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

Lose their valence electrons easily

Form oxides that are basic

Form basic hydroxides

Are good reducing 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

Usually have 4-8 electrons in their outer shell

Gain or share (only in covalent bonds)

valence electrons

Form oxides that are acidic

Are good oxidizing agents

Physical Properties

Metals Non-metals

Good electrical and heat conductors

Malleable—can be hit and shaped

Ductile—can be stretched into wire

Possess metallic luster (shiny)

Opaque as thin sheet

Solid at room temperature (except Mercury [Hg]-

liquid)

Poor conductors of heat and electricity

Brittle- if a solid

Non-ductile

Do not possess metallic luster

Transparent as a thin sheet

Solids, liquids or gases at room temperature

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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

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:

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Positive ions/Cations Negative ions/AnionsCharge Name of ion Formula Charge Name of ion Formula1+ Ammonium

Copper (I)HydrogenLithiumPotassiumSilverSodium

NH4+

Cu+

H+

Li+

K+

Ag+

Na+

1- BromideChlorideHydroxideFluorideIodideNitrateHydrogencarbonate

Br-

Cl-

OH-

F-

I-

NO3-

HCO3-

2+ BariumCalciumCopper (II)Iron (II)Lead (II)MagnesiumNickel (II)StrontiumZinc

Ba2+

Ca2+

Cu2+

Fe2+

Pb2+

Mg2+

Ni2+

Sr2+

Zn2+

2- CarbonateSulphateSulphiteSulphideOxideCarbonateSulphate

CO32-

SO42-

SO32-

S2-

O2-

CO32-

SO42-

3+ AluminiumIron (III)

Al3+

Fe3+3- Nitride

PhosphateN3-

PO43-

Anion Test

Chloride

Cl-Dissolve in dilute nitric acid (HNO3), then a white precipitate forms when silver nitrate solution added. (The

precipitate dissolves in dilute ammonia solution.)

Bromide

Br –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)

Iodide

I-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.

Ammoniu

m

NH4+

Add sodium hydroxide solution to the ammonia (NH3 right now) and heat with Bunsen burner, and then test

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).

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 Result of adding sodium hydroxide solution

copper(II)

Cu2+pale blue precipitate

iron(II)

Fe2+dirty green precipitate

iron(III)

Fe3+orange brown precipitate

2.40 describe simple tests for the gases:

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• 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)

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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

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

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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 (Ar)

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

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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

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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.)

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.)

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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:

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

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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

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

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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.

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

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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

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

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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.

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.

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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.

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 Symbo

l

Reaction with acids Reaction with water Reaction with HCl acid Air

Lithium Li As you can see these metals

(excluding carbon) are above

hydrogen in the reactivity

series so they react with acids

and displace hydrogen gas.

Metal + acid metal salt +

hydrogen

Reacts with cold water

Reacts with cold water

Violent reaction Burns to

form oxide

but getting

less vigorous

Potassium K

Sodium Na

Calcium Ca

Magnesium Mg Reacts, but getting less

vigorousAluminium Al Protected by oxide layer

Carbon C Reacts with steam

Zinc Zn

Iron Fe Reacts

slowlyTin Sn Reacts only slowly with

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steamLead Pb

Hydrogen H H+ ions are responsible for

acidic properties.

Copper Cu These elements are below

hydrogen so they do not

react with acids. (Acids

contain H+ ions)

Doesn’t react with water

nor steam

No reaction

Silver Ag Doesn’t

react with

airGold Au

Platinum Pt

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

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

The metal compound is:

• Melted, then

• Has electricity passed through it

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.

Sodium Na

Lithium Li

Calcium Ca

Magnesium Mg

Aluminium Al

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Zinc Zn Reduction by carbon

e.g. ZnO + Ca Zn + CO

Or sometimes the carbon monoxide is the reducing agent-here think of reduction as

‘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.

Iron Fe

Tin Sn

Lead Pb

Copper Cu These metals can be found uncombined, as the metal itself because they are very

unreactive. We say they are found native. (Copper and silver are often found as ores but

they are easy to extract by roasting the ore.)Silver Ag

Gold Au

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

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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

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) + 2e-

2H+ 2e- + H2

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25

26

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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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