Complete Chemistry Revision Guide IGCSE
Complete Chemistry Revision Guide IGCSE
Complete Chemistry Revision Guide IGCSE
ENDORSED BY
THE ROYAL SOCIETY OF CHEMISTRY AND CAMBRIDGE EXAMINATIONS BOARD.
CONTENTS
[ANDREW RICHARD WARD ALL RIGHTS RESERVED A WARD TUITION - 2009]
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PRINCIPLES OF CHEMISTRY
1.
2.
3.
4.
PHYSICAL CHEMISTRY
5.
6.
7.
8.
INORGANIC CHEMISTRY
9. THE PERIODIC TABLE
10. METALS
11. AIR AND WATER
12. SULPHUR
13. CARBONATES
ORGANIC CHEMISTRY
14.ORGANIC CHEMISTRY
INTRODUCTION
[ANDREW RICHARD WARD ALL RIGHTS RESERVED A WARD TUITION - 2009]
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CHANGES OF STATE
Solid to liquid = melting
Liquid to gas = boiling or evaporation
Gas to liquid = condensation
Liquid to solid = freezing
Sometimes a solid may change directly into a gas missing out the liquid stage.
This is called SUBLIMATION.
Iodine is a black solid. It sublimes to form a purple gas.
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Here is a table to show the melting point and boiling point of some chemical
substances
SUBSTANCE
Aluminium
MELTING
BOILING
POINT(CELSIUS) POINT
(CELSIUS)
661
2467
Ethanol
-117
79
Magnesium oxide
2827
3627
Mercury
-30
357
Methane
-182
-164
Oxygen
-218
-183
EXPLANATION
SOLID IT HAS
NOT MELTED
YET
LIQUID
MELTED AT
LOW
TEMPERATURE
SOLID NOT
YET MELTED
LIQUID
ALREADY
MELTED
GAS
ALREADY
MELTED AND
BOILED
GAS
ALREADY
MELTED AND
BOILED
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BROWNIAN MOTION
Robert Brown discovered this in 1827. This theory explains movement of particles
in liquids. Brown discovered that pollen grains moved on the surface of water
when he looked at them through a microscope. The grains were moving in RAPID
RANDOM MOTION. This was later called BROWNIAN MOTION.
Here is a photograph of the Brownian Motion of particles when photographed
under a microscope. You will see that the motion is rapid and random.
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Shows a pattern:
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GAS LAWS
This topic is studied in detail in Physics but is shown here for you to understand
the behavior of gases in more detail:
There are 2 gas laws which are important. They are named after the scientists who
made their discoveries:
BOYLES LAW
When we keep temperature the same, the volume of a fixed mass of gas is
inversely proportional to the pressure.
This means
Large volume of gas = low pressure of gas
Small volume of gas = high pressure of gas
CHARLES LAW
When we keep the pressure the same, the volume of a given mass of gas is directly
proportional to the temperature.
This means
Large volume of gas = high temperature of gas
Small volume of gas = low temperature of gas
This is the end of chapter one.
A checklist of definitions is shown on the next page.
You must be able to write them down by memory for your examinations.
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TOPIC 1 CHECKLIST
ABSOLUTE TEPERATURE
This is the temperature measured with respect to absolute zero (zero Kelvin) on the
Kelvin Scale.
Temperature in Kelvin = Temperature in Celsius + 273
ATMOSPHERIC PRESSURE
This is the pressure of the atmosphere on the surface of the Earth due to the weight
of the air.
BOILING POINT
The temperature at which gas pressure above a liquid equals atmospheric pressure
BOYLES LAW
At a constant temperature, the volume of a given mass of gas is inversely
proportional to the pressure.
V=1/p
CHARLES LAW
At constant pressure, the volume of a given mass of gas is directly proportional to
absolute temperature.
CONDENSATION
This is the change of vapour of a gas into a liquid. This process involves heat being
produced.
DIFFUSION
This is the process by which different substances mix because of the rapid random
motion of their particles
EVAPORATION
This is a process that occurs at the surface of a liquid and involves the change of
state of a liquid into a vapour at the temperature below boiling point.
[ANDREW RICHARD WARD ALL RIGHTS RESERVED A WARD TUITION - 2009]
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KINETIC THEORY
A theory which accounts for the properties of materials in terms of constituent
particles
MATTER
Anything that occupies space and has a mass
MELTING POINT
This is the temperature at which a solid starts to turn liquid. Pure substances have a
sharp, defined melting point.
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METAL
Usually solid
(occasionally liquid like
mercury)
God
Good
Shiny
Usually high
Usually high
Usually high
Good
NON-METAL
Solid, liquid or gas
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ATOMS
Everything in the universe is made from billions of atoms.
Atoms are too small to be seen by the eye.
The smallest atom- hydrogen atoms are 0.00000007 mm wide.
Chemists use shorthand symbols to label elements and their atoms.
Usually the first or first two letters of the name of the element are used.
Some elements that were discovered many years ago still have Latin names like
Sodium Na Latin name Natrium
Lead Pb Latin name Plumbum
MOLECULES
The atoms of some elements are joined together in small groups called molecules.
Hydrogen, oxygen, nitrogen, fluorine, chlorine, bromine and iodine have atoms
that are joined in pairs. They are known as DIATOMIC MOLECULES.
A phosphorus has 4 atoms joined together, a sulphur molecule has 8 atoms joined
together.
Gases like helium, neon, argon, bromine, krypton and xenon are composed of
separate individual atoms MONOATOMIC MOLECULES.
COMPOUNDS
Compounds are pure substances formed when two or more elements chemically
combine together. Water is an example of a compound.
Water is made from two elements hydrogen and oxygen.
Hydrogen (a pure element) + oxygen (a pure element)
forms water (a pure compound formed from hydrogen gas burning in oxygen gas)
Water contains two atoms of hydrogen and one atom of oxygen to give water a
chemical formula .
Elements other than hydrogen will react with oxygen gas to form chemical
compounds called OXIDES.
[ANDREW RICHARD WARD ALL RIGHTS RESERVED A WARD TUITION - 2009]
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In this case, ethanol (which is alcohol, basically) reacts with ethanoic acid (the
main constituent of vinegar) to form ethyl ethanoate and water. However, the ethyl
ethanoate produced reacts with the water produced to recreate the ethanol and
ethanoic acid again. In practice, the chemicals reach a balance point, called
equilibrium where all four chemicals are present.
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2O3
Note that oxygen gas is diatomic, which means that the oxygen atoms, like
policemen, go around in pairs. A molecule of aluminium oxide consists of two
aluminium atoms combined with three oxygen atoms. Actually, technically the
word "molecule" is inappropriate in that previous sentence. The formula simply
tells us the ratio of aluminium atoms to oxygen atoms in the compound. In the
solid state, the atoms form a giant structure called a crystal lattice rather than
individual discrete molecules. When balancing chemical equations, people often
refer to the number of species on each side to avoid this problem.
You can see by looking at it that there is something wrong with this equation. If
you count the number of atoms of each type on each side, you will see that there is
only one aluminium atom on the left side whereas there are two on the right. There
are two oxygen atoms on the left side, as compared to three on the right side. This
clearly doesn't match.
Left side:
Right side:
We can balance the equation by mutiplying the different atoms and molecules on
each side by different amounts. Firstly, multiply the aluminium atoms on the left
side by 2:
Al + O2
Left
side:
2O3
Right
side:
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Now there are the same number of aluminium atoms on each side of the equation.
We could also multiply the number of oxygen molecules on each side by one and a
half (1.5), which would give three oxygen atoms on the left side (1.5 x 2 = 3) to
match the three oxygen atoms on the right side:
2 Al + 1.5 O2
Left
side:
2O3
Right
side:
This is now balanced, but that 1.5 is a horrible thing to have in an equation - how
can you have one and a half molecules? We can solve this problem by multiplying
everything throughout by 2:
Al +
Left
side:
O2
Al2O3
Right
side:
If you count the number of atoms on each side, you will find that there are four
aluminium atoms on each side and six oxygen atoms. Sorted!
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Another Example
Here's another equation:
Ethane is a gas similar to methane (town gas or natural gas) which burns in oxygen
to give carbon dioxide gas and steam. The steam is simply water in gaseous form
and condenses to form water droplets. Here is the chemical equation rewritten with
the chemical symbols:
C2H6 + O2
+ H2O
Neither the carbon, nor the oxygen atoms nor the hydrogen atoms match. Let's look
at the carbon atoms first. There are two carbon atoms on the left side, but only one
on the right, so we need to put a 2 in front of the carbon dioxide molecule to give
two carbons on each side:
C2H6 + O2
CO2 + 3 H2O
Now we will look at the hydrogen atoms. There are six hydrogen atoms on the left
side and two on the right side, so we treble the number of water molecules on the
right side:
C2H6 + O2
H2O
Now there are two carbon atoms on each side, and six hydrogen atoms on each
side, but the oxygen atoms don't match. There are 2 of them on the left side and 7
on the right side. This is easily solved by multiplying the oxygen molecule on the
left side by 3.5 (as 2 x 3.5 = 7):
C2H6 + 3.5 O2
+ 3 H2O
This gives 2 carbons, 6 hydrogens and 7 oxygens on each side of the equation. The
equation is balanced, but rather inelegant since it contains a decimal. Just double
all the figures in the equation:
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C2H6 +
O2
CO2 +
H2O
The equation has been balanced. You will notice that we left the oxygen atoms
until last. This was deliberate, as oxygen was present on one side of the equation as
an element (i.e. on the left side of the equation there is oxygen present in an
element, not in a compound).
Treat standard groups as an item
You may recognise some standard parts of molecules, ... erm, sorry, species, ... as
being a unit. For instance all sulphates contain the group of atoms SO4. These may
be doubled (or even trebled) if necessary. Some examples of sulphates are shown
below:
Iron(II) sulphate
FeSO4
Sodium sulphate
Na2SO4
Iron (III) sulphate
Fe2(SO4)3
Lead (IV) sulphate
Pb(SO4)2
You will notice that iron forms two sulphates, depending on its oxidation state.
Being a transition metal, it can form different types of ion, Fe2+ and Fe3+ in this
case. Lead also forms different ions, but I have just quoted one of its sulphates. To
show that the sulphate ion is a single group, it is usually included in brackets when
it has to be doubled, so iron (III) sulphate is generally written as Fe2(SO4)3 rather
than Fe2S3O12.
If you can recognise a standard group, such as suphate, phosphate, nitrate etc., then
you should treat it as an indivisible item. It isn't essential to do this, i.e. you can
still balance the equation successfully even if you treat each atom individually, but
treating groups as special items makes life a little easier.
If you are in any doubt, you could temporarily replace the group with a neutral
letter such as X (which is not the symbol of a chemical element). Once the
equation is balanced, put the group back into place, remembering to insert brackets
if necessary. Take the reaction where iron (III) oxide is put in sulphuric acid:
Fe2O3 + H2SO4
2(SO4)3
+ H2O
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2X3
+ H2O
2X3
+ 3 H2O
2(SO4)3
+ 3 H2O
N.B. The approach of treating standard groups as an item only works if those
groups remain unscathed throughout the reaction. If you find that a sulphate
species is broken up (perhaps into an oxide of sulphur), then you can't use this
approach. This is why balancing chemical equations is so much easier if you have
some knowledge of the reactions going on.
Balance the elements last!
You should leave the elements that appear as elements anywhere in the equation
until last. This is because you can balance these elements without affecting any
other elements. Here's an example:
Under certain circumstances, carbon dioxide can be made to react with hydrogen
gas to produce methane and water vapour (which can be electrolysed to produce
oxygen and hydrogen - what a way to produce fuel!)
CO2 + H2
+ H2O
Let's do this the wrong way - let's balance the hydrogen first! There are two
hydrogen atoms on the left (present in the hydrogen molecule) and six on the right,
so we put a 3 in front of the hydrogen molecule on the left:
CO2 +
H2
+ H2O
Now there are six hydrogen atoms on each side. The carbon atoms are balanced,
one on each side, so we only have to balance the oxygen atoms.
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There are two on the left side, and one on the right side. Better put a 2 in front of
the water vapour molecule on the right side:
CO2 + 3 H2
H2O
But now the hydrogens are unbalanced again! We either have to increase the
number in front of the hydrogen molecule on the left side or add more methane
molecules on the right side. Either way, putting a number in front of the water
vapour has changed both the hydrogen and the oxygen.
The proper way to do it would be to balance the carbons and oxygens and then the
hydrogens. Here's the original equation:
CO2 + H2
+ H2O
The carbons are balanced so let's concentrate on the oxygens. There are two on the
left and one on the right, which is easily remedied:
CO2 + H2
H2O
The only element which isn't balanced is hydrogen, which can be balanced without
affecting any other elements. There are now eight hydrogen atoms on the right side
and only two on the left, so we need to multiply the hydrogen on the left by 4:
CO2 +
H2
+ 2 H2O
Now all the elements are balanced, and we didn't have to rebalance anything we
had previously balanced.
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No! No!
2. Don't worry if the numbers turn out to be fractions - you can always double
or treble all the numbers at a later stage.
1
/ 3 H2O
(s)
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Hydrogen and nitrogen react together to produce ammonia gas (note that the
reaction is a reversible one - ammonia also breaks up to form hydrogen and
nitrogen):
H2 +
N2
NH3
O2
CO2 +
H2O
When heated, aluminium reacts with solid copper oxide to produce copper metal
and aluminium oxide:
Al +
Cu
When sodium thiosulphate solution is mixed with brown iodine solution, the
mixture rapidly becomes colourless as the iodine is converted to colourless
sodium iodide:
I2 +
Al2O3 +
Na2S2O3
NaI +
Na2S4O6
Potassium oxide is not a stable compound. In the presence of water (or even
water vapour in the air), it readily converts into potassium hydroxide:
K2O +
H2
KOH
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a double salt consisting of a trivalent metal ion and a group I metal ion. In
this case the alum is potassium iron(III) thiocyanate. The thiocyanate ion is
formed from a carbon and a nitrogen atom (the standard cyanide ion)
together with a sulphur atom:
Fe2(SO4)3 +
KSCN
CO2 +
H2O
O2
H2O
NH3 +
Here's an easy one, the burning of hydrogen gas in oxygen to form steam
(which then condenses to form water).
H2 +
K2SO4
K3Fe(SCN)6 +
P4
Mg3P2
When calcium chloride reacts with silver nitrate solution, a white precipitate
(solid) of silver chloride appears. This is because silver chloride is insoluble:
when silver ions and chloride ions find themselves together in solution, they
immediately react together to form the solid. This leaves the calcium ions
and nitrate ions in solution, effectively forming calcium nitrate solution.
CaCl2 +
AgNO3
AgCl +
Ca(NO3)2
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+ 4 H2C3H8 + 5 O2
2O3
+ 4 H2O
+ 3 Cu
I2 + 2 Na2S2O3
2S4O6
K2O + H2
Fe2(SO4)3
3Fe(SCN)6
(NH4)2CO3
2 H2 + O2
6 Mg + P4
CaCl2 + 2 AgNO3
+ 3 K2SO4
+ CO2 + H2O
2O
3P2
3)2
MIXTURES
A mixture contains more than one substance such as elements and compounds or
elements or compounds.
It is important to understand the DIFFERENCE between MIXTURES AND
COMPOUNDS
We can understand this by looking at the reaction between iron powder and
sulphur.
A mixture of iron powder an sulphur looks completely different from the
individual elements.
The mixture has the individual properties of iron and sulphur for instance we
can use a mangnet to remove all the iron powder from the iron and sulphur
mixture.
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APPEARANCE
EFFECT OF A
MAGNET
IRON
Attracted to it
SULPHUR
Yellow powder
None
IRON/SUPLHUR
MIXTURE
IRON (II)
SULPHIDE
Dirty yellow
powder
Dark grey solid
Iron powder
attracted
No effect
EFFECT OF
DILUTE
HYDROCHLORIC
ACID
Very little when
cold. When warm,
gas made with lots
of bubbles
No effect when hot
or cold
Iron powder reacts
as above
A foul smelling gas
is produced
Here is a very important table that summarizes the differences between mixtures
and compounds
MIXTURE
Contains two or more substances
Composition can vary
No chemical change takes place when
mixture is formed
The properties are those of individual
elements
The components can be easily separated
by physical means
COMPOUND
It is a single substance
The composition is always the same
When the new substance is formed it
always involves a chemical change
The properties are very different to
those of the component elements
The components can only be separated
by one or more chemical reactions
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SEPARATING MIXTURES
Many mixtures contain useful substances mixed with unwanted material such as
impurities. We therefore have to separate useful substances from unwanted
substances.
The separation technique that we use depends on what is in the mixture and also
the properties of the substances present. It also depends on whether the substances
to be separated are solid, liquid or gas.
SEPARATING SOLID/LIQUID MIXTURES
If a SOLID substance is SOLUBLE it will DISSOLVE in a LIQUID and form a
SOLUTION. The solid that an dissolve in the liquid is called a SOLUTE.
The liquid that has the power to dissolve the solid is called the SOLVENT.
Example: sugar dissolves in water when you make a cup of tea or coffee.
Sometimes, the solid does not dissolve in the liquid and is INSOLUBLE.
Example: tea leaves themselves do not dissolve in water.
FILTRATION
This is used when an insoluble solid needs to be separated from a liquid. Sand can
be separated from a mixture with water by filtering through a filter paper.
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The filter contains microscopic holes that allow the small water molecules through
but trap all the larger sand molecules. It acts like a sieve. The sand is called the
RESIDUE on the filter paper and the FILTRATE is the liquid that is allowed to
pass through the filter paper.
DECANTING
Let us think about cooking rice. It is easy to separate the cooked rice from the
water by puring off all the water. This is called decanting.
Decanting is often used to separate an insoluble solid (that may have settled at the
bottom of a container) from a liquid.
Here you can see a clear liquid being decanted from an undissolved solid in the
conical flask.
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CENTRIFUGING
This can separate a solid from a liquid also. The technique is often used instead of
filtration. It is often used when the SOLID Particles are too small that they
SPREAD OUT in the solution and form a SUSPENSION.
They DO NOT SETTLE to the bottom of the container (as heavier particles would
do under the force of gravity).
The technique of centrifuging involves the suspension being SPUN AROUND
VERY FAST in a centrifuge so that the SOLIDS GET FLUNG TO THE
BOTTOM OF THE TUBE.
The pure liquid is decanted after the solids have been forced to the bottom of the
tube. This method is extensively used to separate PLASMA FROM BLOOD
CELLS.
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EVAPORATION
If the solid has dissolved in the liquid we cannot filter or use a centrifuge.
We heat the liquid so that the liquid evaporates and leaves the solid behind.
This technique is commonly used to obtain salt from salty water.
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CRYSTALLISATION
In Jordan, the Dead Sea contains enormous amounts of dissolved salt in water so
much that you can float in the Dead Sea.
We can obtain salt from salt water by using the heat of the sun to evaporate the
water to leave a SATURATED SALT SOLUTION called BRINE.
This is good at preserving food. Tuna fish is often stored in brine.
A SATURATED SOLUTION IS A SOLUTIO THAT CONTAINS AS MUCH
DISSSOLVED SOLUTE AS POSSIBLE AT ANY GIVEN TEMPERATURE.
When the solution is saturated, the salt begins to CRYSTALLIZE and can be
REMOVED.
Here is a salt evaporation pond
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SIMPLE DISTILLATION
If we want to get a solvent from a solution, we carry out simple distillation using
the above apparatus. We can use simple distillation to obtain pure water from salt
water.
The solution is heated in the flask.
The steam rises to the condenser where it condenses back into water again.
The salt is left behind in the flask.
This is done on a large scale in desert countries like Saudi Arabia to obtain pure
water for drinking. This is called DESALINATION.
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The water condenses and moves back down into the conical flask. The ethanol
vapour moves up the column and into the condenser. The ethanol vapour returns to
ethanol liquid and isc collected in the conical flask at the end.When all of the
ethanol has been separated, the temperature steadily rises to 100 Celsius. This
means steam now enters the condenser. We changed the conical flask and collect
the pure water that condenses over.
Fractional distillation can also be used to obtain pure gases from liquid air.
In the diagram, the iodine sublimes and leaves pure salt in the bottom of the
beaker.
CHROMATOGRAPHY
This process is commonly used when we have to separate a mixture of coloured
materials like inks and dyes.
There are many types of chromatography. The simplest type of chromatography is
PAPER CHROMATOGRAPHY.
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PAPER CHROMATOGRAPHY
Here is an example separating the colours that make black ink.
A spot of the ink is placed onto a piece of chromatography paper . the paper is then
put into a suitable solvent such as water.
As the solvent moves up the paper, the dyes become carried with it and begin to
separate. They separate due to them having differences in solubility in the solvent.
They are absorbed in different amounts by the chromatography paper. They
separate as they move up the chromatography paper. The end product of
chromatography is called a CHROMATOGRAM.
The substances on which you perform chromatography do not need to be coloured.
Colourless substances are made visible by covering them in a LOCATING
AGENT. The locating agent will react with the colourless substances to form a
coloured product.
Sometimes a type of chromatography is used which separates out substances due to
differences in CHARGE. This process is known as ELECTROPHORESIS and can
be used in forensic science to separate samples of proteins.
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SOLVENT EXTRACTION
Sugar can be obtained from crushed sugar cane by adding water. The water
dissolves the sugar from the sugar cane. This is called SOLVENT EXTRACTION.
Also some substances present in grass such as chlorophyll can be removed
from crushed grass by using a powerful solvent called ethanol.
CRITERIA OF PURITY
Drugs and pharmaceuticals must be made with an extremely high degree of purity.
To do this, the drugs are dissolved in a suitable solvent and then have fractional
distillation performed on them.
Also, it is illegal to put anything harmful into food.
To make sure that a susbstance is pure we use the following things:
1. MELTING POINT If the substances is pure, it will have a sharp,
defined melting point.
2. BOILING POINT If the substance is pure, the substance will remain
steady at its boiling point and the temperature will not rise.
3. CHROMATOGRAPHY If it is a pure substance, it will produce
only one well-defined spot on the chromatogram.
GELS,SOLS,FOAMS AND EMULSIONS
These are all examples of mixtures which are formed by mixing two substances
which cannot actually mix. These mixtures are often referred to as COLLOIDS.
Colloids are formed by millons of suspended particles.
Generally, colloids cannot be separated by filtration as the size of the dissolved
particles is usually smaller than the holes/pores in the filter paper.
Fruit jelly and custard are examples of gels
Emulsion paint is an example of a sol
Foam could be shaving foam
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The GRP is very light (like plastic) but has strength and flexibility like the glass
fibres.
Bones in the human body are also composite materials.
They are formed from protein strands called collagen mixed with a hard mineral
called calcium phosphate.
Wood is also a composite material. It is made from cellulose fibres mixed with
lignin. Lignin makes the wood strong.
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SYMBOL
Proton
Neutron
Electron
e
n
e
RELATIVE
MASS
1
1
1/1840
CHARGE
+1
0
-1
NOTE: the masses of All the sub-atomic particles are measured in atomic mass
units (amu). This is because they are so light that their masses cannot be measured
usefully in grams.
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Here is the information about the element Krypton in the periodic table.
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ISOTOPES
Isotopes are ATOMS OF THE SAME ELEMENT THAT HAVE DIFFERENT
NUMBERS OF NEUTRONS.
Isotopes of Carbon
Isotopes of an element are atoms of the element that have different numbers of
neutrons in their nuclei. Carbon has three naturally occurring isotopes, which are
shown here with the isotopes of hydrogen. The isotopes of carbon are carbon-12,
which constitutes 98.89 of all carbon atoms and serves as the standard for the
atomic mass scale; carbon-13, which is the only magnetic isotope, making it very
important for structural studies of compounds containing carbon; and carbon-14,
which is produced by cosmic rays bombarding the atmosphere. Carbon-14 is
radioactive, with a half-life of 5760 years. The amount of carbon-14 remaining in
historical artifacts can be used to estimate their age.
In the above 3 examples, all the isotopes have 6 protons and 6 electrons.
Carbon-12 has 6 neutrons
Carbon-13 has 7 neutrons
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Most periodic tables contain a zigzag line which allows you to identify which
elements are metals, nonmetals, and metalloids.
Metals
Most elements are metals. 88 elements to the left of the stairstep line are metals or
metal like elements.
Physical Properties of Metals:
Luster (shininess)
Good conductors of heat and electricity
High density (heavy for their size)
High melting point
Ductile (most metals can be drawn out into thin wires)
Malleable (most metals can be hammered into thin sheets)
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Nonmetals
Nonmetals are found to the right of the stairstep line. Their characteristics are
opposite those of metals.
Physical Properties of Nonmetals:
Since metals tend to lose electrons and nonmetals tend to gain electrons, metals
and nonmetals like to form compounds with each other. These compounds are
called ionic compounds. When two or more nonmetals bond with each other, they
form a covalent compound.
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Metalloids
Elements on both sides of the zigzag line have properties of both metals and
nonmetals. These elements are called metalloids.
Solids
Can be shiny or dull
Ductile
Malleable
Conduct heat and electricity better than nonmetals but not as
well as metals
ALLOYS
An alloy is formed when a metal is mixed with other elements
Alloy
Typical
composition
Special
properties
Some Uses
Steel
Fe(99%),
C(1%)
Duralumin
Al(95%),
Cu(4%),
Mn(0.5%),
Mg(0.5%)
Increased
hardness and
tensile strength,
light weight
Aircraft industry
Bronze
Cu(85%),
Sn(15%)
Stronger than
copper, corrosion
resistant
Weapons and
tools, coinage,
decorations
Fe(85%),
Cr(14%),
Ni(1%)
Cutlery,
domestic
Corrosion resistant appliances,
furnace parts,
nuclear reactors
Stainless
Steel
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Pewter
Sn(85%),
Stronger than tin,
Cu(7%),
but still easy to
Bi(6%), Sb(2%) etch and engrave
Manganin
Cu(81%),
Mn(15%),
Ni(4%)
High resistivity
andlow
temperature
coefficient
Nichrome
Cr(80%),
Ni(20%)
Cupronickel
Cu(75%),
Ni(25%)
Attractive
appearance for
coins, very
ductile(looks like
silver)
Solder
Pb(67%),
Sn(33%)
Brass
Cu(60%),
Zn(40%)
Dental
amalgam
Sn(44%),
Hg(33%),
Ag(22%)
Resistant to
corrosion from the
Braces, dental
acidic products
work
excretedby mouth
bacteria
Domestic
utensils, jewelry
Resistors
Coinage
Page 57
Only atoms with complete electron shells tend to be unreactive like the noble gases
in group 8 of the periodic table.
When atoms combine, they try to achieve full outer electron shells. They do this
either by gaining electrons to fill the gaps in their outer shell to make a full outer
shell or lose electrons to leave a full shell behind.
IONIC BONDING
Ionic bonding involves ELECTRON TRANSFER between metals and non-metals
to form FULL ELECTRON SHELLS.
Metal elements LOSE ELECTRONS and form POSITIVE METAL IONS. This is
called OXIDATION.
Non-metal elements GAIN ELECTRONS and form NEGATIVE NON-METAL
IONS. This is called REDUCTION.
The ions are HELD TOGETHER by strong electrical forces called
ELECTROSTATIC FORCES.
The bonding process can be represented by DOT AND CROSS DIAGRAMS.
Here is the ionic bonding diagram for sodium chloride.
+
sodium metal
chlorine gas
table salt
Page 58
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The chlorine is no longer an atom. It is also an ION. It now does not have an equal
number of protons and electrons. It now has ONE MORE ELECTRON than
protons and is a CHLORIDE ION WITH A -1 CHARGE.
Metals can also transfer more than one electron to a metal.
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Again, noble gas structures are formed, and the magnesium oxide is held together
by very strong attractions between the ions. The ionic bonding is stronger than in
sodium chloride because this time you have 2+ ions attracting 2- ions. The greater
the charge, the greater the attraction.
The formula of magnesium oxide is MgO.
calcium chloride
This time you need two chlorines to use up the two outer electrons in the calcium.
The formula of calcium chloride is therefore CaCl2.
potassium oxide
Again, noble gas structures are formed. It takes two potassiums to supply the
electrons the oxygen needs. The formula of potassium oxide is K2O.
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+2
+3
-3
-2
-1
COVALENT BONDING
This type of bonding involves the sharing of electrons. It only occurs in atoms of
non-metal elements. It usually forms a molecule. The non-metal atoms SHARE
electrons to get a stable FULL OUTER SHELL.
A single covalent bond is formed between TWO NON-METAL ELECTRONS.
Hydrogen gas exists as TWO HYDROGEN ATOMS JOINED TO FORM A
MOLECULE.
Each hydrogen atom wants to fill its electron shell. They do this by sharing
electrons.
Hydrogen
Hydrogen atoms only need two electrons in their outer level to reach the noble gas
structure of helium. Once again, the covalent bond holds the two atoms together
because the pair of electrons is attracted to both nuclei.
A single covalent bond can be shown by a single line. The formula for a hydrogen
molecule can be written as H-H.
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The hydrogen and oxygen atoms in water are also held together by covalent bonds.
Some molecules can form DOUBLE COVALENT BONDS like in the covalent
molecule CARBON DIOXIDE
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On the following page, you will see the covalent bonding in an alcohol called
methanol. Things et slightly more complicated, but the rules of covalent bonding
remain the same.
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GROUP IN 1
PERIODIC
TABLE
NUMBER
X
OF
COVALENT
BONDS
FORMED
Page 67
The properties of sodium chloride can be explained just by looking at the diagram
above.
PROPERTIES OF SODIUM
CHLORIDE
Hard crystals
High melting point of 801 Celsius
Dissolves in water
EXPLANATION IN TERMS OF
STRUCTURE
Strong forces between the ions
Strong forces between the ions
The water is also able to form strong
electrostatic attractions with the ions
Does not conduct electricity when solid Strong forces between the ions stop
them from moving about
Does conduct electricity when molten or The strong forces between the ions have
dissolved in water
been broken down so now the ions are
free to move about
COVALENT MOLECULES
Covalent bonds are also strong bonds. They are INTRAMOLECULAR BONDS
that are formed WITHIN each molecule. Much weaker INTERMOLECULAR
FORCES attract the individual molecules towards each other.
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In the diagram on the other page, the hydrogen atoms have a SLIGHTLY
POSITIVE CHARGE. The oxygen atoms have a SLIGHT NEGATIVE CHARGE.
The bonds that form BETWEEN the water molecules are called HYDROGEN
BONDS and are INTERMOILECULAR forces.
MACROMOLECULES
Some covalently bonded compounds do not exist as simple molecular structures in
the way that hydrogen does. Diamond exists as a GIANT COVALENT
STRUCTURE with each carbon atom bonded to four others.
Diamond is a FORM of the element CARBON.
DIFFERENT FORMS of the SAME ELEMENT are called ALLOTROPES.
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In graphite, carbon atoms form layers of hexagons in the plane of their strong
covalent bonds. The weak bonds are BETWEEN the layers. The layers can SLIDE
OVER EACH OTHER. Graphite is flaky and can be used as a LUBRICANT.
Graphite can be used to conduct electricity (as electrodes in electrolysis as we shall
see later) because the electrons are free to move along the layers.
Here is a summary of the properties of graphite:
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Has a lower density than diamond. This is because of the relatively large
amount of space that is "wasted" between the sheets.
Is insoluble in water and organic solvents - for the same reason that diamond
is insoluble. Attractions between solvent molecules and carbon atoms will
never be strong enough to overcome the strong covalent bonds in graphite
Conducts electricity. The delocalised electrons are free to move throughout
the sheets. If a piece of graphite is connected into a circuit, electrons can fall
off one end of the sheet and be replaced with new ones at the other end.
Finally, another giant atomic structure is silicon (IV) oxide, otherwise known as
silicon dioxide or sand.
There are three different crystal forms of silicon dioxide. The easiest one to
remember and draw is based on the diamond structure.
Crystalline silicon has the same structure as diamond. To turn it into silicon
dioxide, all you need to do is to modify the silicon structure by including some
oxygen atoms.
Notice that each silicon atom is bridged to its neighbours by an oxygen atom. Don't
forget that this is just a tiny part of a giant structure extending on all 3 dimensions.
Silicon dioxide
has a high melting point - varying depending on what the particular structure
is (remember that the structure given is only one of three possible
structures), but around 1700C. Very strong silicon-oxygen covalent bonds
have to be broken throughout the structure before melting occurs.
is hard. This is due to the need to break the very strong covalent bonds.
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doesn't conduct electricity. There aren't any delocalised electrons. All the
electrons are held tightly between the atoms, and aren't free to move.
is insoluble in water and organic solvents. There are no possible attractions
which could occur between solvent molecules and the silicon or oxygen
atoms which could overcome the covalent bonds in the giant structure.
ATOM
Diamond graphite,
metals, high
melting points
MOLECULE
Sand molecules,
high melting point
Carbon dioxide,
water. Low
melting points
ION
All ionic
compounds like
sodium chloride.
High melting
points
None exist
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METALLIC BONDING
Metals are giant structures thatb have high melting points and boiling points
(usually).
Metallic bonds are quite strong so that metal atoms are quite hard to pull apart.
Metals can be beaten into sheets (malleable) due to the nature of the forces holding
the metal together. Since there is no overall force or attraction between the positive
cations, they can be moved from one lattice site to the other as the metals are bent
or shaped.
Metal ions are closely packed in the giant structure and this accounts for the
property of the metal. Metal atoms give away one or more of their pouter shell
electrons to form positive ions called cations. The electrons that have been given
away make up a sea of electrons that surround these positive metal cations.
The negative electrons are attracted to the positive cations and hold them together.
The electrons are free to move all of the way through the whole structure.
The electrons are DELOCALISED which means that they are NOT IN ONE
FIXED POSITION.
[ANDREW RICHARD WARD ALL RIGHTS RESERVED A WARD TUITION - 2009]
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TOPIC 4: STOICHIOMETRY
This is a table of all the POSITIVE METAL IONS
Valency 1
Valency 2
+
lithium
Li
sodium
Na
potassium
silver
Ag
hydronium
H3O
(or hydrogen)
ammonium
NH4
copper I
Cu
mercury I
Hg
Valency 3
2+
aluminium
Al
2+
iron III
Fe
chromium
Cr
magnesium
Mg
calcium
Ca
strontium
Sr
barium
Ba
copper II
Cu
lead II
Pb
zinc
Zn
manganese II
Mn
iron II
Fe
tin II
Sn
2+
3+
3+
3+
2+
2+
2+
2+
2+
2+
2+
Valency 2
-
fluoride
chloride
Cl
bromide
Br
iodide
hydroxide
OH
nitrate
NO3
bicarbonate
HCO3
bisulphate
HSO4
nitrite
NO2
chlorate
ClO3
permanganate
MnO4
hypochlorite
OCl
Dihydrogenphosphate
H2PO4
Valency 3
2-
oxide
sulphide
carbonate
CO3
sulphate
SO4
sulphite
SO3
dichromate
Cr2O7
chromate
CrO4
oxalate
C2O4
thiosulphate
S2O3
tetrathionate
S4O6
monohydrogen
phosphate
HPO4
phosphate
3-
PO4
22-
22-
22-
22-
2-
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Move the mouse cursor, Hints display the name of the ion and the cursor shows its
charge.
Double Click or Drag ions to the balance, Hints will tell you what you need to do to
balance the charges.
When the charges are being balanced, the formula of the substance will be displayed
in the formula frame and an atom counter shows numbers of each type of atom.
To find out the name of the substance move the mouse cursor to the formula frame.
To immediately show the formula Drag the ions directly to the formula frame.
The Test Mode randomly selects ions from the Level you have chosen
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Top Scores
Page 78
The Editor
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Apart from using the free Chemical Formula Tutor software (by kind permission of
Chemserve New Zealand), here is a simple guide to writing chemical formulae.
All subsyances are made from simple building blocks called chemical elements.
Each element has its own unique chemical symbol containing one or two letters.
Elements that were discovered many years ago have Latin names such as
Silver is Ag from the Latin Argentum
Lead id Pb from the Latin Plumbum
When chemical elements combine together they form COMPOUNDS. A
compound can be simplified into a chemical formula
Simple compounds
Many compounds contain just 2 elements.
Here is an example.
When magnesium burns in oxygen gas, a white ash of magnesium oxide is formed.
Here are the steps you use to work out the formula of magnesium oxide:
1. Write down the name of the compound (magnesium oxide)
2. Write down the chemical symbols for the elements in the compound
(Mg and O)
3. Use the Periodic Table to find the combining power of each element
from the table below. Write the combining power of each element
under its symbol (Mg 2, O 2)
4. If the numbers can be cancelled, cancel them (Mg O)
5. Swap over the combining powers and write down the symbols
GROUP
NUMBER
COMBINING
POWER
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There you go, the compound magnesium oxide has the chemical formula MgO but
you probably would have guessed that anyway!
What about Calcium Chloride?
1. Write down the name of the compound (Calcium Chloride)
2. Write down the chemical symbols for the elements in the compound
(Ca and Cl)
3. Use the Periodic Table to find the combining power of each element.
Write down the combining power for each element under its symbol
(Ca1 Cl2)
4. If the numbers can be cancelled down, cancel them (they cant be)
5. Swap over the combining powers. Write them below the symbols as a
subscript.
6. If the number is one please dont write it!
7. The correct formula of calcium chloride is CaCl2
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2H2O + O2 = 2H2O
This means TWO molecules of hydrogen gas react with ONE molecule of oxygen
gas to make TWO molecules of water.
Page 83
2. What is the equation when natural cooking gas at home (methane) burns in
air to form carbon dioxide and water?
WORD EQUATION:
METHANE + OXYGEN GAS = CARBON DIOXIDE GAS AND WATER
SYMBOLS AND FORMULAE CH4 + O2 = CO2 + H2O
BALANCE THE EQUATION CH4 + 2O2 = CO2 + 2H2O
Sample Problem
The work of balancing a chemical equation is in many ways a series of trials and
errors. As a sample exercise, consider the equation given below. Does this
represent a balanced chemical equation?
The work of balancing a chemical equation is in many ways a series of trials and
errors. As a sample exercise, consider the equation given below. Does this represent a
N2 + H2
NH3
To determine whether this reaction is balanced you must first determine how many
atoms of each type are on the reactant side (left-hand side) of the equation and how
many atoms of each type are on the product side (right-hand side). In this example,
you have two blue nitrogen atoms and two gray hydrogen atoms on the reactant
side but only one nitrogen atom and three hydrogen atoms on the product side. For
the purpose of balancing the equation we are not concerned what molecules these
atoms are in, just the number of atoms of each type.
To balance this reaction, it is best to choose one kind of atom to balance initially.
Let's choose nitrogen in this case. To obtain the same number of nitrogen atoms on
[ANDREW RICHARD WARD ALL RIGHTS RESERVED A WARD TUITION - 2009]
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the product side as on the reactant side requires multiplying the number of product
NH3 molecules by two to give:
N2
H2
2NH3
As you can see above, once we know what the molecules are (N2, H2, and NH3 in
this case) we cannot change them (only how many of them there are). The nitrogen
atoms are now balanced, but there are 6 atoms of hydrogen on the product side and
only 2 of them on the reactant side. The next step requires multiplying the number
of hydrogen molecules by three to give:
N2
3H2
2NH3
As a final step, make sure to go back and check whether you indeed have the same
number of each type of atom on the reactant side as on the product side. In this
example we have two nitrogen atoms and six hydrogen atoms on both sides of the
equation. We now have a balanced chemical equation for this reaction.
Here are 5 examples for you to try yourself
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Reaction #1
NO2 = N2O4
Reaction #2
CO + H2 = CH3OH
Reaction #3
NO + O2 = NO2
[ANDREW RICHARD WARD ALL RIGHTS RESERVED A WARD TUITION - 2009]
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Reaction #4
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Reaction #6
This is the first step in the industrial production of nitric acid, HNO3.
NH3 + O2 = NO + H2O
There are some other reactions called combustion where the chemical reacts in
OXYGEN GAS.
These are very common chemical reactions and the4 equations are commonly
found on IGCSE Exam Papers!!!
Reaction #7
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Reaction #8
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ANSWERS
2NO2 = N2H4
CO + 2H2 = CH3OH
2NO + O2 = 2NO2
4NO2 + H2O + O2 = 4NO3H
2NH3 + 2ClF3 = N2 + Cl2 + 6HF
2NH3 + 2.5O2 = 2NO + 3H2O (OR DOUBLE EACH NUMBER)
C3H8 + 5O2 = 3CO2 + 4H2O
C5H12 + 8O2 = 5CO2 + 6H2O
C2H6 + 3.5 O2 = 2CO2 + 3H2O (OR DOUBLE EACH NUMBER)
BALANCING IONIC EQUATIONS
These equations show reactions involving IONS (THESE ARE PARTICLES
WHICH HAVE A + OR CHARGE EITHER BY GAINING OR LOSING
ELECTRONS). The size of the charge on an ion is the same as the combining
power. Whether it is positive or negative depends on whether it is a metal or a nonmetal. METALS HAVE POSITIVE ION AND NON-METALS HAVE
NEGATIVE IONS.
In many ionic reactions, some ions DO NOT play any part in the reaction itself.
We call these ions SPECTATOR IONS.
We can therefore write a very simple ionic equation using only the IMPORTANT
IONS THAT REACT TOGETHER.
In these equations we always use STATE SYMBOLS. Aq means in water.
WORKED EXAMPLE OF AN IONIC EQUATION
Lead (II) nitrate + potassium iodide = lead(II) iodide and potassium nitrate
Pb(NO3)2 + KI = PbI2 + KNO3
BALANCED Pb(NO3)2 + 2KI = PbI2 + 2KNO3
In the above reaction, the potassium ions and nitrate ions ARE SPECTATOR
IONS and DO NOT actually take part in the chemical reaction.
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We should really state that the Relative Molecular Mass (RMM) of carbon dioxide
is 44 since carbon dioxide is a molecule.
The mass of one mole of calcium carbonate (CaCO3) is
(1 x RAM of calcium) + (1 x RAM of carbon) + (3 x RAM of oxygen)
= (1 x 40) + (1 x 12) + (3 x 16) = 100 g.
So, one mole of calcium carbonate has a mass of 100 g.
The Relative Formula Mass of calcium carbonate is 100.
If you are not sure what CaCO3 stands for. It stands for and
means 1 calcium + 1 carbon + 3 oxygens.
The mass of one mole of magnesium oxide (MgO) is
(1 x RAM of magnesium) + (1 x RAM of oxygen)
= (1 x 24) + (1 x 16) = 40 g.
So, one mole of magnesium oxide has a mass of 40 g.
The Relative Formula Mass of magnesium oxide is 40.
MOLES AND MOLECULES
One mole of hydrogen (H2) molecules
has a mass of 2 x the relative atomic mass (RAM) of hydrogen.
This is because each molecule of hydrogen contains two atoms
and the "relative atomic mass" is the relative mass of one atom.
So, one mole of hydrogen molecules has a mass of 2 g (2 x 1).
This is called the relative molecular mass (RMM).
One mole of nitrogen (N2) molecules (RAM = 14)
has a mass of 2 x 14 grams, = 28 g.
One mole of oxygen (O2) molecules (RAM = 16)
has a mass of 2 x 16 grams, = 32 g.
One mole of chlorine (Cl2) molecules (RAM = 35.5)
has a mass of 2 x 35.5 grams, = 71 g.
One mole of bromine (Br2) molecules (RAM = 80)
has a mass of 2 x 80 grams, = 160 g.
Page 94
products.
Example 1.
What mass of magnesium oxide will be formed when 10 g of magnesium is
burned in air?
1) Find how many moles of magnesium are present
in 10 g of magnesium.
The RAM of magnesium is 24.
moles = mass RFM
moles = 10 24 = 0417 moles of magnesium.
2) Write the equation for the reaction between magnesium and oxygen (burning in
air means reacting with oxygen).
magnesium + oxygen
2Mg(s) + O2(g)
magnesium oxide.
2MgO(s)
If you don't know why the formula for magnesium oxide is MgO, see Combining
Power.
Look at the big numbers in front of the elements or compounds.
The big numbers tell you that 2 magnesiums will give 2 magnesium oxides.
This means that 2 moles of Mg give 2 moles of MgO,
or one mole of Mg gives one mole of MgO.
The big numbers in the equation tell you that the proportion of Mg to MgO is one
to one.
Since the proportion of Mg to MgO is 1 to 1, then 0417 moles of magnesium will
give 0417 moles of magnesium oxide.
[ANDREW RICHARD WARD ALL RIGHTS RESERVED A WARD TUITION - 2009]
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sodium oxide.
2Na2O(s)
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For a molecular compound, a reaction might show that the proportion of carbon to
hydrogen is 3 to 6 (divide by 3).
The empirical formula is CH2, not C3H6.
CH2 does not exist as a molecule.
The molecular formula could be any multiple of CH2,
CnH2n where n is a whole number.
The molecule is an alkene, but to know which particular alkene it is,
you must also know the relative molecular mass (RMM).
Here is an example
336 g of iron was found to react with 144 g of oxygen.
What is the empirical formula of iron oxide?
Method
1) Find how many moles of iron react with how many moles of oxygen.
RAM of Fe = 56, RAM of O = 16.
moles = mass RAM
for iron moles = 336 56
= 06 moles.
for oxygen moles = 144 16
= 09 moles.
2) The proportion of moles of iron to moles of oxygen
is reduced to the lowest whole number.
06 moles of Fe to 09 moles of O.
Divide by 3, multiply by 10.
2 moles of Fe to 3 moles of O.
The empirical formula is Fe2O3.
Since iron oxide is an ionic compound,
the actual formula is the same as the empirical formula. Try this one:
20g of calcium was found to react with 355 g of chlorine.
What is the empirical formula of calcium chloride?
Page 100
Find how many moles of calcium react with how many moles of chlorine.
RAM of Ca = 40, RAM of Cl = 355.
moles = mass RAM
for calcium moles = 20 40 = 05 moles.
for chlorine moles = 355 355 = 10 moles.
2) The proportion of moles of calcium to moles of chlorine is reduced to the lowest
whole number.
05 moles of Ca to 10 moles of Cl.
Multiply by 2.
1 mole of Ca to 2 moles of Cl.
The empirical formula is CaCl2.
Since calcium chloride is an ionic compound,
the actual formula is the same as the empirical formula.
MOLECULAR FORMULA FROM A COMPOUND
Example 1.
A compound was found to contain 30 g of carbon and 05 g of hydrogen.
The relative molecular mass of the compound is 42.
What is the molecular formula of the compound?
Method.
1) Find how many moles of carbon react with how many moles of hydrogen.
RAM of C = 12, RAM of H = 1.
moles = mass RAM
for carbon moles = 3 12 = 025 moles.
for hydrogen moles = 05 1 = 05 moles.
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Method.
1) Find how many moles of carbon are present in 6 g of carbon.
RAM of C = 12.
moles = mass RAM
moles = 6 12 = 05 moles of carbon.
2) Write the equation for the reaction.
carbon + oxygen
C(s) + O2(g)
carbon dioxide.
CO2(g)
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Example 2.
The electrolysis of sodium chloride dissolved in water
produces hydrogen at the cathode and chlorine at the anode.
What mass and volume of hydrogen is obtained
if 142 g of chlorine are discharged at the anode?
Method.
1) Find how many moles of chlorine are discharged at the anode.
RMM of Cl2 = 71.
moles = mass RMM
moles = 142 71 = 2 moles of chlorine.
2) Write the half equations for the electrolysis.
(i) 2H+ + 2eH2
(ii) 2Cl - 2e
Cl2
Balance the half equations, to give the same number of electrons on each side.
(the equations are already balanced).
Find the proportion of how many moles of product are discharged at each
electrode.
1 mole of H2 at the cathode gives 1 mole of Cl2 at the anode.
The proportion is 1 to 1.
2 moles of Cl2 gives 2 moles of H2.
3) Find the mass and volume of 2 moles of hydrogen.
The RMM of hydrogen is 2.
mass = moles x RMM
mass = 2 x 2 = 4 grams of hydrogen.
volume = moles x 24,000 cm3
volume = 2 x 24,000 cm3 = 48,000 cm3 of hydrogen.
[ANDREW RICHARD WARD ALL RIGHTS RESERVED A WARD TUITION - 2009]
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Example 4.
The electrolysis of molten aluminium oxide - Al2O3
produces aluminium at the cathode and oxygen at the anode.
See the extraction of aluminium for more detail.
A current of 10 amps is allowed to flow
through molten aluminium oxide for 5 hours.
What mass of aluminium is deposited at the cathode?
Method.
1) Write the half equation for the electrolysis.
Al3+ + 3eAl
3 moles of electrons (3 Faradays)
are required to deposit 1 mole of aluminium.
2) Find how many Faradays have passed through the aluminium oxide in 5 hours.
Q=Ixt
5 hours contains (5 x 60 x 60) seconds
= 18,000 seconds.
Q = 10 x 18,000
= 180,000 Coulombs.
1 Faraday = 96,500 coulombs.
180,000 coulombs = 180,000 96,500 Faradays
= 1865 Faradays.
[ANDREW RICHARD WARD ALL RIGHTS RESERVED A WARD TUITION - 2009]
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Begin by preparing your burette. Your burette should be conditioned and filled
with titrant (liquid in the burette) solution. You should check for air bubbles and
leaks, before proceding with the titration.
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STEP 2
Take an initial volume reading and record it in your notebook. Before beginning a
titration, you should always calculate the expected endpoint volume.
STEP 3
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STEP 4
Approach the endpoint more slowly and watch the color of your flask carefully.
Use a wash bottle to rinse the sides of the flask and the tip of the buret, to be sure
all titrant is mixed in the flask.
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STEP 6
As you approach the endpoint, you may need to add a partial drop of titrant. You
can do this with a rapid spin of a teflon stopcock or by partially opening the
stopcock and rinsing the partial drop into the flask with a wash bottle. Ask your
Teacher to demonstrate these techniques for you, in the lab.
STEP 7
Make sure you know what the endpoint should look like. For phenolphthalein, the
endpoint is the first permanent pale pink. The pale pink fades in 10 to 20 minutes.
[ANDREW RICHARD WARD ALL RIGHTS RESERVED A WARD TUITION - 2009]
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If you think you might have reached the endpoint, you can record the volume
reading and add another partial drop. Sometimes it is easier to tell when you have
gone past the endpoint.
STEP 8
If the flask looks like this, you have gone too far!
STEP 9
When you have reached the endpoint, read the final volume in the buret and record
it in your notebook.
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Subtract the initial volume to determine the amount of titrant delivered. Use this,
the concentration of the titrant, and the stoichiometry of the titration reaction to
calculate the number of moles of reactant in your analyte solution.
HERE ARE SOME WORKED EXAMPLES!
Example 1.
In a titration, 50 cm3 of 2 mol dm-3 sodium hydroxide was exactly neutralised by
30 cm3 of hydrochloric acid.
What is the concentration of the hydrochloric acid?
Method.
1) Write the equation for the reaction.
sodium hydroxide + hydrochloric acid
NaOH(aq)
+
HCl(aq)
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The tiny cells that look like a button in watches, calculators or cameras are also a
type of dry cell. They contain silver metal and zinc oxide. Silver is expensive and
so is the cost of these cells. They last for a long time because the watches and
batteries only run on a very small electrical current.
ALKALINE CELLS
Alkaline cells are very similar o zinc-carbon dry batteries but the electrode
contains ALKALI called potassium hydroxide. The zinc electrode has holes in it
that makes it SLIGHTLY POROS a bit like a sponge. Alkaline cells provide a
larger electric current and last longer than zinc-carbon dry cells.
RECHARGEABLE CELLS
One type of rechargeable cell contains NICKEL AND CADMIUM compounds. A
redox reaction takes place when the cell is being used. When the cell is placed into
a recharger, an electric current passes through it. This REVERSES THE
REACTION and RECHARGES the cell.
FUEL CELLS
These use the energy from fuels to provide electrical energy. A fuel like hydrogen
or methane is passed into the cell along with a stream of oxygen gas. These are
made to react via the external circuit to provide the electric current. Fuel cells are
expensive to make but provided that the fuel is continuously passed in, the cell will
continue to provide an electric current. Fuel cells are used in NASA spacecraft to
provide a constant source of electrical energy.
[ANDREW RICHARD WARD ALL RIGHTS RESERVED A WARD TUITION - 2009]
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Lead (II) bromide PbBr2 is made from lead ions Pb2+ and 2 bromide ions Br-.
When the compound is molten (hot and runny), the ions are free to move and drift
towards the oppositely charged electrode where the ions are turned into either
atoms or molecules.
The equations for the electrolysis are:
NEGATIVE CATHODE: Pb2+ + 2e- = Pb metal
POSITIVE ANODE: 2Br- = Br2 + 2eTHE ELECTROLYSIS OF CONCENTRATED HYDROCHLORIC ACID
Hydrochloric acid is a strong acid and ionizes like so
HCl = H+ + ClThese ions drift to the oppositely charged electrode and get turned into molecules
The equations for the electrolysis are:
NEGATIVE CATHODE 2H+ + 2e- = H2
POSITIVE ANODE 2Cl- = Cl2 + 2e-
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Laboratory tests
The table shows the common laboratory tests for the products from the electrolysis
of brine.
Substance
hydrogen gas
chlorine gas
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The oxygen reacts with the carbon anodes and makes CO2 gas. Anodes have to be
replaced periodically.
ADDITIONA NOTES FOR A* STUDENTS
The aluminium ions are reduced by the addition of electrons to form aluminium
atoms. The oxide ions are oxidized as they lose electrons to form oxygen
molecules. This process often uses massive amounts of electricity and needs
economical low-price electricity. Aluminium is often extracted in countries which
have an abundance of hydro-electric power.
THE PURIFICATION OF COPPER BY ELECTROLYSIS
Copper is extracted from its ore (usually malachite) by reduction with carbon (like
iron ore is during the blast furnace) but the copper that is made is not pure enough
for some o its uses such as electrical wiring and cables. It can be purified by
using electrolysis. The cell used (shown above) has an anode of impure copper.
Copper (II) SULPHATE SOLUTION IS USED AS THE ELECTROLYTE. The
cathode is made from a THIN PIECE OF PURE COPPER.
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IONS
CATHODE (-)
ANODE (+)
Na+ Cl-
2Na
Cl2
Cu2+ Br-
Cu
Br2
Zn2+ I-
Zn
I2
Aqueous solutions are much more complicated as they contain the ions from the
compound and also the ions from the water.
For example, sodium chloride solution
NaCl = Na+ and ClH2O = H+ and OHBoth positive ions will go to the negative cathode but only the H+ ion is discharged
as H2.
Both negative ions go to the positive anode but only the Cl- is discharged as Cl2.
GENERAL RULE IN ELECTROLYSIS
[ANDREW RICHARD WARD ALL RIGHTS RESERVED A WARD TUITION - 2009]
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This means the mercury cell is a more INDIRECT method than the membrane cell.
CHLORINE
Killing bacteria in water, making bleaches, paints, dyes, hydrochloric acid, drugs
and pharmaceuticals
HYDROGEN
Making nylon, margarine, hydrogen peroxide (hair bleach), ammonia
SODIUM HYDROXIDE
This is used for purifying natural gas, sewage treatment, degreasing materials
(oven cleaner), making paper, soaps and detergents.
The chlor-alkali process is also responsible for the following important reactions.
Manufacturing Sodium Chlorate
Sodium hydroxide and chlorine are mixed at room temperature to form sodium
chlorate.
2NaOH(aq) + Cl2(g) NaOCl(aq) + NaCl(aq) + H2O(l)
Uses of the Industrial Products
Sodium
Manufacture of titanium:
TiCl4(l) + 4Na(s) 4NaCl(s) + Ti(s)
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Sodium Hydroxide
Manufacture of soap
Chlorine
Sodium Chlorate
NaOCl is used in domestic bleach. The NaCl impurity is usually left in the
bleach as it is not important unless an acid is added, in which case the two
react together to form chlorine (the solution proportionates).
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NON-METALS
Most have low melting points and
boiling points
Have a dull surface
Are brittle when solid and break/snap
easily
Solids, liquid or gases
Poor conductors of heat and electricity
Glasses have the same properties as ceramics but they are transparent (clear)
GLASS IS TRANSPARENT CERAMIC.
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PLASTICS
Plastics are synthetic (man-made) materials made from large molecules.
Plastics have many useful properties
1.
2.
3.
4.
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The rise in temperature of the water is a measure of the energy transferred to the
water. This method above will NOT give a very accurate answer as heat can be lost
in many ways such as out of the metal can and from the sides of the metal can.
Heat energy will also be lost from the top of the spirit burner And out into the
laboratory.
Nevertheless, the technique can be used to compare the energy released by the
same amount of different fuels.
The energy change can be calculated from the equation:
H (amount of heat given to water) =
Mass of water (grams or kilograms) x specific heat capacity of water x temp. rise
The specific heat capacity of water is a FIXED TERM of 4.2 J /g / 0C
This term means it takes 4.2 JOULES OF ENERGY TO HEAT 1 GRAM OF
WATER BY 1 CELSIUS.
Units are very important in this equation. When you use g for mass of water, your
answer for heat given to water will be in J. However, if the mass is in KILOgrams,
the answer for energy will be in KILOjoules.
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Since 1cm3 of water weighs 1g, the mass of water in the beaker is the SAME as the
volume of water in cm3.
WORKED EXAMPLE
2.0grams of paraffin were burned in a spirit burner under a metal can containing
400cm3 of water. The temperature of the water rose from 20 0C to 700C.
Calculate the energy produced by the paraffin in j g-1 and Kj g -1.
EQUATION
H = 400 X 4.2 X 50
H = 84000 J for every 2 g of paraffin
H = 42000 J g-1
H = 42 KJ -1
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In an endothermic reaction, more heat is needed to break bonds than the energy
that is released when new bonds are formed. The energy changes in endothermic
reactions are usually small.
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NUMBER OF
CARBON
ATOMS
1
2
3
4
5
6
CHEMICAL
FORMULA OF
THE ALKANE
CH4
C2H6
C3H8
C4H10
C5H12
C6H14
ENERGY OF
COMBUSTION
(KJ mol-1)
882
1542
2202
2877
3487
4141
HYDROGEN AS A FUEL
The reaction between hydrogen gas and oxygen gas is VERY EXOTHERMIC. It
produces lots of heat energy.
2H2(g) + 02(g) = 2H20(l)
Burning only 2 grams of hydrogen gives out 485 joules of energy.
This reaction is used for powering space rockets using hydrogen as ROCKET
FUEL.
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Metals and solutions of their own salts can be used to generate electricity. If the
above experiment is set up, a bulb will glow showing that electricity has een
produced in the zinc and copper half-cells.
Zinc is higher than copper in the reactivity series, so is the producer of electrons at
THE CATHODE. The copper takes the electrons at THE ANODE.
The reactions are:
CATHODE (-) Zn(s) = Zn2+(aq) + 2e- (OXIDATION LOSS OF ELECTRONS)
ANODE (+)
Please refer to the section in the next chapter (chapter 7) refer to the section
REDOX and see OIL RIG.
Although zinc/copper is used here as the example, you can get electricity from any
pair of metals set up in a diagram like the one shown above.
The amount of electricity produced depends on the position of the metals in the
reactivity series. The rule is:
[ANDREW RICHARD WARD ALL RIGHTS RESERVED A WARD TUITION - 2009]
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BATTERIES
The words battery and cell are often used to mean the same thing. A battery is
MORE THAN ONE CELL WORKING TOGETHER.
Batteries range from the large lead/acid batteries used in cars, lorries and buses to
tiny little lithium batteries used in hearing aids and watches.
Batteries are now used in mobile phones and i-pods and are now known as mobile
electricity carriers.
Their impact on society is becoming a fascinating study.
Without batteries we would have no mobile phones or laptop computers which are
now part of our everyday lives.
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concentration of a solution
temperature
surface area of a solid
a catalyst
pressure of a gas
light
Simple collision theory can be used to explain how these factors affect the rate of a
reaction.
Two important parts of the theory are:
1. THE REACTING PARTICLES MUST COLLIDE WITH EACH OTHER
2. THERE MUST BE ENBOUGH ENERGY TO OVERCOME TE
ACTIVATION ENERGY THAT IS NEEDED TO START THE
CHEMICAL EACTION.
EFFECT OF CONCENTRATION ON REACTION RATE
Increasing the CONCENTRATION of a reactant will INCREASE the rate of a
chemical reaction.
When a piece of magnesium ribbon is added to a solution of hydrochloric acid, the
following reaction takes place:
Magnesium ribbon + hydrochloric acid = magnesium chloride + hydrogen gas
Mg(s) + 2HCl(aq) = MgCl2(aq) + H2(g)
As the magnesium comes into contact with the hydrochloric acid, there is
EFFERVESCENCE OR FIZZING. Hydrogen gas is made.
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Some experiments were done using the same length of magnesium ribbon to
make the experiment a FAIR TEST. Different concentrations of acid were used.
Four different experiments were done. The volume of hydrogen gas produced was
recorded every minute for 6
minutes. A graph was plotted
and is shown below:
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As the sodium thiosulphate solution is diluted more and more, the precipitate takes
longer and longer to form.
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Catalytic converters
Catalytic converters use metals like platinum, palladium and rhodium to convert
poisonous compounds in vehicle exhausts into less harmful things. For example, a
reaction which removes both carbon monoxide and an oxide of nitrogen is:
Because the exhaust gases are only in contact with the catalyst for a very short
time, the reactions have to be very fast. The extremely expensive metals used as
the catalyst are coated as a very thin layer onto a ceramic honeycomb structure to
maximise the surface area.
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The explanation
You are only going to get a reaction if the particles in the gas or liquid collide with
the particles in the solid. Increasing the surface area of the solid increases the
chances of collision taking place.
Imagine a reaction between magnesium metal and a dilute acid like hydrochloric
acid. The reaction involves collision between magnesium atoms and hydrogen
ions.
Increasing the number of collisions per second increases the rate of reaction.
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You might have a "catalytic converter" in the exhaust pipe of your car. If the
catalyst was Platinum you might expect silly people to start stealing car exhaust
pipes; but there is so little catalyst there that it would not be worthwhile for them.
A little bit of catalyst goes a long way! What does the catalytic converter do? Well
without it the fumes from your car would cause too much pollution and the car
might fail its MOT.
Perhaps you don't think that catalysts are very important.
"..... alters the rate of a chemical reaction ....." This means that catalysts make
chemical reactions go faster. I am still looking for one which will make you do
your homework faster, and another which will make me mark it faster. What about
chemical reactions. Some of them go very slowly, your chemistry experiment
might take hours, days, weeks, or ever years. Imagine if your chemistry teacher
asked you to find out what gas is released from Hydrogen Peroxide: you might
have to sit there watching your test tube for weeks; your chemistry teacher would
keep on asking why you had not finished your work. Eventually you would have
enough gas to test; so weeks later you would say "Oh, it is Oxygen Miss." If you
had put a little pinch of Manganese Dioxide into the test tube, the gas would be
produced in a few minutes. So, you would be able to go long before the end of the
lesson. Even better, you would still have the Manganese Dioxide catalyst which
you would be able to sell back to your teacher to use with another class.
How about the chemical industry. Well they will make much more money if they
can make their products quickly. The manufacturers of Nitric Acid use Platinum as
a catalyst. Even though this is a very expensive metal, it does not cost too much to
use it because they are only using small amounts of it.
"A catalyst is a substance ....." This means that it is some kind of chemical
substance! It could be a pure element; e.g. Platinum, Nickel; or it could be a pure
compound, e.g. Manganese Dioxide, Silica, Vanadium V Oxide, Iron III Oxide; it
coulb be dissolved ions, e.g. Copper ions, Cobalt II ions; or it could be a mixture,
e.g. Iron-Molybdenum, or it could be a much more complicated compound such as
protein (all enzymes are proteins; you learn about them in your biology; they are
special cases.)
Enzymes are biological catalysts. They are slightly different in that they are easily
denatured by heat
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Most catalysts make chemical reactions go faster. Chemists call such catalysts
"positive catalysts" or "promoters". However, sometimes we want a chemical
reaction to go more slowly. So we choose a "negative catalyst"; we could call this
an "inhibitor". My wife put a negative catalyst in our central heating system. She
did this to stop the iron bits from rusting. We did not have a problem with the
Copper pipes (Copper does not rust), but we might have had a problem with the
old Iron radiators: we wanted to stop them from rusting so we used an inhibitor. I
think that we also have an inhibitor in the water cooling system of our car so that
the car radiator does not rust. This is cheaper than buying a new car every year
when the old one has got too rusty.
My baker puts an inhibitor into the bread he makes. This slows down the chemical
reactions which make bread go stale. This is important since we only go shopping
once a week. We used to put Lead in our petrol; this stops the engine from
"knocking". Now we have a better car which uses lead free petrol but the engine
can burn it without knocking.
You might wonder how catalysts work.
There are two ways in which catalysts work. You already know that when two
different molecules bump into each other, they might react to make new chemicals.
We usually talk about "collisions" between molecules, it would be much simpler to
say that the molecules bumped into each other. How fast a chemical reaction is
depends upon how frequently the molecules collide. You have probably been told
about the "kinetic theory" which is all about heat and how fast molecules move
around. What catalysts are doing when they make a chemical reaction go faster is
to increase the chance of molecules colliding. The first method is by "adsorption",
the second method is by the formation of intermediate compounds.
Adsorption This occurs when a molecule sticks onto the surface of a catalyst.
Make sure that you spell this word correctly; it is not the same as absorption. Here
is an example: it is possible to use Platinum as a catalyst to make sulphur Trioxide
from Sulphur Dioxide and Oxygen. Sulphur Trioxide is very important because it
is used to make Sulphuric acid which is needed for car batteries. The molecules of
the two gases (Sulphur Dioxide and Oxygen) get adsorbed (stuck onto) the surface
of a Platinum catalyst. Because the two molecules are held so close together, it is
more likely that they will collide and therefore react with each other. The Sulphur
Trioxide easily falls off the catalyst leaving space for more Sulphur Trioxide and
Oxygen.
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An example
In the manufacture of ammonia by the Haber Process, the rate of reaction between
the hydrogen and the nitrogen is increased by the use of very high pressures.
In fact, the main reason for using high pressures is to improve the percentage of
ammonia in the equilibrium mixture, but there is a useful effect on rate of reaction
as well.
The explanation
The relationship between pressure and concentration
Increasing the pressure of a gas is exactly the same as increasing its concentration.
If you have a given mass of gas, the way you increase its pressure is to squeeze it
into a smaller volume. If you have the same mass in a smaller volume, then its
concentration is higher.
You can also show this relationship mathematically if you have come across the
ideal gas equation:
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Because "RT" is constant as long as the temperature is constant, this shows that the
pressure is directly proportional to the concentration. If you double one, you will
also double the other.
The effect of increasing the pressure on the rate of reaction
Collisions involving two particles
The same argument applies whether the reaction involves collision between two
different particles or two of the same particle.
In order for any reaction to happen, those particles must first collide. This is true
whether both particles are in the gas state, or whether one is a gas and the other a
solid. If the pressure is higher, the chances of collision are greater.
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The darker areas contain the most silver, the lighter areas contain the least silver.
Here is the positive image that forms when the film has been developed.
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The silver salts need to be kept in the dark but even then they do slowly change
into silver metal over a period of time. This is why photographic film has a use
by date. If we use it after this time, all the silver salts will actually be silver and
the film will be completely useless.
PHOTOSYNTHESIS is also started by ultraviolet (UV) light being absorbed by
the green pigment chlorophyll in the leaves of green plants.
Carbon dioxide + water = glucose + oxygen
6CO2(g) + 6H2O (l) = C6H12O6(aq) + 602(g)
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REVERSIBLE REACTIONS
What is a reversible reaction?
Many reactions, such as fuels burning, are irreversible - they go to completion and
cannot easily be reversed. Reversible reactions are different. In a reversible
reaction the products can react to produce the original reactants again.
When writing chemical equations for reversible reactions, we do not use the usual
one-way arrow. Instead, we use two arrows, each with just half an arrowhead - the
top one pointing right and the bottom one pointing left. For example
ammonium chloride
The equation shows that ammonium chloride (a white solid) can break down to
form ammonia and hydrogen chloride. It also shows that ammonia and hydrogen
chloride (colourless gases) can react to form ammonium chloride again.
A very important example of a reversible reaction involving white anhydrous
copper(II) sulphate and blue hydrated copper(II) sulphate, the equation for which is
anhydrous copper(II) sulphate + water
WHITE
BLUE
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Look at the graph. You can see that for any given temperature the yield of
ammonia increases as the pressure increases. You can also see that for any
given pressure, the yield goes down as the temperature increases. This is
because the forward reaction is exothermic.
CHANGING THE POSITION OF EQUILIBRIUM
Reversible reactions can be a nuisance t a Chemist working on an industrial
(massive) scale.
You are trying to make as much product as you can but as soon as it starts to form,
it starts going back into reactants!
Scientists have found ways of increasing the amount of YIELD (PRODUCT)
FORMED by moving the position of the equilibrium (balance) to make more
products and not more reactants. They can do this by
1. Changing concentrations
2. Changing pressure for reactions only involving gases
3. Changing temperature for reactions only involving gases
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Here is an example
A (g) + 2B
2C(g)
Because both reduction and oxidation are going on side-by-side, this is known as a
redox reaction.
Oxidising and reducing agents
An oxidising agent is substance which oxidises something else. In the above
example, the iron(III) oxide is the oxidising agent.
A reducing agent reduces something else. In the equation, the carbon monoxide is
the reducing agent.
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Notice that these are exactly the opposite of the oxygen definitions.
For example, ethanol can be oxidised to ethanal:
You would need to use an oxidising agent to remove the hydrogen from the
ethanol. A commonly used oxidising agent is potassium dichromate(VI) solution
acidified with dilute sulphuric acid.
Ethanal can also be reduced back to ethanol again by adding hydrogen to it. A
possible reducing agent is sodium tetrahydridoborate, NaBH4. Again the equation
is too complicated to be worth bothering about at this point.
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It is essential that you remember these definitions. There is a very easy way to do
this. As long as you remember that you are talking about electron transfer:
A simple example
The equation shows a simple redox reaction which can obviously be described in
terms of oxygen transfer.
Copper(II) oxide and magnesium oxide are both ionic. The metals obviously aren't.
If you rewrite this as an ionic equation, it turns out that the oxide ions are spectator
ions and you are left with:
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Electron-half-equations
What is an electron-half-equation?
When magnesium reduces hot copper(II) oxide to copper, the ionic equation for the
reaction is:
You can split the ionic equation into two parts, and look at it from the point of
view of the magnesium and of the copper(II) ions separately. This shows clearly
that the magnesium has lost two electrons, and the copper(II) ions have gained
them.
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The first thing to do is to balance the atoms that you have got as far as you possibly
can:
ALWAYS check that you have the existing atoms balanced before you do anything
else. If you forget to do this, everything else that you do afterwards is a complete
waste of time!
Now you have to add things to the half-equation in order to make it balance
completely.
All you are allowed to add are:
electrons
water
hydrogen ions (unless the reaction is being done under alkaline conditions in which case, you can add hydroxide ions instead)
In the chlorine case, all that is wrong with the existing equation that we've
produced so far is that the charges don't balance. The left-hand side of the equation
has no charge, but the right-hand side carries 2 negative charges.
That's easily put right by adding two electrons to the left-hand side. The final
version of the half-reaction is:
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Now you repeat this for the iron(II) ions. You know (or are told) that they are
oxidised to iron(III) ions. Write this down:
The atoms balance, but the charges don't. There are 3 positive charges on the righthand side, but only 2 on the left.
You need to reduce the number of positive charges on the right-hand side. That's
easily done by adding an electron to that side:
Combining the half-reactions to make the ionic equation for the reaction
What we've got at the moment is this:
It is obvious that the iron reaction will have to happen twice for every chlorine
molecule that reacts. Allow for that, and then add the two half-equations together.
But don't stop there!! Check that everything balances - atoms and charges. It is
very easy to make small mistakes, especially if you are trying to multiply and add
up more complicated equations.
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You will notice that I haven't bothered to include the electrons in the added-up
version. If you think about it, there are bound to be the same number on each side
of the final equation, and so they will cancel out. If you aren't happy with this,
write them down and then cross them out afterwards!
OXIDATION STATES OXIDATION NUMBERS
Oxidation states simplify the whole process of working out what is being oxidised
and what is being reduced in redox reactions. However, for the purposes of this
introduction, it would be helpful if you knew about:
We are going to look at some examples from vanadium chemistry. If you don't
know anything about vanadium, it doesn't matter in the slightest.
Vanadium forms a number of different ions - for example, V2+ and V3+. If you
think about how these might be produced from vanadium metal, the 2+ ion will be
formed by oxidising the metal by removing two electrons:
The vanadium is now in an oxidation state of +4. Notice that the oxidation state
isn't simply counting the charge on the ion (that was true for the first two cases but
not for this one).
The positive oxidation state is counting the total number of electrons which have
had to be removed - starting from the element.
[ANDREW RICHARD WARD ALL RIGHTS RESERVED A WARD TUITION - 2009]
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It is also possible to remove a fifth electron to give another ion (easily confused
with the one before!). The oxidation state of the vanadium is now +5.
Every time you oxidise the vanadium by removing another electron from it, its
oxidation state increases by 1.
Fairly obviously, if you start adding electrons again the oxidation state will fall.
You could eventually get back to the element vanadium which would have an
oxidation state of zero.
What if you kept on adding electrons to the element? You can't actually do that
with vanadium, but you can with an element like sulphur.
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element
usual oxidation
state
exceptions
Group 1
metals
always +1
Group 2
metals
always +2
Oxygen
usually -2
Hydrogen
usually +1
Fluorine
always -1
Chlorine
usually -1
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Warning!
Don't get too bogged down in these exceptions. In most of the cases you will come
across, they don't apply!
[ANDREW RICHARD WARD ALL RIGHTS RESERVED A WARD TUITION - 2009]
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You might recognise this as an ionic compound containing copper ions and
sulphate ions, SO42-. To make an electrically neutral compound, the copper
must be present as a 2+ ion. The oxidation state is therefore +2.
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You might recognise the formula as being copper(II) sulphate. The "(II)" in
the name tells you that the oxidation state is 2 (see below).
You will know that it is +2 because you know that metals form positive ions,
and the oxidation state will simply be the charge on the ion.
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Example 1:
This is the reaction between magnesium and hydrochloric acid or hydrogen
chloride gas:
Have the oxidation states of anything changed? Yes they have - you have two
elements which are in compounds on one side of the equation and as uncombined
elements on the other. Check all the oxidation states to be sure:.
The magnesium's oxidation state has increased - it has been oxidised. The
hydrogen's oxidation state has fallen - it has been reduced. The chlorine is in the
same oxidation state on both sides of the equation - it hasn't been oxidised or
reduced.
Example 2:
The reaction between sodium hydroxide and hydrochloric acid is:
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Example 3:
This is a sneaky one! The reaction between chlorine and cold dilute sodium
hydroxide solution is:
Obviously the chlorine has changed oxidation state because it has ended up in
compounds starting from the original element. Checking all the oxidation states
shows:
The chlorine is the only thing to have changed oxidation state. Has it been oxidised
or reduced? Yes! Both! One atom has been reduced because its oxidation state has
fallen. The other has been oxidised.
This is a good example of a disproportionation reaction. A disproportionation
reaction is one in which a single substance is both oxidised and reduced.
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TYPE OF SOLUTION
Acidic
Neutral
Alkaline
Universal indicator (UI) can show the strengths of solutions of acids and alkalis
because it has more colours. Each colour is linked to a number on a scale called the
pH scale. The range of numbers is from 1 to 14.
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SUMMARY
pH BETWEEN 1 AND 6 = ACIDIC
pH BETWEEN 8 TO 14 = ALKALINE
pH 7 IS NEUTRAL
pH OF 1 TO 4 IS STRONG ACD
pH OF 5 AND 6 = WEAK ACID
pH OF 8 TO 10 = WEAK ALKALI
pH OF 11 to 14 = STRONG ALKALI
[ANDREW RICHARD WARD ALL RIGHTS RESERVED A WARD TUITION - 2009]
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BASICITY OF ACID
Monobasic (one H+)
Monobasic (one H+)
Dibasic (two H+)
Tribasic (three H+)
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If they dont react with hydrochloric acid or sodium hydroxide they are neutral
oxides.
SUMMARY
BASIC OXIDES ARE METAL OXIDES. THEY HAVE GIANT IONIC
STRUCTURES. THEY REACT WITH ACIDS IN NEUTRALIZATION.
ACIDIC OXIDES ARE FROM NON-METALS. THEY HAVE SIMPLE
MOLECULAR STRUCTURES. THEY REACT WITH BASES OR ALKALIS IN
NEUTRALIZATION. THEY DISSOLVE IN WATER TO FORM ACID
SOLUTIONS.
AMPHOTERIC OXIDES
These are formed by less reactive metals like zinc, aluminium and lead. They can
behave as both an acidic oxide and a basic oxide. They can react with acids and
bases.
NEUTRAL OXIDES
These do not react with acids or alkalis. N example of a neutral oxide is
NITROGEN MONOXIDE GAS, NO.
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The method used to produce a particular salt depends on two factors (i) the solubility of the base used and
(ii) the solubility of the salt to be made.
Since all acids are aqueous solutions the acid needed does not directly affect the
method of preparation.
The following table summarises the solubilities of the various bases and salts,
compound
insoluble in
water
hydroxide
all others
oxide
all others
carbonate
all others
nitrate
all possible
none
sulphate
all others
lead, barium
chloride
all others
silver, lead
There are three main methods of preparing salts (i) for a metal or insoluble base reacting with an acid to produce a soluble salt
filtration is used, e.g. reacting copper(II) oxide with sulphuric acid to make
copper(II) sulphate.
(ii) for a soluble base reacting with an acid to produce a soluble salt titration is
used, e.g. reacting sodium hydroxide with hydrochloric acid to make sodium
chloride.
(iii) for an insoluble base reacting with an acid to produce an insoluble salt a two
stage process involving filtration and then precipitation is used, e.g. reacting
lead(II) oxide firstly with nitric acid to produce lead(II) nitrate and then reacting
the lead(II) nitrate with aqueous iodide ions to produce lead(II) iodide.
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The data for this graph can be gathered by performing a series of experiments
at different temperatures.
A beaker of water is heated to the required temperature and an excess of solid
is added to the eater, with stirring. The excess is filtered off, dried and weighed
and the mass dissolved can be calculated by subtraction.
The graph can be used to calculate how much solid is needed to make a
saturated solution at a particular temperature.
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For example -
Introduction
There are a number of different methods of making salts, such as the reaction of a
metal with an acid. Copper metal, however, does not react with sulphuric acid and
so another method must be used. In this experiment a basic copper compound
(copper(II) oxide) will be reacted with sulphuric acid giving copper(II) sulphate as
one of the products.
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Method
Place 25 cm3 of dilute sulphuric acid in a 100 cm3 beaker and warm it gently over a
Bunsen burner flame to about 50 oC. Add a small spatula measure of copper(II)
oxide and stir, with a glass rod, until it dissolves. Then add a further spatula
measure, with stirring and continue until a small quantity of copper(II) oxide is
present in excess (How can you tell the copper(II) oxide is in excess?). Filter off
the excess copper(II) oxide, allowing the filtrate to pass into an evaporating basin
on a tripod and gauze, and heat it gently until half the liquid has evaporated. Leave
the remaining liquid to crystallise.
Questions
(i) Write up the experiment fully as you do it, being careful to include all
observations at the various stages.
(ii) Write a word and a balanced formula equation for this reaction.
(iii) How would you go about making copper(II) chloride crystals? Write a
balanced formula equation for this reaction.
As shown by the diagrams above, the first stage is the addition of black
copper(II) oxide to sulphuric acid. Mild heating is required for a full reaction
to occur; however, care must be taken to ensure that the acid does not boil as
this would be a great safety hazard.
The copper(II) oxide is added until no more visible reaction can be seen, i.e.
the base no longer dissolves and a black solid is seen in the blue solution.
The mixture is then filtered (stage 2 above) to remove the excess black solid
and leave a clear blue solution in the evaporating bowl. If the blue solution is
heated gently, to remove some of the water and allowed to cool down slowly,
crystals will appear. The slower this crystallization is allowed to occur, the
larger the crystals that will be produced
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Questions
(i) Write up all your observations and results as they occur.
(ii) Write a full balanced chemical equation for this reaction and the reaction of
sodium hydroxide with
a) Sulphuric acid
b) Nitric acid
c)
N.B.: This method of preparation is used when any and all sodium or
potassium salts are made.
The problem with the reaction of a soluble base reaction with acid is that
once all the acid has reacted any excess base will not be visible. This
problem is overcome by adding a third chemical into the reaction mixture
called an indicator.
Indicators are chemicals that change colour with a change in pH. So, if
the indicator is added to an acid it will be one colour. As base is added the
pH of the solution is raised until, once all the acid has reacted, i.e. been
neutralized, and the base is now in excess, the indicator changes colour.
Exemplar indicators Indicator
phenolphthalein
methyl orange
colour in acid
colourless
red
colour in base
purple
yellow
The skill involved in this technique is to add the base to the acid slowly
enough so that the indicator just changes colour with one drop of excess
base.
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ION TESTS
THIS SECTION IS VITALLY IMPORTANT ANDCAN BE UP TO 15
MARKS ON A PAPER 6 ALTERNATIVE TO PRACTICAL PAPER
WHICH IS TAKEN BY BOTH CORE AND EXTENDED STUDENTS
Experiment - Tests on Inorganic Ions
Introduction
You are provided with a series of labeled aqueous solutions containing various
cations (positive ions) and anions (negative ions) as well as sodium hydroxide(aq),
ammonium hydroxide(aq), silver nitrate(aq), lead nitrate(aq), barium chloride(aq),
hydrochloric acid(aq), nitric acid(aq) and Devarda's alloy. Perform the reactions as
detailed below in stages 1, 2 and 3 completing the tables where appropriate.
Method - Stage 1
Take 2 test tubes and pipette one of the cation solutions into each of the test tubes,
to a depth of about 1 cm.
To one of the test tubes add 5 drops of sodium hydroxide(aq) from the bottles in the
lab. Note all observations in the table below.
Then fill the tube up to a depth of about 5 cm with the sodium hydroxide (aq) and
again note your observations down in the table.
Then repeat the above procedure in the second test tube, but use ammonium
hydroxide(aq) in place of the sodium hydroxide(aq).
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Stage 2
As a revision exercise perform the tests for various anions as detailed in the table
below. Complete the observations for the table.
Stage 3
The final part of the practical work involves the use of all the observations you
have made so far.
There are 3 solutions labelled X, Y and Z in the laboratory. Each of them contains
one cation and two anions. To find out what the components of each solution
perform the following routine.
Take 6 test tubes and pipette about a cm depth of one of the unknown solutions
into each of them. To one of the test tubes perform the sodium hydroxide (aq) test
from stage 1 and note down your observations in a suitable table. To the second
tube perform the ammonium hydroxide(aq) test from stage 1. In the remaining test
tubes perform the various anion tests and again note down all positive and negative
observations.
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From your observations in stages 1 and 2 try to determine the composition of the
solution and then repeat for the other unknown solutions.
DETAILED NOTES ON ION TESTS
(1) Metal ions :
When metal ions are reacted with hydroxide ions a displacement reaction
generally occurs, since most metal hydroxides are insoluble in water. Group I
hydroxides, e.g. sodium hydroxide, are all soluble in water and so if a solution
of sodium hydroxide is mixed with a solution of another metal ion a precipitate
is formed.
The transition metal hydroxides produced tend to be uniquely coloured and so
they allow easy identification of the particular metal ion that is present, e.g.
iron(II) hydroxide is green, iron(III) hydroxide is red and copper(II) hydroxide
is blue.
Below is a table summarizing the precipitates formed with common IGCSE
metal ions Cation
aluminium, Al3+(aq)
ammonium, NH4+(aq)
calcium, Ca2+(aq)
copper, Cu2+(aq)
iron(II), Fe2+(aq)
iron(III), Fe3+(aq)
zinc, Zn2+(aq)
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Na2SO4(aq)
+ Cu(OH)2(s)
These equations can also be represented only by the reacting ions, simplifying
the equation 2-OH(aq)
+ Cu2+(aq)
Cu(OH)2(s)
sodium aluminate
Na3Al(OH)6(aq)
Ammonium hydroxide is too weak a base to allow this reaction to occur, so the
aluminium hydroxide does not dissolve in excess ammonium hydroxide.
With zinc hydroxide and copper(II) hydroxide the precipitates do dissolve in
excess ammonium hydroxide, not because of an acid-base reaction, but
because the molecules of ammonia surround the metal ions in solution, making
them soluble.
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+ -OH(aq)
NH3(g)
+ H2O(l)
Below is a table summarizing the observations for the common non-metal ion tests
Anion
Reaction
carbonate, CO32-(aq)
chloride, Cl-(aq)
iodide, I-(aq)
nitrate, NO3-(aq)
sulphate, SO42-(aq)
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(3) Gases :
Below is a table summarizing the various gas tests encountered in GCSE Gas
ammonia, NH3
carbon dioxide, CO2
chlorine, Cl2
hydrogen, H2
oxygen, O2
Test
turns damp red litmus paper blue (+ SMELL !)
gives a white precipitate with limewater
bleaches damp litmus paper
explodes with a 'pop' with a lighted splint
relights a glowing splint
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The modern periodic table, based on atomic number and electron configuration,
was created primarily by a Russian chemist, Dmitri Ivanovich Mendeleev, and a
German physicist, Julius Lothar Meyer, both working independently. They both
created similar periodic tables only a few months apart in 1869.
Mendeleev created the first periodic table based on atomic weight. He observed
that many elements had similar properties, and that they occur periodically. Hence,
the tables name.
His periodic law states that the chemical and physical properties of the elements
vary in a periodic way with their atomic weights. The modern one states that the
properties vary with atomic number, not weight.
Elements in Mendeleev's table were arranged in rows called periods. The columns
were called groups. Elements of each group had similar properties.
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The Periodic table can be divided into nine families of elements each having
similar properties. The families include:
Alkali metals
The alkali metals, found in group 1 of the periodic table, are highly reactive metals
that do not occur freely in nature. These metals have only one electron in their
outer shell. Therefore, they are ready to lose that one electron in ionic bonding
with other elements. As with all metals, the alkali metals are malleable, ductile,
and are good conductors of heat and electricity. The alkali metals are softer than
most other metals.
Alkaline metals
The alkaline earth elements are metallic elements found in the second group of the
periodic table. All alkaline earth elements have an oxidation number of +2, making
them very reactive.
The Transition metals
The 38 elements in groups 3 through 12 of the periodic table are called "transition
metals." As with all metals, the transition elements are both ductile and malleable,
and conduct electricity and heat. Their valence electrons are present in more than
one shell. This is why they often exhibit several common oxidation states.
Other metals
The "other metals" elements are located in groups 3.4.5.6. While these elements
are ductile and malleable, they are not the same as the transition elements. These
elements, unlike the transition elements, do not exhibit variable oxidation states,
and their valence electrons are only present in their outer shell. All of these
elements are solid, have a relatively high density, and are opaque. They have
oxidation numbers of +3, 4, and -3.
Metalloids
Metalloids are the elements found between the boundary that distinguishes metals
from non-metals. Metalloids have properties of both metals and non-metals. Some
of the metalloids, such as silicon and germanium, are semi-conductors.
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Non-metals
Non-metals are the elements in groups 3 to 8 of the periodic table. Non-metals are
not able to conduct electricity or heat very well. As opposed to metals, nonmetallic elements are very brittle. The non-metals can be gases, such as oxygen
and solids, such as carbon. The non-metals have no metallic luster, and do not
reflect light. They have oxidation numbers of 4, -3, and -2.
Halogens
The halogens are five non-metallic elements found in group 7 of the periodic table.
All halogens have 7 electrons in their outer shells, giving them an oxidation
number of -1.
Noble gases
The noble gases are found in group 8 of the periodic table. These elements have an
oxidation number of 0. This prevents them from forming compounds readily. All
noble gases have 8 electrons in their outer shell, making them stable.
Rare Earth
The 30 rare earth elements are composed of the lanthanide and actinide series. One
element of the lanthanide series and most of the elements in the actinide series are
synthetic, that is, human-made. All of the rare earth metals are found in group 3 of
the periodic table, and the 6th and 7th periods.
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GROUP
SODIUM
MAGNESIUM
ALUMINIUM
SILICON
PHOSPHORUS
SULPHUR
CHLORINE
ARGON
1
2
3
4
5
6
7
8
ELECTRON
ARRANGEMENT
2.8.
2.8.2
2.8.3
2.8.4
2.8.5
2.8.6
2.8.7
2.8.8
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The first period contains only 2 elements hydrogen and helium. We dont usually
study this period. Hydrogen is a very reactive gas. Helium is an unreactive gas but
is placed into group 8 as it has a FULL OUTER SHELL OF ELECTRONS.
The elements in the middle block of the periodic table in periods 4,5 and 6 are
called TRANSITION METALS. On of the typical properties of transition metals is
that their compounds are catalysts and speed up the rate of a chemical reaction by
providing an alternative pathway to a reaction by providing it with a lower
activation energy.
Vanadium (V) oxide is used in the Contact Process to manufacture sulphuric acid.
Iron powder is used in the Haber Process to manufacture ammonia.
GROUPS
Columns containing elements with an atomic number increasing down the column
are called groups. The groups are numbered 1 to 8.
Elements in a group have similar properties. You can think of them as a chemical
family.
Groups have family names such as
Group 1 = alkali metals; Group 2 = alkaline earth metals; Group 7 = halogens;
Group 8 = noble gases
IN THE SAME GROUP, ELEMENTS HAVE THE SAME NUMBER OF
OUTER SHELL ELECTRONS AHD HAVE THE SAME CHEMICAL
PROPERTIES
METALS AND NON-METALS
Most elements can be classed as either metals or non-metals. In the periodic table,
metals are found to the left and middle block. Non-metals are found on the right.
METALLOID elements Are between metals and non-metals. They have some
properties of metals and some properties of non-metals. Examples of metalloids
are silicon and germanium. They are often called SEMICONDUCTORS and are
used as computer microchips.
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GROUP PROPERTIES
MOST ELEMENTS ARE METALS. SOME METALS ARE VERY HIGHLY
REACTIVE WHILST OTHERS ARE COMPARATIVELY UNREACTIVE.
THE TWO TYPES OF METALS ARE FOUND IN DIFFERENT PARTS OF
THE PERIODIC TABLE.
GROUP 1 ELEMENTS THE ALKALI METALS
These very reactive metals only have one outer shell electron. This electron is
given away and donated when the metal reacts with a non-metal in ionic bonding.
The more electrons a metal atom has to lose in a reaction, the more energy is
needed to start the chemical reaction.
This is why group 2 elements are less reactive. They have to lose two electrons
when they react with a non-metal to form an ionic bond.
REACTIVITY INCREASES DOWN THE GROUP. As the atoms get bigger, the
outer electrons are far away from the magnetic pull of the nucleus. They can be
removed more easily. The atoms react and form CATIONS POSITIVE IONS
THAT ARE ATTRACTED TO THE NEGATIVE CATHODE IN
ELECTROLYSIS.
PROPERTIES OF GROUP 1 METALS
1. SOFT
2. EASY TO CUT
3. SHINY WHEN CUT BUT QUICKLY GO DARK (TARNISH) IN
AIR
4. VERY LOW MELTING POINTS COMPARED TO MOST OTHER
METALS. SODIUM MELTS AT 98 CELSIUS, POTASSUM AT 63
CELSIUS.
5. VERY LOW DENSITIES COMPARED TO MOST OTHER
METALS. LITHIUM, SODIUM AND POTASSIUM FLOAT ON
WATER.
6. RECAT VERY EASILY WITH AIR, WATER AND HALOGENS.
THE ALKALI METALS ARE STORED IN OIL TO STOP THEM
REACTING WITH AIR AND WATER.
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The chlorine is REDUCED by getting extra electrons from the bromide ions
2Br- = Br2 + 2eThe bromide ions are OXIDISED by losing electrons to the chlorine.
USES OF HALOGENS
Fluorides are used in toothpaste to prevent tooth decay
Fluoride compounds make plastics like TEFLON which is non-stick pan base
Chlorofluorocarbons (CFCS) propellants in aerosols and refrigerants
Chlorine is a bleach
Chlorine compounds are used to kill germs in water and swimming pools. They are
also used as disinfectants (like Dettol) and antiseptics like TCP.
Hydrochloric acid is widely used in industry to make chloride compounds
Bromine compounds make pesticides like cockroach killer
Silver bromide is the light sensitive film coating on photographic film
Iodine solution is an antiseptic in hospitals before you have an operation
PREDICTING PROPERTIES OF ELEMENTS
Metallic elements are on the left of the periodic table. Non-metallic elements are
on the right of the periodic table. In the middle is group 4 a mixture of metal and
non-metal elements.
As you go down any group, the properties change as the ATOMS GET LARGER.
More electron shells exist and there are more protons and neutrons in the nucleus.
The ATOMIC RADIUS of the atom increases as the RADIUS from the centre of
the nucleus to the outer shell electrons INCREASES.
As you go down a group of metals, the elements become more reactive as the
atoms get larger. In Group 1, lithium is least reactive and potassium is most
reactive. Francium is actually the most reactive as it is radioactive. We are not
allowed to keep rubidium, caesium and francium in schools.
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Metallic elements lose electrons when they react. The larger atoms lose electrons
more easily than smaller ones because in smaller atoms, the outer electrons are
closer to the pull of the central nucleus. The easier it is for a metallic element to
lose its electrons, the MORE reactive it is.
As you go down a non-metallic element group like group 7, the elements are LESS
REACTIVE as the atoms get larger. In group 7, fluorine is the MOST REACTIVE
element.
Fluorine is the most reactive non-metal element in the periodic table.
Non-metallic elements GAIN electrons when they react. The smaller atoms nearer
the top of the group do not have the attractive pull of the protons reduced by the
larger number of electron shells between them and the nucleus.
TRENDS IN OTHER GROUPS
The trend in INCREASING REACTIVITY down a metallic group is the same in
groups 1,2 and 3. These groups are also mainly metallic groups.
The trend in decreasing reactivity down a non-metallic group like group 7 is the
same for groups 5,6 and 7 which are also mainly non-metallic.
Group 4 shows a change from metallic properties to non-metallic properties as you
go down the group
C CARBON IS NON-METAL
Si SILICON IS NON-METAL
Ge GERMANIUM IS METALLOID
Sn TIN IS METAL
Pb LEAD IS METAL
Metalloid means that sometimes germanium behaves like ametal and sometimes
it behaves like a non-metal.
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GROUP 1 METAL
Low
Low
White
Very fast
Violent and dangerous
TRANSITION METAL
High
High
Coloured
Slow or no reaction
Slow or no reaction
Transition metals have more than one oxidation state. Copper forms Cu+ and Cu2+
ions. We give names to these as copper (I) and copper (II) ions.
This is why CuO is called copper (II) oxide and Cu2O is copper (I) oxide.
Iron forms Fe2+ and Fe3+ ions.
This is why FeO is called iron (II) oxide and Fe2O3 is iron (III) oxide.
TRANSITION METAL COMPOUNDS AE USUALLY COLOURED. Copper
compounds are usually blue or green. Iron compounds are green or brown. When
sodium hydroxide solution is added to a solution of a transition metal compound, a
coloured precipitate of the metal hydroxide is formed. The colour of the precipitate
helps you to identify the transition metal that is present. For example:
Copper sulphate + sodium hydroxide = copper (II) hydroxide + sodium sulphate
Blue solution + clear solution = blue solution + clear solution
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COLOUR OR METAL
HYDROXIDE
Blue
Green
Green turning to brown
Orange/brown
FLAME COLOUR
Red
Orange/yellow
Lilac (pale pink)
Blue/green
Brick red
Crimson (intense dark red)
Apple green
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TOPIC 10 METALS
THERE ARE 42 KEY WORDS TO LEARN IN THIS TOPIC
KEYWORD
1
Ore
Compounds
Mineral
Oxides
The compounds that metals in ores are most often found in are
__________
Native Metals
Gold, Silver
Native Metals
Extraction
Chemical, Electrical
10
Displacement
11
Reduction
12
Electrolysis
13
Reduction
14
Carbon
15
Blast Furnace
16
17
18
19
Iron oxide
20
Coke
21
Carbon
Found in Coke
22
Carbon monoxide
23
Limestone
24
Slag
25
Calcium Carbonate
Found in Limestone
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26
Too soft
27
28
Alloy
29
Carbon
30
Nickel
31
Tungsten
32
Chromium-Nickel
33
Smart-Alloys
34
Transition metal
35
Smelting
36
Electrolysis
37
Electrolyte
38
Cathode
Negative electrode
39
Anode
Positive electrode
40
Al2O3
41
42
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Here is a table showing the percentage abundance of some of the most common
elements,
metal element
% abundance
non-metal element
Fe
Ca
Cl
% abundance
Al
Si
Metal elements are commonly found as metal oxides and the metal itself can be
extracted using a one of a variety of chemical reactions. The process of changing
the metal ions in a compound to the neutral metal is called reduction. At its
simplest this means that oxygen has been removed from the compound; however,
the term reduction also applies to the addition of electrons to metal ions (see
electrolysis in middle part).
The first of the two most common methods of extraction are a reaction with carbon
(or carbon monoxide) in a Blast Furnace. This method is used mainly to extract
iron from iron oxide ores and zinc from zinc blende.
The other common method of extraction involves passing a great amount of
electricity through the molten ore in process called electrolysis.
METALS AND THE REACTIVITY SERIES
Metals such as sodium and potassium (the alkali metals) react violently with water,
too violently to do experimentally. The group II metals (also called alkaline earth
metals) react less readily and can be used in the laboratory.
Whilst magnesium does react with water it reacts so slowly as to be barely
noticeable. As with the alkali metals the reactivity of the group II metals
increases as the group is descended, with calcium, strontium and barium all
reacting at a reasonable rate.
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However, if the water is turned into steam and passed over the metal the
reaction becomes faster. The reaction with metals such as magnesium and zinc
becomes noticeable as a white smoke is given off. This smoke is the metal
oxide, instead of the metal hydroxide formed with water. Hydrogen gas is also
produced as with water.
Exemplar equation zinc + steam
Zn(s) + H2O(g)
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Take a test tube rack and fill the correct number of test tubes with 5 cm3 of
each metal salt solution. Add a small piece of a metal to each tube, and note
down all the observations you see, both before and after the metal is added,
with special note being taken of the observations of the metal added to the
solution.
Note down what solutions are available, what metals are available, and
construct a table using a full page of A4 paper turned on its side, similar to the
table below. Enter your observations as you make them into the table.
Questions
(i) Which of the metals is the most reactive? Explain your answer.
(ii) Which metal is the least reactive? Explain your answer.
(iii) Arrange the metals in order of their decreasing reactivity. Does this order
agree with any reaction observations you have come across so far?
[ANDREW RICHARD WARD ALL RIGHTS RESERVED A WARD TUITION - 2009]
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(iv) Write word and formulae equations for four different combinations of
metal and metal salt solution that react together.
When a metal is reacted with an aqueous solution of another metal ion a reaction
will occur only if the metal is higher in the reactivity series than the metal ion. This
provides a relatively easy way to obtain a Reactivity Series for any group of
metals.
There is another way which involves forming electrical cells with different metals
and measuring the potential differences produced, though this can give a slightly
different series than the chemical reactions do.
Exemplar equation magnesium + silver nitrate
Mg(s) + 2AgNO3(aq)
The method used to extract a metal from its ore is linked very closely to that
metal's position in the Reactivity Series. The more reactive a metal is the more
it wants to form compound and therefore the harder it is to isolate the pure
metal from its compounds.
The most reactive metals require electricity to extract the metal; less reactive
metals can be reduced with coke in a Blast furnace; the least reactive metals
are found as the pure metal and require very little purification.
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Iron(II) and iron(III) compounds can be separated by their colours most of the
time. Iron(II) compounds are generally green in colour and iron(III)
compounds are generally red in colour (there are plenty of exceptions though).
There is also a chemical test that can distinguish between the two sets of
compounds.
When an aqueous solution of iron ions is reacted with an aqueous solution of
hydroxide ions (e.g. sodium hydroxide or ammonium hydroxide) a coloured
precipitate will be formed. A green precipitate is formed with iron(II) ions and
a red-brown precipitate is formed with iron(III) ions.
The precipitates are insoluble iron hydroxides - Iron(II) hydroxide (green) and
iron(III) hydroxide (red-brown).
General ionic equations are given below for the two reactions Fe2+(aq) + 2-OH(aq)
Fe(OH)2(s)
Fe3+(aq) + 3-OH(aq)
Fe(OH)3(s)
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The carbon monoxide reduces the iron in the ore to give molten iron:
The limestone from 2, reacts with the sand to form slag (calcium silicate):
Both the slag and iron are drained from the bottom of the furnace.
The slag is mainly used to build roads.
The iron whilst molten is poured into moulds and left to solidify - this is called cast
iron and is used to make railings and storage tanks.
The rest of the iron is used to make steel.
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2CO(g)
2CO(g)
Thus the iron ore is reduced to iron metal which is produced at such a
temperature that it is molten. The carbon dioxide gas is simply vented off.
Fe2O3(s) + 3CO(g)
2Fe(l) + 3CO2(g)
There are quite a lot of impurities present in the ore, mainly sand (silicon
dioxide, SiO2). This is removed by the limestone. When the limestone is
heated in the furnace it decomposes to lime (calcium oxide) and oxygen. This
calcium oxide reacts with the sand to produce a slag, calcium silicate, CaSiO3.
CaCO3(s)
CaO(s) + SiO2(s)
CaO(s) + CO2(g)
CaSiO3(l)
The molten iron sinks to the bottom of the furnace with the liquid slag floating
on top of it. It is then a simple matter to allow the slag to flow away and obtain
the iron from beneath it.
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RUSTING OF IRON
Rusting is a very specific reaction of iron, where the metal is turned into
hydrated iron(III) oxide, known as rust.
For iron to rust the metal must be in contact with air and water. Without either
one of these the metal will remain intact and not corrode.
There are a number of separate stages involved, such as the oxidation of iron
to iron(II) ions and then the oxidation of iron(II) ions into the iron(III) oxide,
Equation 4Fe2+(aq) + O2(g) + 2H2O(l)
4Fe(OH)3(s) or 2Fe2O3.3H2O(s)
These processes involve the movement of electrons and as such the rusting
process can be speeded up by having salt present in the water, as this allows
greater conductivity.
RUST PREVENTION
Anything that prevents air and/or water from contacting the surface of the iron
will prevent the metal from rusting; as does anything that reacts faster than the
iron, i.e. is higher in the reactivity series.
In the latter case the oxygen and water will preferentially react with the more
reactive metal, to form metal oxides, before the iron can react. This allows the
overall strength of the metal to remain the same.
Common methods for rust prevention involve painting the metal; covering the
metal in oil or grease; coating the metal in zinc, called galvanising; attaching
blocks of a more reactive metal, such as magnesium or zinc, called sacrificial
protection; alloying the iron with other metals, such as chromium and nickel,
or changing the carbon content of the iron to create steel.
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ALLOYS
Metals are mixed together to create alloys. These alloys have better physical
properties then the individual metals, such as higher melting points, greater
mechanical strength or increased resistance to corrosion.
Steel Steel is an alloy created by blowing oxygen through molten iron produced in
the Blast furnace. The oxygen reacts with the carbon impurity in the iron,
turning it into carbon dioxide, which is vented off. The amount of carbon can
be very carefully controlled, giving a wide range of different steels. For
example, mild steel has about 0.25% carbon and is hard and strong; whilst
high carbon steel has about 1.5% carbon and is harder but more brittle.
Other metals can also be added to the steel as it is made to make an alloy such
as stainless steel. For this chromium and nickel are added which form
unreactive oxides on the surface of the iron and prevent the rusting process
starting.
Others A few other common alloys are bronze, which is a mixture of copper and tin;
brass, which is a mixture of copper and zinc; solder and pewter, which are
mixtures of tin and lead.
Even gold used for jewelry is alloyed with other metals such as zinc or nickel
to produce normal gold as well as white gold.
HEATING HYDROXIDES
Sodium hydroxide and potassium hydroxide are NOT CHANGED by heating
them.
Calcium hydroxide and magnesium hydroxide behave in the same way.
HYDROXIDE + HEAT = OXIDE + WATER
Ca(OH)2(aq) = CaO(s) + H2O(l)
[ANDREW RICHARD WARD ALL RIGHTS RESERVED A WARD TUITION - 2009]
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HEATING NITRATES
Group 1 nitrate + heat = group 1 nitrite + oxygen gas
2NaNO3(s) = 2NaNO2 + O2(g)
2KNO3(s) = 2NKNO2 + O2(g)
All other nitrates behave in the same way when heated
NITRATE + HEAT = OXIDE + NITRITE + OXYGEN GAS
2Ca(NO3)2 = 2CaO(S) + 4NO2(g) + O2(g)
NOTE: NITROGEN DIOXIDE IS A POISONOUS BROWN GAS THAT IS
A MAJOR AIR POLLUTANT. ALL THESE REACTIONS MUST BE
DONE IN A FUME CUPBOARD.
2Mg(NO3)2 = 2MgO(S) + 4NO2(g) + O2(g)
2Fe(NO3)2 = 2FeO(S) + 4NO2(g) + O2(g)
2Cu(NO3)2 = 2CuO(S) + 4NO2(g) + O2(g)
2Zn(NO3)2 = 2ZnO(S) + 4NO2(g) + O2(g)
UNREACTIVITY OF ALUMINIUM
Aluminiums position in the middle of the reactivity series should mean that it
is quite a reactive metal. It isnt.
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USES OF METALS
The uses of aluminium are based on the fact that it has a low density, is strong as
well as being unreactive (due to a layer of aluminium oxide on the surface)
It is used in the manufacture of aircraft (light and strong), drinks cans
(doesnt corrode and can be recycled).
Zinc is used to make the alloy brass by reacting it with copper. Covering iron
with zinc is called GALVANIZING. It stops the iron from rusting underneath.
Copper is an excellent conductor of heat with a high melting point. It is also an
excellent conductor of electricity and is used for electrical wires and cables.
MAKING STEEL FROM IRON
Iron from the blast furnace is brittle (snaps) and corrodes very easily because it
contains a large amount of carbon (from the coke). The corrosion of iron is
called RUSTING. It is a chemical reaction between IRON, OXYGEN AND
WATER. SALT ALSO SPEEDS UP RUSTING.
Common ways to stop rusting are:
1. GALVANIZING Cover the iron with zinc. More reactive zinc
corrodes and is SACRIFICED to PROTECT the less reactive iron
underneath. This is called SACRIFICIAL PROTECTION. This is
why ZINC BLOCKS are found at the front of ships to stop the iron
hull of the ship from corroding.
2. ALLOYING Mixing other metals to the iron to make STEEL.
In steel making, molten iron (from the blast furnace) is mixed with smaller
amounts of scrap (WASTE) iron and steel. Limestone is also added and the
whole mixture is melted. Oxygen is blasted into the mixture which reacts with
carbon in the iron to make carbon monoxide gas.
THIS PROCESS OF MAKING STEEL IS CALLED THE BASIC OXYGEN
PROCESS.
The carbon monoxide gas is recycled back into the TOP of the blast furnace.
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The recirculation of water all over the Earths surface is called the WATER
CYCLE.
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The pattern of rainfall determines where there is desert, rainforests and areas of
land that can or cannot be used for growing things. Rainfall therefore determines
the ECONOMIC WEALTH of the countries of the world.
GLOBAL WARMING
This is thought to be responsible for CLIMATE CHANGES that are affecting
where there is rainfall and also how much of it that there is. Al Gore thinks
global warming is a problem (watch his video An Inconvenient Truth) but
George Bush, U.S president seems to think it is make believe. What do you
think?
Water is essential for life on Earth. The demand for drinking water is increasing
as the earths population has grown from 2 billion people to 6 billion people
since the end of World War 2 in 1945.
Most INDUSTRIAL PROCESSES use water as either a RAW MATERIAL or
for cooling processes. Two-thirds of the water is used in the home. The rest is
used by industry.
It takes, for example, 250,000 tonnes of water to make just one tonne of steel.
Water stored in reservoirs needs to be purified to make drinkable water.
Additionally, tap water in certain areas is treated with SODIUM FLUORIDE
NaF to help to prevent tooth decay.
WATER PURIFICATION PROCESS
1. Water from a reservoir goes to a water treatment plant
2. Water is filtered through course gravel to remove large dirt particles
3. Water is filtered through beds of fine gravel and sand to remove the
smallest particles
4. Chlorine is passed through to kill bacteria and germs
5. Water is supplied directly to your home and industry
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PERCENTAGE PRESENT
78
21
0.9
0.03
0.07
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Carbon dioxide, methane and CFCs are GREENHOUSE GASES. The levels of
these gases in the atmosphere is increasing all the time due to Earths increasing
use of fossil fuels, pollutants from animal waste and the increased use of aerosols
and refrigerants.
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Short wave radiation from the sun warms the ground. The warm Earth gives off
heat as long-wave radiation. Much of this radiation cannot escape the Earth as it
is trapped by the greenhouse gases in the atmosphere. This is called THE
GREENHOUSE EFFECT.
The greenhouse effect is responsible for GLOBAL WARMING this means the
Earth is warmer than what it would be normally. Increasing levels of greenhouse
gases are stopping heat from escaping from Earth and the Earths average
temperature is increasing all of the time.
If global warming continues, the Earths climate will change forever. Polar ice
caps are starting to melt and raise sea levels.
In 30 years if this continues, low-lying countries like Holland, Bangladesh and
Kuwait will be under water.
We do not know for certain that greenhouse gases are the only thing responsible
for the average Earth temperature rising steadily. It may be that these
temperature increases may be part of a natural cycle in the past there have been
Ice Ages followed by very warm periods of time on Earth. Many people are
concerned, however, that it is not part of a natural cycle o we should ACT NOW
TO STOP GREENHOUSE GASES BEING MADE AND PREVEN GLOBAL
WARMING.
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ACID RAIN
Burning fossil fuels gives off many gases including SULPHUR DIOXIDE and
various NITROGEN OXIDES.
SULPHUR + OXYGEN = SULPHUR DIOXIDE GAS
S(s) + O2(g) = SO2(g)
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Sulphur dioxide combines with water in clouds and in the air to form
SULPHURIC ACID.
SULPHUR DIOXIDE + OXYGEN + WATER = SULPHURIC ACID
2SO2(g) + O2(g) + H2O(l) = 2H2SO4(aq)
Nitrogen oxide combines with water in a similar way to make NITRIC ACID.
These 2 gases in water form a combination of acids called ACID RAIN.
Limestone and marble buildings (both made of calcium carbonate CaCO3) are
damaged by acid rain. Metal structures are also attacked by sulphuric acid.
Acid rain HARMS PLANTS AND TREES that take in acidic water.
Here are some trees in Norway that have been killed by acid rain.
Acid rain can also kill life in rivers, streams and lakes.
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In Norway and Sweden, rivers are treated with an alkali called LIME which is
sprayed onto the rivers to NEUTRALIZE the acidity of the RAINWATER.
Acid rain also washes out ions such as calcium and magnesium out of the soil.
These ions are ESSENTIAL IONS for plan growth. Plants need magnesium to
make the green pigment CHLOROPHYLL which traps the sunlight when a plant
makes food by PHOTOSYNTHESIS.
Reducing emission of acid rain gases is expensive. Part of the problem is that
strong winds blow the gases far away from where they were made. Much of the
acid rain in Norway and Sweden comes from power stations in the UK.
Power stations are now fitted with DESU\LPHURIZATION FILTERS to stop
and greatly reduce the amount of sulphur dioxide pollution into the atmosphere.
ULTRA LOW SULPHUR PETROLS are now used for the same reason.
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CATALYTIC CONVERTERS
These are now fitted to ALL cars as STANDARD
Catalytic converters are fitted to car exhaust systems of vehicles to reduce the
level of pollutants being released into the air from burning petrol.
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The catalytic converters change emissions such as carbon monoxide (CO) and
nitrogen oxides like Nitrogen Monoxide NO into less harmful nitrogen gas,
carbon dioxide gas and water vapour.
The catalytic converter is usually made from a mixture of transition metals as an
alloy. If leaded petrol is added by mistake, the catalytic converter stops working
as it is POISONED by the lead compounds which stick or ADSORB to the
surface. The car will stop and the catalytic converter becomes totally useless.
USES OF OXYGEN GAS
In medicine, oxygen is used to help people breathe. It is used in OXYGEN
TENTS. Oxygen is also needed for combustion and burning. It is also used in
WELDING to burn with acetylene (ETHYNE) gas C2H2 to make the acetylene
burn at such a high temperature that it can melt metal.
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The liquid nitrogen and oxygen are then separated by fractional distillation.
Fractional distillation
The liquefied air is passed into the bottom of a fractionating column. Just as in the
columns used to separate oil fractions, the column is warmer at the bottom than it
is at the top.
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The liquid nitrogen boils at the bottom of the column. Gaseous nitrogen rises to the
top, where it is piped off and stored. Liquid oxygen collects at the bottom of the
column. The boiling point of argon (the noble gas that forms 0.9% of the air) is
close to the boiling point of oxygen, so a second fractionating column is often used
to separate the argon from the oxygen.
Uses of nitrogen and oxygen
NITROGEN COMPOUNDS
Fertilizers contain minerals that make a plant grow quick and healthy.
THEY ARE OFTEN CALLED NPK FERTILIZERS. They contain the elements
nitrogen (N) phosphorus (P) and potassium (K). Nitrogen compounds make plant
proteins, phosphorus makes plant roots grow and potassium makes the flowers and
fruits of plants.
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Recycling
At each pass of the gases through the reactor, only about 15% of the nitrogen and
hydrogen converts to ammonia. (This figure also varies from plant to plant.) By
continual recycling of the unreacted nitrogen and hydrogen, the overall conversion
is about 98%.
Explaining the conditions
The proportions of nitrogen and hydrogen
The mixture of nitrogen and hydrogen going into the reactor is in the ratio of 1
volume of nitrogen to 3 volumes of hydrogen.
Avogadro's Law says that equal volumes of gases at the same temperature and
pressure contain equal numbers of molecules. That means that the gases are going
into the reactor in the ratio of 1 molecule of nitrogen to 3 of hydrogen.
That is the proportion demanded by the equation.
In some reactions you might choose to use an excess of one of the reactants. You
would do this if it is particularly important to use up as much as possible of the
other reactant - if, for example, it was much more expensive. That doesn't apply in
this case.
There is always a down-side to using anything other than the equation proportions.
If you have an excess of one reactant there will be molecules passing through the
reactor which can't possibly react because there isn't anything for them to react
with. This wastes reactor space - particularly space on the surface of the catalyst.
The temperature
Equilibrium considerations
You need to shift the position of the equilibrium as far as possible to the right in
order to produce the maximum possible amount of ammonia in the equilibrium
mixture.
The forward reaction (the production of ammonia) is exothermic.
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Notice that there are 4 molecules on the left-hand side of the equation, but only 2
on the right.
According to Le Chatelier's Principle, if you increase the pressure the system will
respond by favouring the reaction which produces fewer molecules. That will
cause the pressure to fall again.
In order to get as much ammonia as possible in the equilibrium mixture, you need
as high a pressure as possible. 200 atmospheres is a high pressure, but not
amazingly high.
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Rate considerations
Increasing the pressure brings the molecules closer together. In this particular
instance, it will increase their chances of hitting and sticking to the surface of the
catalyst where they can react. The higher the pressure the better in terms of the rate
of a gas reaction.
Economic considerations
Very high pressures are very expensive to produce on two counts.
You have to build extremely strong pipes and containment vessels to withstand the
very high pressure. That increases your capital costs when the plant is built.
High pressures cost a lot to produce and maintain. That means that the running
costs of your plant are very high.
The compromise
200 atmospheres is a compromise pressure chosen on economic grounds. If the
pressure used is too high, the cost of generating it exceeds the price you can get for
the extra ammonia produced.
The catalyst
Equilibrium considerations
The catalyst has no effect whatsoever on the position of the equilibrium. Adding a
catalyst doesn't produce any greater percentage of ammonia in the equilibrium
mixture. Its only function is to speed up the reaction.
Rate considerations
In the absence of a catalyst the reaction is so slow that virtually no reaction
happens in any sensible time. The catalyst ensures that the reaction is fast enough
for a dynamic equilibrium to be set up within the very short time that the gases are
actually in the reactor.
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The temperature and pressure are chosen to obtain a good yield of ammonia in a
short time. This involves a compromise between maximum yield and speed of
reaction. As you can see from the graph, ammonia yield is increased by very high
pressures, but decreased by very high temperatures.
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However, lowering the temperature (apart from being difficult at high pressures)
slows down the reaction. The conditions chosen are a compromise between these
conflicting factors:
1. Increasing the pressure above 200 atmospheres would improve yield, but
would also raise temperature, and make the plant much more expensive to
build.
2. Decreasing the temperature below 450C would slow the reaction down too
much and make it hard to maintain sufficiently high pressure.
3. So a balance is struck at 200 atmospheres and 450C.
In addition, two other steps are taken to maximise yield. Firstly, the ammonia is
cooled until it liquefies and is then removed, thus causing more of the nitrogen and
hydrogen to react. Secondly, any unreacted nitrogen and hydrogen is recycled to
give it another chance to react.
AMMONIUM SALTS
These salts will react with a base to produce AMMONIA GAS. An example is the
reaction of ammonium chloride with calcium hydroxide to make ammonia gas.
The gas is tested and is THE ONLY ALKALI GAS. IT TURNS RED LITMUS
BLUE
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Carbon dioxide gas is also formed from combustion. Most of the common fuels
today are HYDROCARBONS They contain only HYDROgen and CARBON.
When a hydrocarbon is burnt in a large amount of air (complete combustion), I
reacts with the oxygen in the air (IT IS OXIDISED) to form CARBON DIOXIDE
AND WATER.
This is an EXAMPLE OF COMBUSTION
FUEL + OXYGEN GAS = CARBON DIOXIDE AND WATER
BURNING METHANE (NATURAL GAS)
Methane + oxygen = carbon dioxide gas + water
CH4(g) + 2O2(g) = 2CO2(g) + 2H2O(l)
Carbon dioxide is also a PRODUCT OF RESPIRATION.
Plants take in carbon dioxide and combine it with water to make sugar + oxygen.
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. . . or by heating sulphide ores like iron pyrites (fools gold) or zinc sulphide (zinc
blende) in a furnace in an excess of air:
In either case, an excess of air is used so that the sulphur dioxide produced is
already mixed with oxygen for the next stage.
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This can then be reacted safely with water to produce concentrated sulphuric acid twice as much as you originally used to make the fuming sulphuric acid.
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The temperature
Equilibrium considerations
You need to shift the position of the equilibrium as far as possible to the right in
order to produce the maximum possible amount of sulphur trioxide in the
equilibrium mixture.
The forward reaction (the production of sulphur trioxide) is exothermic.
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The pressure
Equilibrium considerations
Notice that there are 3 molecules on the left-hand side of the equation, but only 2
on the right.
According to Le Chatelier's Principle, if you increase the pressure the system will
respond by favouring the reaction which produces fewer molecules. That will
cause the pressure to fall again.
In order to get as much sulphur trioxide as possible in the equilibrium mixture, you
need as high a pressure as possible. High pressures also increase the rate of the
reaction. However, the reaction is done at pressures close to atmospheric pressure!
Economic considerations
Even at these relatively low pressures, there is a 99.5% conversion of sulphur
dioxide into sulphur trioxide. The very small improvement that you could achieve
by increasing the pressure isn't worth the expense of producing those high
pressures.
The catalyst
Equilibrium considerations
The catalyst has no effect whatsoever on the position of the equilibrium. Adding a
catalyst doesn't produce any greater percentage of sulphur trioxide in the
equilibrium mixture. Its only function is to speed up the reaction.
Rate considerations
In the absence of a catalyst the reaction is so slow that virtually no reaction
happens in any sensible time. The catalyst ensures that the reaction is fast enough
for a dynamic equilibrium to be set up within the very short time that the gases are
actually in the reactor.
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For hundreds of years, limestone has been heated in a kiln or furnace to make
QUICKLIME or CALCIUM OXIDE (CaO).
LIMESTONE HEATED AT 12000C = QUICKLIME + CARBON DIOXIDE
CaCO3(s) = CaO(s) + CO2(g)
This is an example of THERMAL DECOMPOSITION (BREAKDOWN OF A
SUBSTANCE BY HEAT).
When water is added to quicklime, there is a very fast exothermic reaction
(produces heat) and SLAKED LIME (CALCIUM HYDROXIDE) FORMS.
CALCIUM OXIDE + WATER = SLAKED LIME
CaO(s) + H2O(l) = Ca(OH)2(s)
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Slaked lime is an ALKALI. This is the basis for many of its uses. The major uses
of the products of the limestone cycle LIMESTONE, QUICKLIME AND
SLAKED LIME are shown below.
USES OF LIMESTONE CALCIUM CARBONATE - CaCO3
1.
2.
3.
4.
5.
6.
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ethane :
ethene :
ethanol :
ethanoic acid :
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From the graph above, it can be seen that as the number of carbon
atoms in the organic compound increases the boiling points
increase.
Also, the boiling points tend to follow a straight line with the higher
members of each group i.e. the difference between boiling points
tends towards a single value.
The four homologous series studied at IGCSE are alkanes,
alkenes, alcohols and carboxylic acids. The names and
formulae of these compounds will be dealt with in separate
sections.
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The members of each series differ from each other by the number
of carbon atoms contained in the molecule,
alkane
alkene
alcohol
carboxylic
acid
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3-ethylpentane :
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CO2(g) + 2H2O(g)
in general,
CnH2n+2(g) + (1.5n+0.5)O2(g)
nCO2(g) + (n+1)H2O(g)
If there is not enough oxygen present then instead of carbon dioxide, carbon
monoxide, CO, is produced. Carbon monoxide is particularly toxic and absorbed
into blood, through respiration, very easily. For domestic heating systems it is
particularly important that enough air can get to the flame to avoid carbon
monoxide being generated in the home. Car engines also require a lot of air and
there is a lot of research going on to make the internal combustion engine more
efficient, and so put out less carbon monoxide.
Note also that both alkanes and carbon dioxide are green house gases, i.e. they
trap infra-red (i.-r.) radiation inside the Earth's atmosphere, gradually increasing
global temperatures.
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(ii) Halogenation The only other reaction that an alkane will undergo is a reaction with a halogen (
chlorine or bromine typically ) with UV light present as an initiator of the reaction,
e.g. CH4(g) + Br2(g)
CH3Br(g) + HBr(g)
The UV light causes the formation of free radical halogen atoms by providing
enough energy for the bond between the two halogen atoms to break.
A halogen atom attacks the alkane, substituting itself for a hydrogen atom. This
substitution may occur many times in an alkane before the reaction is finished.
A similar process occurs high up in the earth's atmosphere when CFC's and other
organic solvents react with intense sunlight to produce free radicals, chlorine atoms
in this case. These attack molecules of ozone ( O3 ) depleting ozone's concentration
and leading to the "holes".
(3) Crude oil :
(i) Fractional distillation Crude oil is a mixture of many different hydrocarbon compounds, some of
them liquid and some of them gases. These compounds can be separated
because the different length of alkanes will have different boiling points.
The crude oil is heated up to about 350 oC and is fed into a fractionating
column, as in the diagram below,
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The vapours with the lowest boiling points pass all the way up the column and
come off as gases, e.g. methane, ethane and propane. The temperature of the
column gradually decreases the higher up the vapours go, and so various
fractions will condense to liquids at different heights.
The fractions with the highest boiling points do not vaporize and are collected
at the bottom of the fractionating column, e.g. bitumen
Here is a table with some boiling points for the commonest fractions :
Fraction name
Refinery gases
Gasoline
Naphtha
Kerosene
Diesel oil
Fuel oil
Bitumen
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(ii) Cracking In industry the fractions obtained from the fractional distillation of crude oil
are heated at high pressure in the presence of a catalyst to produce shorter
chain alkanes and alkenes.
e.g. C10H22
C5H12 + C5H10
(iii) Reforming This is a process where straight chain alkanes are turned into branched alkanes
and cyclic alkanes are turned into aromatic compounds.
Both these reactions result in the formation of chemicals that improve the
performance of fuels as well as enable more exotic compounds to be made.
Organic Chemistry - Alkenes
(1) Names :
Alkenes all have a C=C double bond in their structure and their names follow this
pattern,
C2H4 - ethene
C3H6 - propene
C4H8 - butene
C5H10 - pentene
The general chemical formula for an alkene is CnH2n.
(2) Addition reactions of alkenes :
(i) Bromination The double bond of an alkene will undergo an addition reaction with aqueous
bromine to give a dibromo compound. The orange bromine water is decolourised
in the process.
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(ii) Hydrogenation Alkenes may be turned into alkanes by reacting the alkene with hydrogen gas at a
high temperature and high pressure. A nickel catalyst is also needed to accomplish
this addition reaction.
e.g. ethene reacts with hydrogen to give ethane,
This reaction is also called saturation of the double bond. In ethene the carbon
atoms are said to be unsaturated. In ethane the carbon atoms have the maximum
number of hydrogen atoms bonded to them, and are said to be saturated.
(iii) Oxidation The carbon-carbon double bond may also be oxidised i.e. have oxygen added to it.
This is accomplished by using acidified potassium manganate(VII) solution at
room temperature and pressure. The purple manganate(VII) solution is
decolourised during the reaction.
e.g. ethene reacts with acidified potassium manganate(VII)(aq) to give ethan-1,2diol,
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ethene
propene
phenylethene
methyl methacrylate
chloroethene
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polymer- a material produced from many separate single monomer units joined up
together.
An addition polymer is simply named after the monomer alkene that it is prepared
from,
e.g. ethene makes poly(ethene)
propene makes poly(propene)
phenylethene makes poly(phenylethene)
chloroethene makes poly(chloroethene)
methyl acrylate makes poly(methyl acrylate)
The structure above shows just 4 separate monomer units joined together. In a real
polymer, however, there could be 1000's of units joined up to form the chains. This
would be extremely difficult to draw out and so the structure is often shortened to a
repeat unit. There are 3 stages to think about when drawing a repeat unit for a
polymer
[ANDREW RICHARD WARD ALL RIGHTS RESERVED A WARD TUITION - 2009]
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____
This rigidity means that once this type of polymer has been
formed, the structure prevents the material from being melted.
This is called a thermosetting polymer and will only char or burn
when heated.
e.g.vulcanised rubber for car tyres
resins for gluing
[ANDREW RICHARD WARD ALL RIGHTS RESERVED A WARD TUITION - 2009]
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(i) Preparation of ethanol by fermentation Ethanol is prepared in the laboratory and in the alcoholic drinks
industry, by the process of fermentation. This involves the use of
an enzyme ( yeast ) that changes a carbohydrate, e.g. sucrose,
into ethanol and carbon dioxide gas,
C6H12O6(aq)
2CH3CH2OH(aq) + 2CO2(g)
The yeast used requires a certain temperature to be active somewhere between 15 and 37 C. Too high a temperature and
the yeast "dies" and too low a temperature causes the yeast to
become dormant.
The production of carbon dioxide gas can be monitored by bubbling
any gases produced during the reaction through limewater (
calcium hydroxide(aq) ). The formation of a white precipitate (
calcium carbonate ) in the limewater shows that carbon dioxide has
been given off.
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In a process similar to that of crude oil, the ethanol/water mixture can be separated
by fractional distillation because of the difference in boiling points.
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Ethanol boils at 79 C and water boils at 100 C, so that ethanol boils first and
therefore comes over through the condenser first. The fractionating column allows
the vapours to condense and drop back down into the round-bottom flask, stopping
water vapour from passing through into the condenser
(ii) Dehydration of ethanol Experimental sheet for the dehydration of ethanol.
All alcohols contain hydrogen and oxygen ( as well as carbon ) and these atoms
can be removed from an alcohol as a molecule of water ( H2O ). This type of
reaction is called dehydration. It can be accomplished by passing alcohol vapour
over a heated aluminium oxide catalyst.
e.g. ethanol can be turned into ethene,
CH3CH2OH(g)
CH2=CH2(g) + H2O(g)
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Diagram -
Method
Set up a boiling tube in a beaker of cold water as in the diagram above. Add about
a few cm's depth of ethanol followed by another cm of ethanoic acid from a bottle.
Then carefully add a few drops of concentrated sulphuric acid (CARE: very
dangerous) to the boiling tube. Add an anti-bumping granule or two, and heat up
the water bath until the reaction mixture in the boiling tube starts to boil gently.
Keep the reaction boiling gently for about 15 minute. Then raise the boiling tube
out of the water bath and leave to cool.
Carefully add some sodium or calcium carbonate to the boiling tube until no more
fizzing is produced. Filter the solution and carefully smell the clear liquid
remaining.
Note down all your observations during this reaction.
Carboxylic acids will react with alcohols to produce organic
compounds called esters.
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(i) Fats :
These natural materials contain the ester link found in the
synthetic polyesters shown above.
They may be hydrolysed ( broken down ) by a reaction with
sodium hydroxide (a strong base) and heat.
Once hydrolysed they form soaps ( sodium salts of carboxylic
acids ) and glycerol ( propan-1,2,3-triol ).
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(ii) Proteins :
These naturally occurring materials contain the amide link found in
the synthetic polyamides shown above.
These compounds may also be hydrolysed by a reaction with
enzymes and/or aqueous acid. Proteins in the food we ingest are
broken down by stomach acids and enzymes which work at body
temperature.
Once hydrolysed they form amino acids which can then be used
by the human body to prepare vital chemicals needed to sustain
life.
(5) Saponification :
Saponification means "soap-making" and is a reaction in which a
fat, or oil, is turned into a salt of a carboxylic acid.
The oil is heated with a concentrated solution of a caustic base,
such as sodium hydroxide. The base breaks down the ester links,
forming alcohol groups and carboxylate ion groups on different
molecules.
FINALLY BEFORE WE GO
THE FINAL PAGE SUMMARIZES ALL TH REACTIONS IN ORGANIC
CHEMISTRY AT IGCSE.
BEST OF LUCK WITH YOUR STUDIES. YOU DESERVE TO DO WELL
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