Transcript Document

Standard Grade Chemistry
• This series of presentations is
designed to help you revise for
Standard Grade Chemistry.
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Chemical Reactions
Speed of Reaction The Periodic Table
How Atoms Combine
Properties of Substances
Acids and Alkalis
Metals
Carbohydrates
Hydrocarbons
Fuels
Chemical Arithmetic
Reactions of Acids
Metals and Electricity
Plastics and Synthetic Fibres
Click here to finish
Corrosion
Fertilisers
Chemical
Reactions
Chemical Reactions
• A chemical
reaction involves
the formation of
new substances.
• How do we know
that a chemical
reaction has
taken place?
• There is a change
in appearance
• A precipitate is
formed
• A gas is given off
• Energy is released
or taken in
Chemical Reactions
• In an exothermic reaction energy
is released.
• In an endothermic reaction energy
is taken in.
Elements and Compounds
• All the substances
in the world are
made from about
100 elements,
each of which has
a name and a
symbol.
• A compound whose
name ends in "ide"
contains two
elements only.
• A compound whose
name ends in "ite" or
"ate" contains three
elements, one of
which is oxygen,
("ite" has less
oxygen than "ate“)
Solutions
• A solution is
formed when a
material dissolves
in a liquid.
• The material
which dissolves is
called a solute.
• The liquid is
called a solvent.
• A material which can
dissolve is soluble.
• A material which
cannot dissolve is
insoluble.
• A saturated solution
is one where no
more solute can
dissolve.
• A table of solubility
is found in the Data
Booklet.
Chemical Reactions
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Speed of
Reaction
Speed of Reaction
The speed of a
chemical reaction
is increased when
• the size of the
reacting particles
is reduced
• the concentration
of the reacting
materials
increases
• the temperature
is increased
Catalysts
• A catalyst is a
substance which
• speeds up the rate of
a chemical reaction.
• is not used up in the
reaction
• is not changed in the
reaction
• is used in industry to
reduce energy costs.
Enzymes
• Enzymes are
biological
catalysts,
produced by
living things.
• Enzymes are use
to make:
• Alcohol
• Medicines
• Yoghurt
• Washing powders
Speed of Reaction
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The Periodic
Table
The Periodic Table
• The elements are classified by
arranging them in the Periodic
Table.
• The atoms in the Periodic Table are
listed in order of their Atomic
Number.
• Elements in the Periodic Table can
be classified in different ways.
• Solid, liquid or gas
• Metal or non-metal.
• Naturally occurring or man-made
• The vertical columns are called
groups.
• All the elements in any group will
show similar chemical properties.
• Group 1 elements are called the Alkali
Metals.
• Group 7 elements are called the
Halogens.
• Group 0 (or 8) gases are called the
Noble Gases.
• The central block of the Periodic Table
contains the Transition Metals.
Rutherford’s Atom
• Elements are made
of small particles
called atoms
• In the centre of the
atom is the nucleus,
containing protons
and neutrons.
• Electrons orbit
around the nucleus,
like planets around
the Sun.
Atoms
• Most of the mass of the atom is
found in the nucleus
• The nucleus contains positively
charged protons.
• The nucleus also contains
neutrons, which have no charge.
• Negatively charged electrons orbit
around the nucleus.
Sub-Atomic Particles
Particle
Charge
Mass
Location
Proton
positive
1 a.m.u.
nucleus
Neutron
none
1 a.m.u.
nucleus
Electron
negative
negligible
In orbit
around the
nucleus
Atoms
• For each atom the Atomic Number
is equal to the number of protons.
• The Mass Number is the number of
protons + neutrons.
• The number of neutrons is Mass
Number minus Atomic Number.
• The atom is neutral because the
positive charge of the nucleus is
balanced by the negative charge of
the electrons.
• Thus the number of electrons is the
same as the number of protons.
Electrons
• The first shell
holds 2 electrons.
Electrons
• The first shell
holds 2 electrons.
• The second shell
holds 8 electrons
Electrons
• The first shell
holds 2 electrons.
• The second shell
holds 8 electron
• The third shell
holds 8 electrons
• Since electrons
are impossible to
track down can
also show them
pear-shaped in
electron pair
clouds
• Each cloud can
hold two
electrons
• The number of outer electrons in
an atom is the same as the number
of its group in the Periodic Table.
• Atoms with the same number of
outer electrons will have similar
chemical properties.
Isotopes
• Not all atoms of the same element
have the same mass. Most
elements are mixtures of isotopes.
• Isotopes are atoms with the same
number of protons but different
numbers of neutrons.
• Relative Atomic Mass is the
average mass number of an atom.
• It is not whole number because
most elements consist of a mixture
of isotopes.
• Different isotopes have different
abundances.
Atoms and the Periodic
Table
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How Atoms
Combine
How Atoms Combine
• Atoms react in such a way as to
achieve a stable electron
arrangement where they have a
full outer electron shell. (Usually 8
electrons)
How Atoms Combine
• They are trying to achieve a Noble
Gas structure.
• This means that they are trying to
get the same electron arrangement
as the nearest Noble Gas.
The Covalent Bond
• As two atoms
come together
the half-filled
electron pair
clouds overlap to
form a new cloud.
The Covalent Bond
• As two atoms
come together
the half-filled
electron pair
clouds overlap to
form a new cloud.
The covalent bond
The Covalent Bond
• In a covalently
bonded molecule
the two atoms are
held together
because both
nuclei are
attracted to the
shared pair of
electrons.
+
e-e-
+
Molecules
• Atoms are held together by bonds.
• A covalent bond is formed between
two atoms when they share a pair
of electrons.
• Covalent bonds are formed
between two non-metal atoms
Molecules
• A molecule is a group of atoms,
held together by covalent bonds.
• The molecular formula gives the
number of atoms of each type in a
covalent molecule.
• A diatomic molecule is one
containing two atoms.
• We can write
formulae by
counting the
atoms in a model
or picture.
• We can write
formulae by
counting the
atoms in a model
or picture.
Diatomic molecules
• Hydrogen, nitrogen, oxygen, the
halogens and carbon monoxide
exist as diatomic molecules.
• We can draw diagrams to show the
formation of diatomic molecules.
• We can draw diagrams to show the
formation of diatomic molecules.
+
H +
H
• We can draw diagrams to show the
formation of diatomic molecules.

+
H +
H

H2
+
Cl +
Cl

+
Cl +
Cl

Cl2
• Some diatomic molecules involve
more than one covalent bond.
O2
N2
• We can represent
these molecules,
using as a
covalent bond.
• H2
H
H
• Cl2
Cl
Cl
• O2
O
O
• N2
N
N
Covalent molecules
• We can draw similar diagrams of
discrete covalent molecules.
ammonia NH3
methane CH4
• Once again we can represent those
more simply:
H
H
N
H
ammonia NH3
H
H
C
H
H
methane CH4
Valency
• Valency is a
number which
helps us work out
molecular
formulae.
• It is the
combining power
of the atom.
• Valency is
• Group Number
• 8 minus Group
Number
• Size of charge on
ion
• Number after
metals name e.g.
copper(II)
Valency
Group
1
2
3
4
5
6
7
0/8
Valency
1
2
3
4
3
2
1
0
Chemical Formulae
1.
2.
3.
4.
5.
Using valency
Write down symbols
Write down valencies
Swap over
Divide (if possible)
Formula
Chemical Formulae
Using valency carbon and oxygen
1. Write down symbols
C O
2. Write down valencies
4 2
3. Swap over
2 4
4. Divide (if possible)
1 2
5. Formula
CO2
Chemical Formulae
Using valency calcium and chlorine
1. Write down symbols
Ca Cl
2. Write down valencies
2
1
3. Swap over
1
2
4. Divide (if possible)
5. Formula
CaCl2
Chemical Formulae
Using valency copper(II) nitrate
1. Write down symbols
Cu NO3
2. Write down valencies
2
1
3. Swap over
1
2
4. Divide (if possible)
5. Formula
Cu(NO3)2
How Atoms Combine
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Fuels
Fuels
• A fuel is a chemical which burns,
releasing energy.
• An exothermic reaction is one in
which heat is released.
• Combustion is the reaction of a
substance with oxygen, in which
energy is given out.
• The test for oxygen is that it
relights a glowing splint.
• The two main gases in air are
oxygen (about 20%) and nitrogen
(about 80%).
Fuels
• In any chemical reaction breaking
bonds takes in energy while
forming bonds releases energy.
• In an exothermic reaction the
energy released by forming the
bonds in the products is greater
than the energy taken in to break
bonds in the reactants.
Fossil Fuels
• Fossil fuels are coal, oil and natural gas
which have been formed by the decay
of natural materials which lived millions
of years ago.
• Coal, oil and gas are finite resources i.e.
the Earth has only limited quantities.
• A fuel crisis will occur when the amount
of these fuels is no longer sufficient to
supply our needs cheaply.
• Coal was formed from the decay of
forests and vegetation which covered
the earth 500-600 million years ago.
• Layers built up until the heat and
pressure changed the organic material
to coal.
• Oil and natural gas were formed in a
similar way, except that they probably
came from marine plants and animals,
compressed by layers of sand on the sea
bed.
• Both coal and oil contain sulphur.
• When the fuels burn the sulphur
produces a gas called sulphur dioxide.
• This causes pollution since it dissolves
in water to form sulphuric acid (acid
rain).
• Oil causes pollution problems if it is
spilled in water because it does not
dissolve in water and is poisonous to
marine life.
Oil
• All substances have there own
particular melting point and boiling
point.
• Crude oil is a mixture of compounds
which can be to split it into fractions.
• A fraction is a group of chemical
compounds, all of which boil within the
same temperature range. Oil can be
separated into fractions by the process
of fractional distillation.
Fractional Distillation of Oil
gases
(gaseous fuel)
petrol (gasoline)
(petrol)
naphtha
(chemicals)
Heated
oil from
furnace
paraffin (kerosine)
(aircraft fuel)
diesel
(fuel for lorries etc.)
residue (wax, tar)
Oil Fractions
Name
Uses
Gases
Carbon atoms
per molecule
1 to 4
Petrol
4 to 9
Fuel for cars
Naphtha
8 to 14
Chemicals
Paraffin
10 to 16
Aircraft fuel
Diesel
15 to 20
Lorry fuel
Residue
More than 20
Lubricating oil,
tar, wax etc.
Fuel
Oil Fractions
• Viscosity is a measure of the
thickness of a liquid.
• Flammability is a measure of how
easily the liquid catches fire.
• Volatility means how easy it is to
turn the liquid into a gas.
• As the boiling point of a fraction
increases then:
• it will not evaporate as easily.
• it will be less flammable
• it will be more viscous (thicker).
• Moving through the fractions from
gases to the residue
• The molecules present in the
fraction are longer and heavier
• They will find it more difficult to
become a gas i.e. they will be less
easy to evaporate.
• Moving up the fractions from gases to
the residue
• Since combustion involves the reaction
of gas molecules with oxygen
flammability will decrease.
• Increased molecular lengths mean that
molecules become more "tangled up",
so the liquid will become thicker (more
viscous).
Tests
• The test for carbon dioxide is that
it turns lime water cloudy.
• The test for water is that it turns
anhydrous copper sulphate from
white to blue.
• Hydrocarbons burn to produce
carbon dioxide and water only.
To pump
Burning candle
Anhydrous
copper
sulphate
(turns blue)
Lime water
(turns cloudy)
Hydrocarbons
• When a hydrocarbon fuel burns to give
carbon dioxide and water then:
• The carbon in the carbon dioxide and
the hydrogen in water must have come
from the fuel.
• Crude oil is mainly made of compounds
called hydrocarbons (i.e. made of
carbon and hydrogen only).
Incomplete Combustion
• When fuels burn in a limited supply
of air then incomplete combustion
takes place and the poisonous gas,
carbon monoxide (CO) is produced.
• Increasing the amount of air used
to burn fuel improves efficiency
and decreases pollution.
Other products of
combustion
• Fossil fuels contain sulphur which
produces sulphur dioxide when the
fuel is burned.
• The oil industry tries to remove
this sulphur from the fuels before
selling them.
Nitrogen does not
react well because of
its strong bonds.
If there is a high
temperature the
nitrogen and oxygen
will combine to
make nitrogen
oxides.
The experiment
opposite shows how a
high voltage spark,
like one provided by
the spark plug or
lightning will do the
same.
High
Voltage
spark
Air
+
-
Nitrogen does not
react well because of
its strong bonds.
If there is a high
temperature the
nitrogen and oxygen
will combine to
make nitrogen
oxides.
The experiment
opposite shows how a
high voltage spark,
like one provided by
the spark plug or
lightning will do the
same.
High
Voltage
spark
Air
Brown
gas
+
-
Atmospheric Pollution
• The sulphur and nitrogen oxides
produced can dissolve in water,
making acid rain.
• Unburnt hydrocarbons escaping
from car exhausts can help cause
the destruction of the ozone layer.
Reducing Pollution
• Air pollution caused by burning
hydrocarbons can be reduced by:
• using a special exhaust system – a
catalytic converter, in which metal
catalysts (platinum or rhodium)
will convert pollutants into
harmless gases.
• altering the fuel to air ratio.
Fuels
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Hydrocarbons
Homologous Series
• An homologous series is a series
of carbon compounds.
• The alkanes, the alkenes and the
cycloalkanes are examples of
homologous series.
• In a homologous series:
• all members can be represented by
a general formula.
• there is a gradation in physical
properties.
• there is a similarity in chemical
properties
Isomers H
• Isomers are
compounds with
the same
molecular formula
but different
structural
formulae
• For example C4H10
H
H
H
H
H
C
C
C
C
H
H
H
H
H
H
H
C
C
C
C
H
H
H
H
H
H
H
Alkanes
• Alkanes
• All members have a name ending
in -ane.
• Alkanes have a general formula
CnH2n+2
• Alkanes are used as gaseous and
liquid fuels, as well as wax and tar.
Alkanes
• Alkanes
• As we move down the alkanes the
boiling point increases.
• This is because the molecular size
increases, making it more difficult
to change a molecule from liquid
into gas.
Alkanes
• The alkanes,
general formula
CnH2n+2
•
•
•
•
•
•
•
•
•
•
methane
ethane
propane
butane
pentane
hexane
heptane
octane
nonane
decane
CH4
C2H6
C3H8
C4H10
C5H12
C6H14
C7H16
C8H18
C9H20
C10H22
Cycloalkanes
• The cycloalkanes,
general formula
CnH2n
• cyclopropane
C3H6
• cyclobutane C4H8
• cyclopentane
C5H10
• cyclohexane C6H12
Alkenes
• The alkenes,
general formula
CnH2n
•
•
•
•
•
ethene
propene
butene
pentene
hexene
C2H4
C3H6
C4H8
C5H10
C6H12
Alkenes contain a carbon to carbon double bond:
C C
H C C H
ethene
H H
H
H C C C H
propene
H H H
H H
H C C C C H
H H H H
butene
Isomers
• The alkenes and the
cycloalkanes are
isomers.
• They both have the
same general
formula CnH2n
• They have different
structural formulae,
as shown.
CH3
CH
CH2
propene C3H6
CH2
CH2
CH2
cyclopropane C3H6
Saturated and
Unsaturated
• Saturated hydrocarbons contain
only carbon to carbon single bonds
• Unsaturated hydrocarbons contain
carbon to carbon double or triple
bonds.
Unsaturated Hydrocarbons
• The test for unsaturation is that
unsaturated hydrocarbons decolourise
bromine water.
• An addition reaction takes place when a
carbon to carbon double bond breaks
and other atoms add on.
• If hydrogen is added to an alkene then
an alkane is formed.
Cracking Hydrocarbons
• Cracking of long-chain
hydrocarbons produces smaller,
more useful molecules.
• These molecules are unsaturated.
mineral wool
soaked in oil
catalyst
heat
gas
Catalytic Cracking
• A catalyst lowers the temperature
at which cracking takes place.
• Cracking produces some
unsaturated hydrocarbons because
there are not enough hydrogen
atoms to produce completely
saturated products.
Hydrocarbons
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Plastics and
Synthetic
Fibres
Plastics and
Synthetic Fibres
• Most plastics and synthetic (i.e. manmade) fibres come from oil.
• Plastics are selected for various uses,
according to their properties e.g.
lightness, durability, electrical and
thermal insulation.
• Biodegradable means "able to rot
away". Most plastics are not
biodegradable and so cause
environmental problems of disposal.
Burning plastics
• Certain plastics burn or smoulder
to give poisonous fumes.
• All plastics can release carbon
monoxide.
• P.V.C. can release hydrogen
chloride
• Polyurethane releases hydrogen
cyanide.
Thermoplastic or
Thermosetting?
• A thermoplastic plastic is one which can
be melted or reshaped (examples
polythene, polystyrene, P.V.C.)
• A thermosetting plastic is one which
cannot be melted and reshaped
(examples bakelite in electrical fittings,
formica in worktops)
Polymerisation
• A monomer is a small molecule which is
able to join together with other,
similar, small molecules.
• A polymer is the large molecule
produced.
• This process is called polymerisation.
• Plastics and fibres (natural and
synthetic) are examples of polymers.
The making of plastics and synthetic
fibres are examples of polymerisation.
Naming polymers
• Many polymers are made from the
small unsaturated molecules,
produced by the cracking of oil.
• The name of the polymer is derived
from its monomer.
Naming polymers
MONOMER
***ene
ethene
propene
styrene
chloroethene
tetrafluoroethe
POLYMER
poly(***ene)
poly(ethene)
poly(propene)
poly(styrene)
poly(chloroethene)
poly(tetrafluoroethene)
Addition Polymerisation
• The small unsaturated molecules add to
each other by opening up their carbon
to carbon double bonds.
• This process is called addition
polymerisation.
CH2=CH2 + CH2=CH2 
Addition Polymerisation
• The small unsaturated molecules add to
each other by opening up their carbon
to carbon double bonds.
• This process is called addition
polymerisation.
CH2=CH2 + CH2=CH2  -CH2-CH2-CH2-CH2-
The repeat unit is (-CH2-)n
I*
H
H
C
C
H
H
The ethene is attacked by an initiator (I*) which
opens up the double bond
I
H
H
H
H
C
C* C
C
H
H
H
H
The ethene is attacked by an initiator (I*) which
opens up the double bond
Another ethene adds on.
I
H
H
H
H
H
H
C
C
C
C*
C
C
H
H
H
H
H
H
The ethene is attacked by an initiator (I*) which
opens up the double bond
Another ethene adds on.
Then another
I
H
H
H
H
H
H
C
C
C
C
C
C*
H
H
H
H
H
H
The ethene is attacked by an initiator (I*) which
opens up the double bond
Another ethene adds on.
Then another
….
Repeat Units
• You should be able to look at the
structure of a polymer and work
out the repeat unit and the
monomer(s) from which it was
formed.
• The repeat unit of an addition
polymer is always only two carbon
atoms long.
Plastics and Synthetic Fibres
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Carbohydrates
Carbohydrates
• Carbohydrates are important food
for animals.
• Carbohydrates contain the
elements carbon, hydrogen and
oxygen.
• There are two hydrogen atoms for
each oxygen atom in
carbohydrates
Photosynthesis
• Photosynthesis is the process by
which plants make carbohydrates
and oxygen from carbon dioxide
and water, using light energy.
6CO2 +
6H2O + energy  C6H12O6 + 602
• Chlorophyll (the green colour in
plants) is used to absorb the light
energy.
Respiration
• Respiration is the process by which
animals AND plants obtain the supply of
energy that they need for growth,
movement, warmth etc.
• They obtain this energy by breaking
down the carbohydrate, glucose, using
oxygen:
C6H12O6 + 602  6CO2 + 6H20 + energy
• Carbohydrates burn, releasing energy
and producing carbon dioxide and water
The Atmosphere
• The combination of respiration and
photosynthesis lead to the balance of
carbon dioxide/oxygen in the
atmosphere.
• The clearing of forests with the loss of
green plants, reduces the amount of
photosynthesis taking place. This could
alter the balance of the atmosphere,
with a consequent danger to life on
Earth.
Glucose
•
•
•
•
Glucose is a carbohydrate
Glucose is sweet
Glucose dissolves well in water
A beam of light can pass through
glucose solution.
• Benedict's solution will give an
orange precipitate with glucose.
Sucrose
•
•
•
•
Sucrose is a carbohydrate
Sucrose is sweet
Sucrose dissolves well in water
A beam of light can pass through
sucrose solution.
• Benedict's solution will NOT give
an an orange precipitate with
sucrose.
Starch
•
•
•
•
Starch is a carbohydrate
Starch is not sweet
Starch does not dissolve in water
A beam of light cannot pass
through starch solution.
• When iodine is added to starch a
blue/ black colour is produced.
Testing Carbohydrates
• Benedict's solution (or Fehling's
solution) gives a positive test (an
orange colour) with glucose, fructose,
maltose and other sugars but NOT
sucrose.
• The pairs of carbohydrate
glucose/fructose (C6H12O6) and
sucrose/maltose (C12H22O11) are
isomers because they both have the
same molecular formula but different
structural formulae.
Testing Carbohydrates
• Starch gives a blue/black colour
when added to iodine.
Types of carbohydrates
• Monosaccharides are simple sugars
with formula C6H12O6.
• Disaccharides are simple sugars
with formula C12H22O11.
• Polysaccharides are complex
sugars with formula (C6H10O5)n.
Condensation
Polymerisation
• Glucose is a carbohydrate made in
photosynthesis.
• Two glucose molecules join to form
sucrose.
• This is a condensation reaction.
2C6H12O6 C12H22O11 + H2O
• Glucose monomers polymerise to form
starch.
• This is a condensation polymerisation.
nC6H12O6  (C6H10O5)n + nH2O
Hydrolysis
• Hydrolysis takes place when large
molecules are broken down into
smaller molecules by the addition
of small molecules, such as water.
• The breakdown of sucrose and
starch are examples of a hydrolysis
reactions.
Digestion
• During digestion starch molecules are
broken down by the body into smaller
glucose molecules that can pass
through the gut wall into the
bloodstream.
• The breakdown of starch is brought
about using acid or the enzyme
amylase.
• Enzymes, such as amylase, are
biological catalysts
Enzymes
• Enzymes, such as amylase, are
biological catalysts
• An enzyme will work most
efficiently within very specific
conditions of temperature and pH.
• The further conditions are removed
from the ideal the less efficiently
the enzyme will perform.
Digestion
• Sucrose and starch molecules break
down by the addition of water:
C12H22O11
+
H2O  C6H12O6
+ C6H12O6
sucrose
glucose fructose
(C6H10O5)n + nH2O  n C6H12O6
starch
glucose
Monosaccharides have formula C6H12O6.
Disaccharides have formula C12H22O11.
Alcohol
• Alcoholic drinks can be made from
any fruit or vegetable source that
is a source of sugars.
• The enzymes in yeast act as
catalysts in the formation of
alcohol.
Fermentation
• Fermentation is the breakdown of
glucose to form carbon dioxide and
alcohol
C6H12O6  2 CO2 + 2 C2H5OH
• The maximum concentration of alcohol
that can be produced is limited because
an increase in alcohol concentration
limits the efficiency of the yeast.
Distillation of alcohol
• Since alcohol boils at 80oC and
water boils at 100oC distillation of
an alcohol solution increases the
alcohol concentration.
• Alcohol is a member of the alkanol
family, called ethanol
Carbohydrates
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Chemical
Arithmetic
Formula Mass
• Formula mass is found by adding
together the relative atomic masses of
all the atoms present in the formula,
e.g. calcium carbonate CaCO3
• Ca
40
• C
12
• O
16x3
48
• Formula Mass
100
Percentage Composition
• The percentage composition is found as
follows.
• Find the formula mass of the compound.
• Find the fraction made up by the
element required.
• Convert that fraction to a percentage.
Percentage = mass of element required x 100
formula mass of compound
Percentage Composition
• Find the percentage of nitrogen in
ammonium nitrate:
• Formula NH4NO3
• Formula mass = 80
• Nitrogen makes up 28 out of 80
• %N = (28/80) x100 = 35%
Empirical Formula
• The empirical (simplest) formula is
found as follows.
• Take the masses (or percentages)
of each element present.
• Divide the mass of each element by
its relative atomic mass.
• Convert these numbers into a
simple, whole number ratio.
Empirical Formula
• Calculate the empirical formula for the
compound which is 54% calcium, 43%
oxygen and 3% hydrogen.
• Symbol % Divide by RAM
Ratio
Ca
54
54/40 = 1.35
1
O
40
43/16 = 2.7
2
H
3
3/1 = 3
2
Formula is CaO2H2 or Ca(OH)2
Moles
• To connect gram
formula mass, mass
in grams and number
of moles use the
triangle opposite
• gfm = mass of 1
mole
• n = number of moles
• m = mass of
substance
m
n
gfm
Chemical Equations
• Reactants are the materials with which
are present at the start of the reaction
and are changed by the reaction.
• Products are the materials produced by
the chemical change.
• These are separated by an arrow (which
means “gives”).
Reactants  Products
Chemical Equations
• Whenever we write a chemical
equation we need to know
• what substances are present at the
start
• what are the new substances
formed in the chemical reaction.
Chemical Equations
• To know the chemical reactants
and products means we can write a
word equation
• Here we are naming the reactants
and products. e.g.
propane + oxygen  carbon dioxide + water
Chemical Equations
• We need to convert the word equation
into symbols:
C3H8 + O2  CO2 + H2O
If we look closely at this equation we
will realise that it is unbalanced – there
are different numbers of atoms on each
side:
3xC + 8xH + 2xO  C + 2xH + 3xO
Chemical Equations
• We must write a balanced chemical
equation where there are equal
numbers of moles of each type of
atom on both sides.
• We can balance the equation we
have been working with.
Balancing Equations
• Propane has 3 carbons so:
C3H8 + O2  3CO2 + H2O
• Propane has 8 hydrogens so:
C3H8 + O2  3CO2 + 4H2O
• To balance out the oxygens:
C3H8 + 5O2  3CO2 + 4H2O
This is a balanced chemical equation.
Using Chemical Equations
• The numbers we use to balance an
equation are the actual numbers of
moles which react.
• This gives us the mole relationship in
the reaction.
• If we look at the example we have been
given:
C3H8 + 5 O2  3CO2 + 4H2O
1 mole + 5 moles  3moles + 4moles
Using Chemical Equations
• Since one mole is the formula weight in
grams we can now work out the masses
which react.
C3H8 + 5 O2  3CO2 + 4H2O
• 1 mole 5 moles  3mole 4moles
• 1x44g 5x32g  3x44g 4x72g
• 44g
160g  132g
72g
• Now by proportion we can work out any
reacting quantities.
Using Chemical Equations
• How much oxygen is needed to burn
0.22g of propane?
C3H8 + 5 O2  3CO2 + 4H2O
1 mole 5 moles 3mole 4moles
• To burn, 1 mole C3H8 needs 5 moles O2
• 44g C3H8 needs 160g O2
• 0.22g C3H8 needs 0.8g O2
Concentration of Solutions
• To connect volume,
concentration and
molarity of a solution
use the triangle
opposite.
• c = concentration
(m/l)
• n = number of moles
• v = volume (l)
n
c
v
Using Chemical Equations
• How much sodium carbonate would
dissolve in 500ml of 0.5 m/l sulphuric
acid?
•
•
•
•
Na2CO3 + H2SO4  Na2SO4 + CO2 + H2O
1 mole
1 mole  1mole
1 mole 1 mole
500ml of 0.5 m/l sulphuric acid contain
0.5 x 0.5 = 0.25 moles of acid.
0.25 moles of sodium carbonate react with
0.25 moles of sulphuric acid.
1 mole sodium carbonate = 106g
• 0.25 moles sodium carbonate = 26.5g
Acid/Alkali Titrations
• Work out unknown
concentrations and
volumes from the
results of volumetric
titrations.
• You use the equation
VH MH NH =
VOH MOH NOH
V = volume
M = molarity
N = number of
H/OH
H = acid
OH = alkali
Chemical Arithmetic
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Properties of
Substances
Conductivity
• An electric current is a flow of
electrons.
• Conductors are materials which
allow an electric current to pass
through.
• Insulators are materials which do
not allow an electric current to
pass through.
Conductors
• Metals
• Graphite (a form of carbon - the
only non-metallic conductor)
• Solutions of ionic metal
compounds
• Molten ionic compounds
Conductivity
• A metal conducts
because of its
metallic bonding.
• In metallic
bonding the outer
electrons can
jump from atom
to atom, and thus
move through the
solid.
Ions and Conductivity
• Ions move through liquids.
• Positive ions are formed when
atoms lose electrons.
• Negative ions are formed when at
atoms gain electrons.
Ions
• Ions are charged particles
• Atoms gain or lose electrons to
achieve the Noble Gas Structure.
• Positive ions are formed when
metal atoms lose electrons.
• Negative ions are formed when at
non-metal atoms gain electrons.
Ions
Gr.
1
2
3
e- to
lose
1
2
3
e- to
gain
ion
+
2+
3+
4
5
6
7
3
2
1
3-
2-
-
0
• An ionic solution or a melt will
conduct because its ions are free to
move to the electrode of opposite
sign.
• An ionic solid does not conduct
because its ions are unable to
move.
Liquid or Gas
• At room temperature
• A liquid or gaseous compound will
be covalent.
• A liquid or gas contains small
discrete molecules between which
there are fairly small forces of
attraction.
Solids
• At room temperature
• A solid compound can be ionic or
covalent
• Solids are a result of very strong
forces holding the particles
together.
• Ionic solids consist of a lattice of
oppositely charged ions.
Types of Solid
• In an ionic solid
these forces are
the ionic bonds
i.e. the forces of
attraction
between the
oppositely
charged ions.
• A covalent
network solid
consists of a huge
number of atoms
held together by a
network of
covalent bonds.
Soluble in water?
• Most ionic substances are soluble
in water, the lattice breaking, to
free the ions
• Most covalent substances are
insoluble in water but can dissolve
in other solvents.
• An electrolyte is a substance which
conducts when molten or in solution.
• While most ions are colourless, some
are coloured. e.g.
• cobalt - pink/purple;
• copper - blue;
• dichromate - orange;
• nickel - green;
• permanganate - purple
Electrolysis
• Electrolysis occurs when d.c.
(direct current) is passed through
a melt, or an ionic solution. This
changes the compound, releasing
new substances at the electrodes.
Products of Electrolysis
• At the positive electrode: chlorine,
bromine, iodine or oxygen (from
water) are released.
• At the negative electrode:
copper, silver or hydrogen (from
water) are released.
Electrolysis
-
+
Cu2+
• Electrolysis of
copper(II)
chloride
• The positive
copper ion moves
to the negative
electrode where:
Cu2+ + 2e-  Cu
Electrolysis
-
+
Cl-
• Electrolysis of
copper(II)
chloride
• The negative
chloride ion
moves to the
positive electrode
where:
2Cl-  Cl2 + 2e-
Properties of Substances
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Acids and
Alkalis
pH
•
•
•
•
pH is a continuous scale of acidity.
Acids have a pH of less than 7
Alkalis have a pH of more than 7.
Water, and other neutral solutions
have a pH of 7.
Oxides
• Non-metal oxides dissolve in
water, giving acidic solutions.
• Metal oxides and hydroxides,
which dissolve in water, give
alkaline solutions.
Acid Rain
• Acid Rain
• This has damaging effects on buildings
and other structures, soil and plant and
animal life.
• Sulphur dioxide gas dissolves in water
in the atmosphere, producing sulphuric
acid.
• Nitrogen oxides dissolves in water in
the atmosphere, producing nitric acid.
Ions
• Acids and alkalis both contain ions.
In water the concentration of ions
is very low.
• The test for hydrogen is that it
explodes with a "pop" when lit.
H+ and OH- ions
• Acids contain more H+ ions than
water.
• Alkalis contain more OH- ions than
water.
• Water, and other neutral solutions,
contain equal numbers of H+ and
OH- ions.
Dilution
• When an acid is diluted its acidity
decreases and its pH increases.
• When an alkali is diluted is
alkalinity decreases and its pH
decreases.
• When an acid (or alkali) is diluted
then the number of H+ (or OH- )
ions per cm3 of solution decrease
and so the acidity (or alkalinity)
decrease.
Acids and Alkalis
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Reactions of
Acids
Neutralisation
• Neutralisation is the reaction of an acid
with a neutraliser.
• Neutralisers are metal oxides,
hydroxides and carbonates.
• Examples of neutralisation involve
adding lime to soil or water to reduce its
acidity
treating acid indigestion with
magnesium hydroxide
the reaction of H+ (aq) to form water.
Neutralisation
• During a neutralisation reaction then
the pH of any acid or alkali involved
move nearer to 7.
• Neutralisation involves the reaction:
H+ + OH-  H2O
A salt is the substance formed when the
hydrogen ion of an acid is replaced by a
metal (or ammonium) ion.
Salts
Acid
Formula
Salt
Ion
hydrochloric HCl
chloride
Cl-
sulphuric
H2SO4
sulphate
SO42-
nitric
HNO3
nitrate
NO3-
carbonic
H2CO3
carbonate
CO32-
Acids and carbonates
• An acid reacts with a metal carbonate to
release carbon dioxide. Thus acid rain
will dissolve rocks or buildings which
contain carbonates.
• The hydrogen ions from the acid react
with the carbonate ions, to form carbon
dioxide and water.
2H+ + CO32-  H2O + CO2
Acids and metals
• Acids react with some metals to
release hydrogen. The hydrogen
ions in the the acid form hydrogen
molecules.
• Acid rain will dissolve iron
structures very slowly, since iron
reacts with acid to produce
hydrogen.
Salts
• When dilute hydrochloric acid
reacts with acid then hydrogen and
a metal chloride are formed.
• When dilute sulphuric acid reacts
with acid then hydrogen and a
metal sulphate are formed.
Acids, Bases and Alkalis
• A base is a substance which
neutralises an acid.
• ACID + BASE 
SALT + WATER
• An alkali is a soluble base.
• ACID + ALKALI  SALT + WATER
Precipitation
• An easy way to prepare salts is to
react an acid with an insoluble
metal oxide or metal carbonate.
Excess can be removed from the
reaction mixture by filtration.
• Precipitation is the reaction in
which two solutions react to form
an insoluble salt.
Remember Moles?
• To connect gram
formula mass, mass
in grams and number
of moles use the
triangle opposite
• gfm = mass of 1
mole
• n = number of moles
• m = mass of
substance
m
n
gfm
Remember solutions?
• To connect volume,
concentration and
molarity of a solution
use the triangle
opposite.
• c = concentration
(m/l)
• n = number of moles
• v = volume (l)
n
c
v
Working out about
neutralisations
• Work out unknown
concentrations and
volumes from the
results of volumetric
titrations.
• You use the equation
VH MH NH =
VOH MOH NOH
V = volume
M = molarity
NH = number of
H+ ions in acid
NOH =number of
OH- ions in alkali
H = acid
OH = alkali
Reactions of Acids
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Metals
Metals
• Metals have a metallic lustre i.e.
they are shiny
• Metals conduct electricity when
solid or liquid.
• The world's metal resources are
finite and so we must recycle used
metals.
Alloys
• An alloy is a mixture of metals, or a
mixture of metal with non-metal.
Examples:
• Brass
• solder
• "stainless steel"
Properties of Metals
•
•
•
•
•
A metal's properties decide its uses
electrical conductivity – electric wiring
thermal conductivity – pots and pans
malleability - shaped into many objects
strength - made into certain objects
Reactions of Metals
Common reactions of metals.
• Metals react with oxygen to form
metal oxides.
• Metals react with water (either as
liquid or steam) to form the metal
hydroxide and hydrogen.
• Metals react with dilute acid to
release hydrogen.
Reactions of Metals
• N.B. Not all metals react as shown
on the previous slide.
• The ease with which these
reactions take place is a measure
of the reactivity of the metal.
• We can build up a Reactivity Series
from the relative reactivity of the
metals.
Oxidation
• Oxidation is the loss of electrons
by a reactant in a chemical
reaction.
• When a metal reacts to form a
compound it is oxidised.
Reduction
• Reduction is the gain of electrons
by a reactant in a chemical
reaction.
• When a metal compound reacts to
form a metal it is reduced.
Oxidation and Reduction
• OIL RIG
Oxidation Is Loss of
electrons Reduction Is Gain
of electrons
• In a redox reaction oxidation and
reduction go on together.
The Reactivity Series
•
•
•
•
•
•
•
•
•
•
•
•
•
•
K
Na
Ca
Mg
Al
Zn
Fe
Sn
Pb
H
Cu
Hg
Ag
Au
Metals are
listed from
most reactive
to least
reactive.
Reactivity
• The Reactivity Series (also called
the Electrochemical Series) lists
the metals in order of their ease of
oxidation.
• The least active metals are those
whose ions are most easily
reduced.
Recovering Metals
• The less active metals do not react
well and so occur uncombined in
the earth's crust.
• Thus they were some of the first
elements discovered.
• Ores are naturally occurring
compounds of a metal.
• The more reactive metals are
found combined in the earth's
crust, as ores.
• The extraction of a metal from its
ore is an example of reduction.
• Very unreactive metals , such as
gold, silver and mercury, can be
obtained from their oxides by heat
alone.
Recovering Metals
• Other metals from the middle of
the Reactivity Series, such as zinc,
iron, copper and lead, can be
obtained from their oxides by
heating the oxide with hydrogen,
carbon (or carbon monoxide).
• Highly reactive metals, such as
magnesium, calcium, sodium and
potassium, have to be obtained
from their oxides by other
electrolysis.
Recovering Metals
• The more reactive a metal is, the
more difficult it is to break down
its compounds.
• Oxides of reactive metals are most
difficult to break down.
• Oxides of unreactive metals are
most easily broken down.
The Blast Furnace
• Iron is produced from iron ore in the
blast furnace.
• There are two reactions
• The formation of carbon monoxide from
coke (carbon):
C(s) + O2 (g) CO2 (g)
C(s) + CO2 (g)  2CO(g)
• The reduction of iron oxide to iron:
Fe2O3(s) + 3CO(g)  2Fe(s) + 3CO2
Metals
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Metals and
Electricity
Cells
• A cell is made by connecting two
different metals together with an
electrolyte.
• An electrolyte is a material, which
conducts electricity in solution (it
contains ions). The electrolyte is needed
to complete the circuit.
• The voltage generated between
different pairs of metals varies, and this
gives rise to the Electrochemical Series.
Displacement
• Any metal, in an Electrochemical
Series, will displace a metal below
it from one of its compounds.
• This reaction will usually produce
some visible signs.
Displacement
• If zinc reacts with copper sulphate
solution the reactions are:
Cu2+ + 2e  Cu
Reduction
Zn  Zn2+ + 2e
Oxidation
Overall
Cu2+ + Zn  Cu + Zn2+ Redox
• By considering the metals with which
acids will react it is possible to place
hydrogen in the Electrochemical Series.
Chemical Energy in Cells
• Chemical changes can bring about
the production of electrical energy.
• A cell or battery will run out (“go
flat”) when the chemicals which
produce electricity are used up.
Mains or Battery?
• Mains Electricity
• Cannot be
transported
• Uses high voltages,
which can be
dangerous.
• Cheap to use
• Made from
renewable energy
sources
• Battery
• Is easily transported
• Uses low voltages, so
is not dangerous.
• More expensive to
use
• Made from finite
energy sources
Metals and Electricity
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Corrosion
Corrosion
• Corrosion is a chemical reaction in
which the surface of a metal changes
from an element to a compound.
• Different metals will corrode at
different rates.
• Corrosion results in metals forming
compounds and so is an example of
oxidation.
Rusting
• Rusting is the word used to
describe the corrosion of iron.
• Rusting requires the presence of
oxygen (from air) and water.
• Ferroxyl indicator, which turns
blue in the presence of Fe2+ ions
can be used to show the extent of
rusting.
Rusting
• The water must contain dissolved
carbon dioxide or some other
electrolyte.
• Salt acts as an electrolyte and so salt,
spread on the roads in winter,
increases the corrosion of car
bodywork.
• Any other electrolyte would increase
corrosion
Rusting
• When iron rusts the iron atom loses
electrons to form iron(II) ions:
Fe  Fe2+ + 2e
• This is followed by a further loss of
electrons to form the iron(III) ion:
Fe2+ Fe3+ + e
• The electrons lost by the iron are taken
by the water and oxygen and used to
form hydroxide ions:
2H20 + O2 + 4e  40H-
Rusting and electrons
• Iron does not rust when connected to
the negative terminal of a battery
because the electrons flowing onto the
iron prevent it from losing electrons
• By using a cell with an iron nail, a
carbon rod, an electrolyte, ferroxyl
indicator and a centre-zero meter it is
possible to show the formation of Fe2+
at the iron nail and the movement of
electrons away from the iron.
Rusting and electrons
• The iron atoms rust,
losing electrons.
Rusting and electrons
• The iron atoms rust,
losing electrons.
• The blue colour
shows Fe2+ has been
formed.
Rusting and electrons
• The iron atoms rust,
losing electrons.
• The blue colour
shows Fe2+ has been
formed
• The centre-zero
meter shows the
movement of
electrons from the
iron nail to the
carbon rod.
Electrons flow to iron
Magnesium
Iron
• When a cell is set up
with iron and a metal
(say Mg) higher in
the Electrochemical
Series then electrons
flow to the iron.
• The reactions taking
place are:
Mg
 Mg2+ + 2e
Fe2+ + 2e  Fe
Electrons flow from iron
Copper
Iron
• When a cell is set up
with iron and a metal
(say Cu) lower in the
Electrochemical
Series then electrons
flow from the iron.
• The reactions taking
place are:
Fe
 Fe2+ + 2e
Cu2+ + 2e  Cu
Electroplating
• Electroplating
+
-
Electroplating
• Electroplating
• The metal to be
plated on is made
the positive
electrode.
Metal to
be plated on
+
-
Electroplating
• Electroplating
• The metal to be
plated on is made
the positive
electrode.
• The object to be
coated is made the
negative electrode.
+
-
Metal to
be plated on
Metal
object
to be plated
Electroplating
• Electroplating
• The metal to be
plated on is made
the positive
electrode.
• The object to be
coated is made the
negative electrode.
• The solution contains
the ions of the metal
to be plated on.
+
-
Metal to
be plated on
Solution of plating ions
Metal
object
to be plated
Metal Plating
• Galvanising occurs when steel (or
iron) is coated with zinc.
• Tin-plating occurs when steel (or
iron) is coated with tin.
Physical Protection
• Putting a barrier
over the surface
of a metal will
provide physical
protection against
corrosion
• It will not allow
air and water to
come in contact
with the metal.
•
•
•
•
•
•
Painting
Greasing
Electroplating
Galvanising
Tin-plating
Coating with
plastic.
Sacraficial Protection
• Sacrificial protection
• If two metals are
connected electrons
will flow from the
more active metal to
the less active.
• The more active
metal will corrode in
preference to the
less active metal.
• On the FinartGrangemouth oil
pipe bags of scrap
magnesium are
connected every
200 meters so
magnesium
corrodes
sacraficially to
protect the iron.
Tin-plating
• Tin-plating
• If it is scratched then the iron and
tin are exposed. Since the iron is
higher in the Electrochemical
Series it will corrode in preference
to the tin.
• The corrosion of the iron increases.
Galvanising
• Galvanising
• If it is scratched then the iron and
zinc are exposed. Since the zinc is
higher in the Electrochemical
Series it will corrode in preference
to the iron.
• The corrosion of the iron is
prevented.
Corrosion
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Fertilisers
Fertilisers
• Increasing world population
means that we need more efficient
means of food production.
• Growing plants take nutrients from
the soil. These nutrients include
compounds of nitrogen,
phosphorous and potassium.
Fertilisers
• Fertilisers are substances which
are added to the soil to replace the
essential elements needed for
plant growth.
• Different plants require fertilisers
containing different proportions of
these nutrient elements.
Fertilisers
• Artificial fertilisers are soaked out of the
soil by rain.
• They are carried into lakes and rivers
where they increase the number of river
plants.
• When these plants die then there is an
increase in the bacteria which digest
them leading to a decrease in oxygen in
the water.
• This results in the death of fish.
Fixed nitrogen
• Certain plants have nitrifying bacteria
present in nodules in their roots. These
bacteria can convert atmospheric
nitrogen (called free nitrogen) into
nitrogen compounds (fixed nitrogen).
These nitrogen compounds increase the
fertility of the soil.
• Bacterial methods of increasing the
nitrogen content of soil are cheaper
than chemical methods.
Fixed nitrogen
• Recycling of nitrogen compounds into
the soil is brought about by the
decomposition of plant and animal
protein by bacteria in the soil.
• Ammonium salts, potassium salts,
nitrates and phosphates make good
fertilisers because:
they contain some of the essential
elements for plant growth (P, N and K).
they are soluble and pass easily into
the soil and up the plant's roots.
The Nitrogen Cycle
• All living things need nitrogen to make
proteins.
• They cannot use free nitrogen from the
atmosphere.
• They need to get fixed nitrogen in their
food.
• The Nitrogen Cycle describes the place
of nitrogen compounds in Nature.
The Nitrogen Cycle
Animals need nitrogen to make
substances called proteins
ANIMALS
The Nitrogen Cycle
They get this nitrogen by eating
protein which has been made by plants.
ANIMALS
PLANTS
The Nitrogen Cycle
ANIMALS
plant protein
eaten
PLANTS
The Nitrogen Cycle
Plants absorb nitrates
through their roots
nitrates
ANIMALS
plant protein
eaten
PLANTS
The Nitrogen Cycle
nitrates
taken in
by roots
ANIMALS
plant protein
eaten
PLANTS
The Nitrogen Cycle
ammonia
Ammonia, NH3 comes
from animal waste
nitrates
taken in
by roots
ANIMALS
plant protein
eaten
PLANTS
The Nitrogen Cycle
ammonia
sewage
and
manure
nitrates
taken in
by roots
ANIMALS
plant protein
eaten
PLANTS
The Nitrogen Cycle
ammonia
bacteria
sewage
and
manure
Bacteria convert
ammonia into
nitrates
ANIMALS
plant protein
eaten
nitrates
taken in
by roots
PLANTS
The Nitrogen Cycle
atmospheric
nitrogen
ammonia
sewage
and
manure
bacteria
Thunderstorms
make nitrates
from N2
nitrates
taken in
by roots
ANIMALS
plant protein
eaten
PLANTS
The Nitrogen Cycle
atmospheric
nitrogen
ammonia
sewage
and
manure
bacteria
thunder
storms
nitrates
taken in
by roots
ANIMALS
plant protein
eaten
PLANTS
The Nitrogen Cycle
atmospheric
nitrogen
ammonia
sewage
and
manure
ANIMALS
bacteria
Death of living
things also
produces
ammonia
plant protein
eaten
thunder
storms
nitrates
taken in
by roots
PLANTS
The Nitrogen Cycle
atmospheric
nitrogen
ammonia
sewage
and
manure
bacteria
decaying animals
after death
ANIMALS
thunder
storms
compost
plant protein
eaten
nitrates
taken in
by roots
PLANTS
The Nitrogen Cycle
ammonia
atmospheric
nitrogen
thunder
storms
bacteria
Some
sewage
plants
and
can
manure
change
nitrates
compost
N2 into
decaying animals
nitrates
taken in
after death
by roots
ANIMALS
plant protein
eaten
PLANTS
The Nitrogen Cycle
atmospheric
nitrogen
ammonia
sewage
and
manure
bacteria
decaying animals
after death
ANIMALS
thunder
storms
bacteria in
legumes
compost
plant protein
eaten
nitrates
taken in
by roots
PLANTS
The Nitrogen Cycle
• Due to the way in which we live, food,
containing nitrogen compounds, is sent
from the country to the city.
• This means that the nitrogen
compounds are not returned to the area
from which they came.
• To replace that nitrogen, artificial
fertilisers are needed.
• The Haber Process is used to make
ammonia from free nitrogen.
The Nitrogen Cycle
atmospheric
Haber Process nitrogen
ammonia
bacteria thunder
storms
sewage
bacteria on
and
legumes
manure
nitrates
compost
decaying animals
taken in
after death
by roots
ANIMALS
plant protein
eaten
PLANTS
The Haber Process
• Ammonia is produced by the Haber
Process.
• Nitrogen and hydrogen react over an
iron catalyst to produce ammonia:
N2 + 3H2  2NH3
Since the formation of ammonia is a
reversible reaction then not all of the
nitrogen and hydrogen are converted
into ammonia.
The Haber Process
• At low temperatures a large amount of
ammonia is produced slowly.
• At high temperatures a smaller amount
of ammonia is produced more quickly.
• This means that the Haber Process is
carried out at a moderately high
temperature to produce ammonia at the
most economical rate.
The Haber Process
Nitrogen (air)
Hydrogen (Oil industry)
Unreacted
Nitrogen and
hydrogen
Reaction
Chamber
With Fe
catalyst
Ammonia
Cooler
Ammonia
• Ammonia:
has no colour.
has a distinctive sharp smell.
is highly soluble in water, producing an
alkaline solution.
can be converted into an ammonium
compound by its reaction with acid:
2NH3 + H2SO4
 (NH4) 2SO4
ammonia
sulphuric acid
ammonium sulphate
Ammonia
• Ammonia can be prepared by the
reaction of ammonium compound
with alkali:
NH4Cl + NaOH  NaCl + H2O + NH3
This is also used as a test for the
ammonium ion NH4+
Nitric Acid
• Nitrogen is a very unreactive gas.
• Nitric acid (HNO3) is formed when
nitrogen dioxide (NO2), in the presence
of air, dissolves in water.
• The presence of nitrogen oxides in the
air means that they will dissolve in rain
to produce a mildly acidic solution. This
has the result that:
– nitrogen compounds are added to the soil.
– the acidity of the soil will be increased.
Nitric Acid
• Nitrogen dioxide is produced by the
passage of a high voltage spark through
air since a large amount of energy is
required to break the bonds between
the nitrogen atoms in the molecules.
• These conditions occur when
– lightning passes through air
– a spark passes in the spark plug of a car
engine
This does not provide an economic way of
making nitrogen dioxide.
Nitric Acid
• Ammonia normally reacts with oxygen
to give nitrogen and water.
• It will only produce nitrogen oxides if a
platinum catalyst is used.
• This is a step in the Ostwald Process, by
which nitric acid is produced. This can
also be carried out in the laboratory.
The Ostwald Process
Ammonia
Reaction Chamber
(Pt catalyst)
Oxygen
(or) air
Oxides of
nitrogen
Water
Nitric acid
The Ostwald Process
• This process is carried out at a
moderately high temperature to allow it
to proceed fairly quickly.
• Since the reaction is exothermic it is not
necessary to continue heating it after
the reaction has started since it will
supply sufficient energy to continue at a
reasonable rate.
Fertilisers
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