Sodium
(Na)
•
•
•
•
Silvery white metal with a
FeSO4(aq) +H2(g)
When reacting with oxygen, sodium burns
waxy appearance
slowly and eventually catches fire to form
The surface is bright and
sodium oxide.
shiny
4Na(s)+O2(g)
•
The melting point is 97.82
2Na2O(s)
Sodium metal reacts rapidly with water to
ºC and its boiling point is
form a colourless solution of sodium
881.4°C
hydroxide (NaOH) and hydrogen gas.
•
The density is 0.968g/cm3
2Na(s)+2H20(l)
•
Sodium is a good conductor
•
•
2NaOH(aq) + H2(g)
The reaction between sodium and
of electricity
hydrochloric acid is violent and overs
It is soft enough to be cut by
quickly producing the salt, sodium
a knife.
chloride. The sodium ignites, producing a
bright flame.
2Na(s)+ 2HCl(aq)
•
2NaCl(aq) + H2(g)
Sodium metal dissolves readily in dilute
sulphuric acid to form solutions containing
the Na2SO4 together with hydrogen gas.
2Na(s)+ H2SO4(aq)
Calcium
(Ca)
•
Na2SO4(aq) + H2(g)
•
Silver white metallic
•
The surface is shiny
spontaneous reaction takes place, when
•
Relatively soft metal
compared to sodium, calcium oxide is
•
The melting point is 851°C
formed.
and the boiling point is
2Ca(s)+ O2(g)
•
1482°C
•
When reacting with oxygen, a less
2CaO(s)
Calcium reacts slowly with water, forming
It is a good conductor of
calcium hydroxide and hydrogen gas. The
electricity
calcium metal sinks in water and after a
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•
The density is 1.55g/cm3
while bubbles of hydrogen are evident,
stuck to the surface of the metal.
Ca(s) + 2H20(l)
•
Ca(OH)2(aq)+ H2(g)
Calcium metal dissolves readily in dilute or
concentrated hydrochloric acid to form
solutions containing the CaCl2 ion together
with hydrogen gas.
Ca(s) + 2HCl(aq)
•
CaCl2(aq) + H2(g)
The reaction of sulfuric acid with calcium
metal produces a coating of calcium
sulphate (CaSO4) on the metal. The
calcium sulphate is insoluble in water, so
this coating acts as a protective layer which
prevents further attack on the metal by the
acid.
Ca(s) + H2SO4(aq)
Magnesium • Silver or grey coloured
(Mg)
metal
•
CaSO4(aq) + H2 (g)
Once ignited, magnesium burns in the
presence of oxygen, with a characteristic
•
Shiny
blinding bright white flame made up of
•
The melting point is 649°C
white magnesium oxide
and the boiling point of is
2Mg(s) + O2(g)
•
1090°C.
2MgO(s)
There is no significant reaction with water.
•
The density is 1.738g/mL
It reacts slowly with hot water, but
•
It has a hardness of 1-2.5
however reacts with steam forming
•
Magnesium conducts
magnesium oxide or hydroxide.
electricity but magnesium is
Mg(s) + 2H2O(l)
•
a bad metal to make
Mg(OH)2(aq) + H2(g)
Magnesium metal dissolved readily with
electrical contacts and wires
hydrochloric acid. This is a single
from.
replacement reaction and generates
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hydrogen gas as well as magnesium
chloride.
Mg(s) + 2HCl(aq)
•
MgCl2(aq) +H2(g)
Magnesium readily reacts with sulfuric
acid and forms hydrogen gas bubbles and
aqueous magnesium sulphate after the
reactants are consumed.
Mg(s) + H2SO4(aq)
Copper
(Cu)
•
•
•
•
Reddish metal; tarnishes to
Copper metal is stable in air under normal
black or green in air
conditions. When strongly heated, copper
Has a melting point of
metal and oxygen react to form black
1356.6 and boiling Point is
Cu2O.
2840ºC
2Cu(s) + O2(g)
•
The lustre of copper is
water.
•
The hardness is 2.5-3
Cu(s) + H2O(l)
•
It is a good conductor of
•
heat and electricity
The density is 8.96g/cm
•
CuO(s) + H2(g)
It does not react with hydrochloric acid
Cu(s) + 2HCl(aq)
3
2CuO(s)
There is no reaction between copper and
bright metallic
•
MgSO4(aq) + H2(g)
CuCl2(aq) + H2(g)
Copper metal dissolves in hot concentrated
sulphuric acid to form solutions containing
the copper sulphate ion together with
hydrogen gas
Cu(s) + 2H2SO4(aq)
+ SO2(g)
7
CuSO4(aq) + 2H2O(l)
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Aluminium • Silvery- white metal
(Al)
• Soft with dull lustre
•
•
When exposed to oxygen, it rapidly
tarnishes in the cold. A thin layer of
The melting point of
aluminium oxide forms.
aluminium is 660.32 ºC and
4Al(s) + 3O2(g)
•
its boiling point is 2579ºC
It reacts with water liberating a lot of heat
•
The density is 2.7g/cm3
and hydrogen.
•
The hardness is 9
2Al(s) + 6H2O(l)
•
Aluminium is an excellent
3H2(g)
•
heat and electricity
conductor
2Al2O3(s)
2Al(OH)3(aq) +
When aluminium is placed in hydrochloric
acid it may initially appear not to react.
This is because a layer of aluminium oxide
forms on the surface of the aluminium due
to prior reaction with the air and acts as a
protective barrier. The acid must remove
this layer before it is able to react with the
aluminium underneath.
2Al(s) + 6HCl(aq)
2AlCl3(aq) +
3H2(g)
•
Aluminium reacts with sulphuric acid) to
produce aluminium sulphate and hydrogen
gas.
2Al(s) + 3H2SO4(aq)
3H2(g)
8
Al2(SO4)3(aq) +
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2. Describe the reactions of metallic oxides, nitrates,
carbonates and hydroxides
Substance Dilute Acid
Metallic
Oxides
Heat
This reaction forms salt
Heat can decompose metal oxides to give oxygen
and water. The metal
and metal. The more reactive the metal is, the higher
replaces the hydrogen.
the temperature in order to decompose. With metals
Hydrogen gas is liberated
like Aluminium, many thousands of degrees Celsius
and the other product
would be needed to carry this out.
which is the salt, is
2Al2O3
4Al + 3O2
formed from the chemical
bonding of a non-metal
with a metal component.
Na2O + 2 HCl2
NaCl + H2O
Nitrates
The reaction between a
Most nitrates tend to decompose on heating to give
metal nitrate and dilute
the metal oxide, brown fumes of nitrogen dioxide,
acid is a double
and oxygen. They are droplets of water from the salt
decomposition. It is
losing water in crystallization. Magnesium and
driven by the formation of lithium nitrate tend to decompose completely:
an insoluble salt by the
interchange of ions
between the salt and acid.
Compared to this, the rest of the group does not
KNO3(aq) + HCl(aq)
decompose completely when using a Bunsen burner.
KCl(s) + HNO3(aq)
All the nitrates from sodium to caesium decompose
in this same way, the only difference being how hot
they have to be to undergo the reaction. As you go
down the Group, the decomposition gets more
difficult, and you have to use higher temperatures.
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Carbonates
Metal carbonates will
Most carbonates decompose on heating resulting in
react with dilute acids to
a metal oxide and carbon dioxide. There are droplets
give salt, water and
of liquid present from the salt losing water in
carbon dioxide gas. The
crystallization. Carbon dioxide is evolved showing
metal replaces the
that the compound is decomposing. Calcium and
hydrogen. The products
Lithium carbonate react similarly:
are salt, water and carbon
dioxide gas. The salt
produce depends upon the
acid used.
The reactions between
carbonates and acids are
called neutralization
reactions because the acid
It is more difficult to decompose the other
carbonates since the reaction does not occur at
Bunsen burner temperatures. The required
decomposition temperatures increase as you go
down the group.
is neutralized. In other
words, the acid and the
base (carbonate) are
neutralized, or their pH
gets close to 7.
MgCO3(aq)+2HCl(aq)
MgCl2(aq)+ H2O(l)+CO2 (g)
Hydroxides Metal hydroxides react
Metal hydroxides decompose on heating to yield
with dilute acid to form
metal oxides and water. Sodium hydroxide
salt and water.
decomposes to produce sodium oxide and water.
HCl(aq) + NaOH(aq)
Droplets of liquid are present showing that the
NaCl(aq) + H2O(l)
compound is decomposing.
2NaOH(s)→Na2O(s)+H2O(g)
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3. Describe the extraction of Aluminium and Iron
Aluminium
Aluminium is the most abundant metal found the Earth’s crust. Bauxite plays an important role
in some economies and is the only important ore of Aluminium. It is expensive because large
amount of electricity are required in its extraction process. The bauxite is purified to produce
aluminium oxide, a powder from which aluminium can be extracted.
The mined bauxite is either:1. Converted to pure alumina (Al2O3), the anhydrous compound, or
2. Heated to 3000 ºC to produce calcined bauxite
Aluminium is obtained from alumina by electrolysis. The ions in the aluminium oxide must be
free to move so that electricity can pass through it. Aluminium oxide has a very high melting
point (over 2000°C) so it would be expensive to melt it. Aluminium oxide does not dissolve in
water, but it does dissolve in molten cryolite (sodium aluminium fluoride). The addition of
cryolite lowers the melting point to 960ºC. The presence of cryolite also gives the melt better
conducting properties. The use or cryolite reduces some energy cost when extracting aluminium.
During the electrolysis of alumina:
•
A molten mixture of cryolite and aluminium oxide is used as the electrolyte
•
5 V and 100000 A are required
•
The positively charger aluminium ions gain electrons from the cathode and form molten
aluminium
Al3+ + 3ē→Al (l)
•
The oxide ions lose electrons at the anode, and form oxygen molecules which is released at
the anode.
2O2+(l) – 4ē → O2 (g)
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Extraction of Aluminium
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Iron
Iron is produced more than any other metal and most of it is converted to steel. The raw
materials for the iron industry include the iron ore, a reducing agent e.g. coke or natural gas, an
energy source and a flux, usually limestone, to form a slag with silicates and other impurities.
Iron is extracted from iron ore in a huge container called a blast furnace. Iron ores such as
haematite contain iron (III) oxide, Fe2O3. The oxygen must be removed from the iron (III) oxide
in order to leave the iron behind. Dried heated iron ore, limestone and coke are fed into the top of
the furnace. Near the bottom, hot air is blown into the furnace. Due to the heat, the coke burns
producing carbon dioxide and generating a great deal of heat. The carbon dioxide formed is then
reduced the carbon monoxide by the hot coke; CO2 (g) + C(S)→ 2CO (g).The carbon monoxide
reduces the hot iron ore to molten iron; Fe2O3(s) + 3CO (g) → 2Fe (l) + 3CO2 (g).
The molten iron runs to the bottom of the furnace. The iron produced in the blast furnace would
have very large amounts of impurities if limestone was not used. The limestone then breaks
down; CaCO3(s)→ CaO(s)+ CO2(g). The calcium oxide then combines with silicon dioxide which is
the main impurity in iron ore, to form a molten slag; CaO(s) + SiO2(s) →CaSiO3(l). The slag does
not mix with the molten iron but floats on it, then they run off separately. The molten iron
formed in the blast furnace, is an impure form of the metal. The molten metal is then allowed to
solidify in shallow trays known as casts. Because of this, the iron is referred to as ‘cast iron’.
Extraction of Iron
THE BLAST
FURNACE
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4 Explain why metal alloys are often used in place of their metals
An alloy is a combination of metals or a combination of one or more metals with non-metallic
elements. Metal alloys are used because they are often harder than the pure metal. When the
combination occurs, a stronger item is created. It is difficult for the layers of the atoms to move
in alloys because the atoms in the atomic arrangement are of varying sizes. This then increases
its durability and stability. Pure metals have a uniform layer of atoms which makes the layers
more vulnerable to movement. The alloys are used instead of pure metals in manufacturing
because of their hardness. The properties of metals are enhanced by alloys which increase their
usefulness.
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5. Relate the properties of the metals; aluminium, lead, iron,
copper and their alloy to their uses.
Aluminium
Pure aluminium is soft, ductile, and corrosion resistant and has a high electrical conductivity. It
is widely used for foil and conductor cables. It is use in cans, kitchen utensils, window frames,
beer kegs and aeroplane parts but alloying with other elements is necessary to provide the higher
strengths needed for other applications Aluminium is not a very strong metal, but its conductive
qualities make it useful for a variety of applications. For this reason, manufacturers mix
aluminium with other metals to strengthen it, forming several different aluminium alloys.
Aluminium alloys are widely used in automotive engines, particularly in cylinder blocks and
crankcases due to the weight savings that are possible. Some alloys include Magnalium and
Duralumin.
Lead
The properties of lead that make it useful in a wide variety of applications are its lubricity,
malleability, density, electrical conductivity, and coefficient of thermal expansion, which are all
high. It is also characterised by its low elastic modulus, elastic limit, strength, hardness, and
melting point. Lead also has good resistance to corrosion under a wide variety of conditions.
Lead is easily alloyed with other metals. The high density of lead makes it effective when
shielding again x-rays and gamma radiation. The primary consumption of lead is for Lead- Acid
batteries, rolled extrusions, cable sheathing and ammunition. Seamless pipes are usually made
from lead alloys because of its corrosion resistance and flexibility. Anodes made from lead
alloys are used in electro-winning and plating of metals e.g. manganese. An example of a lead
alloy is Solder.
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Iron
Iron is very strong but rusts easily. It also forms a wide range of alloys. It is used to manufacture
steel and also used in civil engineering like reinforced concrete, girders etc. Iron is used to make
bridges, electricity pylons, bicycle chains, cutting tools and rifle barrels. Iron alloys are arguably
the most important class of engineering materials. The most well-known alloy of iron is steel
which contains carbon as its supplemental element. The carbon helps prevent the iron from
rusting, and makes it stronger. People use the material widely in construction, such as for making
screws, nails and beams for buildings and bridges. A common alloy of iron is steel.
Copper
Copper is soft, malleable, easy to bend and a good conductor of heat and electricity as well as
being resistant to corrosion. Pure copper is prone to oxidation making its surface a dull, pale
greenish colour. Manufacturers fuse copper with several other elements to prevent oxidation and
increase its strength e.g. brass, a copper alloy, contains about 20% zinc. These alloys are used
jewellery, nuts and bolts. Copper’s most common use is in electrical equipment like wiring and
motors. Due to its resistance to corrosion, it is used in roofing and guttering. It is also used in
plumbing, cookware and cooking utensils.
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6. Explain the importance of metals and their compounds on
living systems and the environment.
Metal
Importance
Iron
Iron is an essential for all forms of life and is non- toxic. About four grams of iron is
contained in the body. Iron is found in haemoglobin, giving blood cells their red
colour. Haemoglobin carries oxygen from our lungs to the cells, where the oxygen is
used to release energy from food. This energy is required for many chemical reactions
to sustain life. A lack of haemoglobin causes anaemia to develop. Haemoglobin is
made up of four protein molecules. Each haem molecule is a coordination complex
containing a Fe2+ ion held by four covalent bonds. An oxygen molecule can attach
reversibly to the Fe2+ ion.
Zinc
Zinc is a common element in human and natural environments and plays an important
part in many biological processes Zinc is found in cells throughout the body. It is
needed for the body’s defensive system to properly work. It has a role in cell division,
cell growth, wound healing and the breakdown of carbohydrates. Senses of smell and
taste rely on zinc as well. Zinc, an essential trace element is essential for the normal
growth and reproduction of humans and plants. It is vital for the functionality of more
than 300 enzymes, for the stabilization of DNA, and for gene expression.
Magnesium Magnesium is important to living systems for various reasons. It is needed for more
than 300 biochemical reactions in the body. Magnesium helps to maintain normal
nerve and muscle function, helps bones remain strong, supports a healthy immune
system and keeps the heartbeat steady. It also helps to adjust blood glucose levels as
well as aiding in the production of energy and protein. It is a part of chlorophyll.
Calcium
Calcium is an important part of structure of plant cell walls and cell membranes. It is
essential when building strong bones and teeth. It is involved in muscle contraction
and prevents muscle cramp. Calcium is also involved in vascular contraction,
vasodilation, nerve transmission, intracellular signalling, and hormonal secretion. If
someone does not take in sufficient amounts of calcium in his or her diet, the body
will demineralize which would make the bones weak and reduce body mass.
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7. Discuss the harmful effects of metals and their compounds on
living systems and the environment
Metal
Harmful Effects
Lead
Exposure to high levels of lead can cause anaemia, hypertension, kidney and brain
damage. Lead attacks the brain and central nervous system to cause coma,
convulsions and even death. Children who survive lead poisoning may be left
with mental retardation and behavioural disorders. The main sources of lead
entering an ecosystem are atmospheric lead (from vehicles), paint chips, used
ammunition, fertilisers and pesticides and lead- acid batteries or other industrial
products. Lead is a particularly dangerous chemical, as it can accumulate in
individual organisms, but also in entire food chains. These animals experience
the health effect from lead poisoning. Shellfish are very susceptible to very small
concentrations of lead. The functions of phytoplankton can be disturbed by lead
interferes. Phytoplankton is eaten by many large sea animals due to its source of
oxygen. Soil functions are disturbed by lead intervention especially near
highways were extreme concentrations may be present, hence the soil organisms
suffer. Lead can be released directly into the air, as suspended particles. Historic
major sources of lead air emissions were motor vehicles and industrial sources.
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Mercury
The inhalation of mercury vapour can cause neurological and behavioural
disorder such as tremors, emotional instability, insomnia, memory loss, neuromuscular changes and headaches. They can also harm kidneys and thyroids. The
inorganic salts of mercury are corrosive to the skin, eyes and gastrointestinal
tract, and may induce kidney toxicity if ingested. A very important factor in the
impacts of mercury to the environment is its ability to build up in organisms and
up along the food chain. The mercury enters freshwater lakes and rivers, then
accumulates in the sediments at the bottom. Methyl mercury is taken up by the
small organisms and enter aquatic food chains. Mercury is released into the
atmosphere from the burning of fossil fuels in power stations and domestic and
industrial wastes in incinerators. Mercury compounds are released directly into
the land from many fungicides.
Nickel
The most serious effects from exposure to nickel are chronic bronchitis, reduced
lung function, cancer in the lung and nasal sinus. Nickel can be released to the
environment from the stacks of large furnaces used to make alloys from power
plants and trash incinerators. The nickel that comes out attaches to small
particles of dust that settle to the ground or is taken out of the air in rain or
snow. Nickel can also be released into the environment by industrial waste
water. The nickel released into the environment end up in soil. Nickel might
even seep into groundwater. Plants can take up and accumulate nickel but it
does not accumulate in small animals on land.
Cadmium
Exposure to cadmium can result in flu like symptoms (chills, fever and muscle
pains) and can damage the lungs. Being exposed to lower levels of cadmium
over a long period of time can cause damage to kidney, lungs and bones. In the
environment, it is toxic to plants, animals and microorganisms. Because of its
high rates of soil-to-plant transfer, cadmium is a contaminant found in most
human foodstuffs.
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8. Describe the physical and chemical properties of non-metals
NonMetal
Physical properties
Chemical properties
Hydrogen
•
Colourless, odourless gas
•
•
The density is 0.09g dm-3
with oxygen when the existing
•
Its melting point is -219ºC and the
molecular bonds break and new
boiling point is -183ºC
bonds are formed between oxygen
•
No hardness nor lustre
and hydrogen atoms. But hydrogen
•
The conductivity is 0.18 W/m-ºC
does not react with oxygen at room
Hydrogen molecules violently react
temperature, a source of energy is
needed to ignite the mixture.
H2+ O2→ H2O
•
Hydrogen reacts with alkali metals
forming a substance known as a
hydride.
2Na(s) + H2(g)→2NaH(s)
•
Hydrogen gas is a reducing agent
when it reacts with non-metals
H2(g)+ FeO → Fe + H2O
0
+2 -2
0
Hydrogen acts as an oxidizing agent
when it reacts with metals.
2Na(s) + H2(g)→2NaH(s)
0
20
0
+1
-1
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Chlorine
•
Pale Green-yellow gas with a sharp
•
When chlorine reacts with oxygen,
odour (Choaking smell)
the oxides formed are highly unstable
•
The density is 3.21 g dm-3
(explosive)
•
The melting point is -101ºC and the
2Cl2(g) + O2(g) → 2Cl2O(g)
boiling point is -34ºC
•
•
No hardness nor lustre
•
The conductivity is 0.0089 W/m-ºC
The type of reaction between chlorine
and most metals is direct
combination.
2Fe(s) + 3Cl2(g) → 2FeCl3(s)
M + Cl2 → MCl
•
Chlorine is a good oxidising agent,
either as a gas or as an aqueous
solution.
H2(g) + Cl2(g) → 2HCl(g)
Oxygen
•
Colourless, odourless gas
•
•
The density is 1.43 g dm-3
form the corresponding metal oxide.
•
The melting point is -219 ºC and the
2Mg(s) + O2(g) → 2MgO(s)
•
boiling point is -183 ºC
When oxygen reacts with metals, to
Oxygen is good oxidising agent,
•
No hardness nor lustre
oxidising most metals to metal oxides
•
The conductivity is 0.034 W/m-ºC
and some non-metals to non-metal
oxides.
5O2(g) + 4P(s) → P4O10(s)
Carbon
•
•
•
Black solid (graphite) conducts
electricity, brittle
there is the formation of carbon
Colourless crystal (diamond),
dioxide.
extremely hard
C(s) + O2(g) → CO2(g)
•
The density of carbon is 2.26g/cm3
•
The melting point 3367ºC and the
•
There is 0.5 hardness but no lustre
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Carbon does not normally react with
metals
•
boiling point 4827ºC
•
When carbon reacts with oxygen,
Carbon is a good reducing agent. At
high temperatures, it reduces metal
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•
oxides below aluminium in the
The conductivity is 2000 W/m-ºC
reactivity series to the metal.
Sulphur
•
Yellow solid (soft)
•
•
The density is 2 gcm3
between sulphur and oxygen
•
The melting point is 115.21ºC and
resulting in sulphur dioxide.
the boiling point is 444.6ºC
S(s) + O2(g) → SO2(g)
•
There is 2 hardness and no lustre
•
The conductivity is 0.205 W/m-ºC
•
There is a non-vigorous reaction
With metals, Sulphur reacts on
heating to produce metallic sulphides
which are ionic compounds
Mg(s)+ S(s) → MgS(s)
•
Sulphur is a good reducing agent.
The reducing properties are
demonstrated in their reactions with
the oxidising acids, concentrated
sulphuric acid and concentrated nitric
acid.
S(s) + 2H2SO4(aq) → 3SO2(g) + H2O(l)
Nitrogen
•
Colourless, odourless gas
•
•
The melting point is -210ºC and the
oxygen at room temperature although
boiling point is 195.8ºC
it will react at higher temperatures
•
There is no hardness nor lustre
and in the presence of an electric
•
The conductivity is 0.060 W/m-ºC
spark.
•
The density is 0.00116cm-3
•
Nitrogen does not combine with
There is a vigorous reaction between
nitrogen and metals resulting in the
formation of a nitride.
3Mg(s) + N2(g) → Mg3N2(s)
•
Nitrogen prevents substances from
being oxidised.
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9. Describing the laboratory preparation of oxygen, carbon
dioxide, ammonia.
Oxygen
Concentrated hydrogen peroxide is dropped slowly from a dropping funnel into a flask
containing manganese (IV) oxide catalyst, 2H2O2 (aq)→O2(g) +2H2O(l). The oxygen is collected in
the gas jar by downward displacement of water. Oxygen is slightly soluble in water so most of
the oxygen produced will be collected in the gas jar. The oxygen collected in the gas jar will
contain some water vapour because it is collected over water. If the oxygen is required dry, it
can be passed over anhydrous calcium chloride in a U-tube.
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Carbon Dioxide
Carbon dioxide is prepared by dropping dilute hydrochloric acid onto marble chips. In this
preparation, carbon dioxide is denser than air, hence it is collected in the gas jar by upward
displacement of air. It is soluble in water so it is preferable not collect it over water. The
carbon dioxide entering from the reaction flask will contain some water vapour. This then
arises from the water in the dilute hydrochloric acid. The carbon dioxide can be passed over
anhydrous calcium chloride in a U-tube to dry it.
Test: Bubble the gas through lime water (calcium hydroxide). Goes from
colourless to milky white
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Ammonia
Ammonia is an alkaline gas which is very soluble in water. It can be prepared by warming any
alkali with ammonium salt. In the diagram below, calcium hydroxide and ammonium chloride is
heated gently.
The ammonia is less dense than air so it is collected in the gas jar by downward displacement of
air. An aqueous solution should not be used since ammonia is soluble in water. The ammonia
then passes over calcium oxide to dry it. Calcium chloride cannot be used because ammonia
reacts with it as well as concentrated sulphuric acid since that as well reacts with ammonia.
NH4Cl + Ca(OH)2 → NH3 + H2O + CaCl2
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10. Explain the use of gases based on their properties
Oxygen
Oxygen is an odourless, colourless and tasteless gas. It is made up of 22% of the air. Oxygen
gas can generate temperatures of 3000ºC making it suitable for oxy-hydrogen and oxyacetylene blow torches. This is used in industry for cutting, welding and melting metals.
Oxygen does not burn but supports combustion. It is used to produce energy in industrial
processes, generators and ships. Oxygen is also used in airplanes and cars. Liquid oxygen is
burnt by spacecraft for thrust. Oxygen is slightly soluble in water; astronauts, mountaineers
and scuba divers use breathing apparatus that contain oxygen gas. Oxygen is used to destroy
bacteria, treat victims of carbon monoxide, and for aerobic respiration.
Hydrogen
Hydrogen is colourless and odourless gas. It is less dense than air, due to this it is used in
weather balloons that are fitted with equipment to record information necessary to study the
climate. Hydrogen is used in fertilisers, food and chemical and paint industries. Hydrogen
fuel cells are used to generate electricity from oxygen and hydrogen. Hydrogen gas is used in
the processing of petroleum products to break down crude oil into fuel oil, gasoline and such.
Hydrogen is important in creating ammonia (NH3) for use in making fertilizer. Hydrogen gas
is used as a hydrogenating agent to for polyunsaturated fats, such as used in margarine.
Carbon Dioxide
Carbon Dioxide is a colourless and odourless gas. Carbon dioxide does not burn and is
denser than air. Some fire extinguishers contain CO2 when sprayed on fire. Since it is denser
than air and the ‘blankets’ of fire, it’s sealed off from the oxygen. Carbon dioxide
extinguishers are especially useful for dealing with fires involving flammable liquids and
electrical equipment. Carbon dioxide is used make drink ‘fizz’ by pumping carbon dioxide
into the drink under pressure. Solid Carbon Dioxide is used to keep materials cold. It is also
used in pressurising oil wells and as an aerosol propellant.
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Nitrogen
Nitrogen gas is used in handling explosive mixtures, to anneal metals at high temperatures.
This is because of the inertness of nitrogen which prevents premature explosions. It is used to
flush out boilers and pipes during non-use periods since it reduces the chance of corrosion.
Nitrogen protects food from spoilage because bacteria cannot survive in an atmosphere of
nitrogen. Used to make ammonia
Chlorine
This green-yellow gas is soluble in water and has bleaching properties and so turns red
litmus, colourless. Chlorine is used to make sodium hypochlorite which is present in many
bleaches because it is a powerful oxidant. It is involved in cotton and paper manufacture of
polychloroethene since it adds readily to alkenes.
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11. List the usage of non-metals, C, S, P, Si, N, Cl and their
compounds
Carbon
•
Carbon, in the form of graphite, is used as a lubricant and is also used as ‘lead’ in
lead pencils. This is possible because there are weak forces between the layers
allowing the plates of graphite to slide past each other.
•
As diamond, it is used in drill tips for high speed drills because of its hardness.
•
Diamond is also used in jewellery because of its lustre.
•
Carbon fibres are used to strengthen some types of plastic as well.
•
Used to cut glass
•
Graphite is used as electrodes
Sulphur
•
Sulphur is used in the production of chemicals e.g. sulphuric acid which is normally used
in the manufacture of fertilisers and detergents.
•
Matches and gun power contain sulphur since it burns easily and quickly.
•
Sulphur forms links between polymer chains hence it is used in the manufacture of tyres to
make the rubber harder, known as vulcanisation.
•
Sulphur powder is used as a fungicide on plant materials.
Phosphorous
•
Phosphorous is used to make flares and fireworks and
•
phosphorous sulphide is used to make the heads of ‘strike anywhere’ matches because of
its spontaneous inflammable nature.
•
Phosphates are used in fertilisers and pesticides, cleaning agents and water softeners.
•
Phosphorous is an essential element for plant growth and contains particle binding
properties.
•
A small amount of phosphorous is used to make the alloy, phosphor bronze.
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Silicon
•
Silicon is used in electronic devices like transistors and calculators due to its semiconducting property.
•
Silicon in its highly purified form is used to make silicon chips for computers.
•
Sand contains silicon oxide which is used in the manufacture of glass. The glass is made
by heating the sand with calcium oxide and sodium carbonate.
•
Glass fibres are silicates which are used to strengthen plastics (fibreglass). Fibreglass has
a low density, it is strong and is used to make pipes and storage tanks.
•
Traces of transition element atoms found in silicates are used for jewellery.
Nitrogen
Done above
Chlorine
Chlorine is also used to sterilise swimming pools and in water treatment. The active
ingredients of some insecticides are chlorine-containing compounds. A major use of
chlorine is to make the monomer for the plastic, PVC. Some dry cleaning and industrial
solvents and refrigerants contain compounds of chlorine.
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12. Describe the harmful effects of non-metals on living systems
and the environment.
Non- Metal
Carbon Dioxide
Harmful Effects
Carbon dioxide is the product of all combustion reactions and it is
used during photosynthesis. The burning of fossil fuels,
population explosions and deforestation leads to an increase in
carbon dioxide concentrations. It is referred to as a greenhouse
gas because they are good absorbers of infrared radiation leading
to the heating of the atmosphere. This heat causes global warming
and changes in rainfall pattern. This then leads to polar ice caps
melting, rise in sea levels, more violent and unpredictable weather
patterns, formation of more desserts and an increase in the
temperature of the oceans. Exposure to CO2 can produce a variety
of health effects. These may include headaches, dizziness,
restlessness, a tingling or pins or needles feeling, difficulty
breathing, increased heart rate, elevated blood pressure, coma,
asphyxia, and convulsions.
Carbon Monoxide
Carbon monoxide is formed from the incomplete combustion of
carbon or carbon compounds e.g. decay of organic matter and
burning fossil fuels. When carbon monoxide is emitted into the
atmosphere it effects the amount of greenhouse gases, which are
linked to climate change and global warming. If CO is inhaled, it
can cause headache, dizziness, vomiting, and nausea. If CO levels
are high enough, you may become unconscious or die. Exposure
to moderate and high levels of CO over long periods of time has
also been linked with increased risk of heart disease.
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Chlorofluorocarbons
Aerosols sprays and some refrigerants contain CFCs which
catalyse the breakdown of ozone into oxygen. Because of it is
colourless, non- toxic, heat resistant, inert, non-flammable and can
easily be liquefied, CFCs are allowed to escape into the air
without being detected and move up the atmosphere. CFCs are
said to be responsible for the destruction of the ozone layer. This
ozone layer is vital since it protects the earth against the harmful
ultra-violet rays. The breakdown of ozone causes the formation of
holes in the ozone layer. This results in the increased risk of
getting skin cancer, eye cataracts and reduced resistance to some
disease. Warming of the earth
Nitrates and
Eutrophication is caused by the excessive amounts of nitrates and
Phosphates
phosphates in lakes and rivers. The nitrates and phosphates from
fertilisers applied to the fields dissolve in groundwater and get
into lakes and rivers. Their presence causes the excessive growth
of algae which then cover the surface of the water. Due to the lack
of sunlight, water plants die and the bacteria feeds on the plant
remains. The aerobic bacteria uses up the oxygen in the water, so
the aquatic animals die. This affects people's opportunity to use
lakes and rivers for leisure activities. These losses can mean that
the value of tourism and properties decreases.
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BIBLIOGRAPHY
•
Chemistry for CSEC- Nelson Thomas Ltd
•
Chemistry for CSEC – Keane Campbell et al
•
https://courses.lumenlearning.com/introchem/chapter/properties-of-carbon/
•
https://www.engineersedge.com/heat_transfer/thermal-conductivity-gases.htm
•
https://www.ncbi.nlm.nih.gov/pmc/articles/PMC2831915/
•
www. https://www.rsc.org/periodic-table/element/6/carbon
•
https://webcam.srs.fs.fed.us/impacts/mercury/index.shtml
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Reactivity Series of Metals
•
•
•
•
The chemistry of the metals is studied by analysing their reactions with water and acids
and oxygen
Based on these reactions a reactivity series of metals can be produced
The series can be used to place a group of metals in order of reactivity based on the
observations of their reactions with water and acids oxygen
The non-metals hydrogen and carbon are also included in the reactivity series as they are
used to extract metals from their oxides
Table of Metal Reactions
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The reactivity series mnemonic
•
•
Observations from the table above allow the following reactivity series to be deduced
The order of this reactivity series can be memorised using the following mnemonic
o “Please send cats, monkeys and cute zebras into hot countries signed Gordon"
You can learn the reactivity series with the help of a silly phrase
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Reactions of Metals
Reaction with cold water
•
•
The more reactive metals will react with cold water to form a metal hydroxide and
hydrogen gas
Potassium, sodium and calcium all undergo reactions with cold water as they are the
most reactive metals:
metal + water → metal hydroxide + hydrogen
•
For example, calcium and potassium:
Ca (s) + 2H2O (l) → Ca(OH)2 (aq) + H2 (g)
K (s) + H2O (l) → KOH (aq) + H2 (g)
Reaction with steam
•
Metals just below calcium in the reactivity series do not react with cold water but will
react with steam to form a metal oxide and hydrogen gas, for example, magnesium:
Mg (s) + H2O (g) → MgO (s) + H2 (g)
Reaction with dilute acids
•
•
•
•
•
•
Only metals above hydrogen in the reactivity series will react with dilute acids
Unreactive metals below hydrogen, such as gold, silver and copper, do not react with
acids
The more reactive the metal then the more vigorous the reaction will be
Metals that are placed high on the reactivity series such as potassium and sodium are
very dangerous and react explosively with acids
When acids react with metals they form a salt and hydrogen gas:
The general equation is:
metal + acid ⟶ salt + hydrogen
•
Some examples of metal-acid reactions and their equations are given below:
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Acid-Metal Reactions Table
Reaction with oxygen
•
•
•
Some reactive metals, such as the alkali metals, react easily with oxygen
Silver, copper and iron can also react with oxygen although much more slowly
When metals react with oxygen a metal oxide is formed, for example, copper:
metal + oxygen → metal oxide
2Cu (s) + O2 (g) → 2CuO (s)
• 4Na (s) + O2 (g) → 2Na2O (s)
• 2Mg (s) + O2 (g) → 2MgO (s)
•
•
Gold does not react with oxygen
Deducing the order of reactivity
•
•
•
•
•
•
The order of reactivity of metals can be deduced by making experimental observations
of reactions between metals and water, acids and oxygen
The more vigorous the reaction of the metal, the higher up the reactivity series the metal
is
A combination of reactions may be needed, for example, the order of reactivity of the
more reactive metals can be determined by their reactions with water
The less reactive metals react slowly or not at all with water, so the order of reactivity
would need to be determined by observing their reactions with dilute acid
Temperature change in a reaction can also be used to determine the order of reactivity
The greater the temperature change in a reaction involving a metal, the more reactive
the metal is
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5 Unique Properties Of Water
Water is one of the most important sources of life for us, and not only is it healthy, but it is also a
unique substance with some interesting properties.
▪
When water begins to evaporate off of a surface, it creates a cooling effect.
▪
Lower density of ice allows for only the tops of lakes to be frozen.
▪
Water is an extremely potent solvent due to its characteristic of high polarity.
Imagining life without water is impossible, literally. It is one of the most important sources of
life for us, and not only is it healthy, but it is also a unique substance with some interesting
properties. This article will deal with those and try to show you a more interesting side to this
substance we mostly take for granted.
The five main properties that will be discussed in this article are its attraction to polar
molecules, its high specific heat, the high heat of vaporization, the lower density of ice, and its
high polarity. So let’s begin!
5. Attraction To Other Polar Molecules
The property of cohesion allows liquid water to
have no tension on the surface.
Cohesion is what we call the ability of water to
attract other water molecules. It is one of its most
important properties. Water has a high polarity, and
it gives it the ability to being attracted to other
water molecules. These molecules are held together
by the hydrogen bonds in water.
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The property of cohesion allows certain insects can walk on water. Also, because of cohesion,
water manages to remain liquid at moderate temperatures and not turn into a gas. There is also
the ability of water to bind with molecules of different substances. This is called adhesion.
Through this property, water can be adhesive to any other molecule it can form a hydrogen bond
with.
4. High Specific Heat
Water manages to stay liquid because of two of its properties, high specific heat, and its high
heat of vaporization. More on the latter in the next paragraph, but here we will focus on the
former. High specific heat refers to the amount of energy that is absorbed or lost by one gram of
a specific substance to change the temperature by 1 degree celsius.
Since water molecules form hydrogen bonds with each other, a lot of energy is needed to break
those bonds. By breaking them, we allow the molecules to move around freely, and they have a
higher temperature. A more straightforward way to describe this would be to say that with many
individual water molecules floating around, more friction is created, which creates more heat and
higher temperatures. The hydrogen bonds absorb this heat. This is why water takes longer to get
heated and holds its temperature longer.
3. High Heat Of Evaporation
When water starts evaporating off of a surface, it
creates an effect of cooling.
This is another unique property that allows water to
maintain its temperature. The high heat of
vaporization refers to the amount of heat energy that
we need in order to be able to change one gram of
water into gas. A lot of energy is required to break
the hydrogen bonds between water molecules.
When water starts evaporating off of a surface, it creates an effect of cooling. This is similar to
humans and sweating. When we get hot, chemical bonds in our bodies begin breaking down, and
we start sweating as a cooling effect for our bodies. This is the same as the evaporation of water
and its cooling of the surface.
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2. Lower Density Of Ice
The reason icebergs are floating on the sea
surface is because of this lower density.
The hydrogen bonds between water molecules
start forming into ice crystals at higher
temperatures. Once they achieve this state, these
bonds become even more stable and will
maintain the shape of ice as long as the
temperature is not changing. Ice is the dense
form of water, and it has a lower density than water. The reason for this is the fact that the
hydrogen bonds are getting more spaced out in ice form. They are further apart from each other
than when in liquid form.
The reason icebergs are floating on the sea surface is because of this lower density. Also, it
allows for only the tops of lakes to be frozen, which is the fact that not many are aware of. While
most people are aware that the reason icebergs float is the lower density, not many know about
the reason why only the tops of lakes are frozen.
1. High Polarity
An excellent example of the high polarity of
water would be the fact that salt dissolves in
water.
Water is a polar molecule, which means it can
attract other polar molecules. The level of
polarity in water is extremely high, uniquely so.
It can form hydrogen bonds with other
elements. This makes water an extremely potent
solvent. The molecules that attract water molecules the most are those with a full charge, as an
ion.
An excellent example of the high polarity of water would be the fact that salt dissolves in water.
The salt molecules get surrounded by water molecules, and it separates the sodium from the
chloride. Water forms special hydration shells around those ions.
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•
Water dissociates salts by separating the cations and anions and forming new
interactions between the water and ions.
•
Water dissolves many biomolecules, because they are polar and therefore
hydrophilic.
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