IGCSE 3RD FORM CHEMISTRY NOTES – ADITYA GOEL
(With credit to  http://askmichellechemistry.blogspot.co.uk & BBC BITESIZE)
Basic Periodic Table Notes:
The periodic table is a list of elements arranged in order of their increasing atomic (proton) number.
A period is a horizontal row of elements--the number of electron shells is the same as the period
number of the element.
A group is a vertical column of elements--the number of valence electrons (outer shell electrons) is
the same as the group number of the element.
Since elements with similar electronic configurations have similar chemical properties, we can
deduce that elements in the same group have similar chemical properties hence all elements in
Group 1 are alike, as are the Halogens too, but in their own manners.
Chemical Properties
Metals
Non-metals
Usually have 1-3 electrons in their outer shell
Usually have 4-8 electrons in their outer shell
Lose their valence electrons easily
Gain or share (only in covalent bonds)
Form oxides that are basic
valence electrons
Form basic hydroxides
Form oxides that are acidic
Are good reducing agents
Are good oxidizing agents
Metal carbonate + acid  salt + water + carbon
dioxide
(Acid reactions with basic substances = salt + water)
acid + metal oxide → salt + water
acid + metal hydroxide → salt + water
(with reactive metals like Zn)
acid + metal → salt + hydrogen
Physical Properties
Metals
Non-metals
Good electrical and heat conductors
Poor conductors of heat and electricity
Malleable—can be hit and shaped
Brittle- if a solid
Ductile—can be stretched into wire
Non-ductile
Possess metallic luster (shiny)
Do not possess metallic luster
Opaque as thin sheet
Transparent as a thin sheet
Solid at room temperature (except Mercury [Hg]-
Solids, liquids or gases at room temperature
liquid)
The block of metals between Groups 2 and 3 are known as the transition metals/elements, and
form colored compounds
Common Ions to be remembered:
1.27 describe the formation of ions by the gain or loss of electrons
Positive ions/Cations
Negative ions/Anions
Charge
Name of ion Formula
Charge
Name of ion
Formula
1+
Ammonium
NH4+
1-
Bromide
Br-
Copper (I)
Cu+
Chloride
Cl-
Hydrogen
H+
Hydroxide
OH-
Lithium
Li+
Fluoride
F-
Potassium
K+
Iodide
I-
Silver
Ag+
Nitrate
NO3-
Sodium
Na+
Hydrogencarbonate
HCO3-
Barium
Ba2+
Carbonate
CO32-
Calcium
Ca2+
Sulphate
SO42-
Copper (II)
Cu2+
Sulphite
SO32-
Iron (II)
Fe2+
Sulphide
S2-
Lead (II)
Pb2+
Oxide
O2-
Magnesium
Mg2+
Carbonate
CO32-
Nickel (II)
Ni2+
Sulphate
SO42-
Strontium
Sr2+
Zinc
Zn2+
Aluminium
Al3+
Nitride
N3-
Iron (III)
Fe3+
Phosphate
PO43-
2+
3+
2-
3-
2
In ionic compounds, electrons have been donated/given. Ions are atoms or molecules, in which the
atoms have opposite charges, as one element donates electrons to achieve a full outer valence,
becoming electrically positive, the other element(s) gain an these electrons, becoming electrically
negative. The elements thus are bonded by a strong electrostatic charge. Metals tend to give
electrons, so they form cations (positive ions), hence normally elements from group 1-3 will form
cations.
If electrons are gained, the ion has a negative charge. Non-metals tend to do this, and they form
anions (A-Negative-ION - ANION). So elements from group 5-7 will form anions. Group 0/8 are the
noble gases and are inert + unreactive, so they do not form ions.
Tests for anions:
Anion
Chloride
ClBromide
Br –
Iodide
I-
Test
Dissolve in dilute nitric acid (HNO3), then a white precipitate forms when silver nitrate solution added. (The
precipitate dissolves in dilute ammonia solution.)
Dissolve in dilute nitric acid (HNO3), then a cream precipitate forms when silver nitrate solution added. (The
precipitate is insoluble in dilute ammonia solution, but will dissolve in concentrated ammonia solution)
Dissolve in dilute nitric acid (HNO3), then a pale yellow precipitate forms with silver nitrate solution. (The
precipitate is insoluble in dilute and concentrated ammonia solution)
Sulfate
SO42-
Dissolve in dilute hydrochloric acid (HCl), then add to barium chloride (, after which white precipitate forms.
Ammonium Add sodium hydroxide solution to the ammonia (NH3 right now) and heat with Bunsen burner, and then test
NH4+
the gas given off (liberated ammonium) with damp red litmus paper, which turns blue if positive for
ammonium.
Carbonate
CO32-
Add dilute HCl acid, and then pass the carbon dioxide gas through limewater, which turns milky (cloudy).
3
We can also do tests for cations too:
We can test for positive ions by adding sodium hydroxide solution and noting the colour of the
precipitate, as shown in the table below.
Cation
copper(II)
Cu2+
iron(II)
Fe2+
iron(III)
Fe3+
Result of adding sodium hydroxide solution
pale blue precipitate
dirty green precipitate
orange brown precipitate
2.40 describe simple tests for the gases:
• Hydrogen: apply a lit splint and you will hear a squeaky pop sound
• Oxygen: apply a glowing splint and the splint relights
• Carbon dioxide: bubble it into limewater and it goes milky white
• Ammonia: use a damp red litmus paper and it turns blue
• Chlorine: use damp blue litmus paper and it goes red (then bleaches it white)
1.28 understand oxidation as the loss of electrons and reduction as the gain of electrons
OILRIG - Oxidation Is Loss, Reduction Is Gain (of electrons)
Atomic Structure
1.8 recall that atoms consist of a central nucleus, composed of protons and neutrons, surrounded
by electrons, orbiting in shells
The center of an atom is the nucleus, which is composed of protons and neutrons. Their masses are
roughly equal and since the mass of the electron is pretty much negligible, most of the mass of an
atom is in the nucleus.
The electrons are found in a series of energy levels which you call shells at IGCSE. Each 'shell' can
only hold a certain number of electrons, these shells can be thought of as getting progressively
further from the nucleus. Electrons will always go into the lowest possible energy level, provided
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there is space. The first shell can only hold 2 electrons, then the shells after that can hold a
maximum of 8. (i.e. Ca  {2,8,8,2})
In a diagram, the electrons are shown on circles around the nucleus. Beware that these circles are
just imaginary lines to help you understand that the electrons orbit around the nucleus, at IGCSE
level you just need to accept that. Due to Heisenberg’s uncertainty principle, truly the location and
direction of the e- cannot actually be determined and this is covered in higher courses of Chemistry.
http://www.chemguide.co.uk/atoms/properties/atomorbs.html#top
Above, there is an example of a “dot and cross” electronic diagram. Dots or crosses are used to represent electrons, in the above;
Carbon has 4 outer shell electrons. They are drawn far apart even though you could draw them close to each other like in the first
shell; this is because electrons would repel each other as they have the same negative charge. (Remember like charges repel). So
only if you have more than 4 OSE, then do you draw them in pairs. Remember, draw the 4 OSE like in the diagram of carbon
above, and then pair up any OSE left. The following diagram might help you understand:
This way of drawing electrons is clear and makes it easy to count too. When you learn about ions, dot-and-cross diagrams are useful
and they help you see how the electrons are transferred. Like in the above diagram, the Chlorine atom gains one electron (the cross)
from the sodium atom to become a
Chloride
ion (Cl-)
whilst the sodium atom becomes a sodium ion (Na+)
It is positive because it lost one electron, so it has one more proton than electron now. :)
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1.9 recall the relative mass and relative charge of a proton, neutron and electron
Relative mass
Relative charge
Proton
1
+1
Neutron
1
0
Electron
1/1836 (negligible)
-1
1.10 understand the terms atomic number, mass number, isotopes and relative atomic mass (A r)
Atomic number is the number of protons there are in the nucleus, it is sometimes called the proton
number, though atomic number should be more accurate because atoms are electrically neutral, the
number of protons and electrons are equal. (Protons have a charge of +1 whilst electrons are -1, so
they cancel each other out.) So the atomic number tells you the number of protons and the number
of electrons.
Mass number is the number of protons and neutrons. It is sometimes known as the nucleon
number, because protons/neutrons are nucleons. So if a question asks you for the number of
neutrons in an atom, mass number - atomic number = no. of neutrons.
Protons
Neutrons
Mass number
Carbon-12
6
6
12
Carbon-13
6
7
13
Carbon-14
6
8
14
6
Isotopes: these are atoms that have the same atomic number but different mass numbers, i.e. same
number of protons but different number of neutrons.
The number of neutrons in an atom can vary a little. For instance, there are three kinds of carbon
atoms 12C, 13C and 14C. Their number of neutrons varies but they have the same number of protons,
because each element's atomic number is unique. If it has a different number of protons, it wouldn't
be the same element anymore. So these atoms are isotopes of carbon. Bear in mind that the fact
that they have varying numbers of neutrons makes no difference whatsoever to the chemical
reactions of the carbon. Though their physical properties may vary.
Relative atomic mass (web definition): the ratio of the average mass per atom of the naturally
occurring form of an element to one-twelfth the mass of an atom of carbon-12.
Symbol Ar Abbreviation r.a.m.
Essentially, the r.a.m. of an element is the average naturally occurring mass you would find the
element naturally. Thus, you must factor in the fact that different isotopes are prevalent in varying
abundances, by multiplying the percent of overall amount of this specific isotope on earth by the
mass number of the specific isotope. Do this for all isotopes, and average, thereby finding the
average mass of the element prevalent on Earth. It is there as, obviously many elements have
multiple isotopes, and therefore not one specific value for atomic mass. Relative atomic mass solves
this.
Remember that the mass number is always higher than the atomic number
1.11 calculate the relative atomic mass of an element from the relative abundances of its isotopes
You multiply the relative abundance of each isotope by its mass number, add these together, and
divide by 100. It's easier to understand through an example, in this case I'll use chlorine, since it's
pretty common.
and
Chlorine consists of 75% Chlorine-35 and 25% Chlorine-37. You can think of the data as 100 atoms,
75 having a mass of 35 and 25 with a mass of 37. So the calculation is:
[(75 x 35) + (25 x 37)] / 100 = 35.5
So the RAM of chlorine, or Ar(Cl) is 35.
(There are tiny percentages of other chlorine isotopes but the two shown above are the most
common, and so the rest are ignored at IGCSE level.)
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The RAM of an element will be closer to the mass number of the more abundant isotope. For
example, the RAM of chlorine is 35.5, which is closer to chlorine-35, because it is the more abundant
isotope. Obviously 75% > 25%!
1.12 understand that the Periodic Table is an arrangement of elements in order of atomic number
The number of protons in the element's atom increases across the Periodic Table as you've probably
noticed. The proton number defines the element.
1.13 deduce the electronic configurations of the first twenty elements from their positions in the
Periodic Table
To work out the electronic arrangement of an atom:
Look up the atomic number in the Periodic Table - making sure that you choose the right number if
two numbers are given. The atomic number will always be the smaller one and tends to be below
the symbol.
This tells you the number of protons, and hence the number of electrons.
Arrange the electrons in levels, always filling up an inner level before you go to an outer one,
remembering the first shell can contain only 2 electrons.
e.g. to find the electronic arrangement in oxygen
the Periodic Table gives you the atomic number of 8.
Therefore there are 8 protons and 8 electrons.
The arrangement of the electrons will be 2,6. (First shell only holds 2 electrons, then there's 6 left
which occupy the second shell.)
1.14 deduce the number of outer electrons in a main group element from its position in the
Periodic Table
If you look at the patterns in the table:
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The number of electrons in the outer level is the same as the group number. (Except with helium,
which has only 2 electrons. The noble gases are also usually called group 0 - not group 8.) This
pattern extends throughout the Periodic Table for the main groups (i.e. not including the transition
elements).
So if you know that barium is in group 2, it has 2 outer shell electrons (btw, outer shell electrons
which I will abbreviate to OSE are also known as valence electrons); iodine is in group 7, so it has 7
OSE, lead is in group 4, so surprise surprise, it has 4 OSE.
Noble gases have full outer shells. Thus they are unreactive, as they do not need to lose or gain
electrons.
Group 1 - Alkali Metals
All alkali metals react vigorously with cold water. In each reaction,
hydrogen gas is given off and the metal hydroxide is produced.
The speed and violence of the reaction increases as you go down
the group. This shows that the reactivity of the alkali metals increases as you go down group 1.
Lithium
When lithium is added to water, it floats. It fizzes steadily and becomes smaller, until it eventually
disappears.
lithium + water → lithium hydroxide + hydrogen
2Li(s) + 2H2O(l) → 2LiOH(aq) + H2(g)
Sodium
When sodium is added to water, it melts to form a ball that moves around on the surface. It fizzes
rapidly, and the hydrogen produced may burn with an orange flame before the sodium disappears.
sodium + water → sodium hydroxide + hydrogen
2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)
Potassium
When potassium is added to water, the metal melts and floats. It moves around very quickly on the
surface of the water. The hydrogen ignites instantly. The metal is also set on fire, with sparks and a
lilac flame. There is sometimes a small explosion at the end of the reaction.
potassium + water → potassium hydroxide + hydrogen
2K(s) + 2H2O(l) → 2KOH(aq) + H2(g)
Strong alkalis
The hydroxides formed in all of these reactions dissolve in water to form alkaline solutions (e.g.
KOH). These solutions turn universal indicator purple, showing they are strongly alkaline. Strong
alkalis are corrosive.
Why does the reactivity increase down the group? – Higher tier
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All alkali metals have one electron in the outer shell. In a reaction, this electron is lost and the alkali
metal forms a +1 ion. As you go down group 1, the number of electron shells increases – lithium has
two, sodium has three etc. Therefore, the outermost electron gets further from the nucleus. The
attraction from the positive nucleus to the negative electron is less. This makes it easier to remove
the electron and makes the atom more reactive.
Group 1 Compounds:
Aim: to investigate the properties of group 1 compounds.
Using NaCl, Na2CO3, NaOH, Na2SO4, we can conclude that:
o
All group 1 compounds are soluble in water.
o
All group 1 compounds form white compounds.
o
NaOH (being a metal hydroxide) was very alkaline - pH14
o
Na2CO (being a carbonate) was slightly alkaline
o
Na2SO4 & NaCl were both neutral – pH7
Group 1 & 2 Flame Tests:
Method:
o
Dip flame test wire into concentrated HCl acid.
o
Dip flame test wire into sample
o
Heat over Bunsen burner
Results:
Lithium (Li+) – Red flame
Sodium (Na+) – Persistent orange flame
Potassium (K+) – Lilac flame
Calcium (Ca2+) – Brick red
Copper (Cu2+) – Blue green
Reactions of Group 2 Alkali Earth Metals
Mg and Ca will burn in the air, being very reactive, and are thus used often for fireworks.
Mg burns in the air with a brilliant white flame, to form MgO.
MgO is BASIC.
Mg will not react with cold water, due to its position in the reactivity series, but with react with
steam.
Mg + H2O  MgO + H2
MgO is basic too.
Ca burns in the air with a red flame to form CaO, known as quick lime.
It reacts with water to form calcium hydroxide (SLAKED LIME – Ca(OH)2) and will dissolve a little in
water to form lime water.
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Group 2 carbonates and hydroxides are insoluble in water.
Group 2 nitrates and chlorides are soluble in water
Group 2 carbonates thermally decompose to form oxides and carbon dioxide.
Group 1 elements:
• Are metals
• Are soft with melting points and densities very low for metals
• Have to be stored out of contact with air or water
• React rapidly with air to form coatings of the metal oxide
• React with water to produce an alkaline solution of the metal hydroxide and hydrogen gas
• Increase in reactivity as you go down the Group
• Form compounds in which the metal has a 1+ ion
• Have mainly white compounds which dissolve to produce colorless solutions
Group 7 elements- chlorine, bromine and iodine
2.9 recall the colours and physical states of the elements at room temperature
2.10 make predictions about the properties of other halogens in this group
Halogen
Colour
State at Room Temperature
F2
Yellow
Gas
Cl2
Green
Gas
Br2
Red-brown
Liquid
I2
Grey
Solid
At2
Dark coloured
Solid
F2 – Used for toothpaste
Cl2 – Used for sanitation and disinfectant for swimming pools
Br2 – Making pesticides
I2 – Cleaning/ sterilizing wounds
The group 7 elements get darker down the group, so we can deduce that astatine is dark
colored, and is a solid too. As atoms get bigger down groups, their intermolecular forces grow
stronger, so astatine can only be a solid.
2.11 understand the difference between hydrogen chloride gas and hydrochloric acid
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Both hydrogen chloride and hydrochloric acid have the formula HCl. Hydrogen chloride is a gas,
and hydrochloric acid is its solution in water. When hydrogen chloride is dissolved in water, it
forms H+ ions which makes it acidic, as it dissociates.
This is because water, the covalent H2O’s individual atoms aren’t aligned, thus the electrons
spend slightly more time with the O2 atom, making the Oxygen marginally negative, and the
Hydrogen is marginally positive (this is a theta charge), thus H2O is known as a dipole. Thus, the
H+ in HCl is attracted slightly to the negative Oxygen and the Cl- is attracted slightly to the
marginally positive Hydrogen in the water molecules. However, as H+ is formed, it is now acidic.
2.12 explain, in terms of dissociation, why hydrogen chloride is acidic in water but not in
methylbenzene
When hydrogen chloride is dissolved in water, it dissociates (basically, just splits up) to
form H+ ions which are responsible for its acidic properties. But when hydrogen chloride is
dissolved in methylbenzene, that's all that happens. It dissolves. It doesn't dissociate, so it
doesn't form ions. This means that it just exists as HCl molecules, not H+ and Cl- ions. So it is not
acidic.
2.13 recall the relative reactivities of the elements in Group 7
The halogens become less reactive as you go down the group, this means that its oxidising ability
falls as you go down the group. (The halogens are good oxidising agents, this means it takes
electrons away. If it takes electrons away from something else, it means it itself gains electrons.
Try to get your head around that. OILRIG-oxidation is loss, reduction is gain.)
So basically, when a halogen oxidises something, it does so by removing electrons from it.
X2 + 2e- à 2X- (halide ion) They gain an electron to have full outer shells, but that means they
have a negative 1 charge.
Each halogen has the ability to oxidise the ions of those underneath it in the Group, but not
those above it. Chlorine can remove electrons from bromide or iodide ions, and bromine can
remove electrons from iodide ions.
Chlorine is a strong oxidising agent because its atoms readily attract an extra electron to make
chloride ions. Bromine is less successful at attracting electrons, and iodine even less successful.
Why? This is because the 'incoming electron' would be further away from the nucleus as you go
down the group, as the atoms get larger. As there are more electron shells, the 'incoming
electron' is further away, and so it doesn't feel the nucleus attraction as much-so it is less
strongly attracted. So the ion is less readily formed.
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2.14 describe experiments to show that a more reactive halogen will displace a less reactive
halogen from a solution of one of its salts
e.g. if you add chlorine to potassium bromide solution, chlorine would displace the bromide
from its salt.
Cl2 + 2KI  2KCl + I2
Remember the Group 7 elements are diatomic, so it must be 2KI so that when iodine is
displaced, it forms I2.
Remember: Each halogen has the ability to oxidise the ions of those underneath it in the Group,
but not those above it.
2.15 understand these displacement reactions as redox reactions
Redox reactions are basically reactions where one species is being oxidised and one is being
reduced. So the more reactive halogen will remove the electrons, so in the above reaction,
Chlorine oxidises iodine and gains an electron each (chlorine is diatomic) and so the iodide ions
become iodine atoms again.
When I say chlorine oxidises iodine, it means that iodine is oxidised as it LOSES an electron
(OILRIG!), but chlorine is reduced, as it gains an electron. Potassium forms K+ ions and chlorine
forms Cl- ions, so they can form KCl.
It's all a bit confusing sometimes, but always refer to OILRIG. Even though chlorine may be
reduced, it's called an oxidising agent because it oxidises other stuff--taking electrons away from
them.
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Displacement reactions and the reactivity series:
2.30 recall that metals can be arranged in a reactivity series based on the reactions of the
metals and their compounds: lithium, potassium, sodium, lithium, calcium, magnesium,
aluminum, zinc, iron, copper, silver and gold.
A more reactive element will displace a less reactive one from a compound, thus being the reducing
agent. However, in reverse, nothing happens.
Element
Symbol
Reaction with acids
Reaction with water
Reaction with HCl acid
Air
Lithium
Li
As you can see these metals
Reacts with cold water
Violent reaction
Burns to
Potassium
K
(excluding carbon) are above
form oxide
Sodium
Na
hydrogen in the reactivity
but getting
Calcium
Ca
series so they react with acids
Magnesium
Mg
and displace hydrogen gas.
Aluminium
Al
Metal + acid metal salt +
Carbon
C
hydrogen
Zinc
Zn
Iron
Fe
Tin
Sn
Reacts only slowly with
Lead
Pb
steam
Hydrogen
H
Reacts with cold water
less vigorous
Reacts, but getting less
Protected by oxide layer
vigorous
Reacts with steam
Reacts
slowly
H+ ions are responsible for
acidic properties.
Copper
Cu
These elements are below
Doesn’t react with water
No reaction
Silver
Ag
hydrogen so they do not react
nor steam
Gold
Au
with acids. (Acids contain H+
react with
Platinum
Pt
ions)
air
Doesn’t
2.34 understand the terms redox, oxidising agent and reducing agent
A redox reaction is a reaction in which both reduction and oxidation are occurring. They always go
together.
An oxidizing agent is a substance that causes another substance to be oxidized. So it causes
something else to lose electrons, and gains these electrons itself. So the oxidizing agent itself is
reduced. *This confuses people!! Remember that oxidizing agent doesn't get oxidized; don't let the
name fool you.
A reducing agent is a substance that reduces something else. So it causes the substance to gain
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electrons, by losing electrons itself. So the reducing agent is said to be oxidized. It can also be taken
as the reducing agent takes away oxygen from the other substance, such as:
Magnesium + copper (II) oxide  magnesium oxide + copper
5.1 explain how the methods of extraction of the metals in this section are related to their
positions in the reactivity series
Order of reactivity
Symbol
Method of Extraction
Potassium
K
Electrolysis
Sodium
Na
The metal compound is:
Lithium
Li
Calcium
Ca
•
Melted, then
Magnesium
Mg
•
Has electricity passed through it
Aluminium
Al
These metals are very reactive and are above carbon in the reactivity series, so they
cannot be reduced by it. As they are very reactive, the make very stable compounds that
requires a lot of energy to separate into its elements. So electrolysis is used.
Zinc
Zn
Reduction by carbon
Iron
Fe
e.g. ZnO + Ca Zn + CO
Tin
Sn
Or sometimes the carbon monoxide is the reducing agent-here think of reduction as
Lead
Pb
‘taking oxygen away’ to leave pure metal. Carbon is cheap and can also be used as the
source of heat. If the ore is a sulphide, it is roasted first to get the oxide. Roasting is a
process where is basically heating the ore in air.
Copper
Cu
These metals can be found uncombined, as the metal itself because they are very
Silver
Ag
unreactive. We say they are found native. (Copper and silver are often found as ores but
Gold
Au
they are easy to extract by roasting the ore.)
Platinium
Pt
Methods of finding a reactivity series:

Reactions with oxygen: metal + oxygen  metal oxide

Reactions with water: metal + cold water  metal hydroxide + hydrogen
metal + hot water  metal oxide + hydrogen

Reactions with acid: metal + acid  salt + hydrogen

Displacement reactions: where a more reactive metal displaces a less reactive one from a
compound
Because aluminium is above carbon in the reactivity series, it has to extract using electrolysis.
Aluminium oxide however, has a very high melting point and it won't be practical to electrolyse
molten aluminium oxide. Instead, it is dissolved in molten cryolite. Cryolite is another aluminium
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compound that melts at a more reasonable temperature. So the electrolyte is a solution of
aluminium oxide in molten cryolite at a temperature of about 1000°C.
Extracting Iron – Blast furnace
5.4 describe and explain the main reactions involved in the extraction of iron from iron ore
(haematite), using coke, limestone and air in a blast furnace
Haematite is basically iron oxide, and the oxygen must be removed to leave the iron behind.
Reactions in which oxygen is removed are called reduction reactions. Since carbon is more reactive
than iron, it can displace the iron from its oxide. Hence the method for extraction of iron is called
'reduction by carbon'.
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The iron ore, coke and limestone (the charge) enter the blast furnace at the top. The hot waste
gases at the top of the furnace are piped away and used to heat the air blast at the bottom.
As the coke, which is impure carbon, enters the furnace, it is oxidized by hot air, causing it to burn
in an exothermic reaction that provides a lot of heat for the furnace itself.
C (s) + O2 (g)  CO2 (g)
At high temperatures in the furnace, the carbon dioxide is reduced by more carbon to give carbon
monoxide.
CO2 (g) + C (s)  2CO (g)
It is the carbon monoxide, which is the main reducing agent in the furnace.
Iron ore, or haematite, Fe2O3 is then reduced by the carbon monoxide, leaving iron and carbon
dioxide as the products:
Fe2O3 (s) + 3CO (g)  2Fe (l) + 3CO2 (g)
Due to the high temperatures, the iron produced melts and flows to the bottom of the furnace,
being denser than slag, where it can be tapped off.
Limestone (CaCO3) is added, thermally decomposes in the heat to form carbon dioxide and calcium
oxide – this is an endothermic reaction, thus not too much limestone should be added to the blast
furnace.
CaCO3 (s)  CaO (s) + CO2 (g)
Calcium oxide is a basic oxide, being a metal oxide, and its function is to react with acidic oxides
such as silicon dioxide, SiO2. Silicon dioxide is the main constituent of sand. The product is calcium
silicate, known also as slag, which melts and floats on top of the iron, being less dense. Slag is used
to make roads.
CaO (s) + SiO2 (s)  CaSiO3 (l)
Rusting:
o
Most metals just form a dull coating when exposed to air, as the metal reacts and forms a
compound – this is known as corrosion.
o
Rusting is the name given to corrosion of iron and steel
o
Rusting occurs only when the metal is in contact with both oxygen and water, causing an
orange-brown rust to form called iron oxide.
o
Iron + oxygen + water  hydrated iron oxide
o
Acid or salt will speed up rusting.
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o
Preventing Rusting:
o
Physical barrier – like paint, grease, plastic coating, electroplating, which acts as a physical
barrier for the iron/steel to the outside world.
o
Sacrificial barrier – where a more reactive metal, like Zinc (galvanizing) is attached to the
iron and corrodes instead of the iron as it is more reactive. Galvanizing is the coating of
iron/steel with zinc.
Writing ionic formulae:
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Electrolysis of BRINE (NaCl water)
Terminology:
Cathode  negative electrode
Anode  positive electrode
Salt water (BRINE) is put in a container, and then an electric current is passed through it, at the
negative electrode (cathode) the Hydrogen in the water is attracted to it, thus it breaks its bond in
H2O, leaving OH- behind and goes to the oppositely charged cathode and evaporates as Hydrogen
gas. The hydrogen becomes H+ also partially due to Na being more reactive than it when Cl- leaves
NaCl to go to the positive anode (before evaporating as chlorine gas) and therefore displacing H+
from H2O to form the alkaline NaOH at the bottom.
2Cl- (aq)  Cl2(g) + 2e2H+  2e- + H2
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Acids
When acids dissolve in water they produce hydrogen ions, H+. For example, looking at hydrochloric
acid:
HCl(aq) → H+(aq) + Cl-(aq)
Remember that (aq) means the substance is in solution.
Alkalis
When alkalis dissolve in water they produce hydroxide ions, OH-. For example, looking at sodium
hydroxide:
NaOH(aq) → Na+(aq) + OH-(aq)
Ammonia is slightly different. This is the equation for ammonia in solution:
NH3(aq) + H2O(l) →
(aq) + OH-(aq)
Be careful to write OH- and not Oh-.
Neutralization reaction
When the H+ ions from an acid react with the OH- ions from an alkali, a neutralisation reaction
occurs to form water. This is the equation for the reaction:
H+(aq) + OH-(aq) → H2O(l)
If you look at the equations above for sodium hydroxide and hydrochloric acid, you will see that
there are Na+ ions and Cl- ions left over. These form sodium chloride, NaCl.
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