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Preliminary Chemistry Topic 2
METALS
What is this topic about?
To keep it as simple as possible, (K.I.S.S.) this topic involves the study of:
1. OUR USE of METALS
2. CHEMICAL ACTIVITY of the METALS
3. PATTERNS of the PERIODIC TABLE
4. QUANTITY CALCULATIONS... the MOLE
5. METALS from their ORES
...all in the context of how Chemistry contributes to cultural development
but first, an introduction...
Chemistry of the Metals
Technology Needs Metals
In the previous topic you learnt about the
Elements of the Periodic Table. In this topic you
will concentrate on the chemistry of the metals,
and some of the chemical patterns that they
show.
... and Speaking of Patterns,
in this topic you will find that
The great sweep of human cultural development
has many aspects... Language, Religion, Art &
Music, and, of course, Technology.
The history of technology is closely linked with
our use of metals; in fact historians have named
some parts of history after the metals that
changed the way people lived.
The Periodic Table
is full of patterns
Dagger from the “Bronze Age”
M
This topic starts with a quick look at the history
of metal use, and ends with a study of how we
get metals from the Earth, and the chemistry of
the extraction process.
s
l
a
et
No
nMe
tal
s
Measuring Chemical Quantities
In this topic you will also be introduced to the
concept of the “Mole”...
not a burrowing mammal!
not a traitor within the group!
not a gangster’s girlfriend!
certainly not a skin blemish!
A Chemical Mole is a clever way to measure
quantities; essential for analysis & chemical
manufacture.
Electrically powered smelter plant
for extracting
Aluminium from its ore
If you know the mass,
you can figure out
how many atoms there are...
thanks to the mole.
Photo courtesy of Comalco Aluminium Ltd
Preliminary Chemistry Topic 2 “Metals”
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CONCEPT DIAGRAM (“Mind Map”) OF TOPIC
Some students find that memorising the OUTLINE of a topic helps them learn and remember
the concepts and important facts. As you proceed through the topic, come back to
this page regularly to see how each bit fits the whole.
At the end of the notes you will find a blank version of this “Mind Map” to practise on.
History of
Metal Use
The Activity Series
of the Metals
Metals
We Use
Today
Our Use of
Metals
Electron Transfer
REDOX
1st Ionisation
Energy
Chemical Activity
of the Metals
Activity & Usage
of Metals
Patterns of the
Periodic Table
METALS
History of the
Periodic Table
Definition of the
Mole.
Avogadro’s
Number
Quantity
Calculations
Extracting Metals
from Ores
the Mole
Molar Ratios in
Reactions
Empirical
Formulas
Case Study:
Extracting
Copper
from its Ore
Minerals
Ores
&
Resources
Mole Quantity
Calculations
Gay-Lussac’s Law
&
Avogadro’s Hypothesis
The Case for
Recycling Metals
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1. OUR USE OF METALS
The First Uses of Metals
The Iron Age (approx. 2,500 to 1,500 years ago)
For most of human existence, people used tools of
stone, wood and bone. Primitive tribes were familiar
with gold which occurs uncombined in nature, but it is
too soft to be useful for anything but jewellery and
decoration.
About 1,000 B.C. the extraction of iron from its ores
was discovered. This requires much higher
temperatures, and the breakthrough was probably the
invention of the bellows, a device to pump air into a
furnace so the wood or charcoal burns hotter.
About 5,000 years ago, in the Middle East, some
people accidentally discovered that if certain rocks
were roasted by fire, small amounts of copper
would be found later in the ashes. Copper is too
soft to be really useful, but there was a brief
“Copper Age” around the eastern end of the
Mediterranean Sea. Copper was used for
decoration, jewellery, small utensils, and
occasionally for knives and spear points.
Iron is stronger and harder than bronze. A warrior
armed with iron weapons will usually beat a bronzearmed man. Iron tools and even the humble nail
allowed new developments in buildings, ships,
wagons... remember that towns, trade and commerce
give wealth and power. An iron plough allows more
land to be cultivated to grow more food, to feed a
bigger army... and so on.
It is no accident that the dominant world power of this
time was ancient Rome, because their technology
was based on iron.
The big breakthrough was the discovery by these
copper-using people that if they roasted copperbearing rocks (ores) with tin ores, the resulting
“alloy” (mixture) of copper and tin produced a
much harder metal, “bronze”, which could be cast
in moulds, and hammered to shape many useful
tools and weapons.
From the Medieval to the Modern
After the collapse of the Roman Empire the various
cultures that dominated the “Dark Ages” still had ironbased technologies.
The next great technological change was the
“Industrial Revolution” which began about 1750 in
England. This had many aspects, but the big change
in technology was the use of coal (instead of wood) for
fuel. As well as steam engines, coal allowed for large
scale smelting of iron and the invention of steel (an
alloy of iron with carbon).
The Bronze Age (approx 4,500 to 2,500 years ago)
It is no accident that the rise of the great
ancient civilizations occurred about this
time. The stone blocks of the pyramids and
temples of ancient Egypt were cut and shaped
with bronze chisels. Egyptians, and later
Greeks, dominated their world because their
soldiers were armed with bronze swords,
spears and arrowheads.
The engines, tools and machinery of the great
factories were based on steel. Transport was
revolutionised by steel locomotives running on steel
rails. Steel ships replaced wooden ones, and steel
weapons (machine guns,
tanks and artillery) achieved
new heights (depths?)
in warfare and mass
destruction.
With bronze tools they built better ships and wagons for
transport and trade, which brought wealth and power.
Sad as it might be, the
facts of human history
are that progress has
been marked by conflict,
war and conquest, and
metals have been a vital
part of that development.
In the 20th century, new metals
and alloys became available... aluminium, titanium,
chromium, and many more.
Metal has many
advantages over stone,
wood, or bone:
This was made possible by electricity, which is
needed in large amounts to extract some metals from
their ores, or to purify and process them once
extracted.
• metal is harder, stronger, and flexible, not brittle.
• metal can be cast, hammered or drawn into shapes not
possible in stone, such as saw blades, swords and armour.
Human Progress has always been linked
to our use of Metals.
• when tools become blunt, metal can be re-sharpened.
Progress in metal usage has always been
linked to the availability of energy
to extract the metals.
Basically, a warrior with a bronze sword always beats a bloke
with a stone axe... we call that progress!
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The Metals We Use Today
Solder is an alloy of 30-50% tin with lead.
In one sense, we are still in the “Iron Age”. Iron
is still the metal we use the most, but nearly
always it is mixed with other elements in a
variety of alloys, notably steel.
Its most notable property is a very low melting
point, around 150-200oC.
Its major use is in plumbing for sealing the joints
between pipes, and in electronics for connecting
small components on a “circuit board”.
Metals That Are Used
in Their Pure State
Although we use a wide range of alloys, there are
some important metals we use in their pure,
elemental state.
Aluminium is very lightweight, yet strong and
corrosion resistant
Steel is used for bridges, tools and machinery,
bolts, screws and nails, reinforcing inside
concrete structures, engines, vehicle bodies,
trains and their rails, ships, and “tin” cans.
Its lightweight strength is
perfect for aircraft
construction.
Why is steel so widely used?
Lightweight and a good
conductor, it is used for
electricity power lines.
• Iron ore occurs in huge deposits, so iron is
common and economical to produce.
• Steel (in its various forms) is hard and strong.
• It can be cast, milled, rolled, worked, bent, cut
and machined into any shape or size.
Malleable and corrosion resistant, it is ideal for
window frames and drink cans.
Copper is used for electrical wiring in buildings and
As always, our usage of the different steel alloys
is linked to their particular properties:
Steel
Alloy
Iron,
with...
Properties
Uses
Mild steel 0.2%
carbon
strong, but
malleable
car bodies,
pipes, roofing
Tool steel
very hard
drills, knives,
hammers
resists
corrosion,
hygenic
food utensils,
medical tools
1-1.5%
carbon
Stainless 20% nickel
Steel
& chromium
appliances, because of its great electrical conductivity
and its ductility for ease of wire-making.
Metal Extraction Needs Energy
Our use of different metals through history can be
linked to the availability of energy.
In topic 1, you learned about the process of
chemical decomposition; where a compound
breaks down into simpler substances.
Decomposition is generally an endothermic
process; energy is absorbed by the reactants
during the reaction. Generally, you must supply
energy to make the process happen.
Brass
Metal ores are mineral compounds. To obtain the
elemental metal involves decomposition, which is
endothermic and requires energy. Some
compounds require more energy than others for
decomposition.
is a common “non-ferrous” (no iron) alloy.
Brass is an alloy
of copper and
zinc (about 50%
each)
Copper and tin ores require little energy. A decent
wood fire can “smelt” the metal from its ore. This why
copper and bronze were used in ancient times.
Iron ore requires more energy for decomposition.
That’s why the “Iron Age” came later.
Brass is very hard, but easily machined for screw
threads, etc. It is more expensive than steel, but is
corrosion resistant, so it is ideal for taps and
fittings for water and gas pipes.
Preliminary Chemistry Topic 2 “Metals”
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Aluminium and other “modern” metals require
even more energy, and electricity works better
than heat, so these only became available in quite
recent times.
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Worksheet 1 Our Use of Metals
Fill in the blank spaces.
Student Name...........................................
Before metals, people used tools mainly
made from a)............................. or
................................. The first metal used
was probably b).................................,
because it occurs in the elemental state
in nature. However, it is too soft to be
used for tools, so was just used for
c)................................
Today, the metal we use most is still
t)...................., in the form of the alloy
u)................... Its widespread use is
because:
• it is common and v)................................
to produce.
• it is very w).................. and .....................
Steel comes in a variety of alloys,
including x).................. steel (car bodies,
pipes, roofing) and y)....................... steel
used for food utensils and medical
tools.
Metallurgy (the technology of metals)
began
with
the
extraction
of
d).............................. from ores that were
simply
e)............................................
.............................................
A big improvement was the mixing of
ores of f)....................... and ......................
This produced the alloy g).......................,
which made tools and weapons with
many advantages over stone:
Other alloys used widely include:
• brass, a mixture of z).................... and
........................
• aa)................................., with a very low
melting point, is an alloy of
ab)........................ and ..............................
and is used in ac).....................................
and ....................................
• metal is h).................... and ....................
and is not i)........................ like stone.
As well as many alloys, there are some
metals commonly used in their pure,
elemental form:
• Aluminium, which has the advantages
of being ad)......................... and resistant
to ae)...........................Uses include
af)..................................... and ..................
...........................
• ag)......................... is used for electrical
wiring
because
of
its
good
ah)............................... and because it is
ai)................................ so it is easy to
draw out into wires.
• metal can made into intricate shapes,
such as j)..........................., not possible
in stone.
Later, bronze was replaced by
k)...................... which is l)................
...................... and....................., but
requires more m)............................ for its
extraction.
During the “Industrial Revolution”, the
use of n)................. for energy led to the
production of o)............................ which
is iron with a small amount of
p)........................... in it. This allowed the
development of machinery, trains and
the modern industrial world.
Chemically, the extraction of metals
from ores involves aj)...............................
reactions, which are ak).............-thermic.
Some
metals,
such
as
al)............................. require very little
energy,
others
such
as
am)...................................... require much
more. In many cases an)...........................
works better than heat in the extraction
and purification processes. The
changes in ao)............................ usage
through history can be directly linked to
society’s changing sources and uses of
ap)......................................
In the 20th century new metals such as
q).............................. became available
because the r)..................................
needed to extract it from its
s)................... was available.
Practice Test Questions are at the end
of the next section
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2. CHEMICAL ACTIVITY OF THE METALS
Metals React With Oxygen
Metals React With Acids
One of the most familiar laboratory reactions is
the burning of magnesium:
The different activity levels of the metals is most
clearly seen when metals are reacted with dilute acids.
Magnesium + Oxygen
2 Mg
+ O2
Magnesium oxide
2 MgO
You may have done experimental work to
observe how vigorously different metals react
with a dilute acid.
In fact, many metals will burn, some a lot more
readily and violently than magnesium:
Sodium + Oxygen
4 Na
+ O2
Metals like calcium and
magnesium react
vigorously.
Sodium oxide
2 Na2O
Zinc and iron are slower.
In these cases there is a violent exothermic
reaction, with light and heat energy produced.
The product is often a powdery, crumbly solid.
Lead is very slow indeed.
Copper does not react at
all.
Other metals, such as aluminium and zinc, react
on the surface and the oxide compound formed
is airtight and prevents further reaction. That’s
why these metals are often dull-looking... the
surface coat of oxide is dull.
Aluminium + Oxygen
4 Al
+ 3 O2
When there is a reaction,
the gas produced is
hydrogen.
Aluminium oxide
2 Al2O3
Examples:
The point is, that different metals
have different chemical activities.
Zinc + Hydrochloric
acid
Zn + 2 HCl
Metals React With Water
Another favourite school reaction is when sodium
reacts with water. This is often done outdoors,
because it results in an exciting little explosion.
Magnesium + Nitric
acid
Mg + 2 HNO3
Iron
What happens is:
2 Na
Water
+ 2 H2O
Hydrogen + Sodium
(gas)
hydroxide
H2
+ 2 NaOH
+
Sulfuric
acid
Fe + H2SO4
Hydrogen + Zinc
chloride
H2
+ ZnCl2
Hydrogen + Magnesium
nitrate
H2
+ Mg(NO3)2
Hydrogen
H2
+ Iron(II)
sulfate
+ FeSO4
The ionic compounds formed are collectively
known as “salts”, so the general pattern of the
reactions is
(In fact this is NOT the explosion reaction. The
explosion is the reaction of the hydrogen with
oxygen, to form water)
Metal + Acid
Once again, some metals react easily and
rapidly and form the metal hydroxide, while
others react slowly if heated in steam, and form
oxides.
Zinc + Water
Zn + H2O
A flame test
gives a “pop”
explosion
The metal is “eaten away”
and dissolves into the liquid.
This is because it forms a
soluble ionic compound. Exactly what the
compound is, depends on which acid is used.
Other metals, such as copper, react with oxygen
very slowly and only if heated strongly. Some,
like gold, will not react at all.
Sodium +
Bubbles of
gas are
produced.
Hydrogen + a Salt
It will help you greatly to learn
the common laboratory acids
Common Name
Chem Name
Hydrochloric
= Hydrogen chloride
Sulfuric
= Hydrogen sulfate
Nitric
= Hydrogen nitrate
Hydrogen + Zinc oxide
H2
+ ZnO
Formula
HCl
H2SO4
HNO3
Metals like copper and gold do not react at all.
WORKSHEET at end of section
There is an “Activity Series” of metals.
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Electron Transfer
in Metal Reactions
The Activity Series of the Metals
From these 3 patterns of reaction, it seems there
is a further, underlying pattern. Certain metals,
like sodium, always seem to react readily and
vigorously. Others, like copper, always react
slowly or not at all.
The chemical reactions that allow us to see the
pattern of the Activity Series are just part of an
even greater pattern in Chemistry... the process
of electron transfer.
From this, and other reaction studies, the
common laboratory metals can be arranged in
an “Activity Series”:
Most
Active
To understand this, look again at the reaction
between a metal and an acid:
Zinc + Hydrochloric
acid
K
Zn
+
Hydrogen + Zinc
(gas)
chloride
2 HCl
H2
+
ZnCl2
Na
HCl and ZnCl2 are both ionic compounds. Here
is the equation re-written to show the individual
ion “species”.
Li
Ba
Zn + 2H+ + 2Cl-
Ca
Study this carefully and make sure you
understand why there have to be 2 of some ions
to agree with the original balanced equation.
Mg
Al
Notice that the chloride ions (Cl-) occur on both
sides of the equation unchanged. Nothing has
happened to them at all. We say they are
“spectator ions”. Like by-standers at a car crash
they are not involved, while other atoms and
ions undergo serious changes.
Zn
Fe
Sn
Since they aren’t actually involved, we can leave
the spectators out. This is called a “net
equation”.
Pb
Cu
Zn + 2H+
Ag
Least
Active
Au
2 6
H2 + Zn2+
Now we can see what really happened;
• a zinc atom became a zinc ion
and
• 2 hydrogen ions became a (covalent)
hydrogen molecule.
If you look for these metals on the Periodic Table
you will notice a further pattern.
3
H2 + Zn2+ + 2Cl-
Positions of the first 6 metals
of the Activity
Series.
To do this, the zinc atom has to lose 2 electrons,
and the hydrogen ions must gain a pair of
electrons to share.
1 5
4
The highly active metals all lie to the extreme left of
the table, AND the higher their activity, the lower down
the table they are within each column.
Zn2+ + 2e-
2H+ + 2e-
H2
Now it should be clear what really happened: the
zinc atom gave a pair of electrons to some
hydrogen ions. Electrons were transferred from
one “species” to another.
This is one of many patterns that allows you to use
the Periodic Table instead of learning many small
facts. For example, instead of memorising the Activity
Series fully, you can remember the pattern above and
always be able to figure out the order of the most
active metals.
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Zn
The equations above are “Half-Equations” and
are often used to describe what is really
happening in a reaction.
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Oxidation and Reduction
First Ionisation Energy
The transfer of electrons from one species to
another is one of the most fundamental and
important general reactions of Chemistry.
Although you’re not yet required to know
about Oxidation and Reduction, this bit you
have to learn.
The reaction between zinc and acid can be
visualised like this:
Definition
The Ionisation Energy of an element is the
energy required to remove an electron from
an atom.
electrons transferred
+
For technical reasons, the measurement of this
energy is carried out for atoms in the gas state.
2 Hydrogen ions
Zn(g)
Hydrogen molecule
The energy required for this to happen is the
“1st Ionisation Energy”
+2
Zinc ion
We know that zinc atoms normally lose 2
electrons to form the Zn2+ ion. However, the
formal definition for this process involves
just the loss of 1 electron.
Covalent bond
(2 electrons being shared)
The zinc atom has lost 2 electrons,
Zn
Every element has its own characteristic
value, even those elements which would not
normally lose electrons, such as non-metals
like chlorine.
Zn2+ + 2e-
Cl(g)
For historical reasons,
the loss of electrons is called “Oxidation”
Even the inert gases, which normally do not
form ions at all, can be forced to lose an
electron if energy is added. They too have a
1st Ionisation Energy value.
H2
The gain of electrons is called “Reduction”
Ionisation Energy
Determines the Activity Series
Neither process can occur alone... they must
occur together
In order for a metal to begin reacting with an
acid, (or with water or oxygen) it must lose an
electron. This will require the input of its 1st
Ionisation Energy.
The zinc oxidation allows the hydrogen to be
reduced, and the hydrogen reduction allows the
zinc to be oxidised.
If the value for 1st Ionisation energy is very
low, the metal will gain this energy easily and
quickly from its surroundings. It will readily
enter the reaction, and the reaction will
proceed vigorously.
The total reaction is an “Oxidation-Reduction”
and is commonly abbreviated to “REDOX”.
Note that the syllabus does NOT require you to
know these definitions yet, but it is worth
knowing about Redox for future topics. You ARE
required to know about electron transfer and its
involvement in metal reactions.
If its value for 1st Ionisation energy is higher,
the atom cannot react so readily or
vigorously... its activity is lower.
The ACTIVITY SERIES of the Metals
is determined by
1st IONIsATION ENERGY
WORKSHEET at the end of section
Preliminary Chemistry Topic 2 “Metals”
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Cl+(g) + e-
Normally a chlorine atom forms a negative ion
by gaining an electron.
Technically though, it is possible for it to lose
an electron if energy is added.
This energy is the “1st Ionisation Energy”
and the hydrogen ions have gained electrons.
2H+ + 2e-
Zn+(g) + e-
Increasing values for 1st Ionisation Energy
Zinc atom
+
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K
Na
Li
Ba
Ca
Mg
Al
Zn
Fe
Sn
Pb
Cu
Ag
Au
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Another example is the choice of metals for
water pipes.
Choice of Metals
Based on Activity
Sometimes which metal is chosen for a
particular application is based on its position in
the Activity Series.
Steel is cheap, but
since iron is about the
middle of the Activity
Series it will corrode
(rust) by contact with
water. Is it better to
choose a lower activity
metal such as copper,
which will not corrode
as quickly, but is more
expensive?
Example
In critical electronic connections, such as
computer network plugs, it is essential that the
electric signals get through without loss or
distortion.
Normally we use copper for electrical wiring, but
in a critical connection plug it is worth the extra
expense of using gold.
The decision is usually
to use cheap steel
pipes for longer,
outdoor uses like your
garden taps.
Copper is a low activity metal, but can slowly
react with oxygen to form a non-conducting
oxide layer in the connection. Gold is lower
down the activity series and will not react at all,
so the plug connection cannot corrode.
Brass fittings
Copper pipe
Indoors, where
distances are shorter,
copper is chosen,
especially for hot water
supply. Indoors a
rusted-out leaking steel pipe would be a
disaster, so it’s worth paying more for copper.
Gold’s extremely low chemical activity (due to
a relatively high 1st
Ionisation Energy)
is part of the
reason it has
always been used
for jewellery.
Interestingly, sometimes the higher activity
metals corrode less. Aluminium and zinc are
higher up the Activity Series than iron. They
react rapidly when exposed to oxygen, but the
surface layer of oxide is airtight and waterproof,
and prevents oxygen or water getting to the
metal underneath. Therefore, these metals can
be used in situations where corrosion needs to
be prevented.
Gold’s low activity
means it will not
tarnish or corrode,
so it retains its
beautiful colour and
lustre.
“Galvanised” steel is coated with a thin layer of
zinc to prevent (or slow down) corrosion of steel
roofing, fence wires, nails, bolts, etc.
Bronze & Gold
have been used throughout history
in Art and Religion
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Worksheet 2 Chemical Activity of the Metals
Fill in the blank spaces.
Student Name...........................................
All these reactions involve the transfer of
n)......................... In the case of the Metal + Acid
reaction,
the
metal
atoms
always
o)......................... electron(s) while a pair of
p)............................ ions gain 2 electrons (which
they share in a q)......................... bond) and form
a r)...................... molecule with formula s)...........
When a metal reacts with oxygen it forms an
a)......................... compound.
METAL + OXYGEN
b) .............................
Some metals will also react with water, forming
c).....................................
gas
and
a
d)...................................... compound.
METAL + WATER
“Oxidation” is the technical term for
t)..................... ................................. The opposite
is “u)...................................”
In the Metal + Acid reaction, the metal is always
v)..............................................
while
w)..............................
ions
are
always
x)..................................................
c).................. + d).................
Most metals will react with acids, forming
e).......................... gas and an ionic compound
called a “f)...........................”
METAL + ACID
e)....................... + f).................
In all these reactions the various metals react at
g)............................... rates, showing an order of
chemical h)......................... From these reactions
and others, the “Activity Series” has been
determined.
The “1st y)........................... Energy” of an
element is defined as the energy required to
z)......................................... ................... from
atoms in the aa)................. state. The very active
metals are like that because they have very
ab).................... (high/low) values for this. Metals
further down the series do not react as
vigorously
because
their
values
are
ac)...........................................
Metals such as i).............................. and
............................. are the most active. These are
the elements located in the j)...........................
columns of the Periodic Table.
Some metals such as k)....................... and
......................... have very low activity, and often
do not react at all. Other common metals like
l).................................. and ....................................
are in the middle of the series. They will react,
but generally do so m).......................................
Sometimes the choice of which metal to use is
determined by the activity level. An example is
ad)............................... .... .......................................
.............................................
Worksheet 3 Practice Problems
Student Name...........................................
a) Lead
(assume lead(IV) ion forms)
4. All the following equations are Metal + Acid
reactions.
Fill in all blank spaces, then re-write in symbols and
balance.
b) Iron
(Assume iron(III) ion)
a) Zinc + Sulfuric acid
1. Write a balanced, symbol equation for the reaction
of each of the following metals with oxygen.
................. +.....................
c) Lithium
b) Calcium + Hydrochloric
acid
2. a) Arrange the metals in Q1 in order of decreasing
chemical activity.
b) Which one(s), if any, might ignite easily and burn
in air with a visible flame?
3. Write a word equation AND a balanced, symbol
equation to describe the reaction of:
a) calcium metal with water (reacts spontaneously at
room temperature)
................. +...................
c).................. +......................
Hydrogen + Barium
nitrate
d).................... + ....................
Hydrogen + iron(II)
chloride
5. For each of the reactions in Q4, which chemical
species
a) lost electrons?
b) gained electrons?
b) Tin metal with water (heated in steam) (Assume
tin(II))
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Worksheet 4 Test Qestions
Multiple Choice
sections 1 & 2
8. (3 marks)
Give a reason why
a) metal tools are superior to stone tools.
1.
Which list shows metals used by humans in the
correct chronological order of their history of usage?
A. bronze, aluminium, iron
B. copper, bronze, iron
C. gold, iron, bronze
D. copper, steel, bronze
b) iron replaced bronze in the history of metallurgy.
c) aluminium did not come into common use until the
20th century.
2.
Which list correctly identifies an alloy, and the
elements it contains?
A. Steel; iron and tin
B. Bronze; tin and zinc
C. Solder; copper and lead
D. Brass; copper and zinc
9. (6 marks)
The most common metal in use today is steel, which
comes in a variety of forms, with different properties
and uses. Compare 3 different types of steel, stating
the composition of each and relating its properties to
a common use.
3.
The metals used by humans have changed over the
course of history. The availability of new metals has
often been dependent on the:
A. availablity of energy to extract metals from ores.
B. discovery of new minerals as people explored
the world.
C. invention of new alloys.
D. development of new technologies to use
the metals.
10. (5 marks)
Give an outline of an experiment you have done to
investigate the relative chemical activity of some
metals. Include the observation(s) you made to
assess metal activity, and state the conclusion(s)
reached.
4.
A metal which reacts readily and vigorously with
oxygen, water and dilute acids would probably:
A. have a high value for 1st ionisation energy.
B. be from the “Transition” block of the
Periodic Table.
C. have a very low 1st ionisation energy.
D. be located at extreme right of the Periodic Table.
11. (6 marks)
Write a balanced symbol equation for the reaction of:
a) magnesium with hydrochloric acid.
5.
If nickel reacted with sulfuric acid, the products of the
reaction would be:
A. hydrogen gas and nickel sulfate
B. carbon dioxide gas and nickel sulfate.
C. nickel sulfide and hydrogen gas.
D. sulfur dioxide gas and nickel hydroxide.
b) calcium with water (reacts at room temperature).
c) potassium with oxygen.
12. (4 marks)
When barium metal reacts with an acid there is an
exchange of electrons such that hydrogen gas and
barium ions are formed. Write 2 “half-equations” to
show clearly the species gaining, and the species
losing, electrons.
6.
During the reaction in Q5, the basic underlying
change occurring is:
A. the breaking covalent bonds.
B. the transfer of electron(s) from one species
to another.
C. chemical changes in “spectator ions”.
D. physical dissolving of metal in the acid.
13. (4 marks)
a) Write an equation (including states) for the first
ionisation of
i) magnesium
Longer Response Questions
Mark values shown are suggestions only, and are to
give you an idea of how detailed an answer is
appropriate. Answer on reverse if insufficient space.
ii) oxygen
7. (5 marks)
Give an example of
a) a metal used in its elemental state, and
b) Describe how the Activity Series of Metals is
related to the values of 1st Ionisation Energy.
b) a non-ferrous alloy (naming its components)
in common use. For each, relate the properties of the
metal to its particular use(s).
Preliminary Chemistry Topic 2 “Metals”
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3. PATTERNS OF THE PERIODIC TABLE
Atomic Structure, Atomic Number and Mass
Here is a quick reminder of some basics about atoms you need to know:
In the Nucleus
are
Protons &
Neutrons
The Periodic Table
is firstly a list of the elements, arranged in order,
and showing all the basic details.
Atomic Number
In orbit around
the nucleus are
the Electrons
18
Ar
Electrons = Protons = “Atomic Number”
Argon
Each element’s atoms have a different,
characteristic,
number of protons (and
electrons). Therefore, each element has a
different Atomic Number.
39.95
Equal to the number of protons
in each atom. (Also equals the
number of electrons in the
neutral atom.)
Chemical Symbol
Element Name
“Atomic Weight”
NOT the “Mass Number”
However, the Periodic Table is far more than a
simple list. Why is it such a complicated shape?
In the Periodic Table the elements are arranged
in order of Atomic Number.
The shape and arrangement of the Periodic
Table is a very clever device to allow many
patterns and groupings to be accommodated.
You have already learnt one pattern in the
position of the most active metals, and their 1st
Ionisation Energies.
There are lots more...
Protons + Neutrons = “Mass Number”
(Electron mass is insignificant)
The Mass Number is always a whole number,
but in the Periodic Table the “Atomic Weight” is
shown instead.
(How and why this is different will be explained in a later topic)
History of the Periodic Table
Mendeleev used many physical and chemical
properties:
• atomic weight
• density
• melting point
• formula of oxide compound
• density of oxide
and many more,
and arranged the elements in order of weight, but
with elements with similar properties under each
other.
The modern concept of a chemical element
developed almost exactly 200 years ago.
By 1830 there were about 40 known elements.
Even with such a small sample, people began to
notice patterns:
Dobereiner (German) pointed out that there
were several
groups of 3 elements with
remarkably similar properties:
Similar elements placed in
vertical columns
Inert Gases had NOT
been discovered
Lithium, sodium & potassium was one “triad”.
Chlorine, bromine and iodine formed another “triad”.
Mendeleev’s vertical “families”
included Dobereiner’s “triads”
and Newland’s “octaves”, but
had one big difference...
By 1860, with over 60 known elements,
Newlands (English) proposed a “Law of
Octaves”.
If the elements were arranged in order of relative
weights, Newlands found that every 8th element
(an “octave”) was similar in properties. These
similar elements included Dobereiner’s triads.
Mendeleev’s genius was to realise that there were
probably missing elements that hadn’t been
discovered yet. He cleverly left gaps in his table for
these undiscovered elements.
The system worked well for the first 20 elements,
but then became confused.
The most famous case was that of the “missing”
element Mendeleev called “eka-silicon”. He used the
patterns in his table to predict, very precisely, the
properties for eka-silicon. Scientists went looking for
such a substance and soon found a new element
(which was named “Germanium”) with properties
exactly as predicted.
The basis of the modern Periodic Table was
developed by the Russian, Dmitri Mendeleev
in 1869.
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Patterns of the Periodic Table
In Mendeleev’s day no-one could explain why these patterns existed.
However, when scientists see patterns in nature like this, they know there must be
underlying “rules” or “laws of nature” causing and controlling the patterns.
Perhaps Mendeleev’s great contribution was not just the Periodic Table itself,
but the stimulus it gave other scientists to investigate the reasons behind the patterns.
Within 40 years Science had unravelled the secrets of atomic structure,
the electron energy levels, and more.
At this stage, your task is to learn some of the patterns.
Melting Point
You learned in topic 1 how melting point is determined by
the bonding within a substance.
At the left side of the table are the very active metals of the
Activity Series. They are also usually soft, and have
relatively low (for metals) melting points.
Electrical Conductivity
As you go across any row (“period”) of the table,
you will move through a number of metals, then one
or two semi-metals, then into the non-metals.
Therefore, the conductivity will start out high, but
rapidly decrease as you encounter a semi-metal,
and become extremely low at the non-metals.
Semi-Metals
NonMetals
Metals
Moving to the right across a period you enter the “Transition
Block” containing typical hard, high melting point metals,
held strongly together by “metallic bonding”.
Further right you hit the Semi-Metals. These often have very
high melting points because of their covalent lattice
structure.
Then you enter the Non-Metals which have covalent
molecular structures and quite low mp’s. At the far right
column, each period ends with an Inert Gas which are all
single-atom molecules, and have the lowest mp of each
period.
This pattern repeats itself along each period.
Conductivity
2,000
decreasing
Melting Points of Elements
Periods 3
V
(oC)
Boiling Points
follow a similar pattern to
Melting Points
Sketch Graph.
0
Melting Point
1,000
Si
Valencies are the same
down each group
Peaks are Transition Metals
Metals
or Semi-M
Period 4
Rb
K
Na
Inert Gases
Ar
Kr
Atomic Number
Chemical Bonding, Valency & Reactivity
What you’ve already learnt about the Activity Series, Ionic and Covalent Bonding and Valency
will help you make sense of the following:
Group 8 Inert Gases
+1
0
+2
+3
4
-3
3
1
-2
2 -1
Activity of Metals
Most active at
bottom-lleft.
Activity (generally)
decreases upwards
and to the right.
Metals
(+ve ions)
Bonding
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Metals
Activity of Non-M
Semi-M
Metals
(Covalent only)
Most active at top-rright
(Fluorine)
Non-M
Metals
(Covalent or (-v
ve) ions)
13
No chemical reactions,
no bonding
Activity (generally)
decreases downwards
and to the left.
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Atomic Radius
The following diagrams
are to scale and show
the relative sizes of the
first 20 elements
The size of an atom is the distance across its outer electron shell.
You might think that the atoms along each period would be the
same size, because it’s the same orbit being added to.
However, the increasing amount of positive charge in the
nucleus pulls that orbit inwards closer and closer to the centre.
H
37
Na
186
K
231
Radius increasing down a group
Li
152
He
50
The numbers given are the atomic radii in picometres.
1 picometre = 1x10-112 metre
Mg
160
Ca
C
B
Be
N
77
88
112
Al
66
P
Si
143
O
70
S
110
118
102
Ne
F
70
68
Ar
Cl
94
99
Radius decreasing across a period
197
Down each group the radius increases.
This is because, as you go down a group, you have added
an entire electron shell to the outside of the previous layer.
“Periodic” means “recurring at
regular intervals”.
This graph shows what a
spreadsheet plot gives for the
radii of the first 37 elements.
Notice how the same graphical
pattern keeps recurring... it is a
periodic pattern.
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300
200
Rb
K
De
acr creasi
oss ng
a p Tren
erio d
d
Li
100
When you do, you can clearly
see how the Periodic Table got
its name.
rend
sing T
Increa a group
down
Na
He
Ne
g Trend
Increasaingroup
n
dow
Ar
Kr
0
The Syllabus requires that you
produce a table and a graph of
the changes in a property
• across a period,
and
• down a group
Atomic Radius (picometre)
Spreadsheet Plot of Atomic Radii
1
10
20
30
Atomic Number
There are a number of irregularities and “glitches”
apparent on the graph. It is beyond the scope of
this course (and way beyond the K.I.S.S. Principle)
to attempt an explanation of these.
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Ionisation Energy
Successive Ionisation Energies
The meaning of the “1st Ionisation Energy” was
explained previously in relation to the Activity
Series of Metals.
Having added the energy of 1st I.E. and removed
an electron from any atom, it is then possible to
add more energy and remove a 2nd electron,
and a 3rd, and so on.
A+(g)
A(g)
+
e-
where “A” stands for any atom
in the gas state
Any atom can lose an electron if enough energy
is supplied... even atoms which do not normally
lose electrons.
You should remember that the very active
metals are the ones with low 1st ionisation
energies. They easily lose their outer electron(s)
and so react readily.
decreasing
Highest value
1st I.E. increases to the right because each atom
across a period has more and more (+ve)
nuclear charge attracting and holding electrons
in the orbit concerned. Therefore, it requires
more energy to remove an electron.
1st I.E. decreases down each group because, at
each step down, an extra whole layer of electrons
has been added to the outside of the atom. The
outer shell is further away from the nucleus, and
is partially “shielded” from nuclear attraction by
the layers of electrons underneath it. Therefore, it
becomes easier and easier to remove an electron.
A+(g)
A2+(g)
+
e-
3rd I.E.
A+2(g)
A3+(g)
+
e-
Element
Electron
Config.
1st
I.E.
2nd
I.E.
3rd
I.E.
4th
I.E.
Sodium
2.8.1
0.5
4.5
6.9
9.6
Magnesium 2.8.2
0.7
1.4
7.7
10.5
Aluminium 2.8.3
0.6
1.8
2.8
11.6
Notice how the values “jump” (underlined data)
as the next ionisation has to remove an electron
from the next lower orbit.
Highest Value
Fluorine
(values decrease down)
Electronegativity
Preliminary Chemistry Topic 2 “Metals”
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2nd I.E.
Patterns in Successive Ionisation Energy Data
(values shown are energy units)
Successive Elements on Period 3
Explanations:
Atoms with a tendency to gain
electrons and form negative ions have
high values.
Atoms with a tendency to lose
electrons easily (low 1st I.E.) and form
(+ve) ions have very low values.
Once again, there is a pattern in these
values in the Periodic Table.
e-
+
Once the entire outer orbit has been stripped
away, the next ionisation must remove an
electron from an underlying orbit, which
requires a huge increase in the next ionisation
energy. This results in an interesting pattern:
increasing
is a value assigned to each element to
describe the power of an atom to
attract electrons to itself.
A+(g)
Once the first electron is removed, the
remaining electrons are pulled in tighter to the
nucleus. Each one experiences increased force
of attraction, so it requires more energy to
remove the next electron. Therefore, each
successive ionisation requires more energy.
The trend for the whole Periodic table is:
Lowest
A(g)
...and so on,
according to how many electrons
the atom has
The Periodic Trend
in 1st Ionisation Energy
1st Ionisation
Energy
1st I.E.
1.0 1.5
0.9
Electronegativity Values
of selected elements
(values decrease to left)
Inert gases
not included
2.0 2.5 3.0 3.5 4.0
3.0
0.8
2.8
0.8
2.5
0.7
2.2
0.7
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Worksheet 5
Patterns of the Periodic Table
Fill in the blank spaces.
Student Name...........................................
As early as 1830, the German
a).......................................
noticed
patterns in the properties of the
elements. In 1860, the English scientist
b)................................ proposed a “Law
of c)..............................” describing the
repeating pattern of properties.
Chemical Reactivity is different for
metals and non-metals. The most active
metals are located at the left
z)......................... (top/bottom) of the
table. Generally, activity decreases
aa)......................
and
to
the
ab)............................. The Inert Gases
show no chemical activity. Apart from
them, the most active non-metals are
located on the right ac)............................
(top/bottom) of the table. Activity
generally decreases as you move
ad).......................... and ...........................
It was the Russian d).................................
who invented the e)......................
......................, in more or less its modern
form. He realised that there were
probably many elements that had not
f)........................, so he g)............
.................. in his table for later
additions. By studying the details of
known elements, he was able to
h)........................ very precisely the
properties of the missing elements.
Atomic Radius ae) ....................................
across a period because each
successive element has af)......................
(more/less) positive charge in the
ag)........................ to attract the electron
shell and pull it inwards. As you go
down
a
group
the
radius
ah)........................... as each new
electron shell is added.
Sure enough when discovered, the
missing elements were found to have
properties i)............................................
The patterns in the Periodic Table
include:
First
ai)........................
Energies
aj)....................... across a period, as the
increasing amount of nuclear charge
makes it more and more difficult to
ak)............................ an electron. The
values al)...................... down a group
because each extra shell of electrons is
am)................. (more/less) strongly held
than the previous.
Conductivity,
which
generally
j)............................ to the right, as you go
from metals to k).................................
and ...............................
Melting Points: tend to l)........................
to about the middle of each period, then
m)............................. The highest value is
usually a n)..................... metal or one of
the o)............................... elements. The
lowest value on each period is always
the p)............................. gas member on
the extreme q)......................... (right/left)
Successive
Ionisation
Energies
measure the energy required to
an)............................
another,
subsequent electron from an atom. The
energy required to remove the next
electron is always ao)...............................
(higher/lower). When the next electron
happens to be in the next lower shell,
the value ap)................................ by a
huge amount.
Valencies are r)............................... down
each vertical group. Bonding follows
the pattern of the main categories of
elements.
s)........................... form
t)........................... bonds when they lose
electrons and become u)....................
ions. The Semi-metal elements form
only v)........................... bonds. The Nonmetals can bond w)................................
or can x).................. electrons to form
y)................. ions.
aq).................................... is a value
which describes the power of an atom to
ar)............................. electrons. The
element with the highest value is
as)............................,
and
values
decrease as you move to the
at)......................... and as you move
au)............................ the Periodic Table.
WHEN COMPLETED, WORKSHEETS
BECOME SECTION SUMMARIES
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Worksheet 6
Patterns of the Periodic Table
Test Questions
Student Name...........................................
6. (5 marks)
a) Sketch a graph (values are not required) to
show the general changes in melting points of
the elements across one period of the Periodic
Table.
Multiple Choice
1.
The scientist most responsible for
development of the Periodic Table was:
A. Avogadro
B. Newlands
C. Gay-Lussac
D. Mendeleev
the
2.
Element “X” is in Group 2 and element “Y” in Group 7.
If X & Y
formed a
compound,
you would
expect it to be
A.
B.
C.
D.
ionic, with formula X2Y
covalent, with formula X2Y
ionic, with formula XY2
covalent, with formula Y2X
b) Briefly explain the general trend shown in
your graph.
3.
If the elements “X” & “Y” in Q2 lie in the same
period of the table, you would expect:
A. X to have a smaller radius than Y.
B. Y to have a higher electronegativity than X.
7.
On each of the following Periodic Table
diagrams label the arrows with the word
“increasing” or “decreasing” to correctly
describe the trend in the direction shown.
C. X to have a higher 1st ionisation energy than Y.
D. Y to have a higher melting point than X.
a) Atomic
Radius
4.
The reason for the trend in atomic radius as you
move across a period to the right, is:
A. increasing nuclear charge.
B. addition of extra electron shells.
C. decreasing attraction of electrons to the nucleus.
i)
ii)
D. increasing mass of the atoms.
b) Electronegativity
Longer Response Questions
Mark values shown are suggestions only, and are to
give you an idea of how detailed an answer is
appropriate. Answer on reverse if insufficient space.
Also indicate
(“H”&“L”) the
position of
elements with
highest &
lowest values.
5. (5 marks)
a) Write equations to represent the 1st, 2nd, 3rd
& 4th ionisations for a calcium atom.
i)
ii)
c) 1st
Ionisation
Energy
i)
Show“H”&”L”
ii)
b) Between which two of these successive
ionisations would you expect a huge increase in
the required energy?
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4. QUANTITY CALCULATIONS & THE MOLE
Defining the Mole
Quantities in
Chemical Calculations
For technical reasons, the “atomic standard”
used to compare the masses of all atoms is the
carbon atom, which contains
Atoms, molecules and ions always react with each
other in fixed, whole-number ratios. That’s why
balancing an equation is so important... it actually
brings the equation into line with what is happening
at the particle level.
6 protons
6 neutrons
6 electrons
For example, when hydrogen and oxygen react to
form water, the balanced equation is
2H2 + O2
Atomic Number = 6
Mass Number = 12
2H2O
This is a true description of what is happening to the
molecules:
2 Molecules
of H2
+
1 Molecule
of O2
6p+
6n0
The mass of this atom is
defined to be exactly 12.000000
atomic mass units (a.m.u.) and all other atoms
are given masses relative to this one.
Since this is the standard of comparison, the
formal definition of the mole is:
“the number of atoms contained in
exactly 12 grams of carbon-12”
2 Molecules
of H2O
However, when we carry out chemical reactions in the
laboratory or in Chemical Industry, we cannot see or
count the molecules. Instead, we measure the mass
or volume of substances.
Note: In Topic 1 it was pointed out that the Mass
Number for any atom is a whole number. It has still
not been explained why the “Atomic Weights” in the
Periodic Table are mostly not whole numbers.
To measure out the correct numbers of particles for a
reaction we need a simple way to convert masses and
volumes to numbers of molecules, and vice-versa.
That’s the purpose of
This WILL be explained in a later topic.
If you cannot wait, go find out about “Isotopes”.
The Mole
Avogadro’s Number
1 mole is a quantity of a chemical substance.
Just how many atoms are in 1 mole?
1 mole of any element or compound contains
exactly the same number of particles.
Obviously, it is a very large number. We now
know that it is about 6,000 billion trillion.
1 mole of each substance has a different mass,
because the atoms and molecules all weigh
differently.
Avogadro’s Number
6.022 x 1023
particles in 1 mole of anything
The really clever and convenient thing about the
mole is its link to the Periodic Table and the
“Atomic Weights” shown.
6
C
This number is named in honour of an Italian
scientist who you will learn about soon.
82
18
207.2 grams of
Lead
contains
6.022 x 1023
Lead atoms
Ar Pb
Carbon
Argon
Lead
12.01
39.95
207.2
1 mole
1 mole
1 mole
= 12.01 grams
= 39.95 grams
= 207.2 grams
39.95 grams of
Argon
contains
6.022 x 1023
Argon atoms
EACH OF THESE HAS THE SAME NUMBER OF ATOMS
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18
12.01 grams
of
Carbon
contains
6.022 x 1023
Carbon atoms
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Calculating Mole Quantities
Moles and Numbers of Particles
You need to be able to calculate mole quantities
in terms of both mass and number of particles.
Since one mole of any substance contains
Avogadro’s Number of particles:
Molar Mass
No. of moles = No. of particles you have
Avogadro’s Number
The “Molar Mass” of any chemical species is the
mass (in grams) of 1 mole of the substance.
n= N
NA
You need to add up all the Atomic Weights of all
the atoms given in the formula.
Examples:
Name
Argon
Sodium
Formula
Ar
Na
Example Calculations
1. How many moles are present in a sample of lead
containing 7.88 x 1024 atoms?
Molar Mass (g)
39.95
22.99
(for elements like these just use Atomic Weight)
n= N
NA
Solution
Oxygen
Chlorine
O2
Cl2
(16.00 x 2) = 32.00
(35.45 x 2) = 70.90
(these elements are diatomic molecules... 2 atoms each)
= 7.88x1024
6.022x1023
= 13.1 mol
2. a) How many atoms of lead are needed to make
0.0250 mole?
b) What would be the mass of this quantity?
Water
H2O (1.008x2 + 16.00) = 18.016
Carbon Dioxide CO2 (12.01 + (16.00x2)= 44.01
Sodium chloride NaCl (22.99 + 35.45) = 58.44
Solution
a) n = N
so N = n x NA = 0.0250 x 6.022x1023
NA
= 1.51 x 1022 atoms
(add up At.weights of all atoms in the formula)
Worksheet at the end of this section
b) m = n x MM = 0.0250 x 207.2 (molar mass of Pb)
= 5.18 g
Number of Moles in a Given Mass
When you weigh a chemical sample you then
need to be able to calculate how many moles
this contains.
Mole Quantities
in Chemical Equations
When you consider an equation like
2H2 + O2
you know it means
No. of moles = mass of substance you have
molar mass
2H2O
n= m
MM
Example Calculations
1. How many moles in
Solution
2 Molecules
of H2
a) 5.23g of magnesium?
b) 96.7g of water?
a)
n = m = 5.23 = 0.215 mol
MM
24.31
b)
n= m
MM
+
1 Molecule
of O2
However, the number of molecules reacting is
really just a ratio. The actual numbers might be
=
96.7
(2x1.008 + 16.00)
= 96.7/18.016
= 5.37 mol
2 million H2 + 1 million O2
2 million H2O
or, 200 zillion H2 + 100 zillion O2
200 zillion H2O
or, (let’s use Avagadro’s number)
(2 x NA) H2 + NA O2
(2 x NA) H2O
2. What mass is needed if you want to have 1.50
moles of salt (sodium chloride)?
= 2 moles H2 + 1 mole O2
n= m
MM
so m = n x MM = 1.50 x (22.99 + 35.45)
= 1.50 x 58.44
= 87.7 g
2 moles H2O
The Balancing Coefficients
in a Chemical Equation
May be Interpreted as
Mole Ratios
Worksheet at the end of this section
Preliminary Chemistry Topic 2 “Metals”
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2 Molecules
of H2O
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Mole Quantities
in Chemical Equations (cont.)
Calculating Mass Quantities
in Reactions
The balancing coefficients of an equation can
be interpreted as the mole ratio of reactants and
products.
Mole calculations allow you to calculate the mass of
products and reactants involved in a reaction.
So,
2 H2
+
O2
Example Problem
Aluminium burns to form aluminium oxide.
If 4.29g of aluminium was burned,
a) what mass of oxygen would be consumed?
b) what mass of aluminium oxide would be formed?
2 H2 O
means 2 mol. reacts with 1 mol. to form 2 mol.
or,
4 mol. reacts with 2 mol. to form 4 mol.
or,
100 mol. reacts with 50 mol. to form 100 mol.
Solution
Always start with the balanced equation:
or any other proportional quantities.
Example Problem
a) If 0.50 mol of sodium reacted completely with
hydrochloric acid, how many moles of products
would be formed?
4 Al
mole
ratio
Solution
a) The balanced equation is
so,
H2
:
0.50 mol : 0.50 mol :
: 2 mol.
0.25 mol : 0.50 mol
:
b) Mass of Hydrogen: m = n x MM = 0.25 x 2.016
= 0.50 g
Mass of salt:
m = n x MM = 0.50 x 58.44
= 29 g
Worksheet at the end of this section
Using Mass & Mole Ratios
to Determine a Formula
ceramic
crucible
A common experiment is to burn a piece of magnesium in a crucible, as
suggested by the diagram.
Magnesium + Oxygen
= 4.29
26.98
= 0.159 mol
∴ mass of Al2O3: m = n x MM = 0.0795 x 101.96
= 8.11 g
Worksheet at the end of this section
Reaction:
2
b) Mass Al2O3 produced:
mole ratio Al : Al2O3 = 4: 2 (i.e. 2:1)
\ moles of Al2O3 = 1/2 x 0.159 = 0.0795 mol
Answer: 0.25 mol of H2 and 0.5 mol of NaCl
Practical Work:
3
a) Mass O2 consumed:
mole ratio Al : O2 = 4 : 3
∴ moles of O2 = 0.159 x 3
= 0.119 mol
4
∴ mass of O2: m = n x MM = 0.119 x 32.00
= 3.81 g
+ 2 NaCl
1 mol
:
No. of moles of Aluminium: n = m
MM
b) What mass of each product would be formed?
2 Na + 2 HCl
mole
ratio 2 mol :
2 mol
4
2 Al2O3
+ 3 O2
Magnesium oxide
Careful measurement of mass allows the empirical formula for magnesium
oxide to be determined.
Typical Measurements
Mass of empty crucible = 42.74 g
Mass of magnesium
= 2.05 g
Mass of crucible
+ product after burning = 46.22 g
∴ Mass of magnesium oxide
formed = 3.48 g
∴ Mass of oxygen in
compound
= 1.43 g
Preliminary Chemistry Topic 2 “Metals”
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Analysis of Results
Remember that to convert any mass to moles:
n = m / MM
Elements
Ratio of masses:
Ratio of moles:
Magnesium :
2.05 g
:
2.05 / 24.31 :
Oxygen
1.43 g
1.43 / 16.00
(divide by Atomic Weight)
= 0.0843
Simplified ratio = 0.0843/0.00843 :
=
1.0 :
Nearest whole number ratio 1 :
∴ Empirical Formula is MgO
20
mol : 0.0894 mol
0.0894/0.0843 (divide both by the
1.06
smaller)
1
There is often a large error
due to incomplete burning
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A Little History... How the Mole was Invented
Avogadro’s Hypothesis
Gay-Lussac’s Law
The Italian, Amadeo Avogadro (1776-1856) was trained in
Law, but became very interested in Science.
Joseph Gay-Lussac was a French scientist with
an unfortunate name by modern standards. He
lived 200 years ago, and was very interested in
flight using balloons, so he investigated the way
gases react chemically.
In 1811, he noticed the similarity between GayLussac’s Law (an empirical “law” based on
experiment) and the concept that atoms must
combine in simple, whole number ratios to form
compounds.
After a series of clever experiments, in which the
volumes of reacting gases were measured, in
1808 he proposed the “Law of Combining
Volumes”:
This led him to make an hypothesis:
Equal Volumes of all Gases
Contain Equal Numbers of Molecules
When measured at constant temperature
and pressure, the volumes of gases in a
chemical reaction show simple, wholenumber ratios to each other.
(when measured at the same conditions
of temperature and pressure)
This was a vital breakthrough in the history of
Chemistry.
The volume of a gas is easily changed by temperature
and pressure, so it is very important that the volumes
are all measured at the same conditions.
For example, consider the reaction:
Examples of Gay-Lussac’s Law
Hydrogen(g) + Chlorine(g)
1 litre
1 litre
Hydrogen(g) + Oxygen(g)
2 litres
1 litre
Hydrogen(g) + Nitrogen(g)
3 litres
1 litre
Hydrogen(g) + Chlorine(g)
Hydrogen chloride(g)
2 litres
Prior to Avogadro, it was assumed that the the
reaction involved single atoms, like this:
Water(g) (vapour)
2 litres
H(g)
Hydrogen(g) + Chlorine(g)
1 volume :
2 H2(g) +
Cl2(g)
2 HCl(g)
O2(g)
2 H2O(g)
3 H2(g) + N2(g)
2 NH3(g)
HCl(g)
1 volume
:
Hydrogen chloride(g)
2 volumes
Now, reasoned Avogadro, gases react in simple,
whole-number volume ratios because each litre
of gas has the same number of molecules in it.
Therefore, to get the volume ratios shown
above, each hydrogen molecule, and each
chlorine molecule, must have 2 atoms!
Now consider the balanced equations for these
three example reactions:
+
+ Cl(g)
but the combining volumes (discovered by
experiment) were
Ammonia(g)
2 litres
Notice that in every case, that the volumes are
always in a simple, whole number ratio to each
other.
H2(g)
Hydrogen chloride(g)
i.e. Hydrogen is H2(g) and Chlorine is Cl2(g), and
the correct equation is
H2(g)
The mole ratios are the same as the volume
ratios discovered by Gay-Lussac!
+
Cl2(g)
2 HCl(g)
Then, for the same reaction, scientists could
measure the masses of these gases as well as
volumes. This showed that chlorine atoms must
be about 35 times heavier than hydrogen...
chemists were on the way to figuring out the
relative atomic weights of elements, and being
able to calculate chemical quantities.
Why should this be?
... enter
Avogadro!
Although he did not invent the concept of the
mole, we name the number of particles in 1 mole
in Avogadro’s honour... Avogadro’s Number.
Preliminary Chemistry Topic 2 “Metals”
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Molar Volume of a Gas
Comparing Mass Changes
When Metals Burn
If 1 mole of any chemical species contains the
same number of particles (Avogadro’s Number)
AND if equal volumes of gases contain equal
number of particles (Avogadro’s Hypothesis),
then it follows that
Atoms always react in simple whole-number
mole ratios, but atoms have different masses,
and compounds have various formulas, so the
result is that chemicals do NOT react in simple
ratios by mass.
1 mole of any gas
must occupy the same volume,
That’s why we need the mole... we measure
quantities by their mass, but this makes no
sense until moles are calculated.
if measured at the same
temperature and pressure.
The syllabus requires that you should consider
the mass changes involved when various
metals combine with oxygen to form their oxide
compound.
This volume is the “Molar Volume” and is the
same for every gas. Usually, the volume is
measured at 25oC and a pressure of 100 kPa.
(kPa = kilopascals, the normal unit for
measuring gas pressures in Chemistry.)
The following table shows the mass changes for
100g of the metal in each case:
100g of
Metal
Formula
of oxide
Lithium
Li2O
Iron
Mass O2
needed(g)
1 mole of any
gas = 24.79 litres
o
at 25 C and 100 kPa
Mass of
Oxide formed
115
215
Fe2O3
43
143
Zinc
ZnO
49
149
Lead
PbO2
15
115
Mole Calculations Involving Gases
This additional knowledge opens up the
opportunity to carry out quantity calculations
which involve mass and volumes of gases.
Example Problems
1.
If 15.65g of calcium carbonate (CaCO3) was
completely decomposed by heat, what volume of
carbon dioxide gas would be produced (if
measured at 25C, 100kPa)?
Empirical & Molecular Formulas
You are reminded that a molecular formula really
does describe the atoms present in a molecule.
Solution
Always begin with the balanced equation for the
reaction.
CaCO3(s)
CO2(g) + CaO(s)
mole ratio = 1
:
1
:
1
Moles of CaCO3: n = m = 15.65 = 0.1564 mol
MM
100.09
Mole ratio is 1 : 1, so moles of CO2 formed = 0.1564
The molecular compound methane,
has formula CH4, because that’s
exactly what each molecule contains...
1 carbon atom and 4 hydrogen atoms.
Lattice structures, either ionic or covalent
are NOT molecular.
Example: sodium
chloride, NaCl
∴ Volume of CO2 = 0.1564 x 24.79
= 3.877 L (at 25C, 100kPa)
Molar Vol. of
all gases at
25C, 100kPa
The formula does NOT
describe a molecule,
but only gives the simplest ratio between the
bonded atoms... this is an empirical formula.
2.
What volume of hydrogen gas (at 25C, 100kPa)
would be produced if 10.00g of lithium metal was
reacted with sulfuric acid?
Earlier was an example of how formulas are
determined by analysing the mass composition
of a compound.
Solution
2 Li(s) + H2SO4(aq)
2
:
1
H2(g) + Li2SO4(aq)
1
:
1
Moles of lithium: n = m
= 10.00 = 1.441 mol
MM
6.941
Mole ratio is 2:1, so moles of H2 = 1/2 x 1.441=0.7204
You should note that this method can only
produce an empirical formula. (In fact, the word
“empirical” means something determined by
experiment, not by theory.)
∴ Volume of H2 = 0.7204 x 24.79
= 17.86 L (at 25C, 100kPa)
If a molecular compound, with molecular
formula X2Y6 was analysed by mass
measurements, its empirical formula would be
calculated to be XY3... simplest ratio of atoms.
Preliminary Chemistry Topic 2 “Metals”
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:
Worksheet at the end of this section
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Worksheet 7
Masses, Moles & Particles
Practice Problems
Student Name...........................................
1. Molar Masses
3. Moles and Number of Particles
Calculate the molar mass of:
a) How many particles (atoms/molecules) in:
a) potassium
c) tin
e) nitrogen gas
g) sodium iodide
i) ammonia
k) aluminium oxide
i) 3 moles of water?
ii) 2.478 mol of CO2?
iii) 5 mol of salt?
iv) 0.007862 mol of copper
v) 1/1000 mol of helium
b) krypton
d) bromine (Br2)
f) magnesium oxide
h) iron(III) sulfide
j) copper(II) sulfate
l) glucose (C6H12O6)
b) Convert between mass, moles and no.of particles.
i) If there are 8.800x1025 atoms of tin, how many moles
is this, and what would be the mass?
2. No. of Moles in a Given Mass
How many moles in:
ii) You have a sample containing 2.575x1024
molecules of water. How many moles is this, and what
is its mass?
a) 100.0g of lead?
b) 100.0g of zinc?
c) 100.0g of water
?
d) 100.0g of copper(II) nitrate?
e) 38.55g of magnesium fluoride?
f) 60.00g of carbon dioxide?
g) 1.000g of zinc oxide?
h) 500.0g of glucose (C6H12O6)?
i) 3.258 x 10-3g of salt (sodium chloride)?
j) 128.6g of ammonium carbonate?
iii) If you weigh out 400.0g of water, how many moles
is this, and how many molecules are present?
iv) If you have 2.569g of pure nickel, how many atoms
are there?
v) What mass of sulfur would contain 2.500x1023
atoms?
Worksheet 8 Mole Ratios & Mass in Reactions
Practice Problems
Student Name .........................................
1. Mole Ratios in Equations
2. Mass Quantities in Reactions
Sodium reacts with water as follows:
2Na + 2H2O
H2 +
2NaOH
a) Calcium burns in oxygen to form calcium oxide:
2Ca + O2
2CaO
If 8.50g of calcium reacted, what mass of calcium
oxide would be formed?
a) If 1 mole of sodium reacted, how many moles of
i) hydrogen formed? ii) water consumed?
b) Silver carbonate decomposes when heated:
2Ag2CO3
2CO2
+
4Ag
+ O2
If 20.0g of silver carbonate was decomposed
i) what mass of silver metal would form?
ii) what mass of CO2 would be produced?
iii) what mass of O2 would be formed?
b) If 0.25 mol of NaOH formed, how many moles of
i) sodium consumed? ii) hydrogen formed?
c) If 0.75 mole of hydrogen formed, how many moles of
i) sodium consumed? ii) NaOH produced?
d) If 0.5 mole of Al used, how many moles of
i) Alum.oxide formed?
ii) oxygen used?
c) Aluminium reacts with hydrochloric acid:
2Al + 6HCl
3H2
+ 2AlCl3
If 6.50g of aluminium reacted
i) what mass of HCl would be consumed?
ii) what mass of hydrogen gas produced?
iii) what mass of aluminium chloride produced?
e) If 0.1 mole of alum.oxide formed, how many
moles of
i) aluminium used? ii) oxygen used?
d) Tin reacts with steam as follows:
Sn(s) + 2H2O(g)
2H2(g) +
Aluminium reacts with oxygen:
4 Al
+ 3 O2
2 Al2O3
FOR MAXIMUM MARKS SHOW
FORMULAS & WORKING,
APPROPRIATE PRECISION & UNITS
IN ALL CHEMICAL PROBLEMS
Preliminary Chemistry Topic 2 “Metals”
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SnO2(s)
If 14.8g of tin reacted
i) what mass of tin(IV) oxide would be formed?
ii) What mass of steam would be consumed?
iii) what mass of hydrogen would be produced?
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Worksheet 9
Practice Problems
Empirical Formulas
Student Name .........................................
1. A compound containing only copper and
chlorine is decomposed, and the masses
measured to find the mass composition:
Mass of copper present = 12.84g
Mass of chlorine present = 7.16g
i) Find the empirical formula.
ii) Name the compound.
3. A compound was found to contain 30%
nitrogen and 70% oxygen by mass.
i) Find the empirical formula.
ii) It is later found that its molecular formula is a
2 times multiple of the empirical. Write the
molecular formula.
2. i) Find the empirical formula of a compound
containing carbon and hydrogen; a sample was
found to contain 1.5g of carbon and 0.5g of
hydrogen.
ii) Name the compound, given that its empirical
and molecular formulas are the same.
iii) Name the compound.
Worksheet 10
Reactions Involving Gases
Practice Problems
Student Name .........................................
1. Volumes of Reacting Gases
2. Mass & Gas Volume Calculations
o
All volumes measured at 25 C, 100kPa
( Assume all gases are measured at the
same temperature & pressure)
2 H2(g) +
O2(g)
a) To “scrub” the air and remove poisonous CO2
on board the Space Shuttle, the air is continually
pumped through canisters containing 5.00kg of
lithium oxide. The reaction is
2 H2O(g)
a) If you used 5 litres of hydrogen, how many
litres:
i) of oxygen consumed?
Li2O(s) + CO2(g)
Li2CO3(s)
ii) of water vapour formed?
i) How many moles of lithium oxide in each
canister?
b) If you used 0.25 litres of oxygen, how many
litres of
i) water vapour formed?
ii) How many moles of CO2 can this absorb?
iii) What volume of CO2(g) is this?
ii) hydrogen consumed?
b) Iron reacts with oxygen:
4Fe(s) + 3O2(g)
2Fe2O3(s)
c) If this reaction produced 20 litres of steam,
how many litres of
i) hydrogen consumed?
i) If 10.0L of O2 reacted, what mass of iron(III)
oxide would be formed?
ii) oxygen consumed?
ii) If 100g of iron reacted, what volume of oxygen
would be needed?
c) The electrolysis
decomposition:
2H2O(l)
Ammonia gas is formed by reaction of hydrogen
with nitrogen
3 H2(g) + N2(g)
2 NH3(g)
water
2H2(g)
+
causes
O2(g)
i) If 1.00g of water was decomposed, what
volumes of each gas (measured at 25C, 100kPa)
would be formed?
d) In order to make 9 litres of ammonia, what
volume
i) of hydrogen needed?
ii) of nitrogen needed?
In an electrolysis experiment, 50mL (0.050 L) of
oxygen was produced. (at 25C, 100kPa)
e) If 0.6 litre of hydrogen reacted, what volume
i) of ammonia formed?
ii) What volume of hydrogen (at 25C, 100kPa) was
produced?
iii) What mass of water must have been
decomposed?
ii) of nitrogen need?
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Worksheet 11
Mole Concept
Fill in the blank spaces.
The formal definition of the mole is “the
a)...................... of atoms in 12.00 grams of
b)..................................”
Student Name...........................................
By converting between the h).......................... of
substances
and
the
number
of
i)..............................., it becomes possible to
calculate all the quantity relationships during a
chemical j).................................... From the mass
composition it is also possible to determine the
k).................................. formula of compounds.
One mole of any substance contains the same
number
of c)............................ The mass
of 1 mole of any substance is equal to its
d)................................... in grams. The actual
number of particles in one mole of anything is
known as “e)...................................’s Number”
and has a value of f).......................
Historically, the mole concept arose from the
work of 2 men: The Frenchman l)..............................................
discovered
that
“the
m)........................... of gases in chemical
reactions
always
show
simple,
n)............................... ratios to each other”. Soon
after, the Italian o)................................. suggested
that “Equal p)......................... of all gases contain
q)....................... numbers of r)...........................
(when measured at the same conditions of
s).............................. and ..........................)
In a balanced chemical equation, the “balancing
numbers” (coefficients) may be interpreted as
being g)........................... .............................. of
reactants and products.
The standard conditions usually used are a
pressure of t)................ and temp. of u)........oC.
Worksheet 12 Calculations & the Mole
Student Name...........................................
Test Questions
5. Carbon monoxide gas reacts with oxygen gas to
Multiple Choice
form carbon dioxide gas as follows:
1.
An atom of argon is about twice as heavy as an atom
of neon. You would expect:
A. a mole of argon to contain about half as many
atoms as a mole of neon.
B. equal masses of each element to contain about the
same number of atoms.
C. 2g of argon to contain about the same number of
atoms as 1g of neon.
D. the molar mass of neon to be about twice the molar
mass of argon.
2CO(g) + O2(g)
If 100mL of carbon dioxide was produced, then the
total volume of reactants (all measured at the same
temp. & pressure) before the reaction would have
been:
A. 100mL
B. 150mL
C. 50mL
D. 250mL
Longer Response Questions
6. (6 marks)
a) Write a balanced equation for the reaction of
aluminium metal with hydrochloric acid.
b) If 6.58g of aluminium reacted fully, calculate:
i) the number of aluminium atoms involved.
ii) the mass of aluminium chloride formed.
iii) the volume of hydrogen gas. (at 25C, 100kPa)
2.
Which line shows correctly the molar mass (to the
nearest gram) of the named substance?
A. water, 18g
B. carbon dioxide, 28g
C. oxygen gas, 16g
D. helium gas, 8g
7. (4 marks)
It was found by experiment that a compound
containing only tin and oxygen, contained 88% tin, by
mass. Showing your working, determine the empirical
formula for this compound, and give its correct
chemical name.
3.
Aluminium reacts with oxygen to form aluminium
oxide.
4 Al
+ 3 O2
2 Al2O3
If 1 mole of aluminium (about 27g) was to be reacted,
you would need how many moles of oxygen gas?
A. 0.75 mol
B. 3 mol
C. 1 mol
D. 1.3 mol
8. (4 marks)
In the reaction of nitrogen and hydrogen gases to
form ammonia gas, it was found by experiment that
300mL of hydrogen reacted completely with 100mL of
nitrogen. 200mL of ammonia gas was produced. All
the gas volumes were measured at a pressure of 10
standard atmospheres and 150oC.
4.
Avogadro’s number can be described by the
abbreviation NA. If you had 2 moles of methane (CH4),
then the number of hydrogen atoms present is:
A. 2 x NA
C. 8 x NA
a) Write a balanced equation for the reaction.
b) Explain how the experimental measurements are in
agreement with Gay-Lussac’s Law.
B. 4 x NA
D. 10 x NA
Preliminary Chemistry Topic 2 “Metals”
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5. METALS FROM THEIR ORES
The Importance of
Predicting Yield from an Ore
Ores and Minerals
... and now back to the metals.
The whole situation of economic feasibility
makes the science of Analytical Chemistry vital
in the mining and metals industry.
Minerals are naturally occurring compounds.
“Rocks” are mixtures of various minerals. Most
minerals are lattice structures, both ionic and
covalent. Some very common minerals include:
Mining operations cost millions of dollars to set
up. To do so, the operators need to be sure that
the ore contains enough metal to be profitable.
Chemical analysis in the laboratory is used to
measure the mineral content of the ore body, to
predict the final yield of the metal.
• Silica, which is chemically silicon dioxide
(SiO2) and is the most common mineral on
Earth. Other compounds are often included in
the silica lattice to make “silicate” minerals.
These occur in virtually all rocks.
• Calcite, which is calcium carbonate (CaCO3) is
the main mineral in limestone and marble.
Some minerals contain significant quantities of
metal(s), chemically combined in the compound.
Ores are rocks and/or minerals from which it is
economically worthwhile to extract a desired metal.
It is the economic part of this definition which is
critical. For example, there are many rocks and
minerals that contain significant amounts of iron and
aluminium. These are not “iron ore” or “aluminium
ore” unless it is economically worthwhile to mine and
process them to get the metal.
Photo courtesy of
Comalco Aluminium Ltd
Ores are Non-Renewable
Resources
Minerals and ores have been formed over
millions and billions of years of geological
processes on Earth.
Because of that time-frame, the ores are nonrenewable in the sense that once we use them
up, they cannot be replaced.
There is no immediate concern for running out
of the most important ores, but unlimited
exploitation of any non-renewable resource is:
What Makes It Economically Worthwhile?
Basically, economic feasibility is the balance
between:
• irresponsible, to future generations.
• unsustainable, because all non-renewable
things must eventually run out.
• economically stupid, because it may be
cheaper to re-use and recycle, than
to constantly extract “new” materials.
• environmentally damaging, because mining
and metal smelting have a history of pollution
and ecosystem destruction.
• the Commercial Price for which the metal can
be sold and
• the Production Costs of mining and
transporting the ore, and chemically extracting
and purifying of the metal.
Another factor is the abundance of the metal
and its ores on Earth. For example, iron is
relatively cheap because it is very common in
huge deposits of iron ores.
In the not-too-distant future it may become
economically worthwhile to begin “mining” the
old rubbish dumps around our cities, to recover
the discarded metals in society’s garbage.
Platinum is very rare, so it commands a high
price. This makes it worthwhile to mine even
very low-grade ores. A low-grade iron ore would
not be worth mining!
Preliminary Chemistry Topic 2 “Metals”
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Case Study:
Extraction of Copper from its Ores
Copper Ores
include a variety of
compounds of copper, including:
Froth Flotation to Concentrate the Ore
• copper(I) sulfide, Cu2S
• copper(II) hydroxide mixed with
copper(II) carbonate, Cu(OH)2.CuCO3
The ore is crushed into a powder and the copper
minerals are separated from the silicates by a process
of “Froth Flotation” which relies on differences in
“wettability” and density.
Compressed air blown in to
create a froth of bubbles
These compounds usually occur as thin
“veins” of blue-green minerals embedded in
masses of worthless silicate rock.
Froth
The copper content of the entire ore body
might be only 3% or less. Therefore, the first
step after mining is to separate the copper
minerals from the “rock”.
Crushed
Ore
in a
slurry of
water and
“wetting
oil”
Waste
Mineral
Slurry
Chemistry of Smelting
The concentrated copper minerals now
undergo Decomposition Reactions.
Compressed air creates a froth of bubbles
in a detergent solution.
In Australia, the main copper ores contain
copper(I) sulfide. If this is heated in a furnace
supplied with plenty of air the reaction is:
Copper(I) sulfide + oxygen
Cu2S + O2
Froth
overflows
for
collection
Copper minerals, sprayed with a special oil,
cling to the bubbles and are carried upwards to
overflow with the froth.
Copper + Sulfur dioxide
2Cu + SO2
Silicate minerals are wetted by the water and,
being denser, sink to the bottom.
The copper collected is about 98% pure.
The collected froth is then treated to separate
the oil and detergent for re-use.
Sulfur dioxide is a serious pollutant if
released from the smelter.
These days it is collected and used to
manufacture sulfuric acid... a useful by-p
product.
The ore concentrate is now about 30% copper.
Final Purification by Electrolysis
The major use of copper is for electrical wiring. For this it needs to be 99.9% pure.
Copper is purified by electrolysis:
Cu
Cu2+ + 2e-
The copper dissolves
into the solution, but
impurities do not.
-
+
The impure copper is
immersed in CuSO4
solution and electrified:
Impure
Copper
dissolves
into
solution
Preliminary Chemistry Topic 2 “Metals”
Copyright © 2005-2
2009 keep it simple science
www.keepitsimplescience.com.au
Cu+2
ions
migrate through
CuSO4 solution
Pure
Copper
deposits
on
electrode
After migrating
through the solution,
the ions are redeposited as pure
copper metal on the
other electrode:
Cu2+ + 2e-
Cu
Impurities
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Producing the electricity usually
involves the burning of coal at a power
station. The burning of fossil fuels like
coal is a major contributor to the
“Greenhouse Effect” which many
scientists are now convinced is causing
massive climate changes to the entire
Earth.
The Case for Recycling
The point that mineral ores are nonrenewable has already been made.
Eventually, any non-renewable resource
must run out, so recycling is inevitable.
There is also a strong environmental
case for recycling of metals, especially
aluminium.
Extracting
aluminium from
its ore requires
about 200kJ
(kilojoules) of
energy per kg
of metal. This
energy is
mainly in the
form of
electricity,
which is
needed in huge
quantities for
the electrolytic
smelting
process.
Recycling aluminium requires about 7kJ
of energy, a saving of about 96% in
energy and environmental impact!
Most local councils now operate “Recycling Centres” which can sort out paper,
glass, plastic, etc from our garbage, as long as we remember to put recyclables in
the correct bin. Aluminium (mainly drink cans) collected this way is returned to
scrap-metal businesses which clean and re-melt the metal to return it to
manufacturing industry for re-use.
Scrap Metal
awaiting recycling.
Photo by Pawel Grabowski
Preliminary Chemistry Topic 2 “Metals”
Copyright © 2005-2
2009 keep it simple science
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Usage & copying is permitted according to the
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keep it simple science
Worksheet 13
Metals from Ores
Fill in the blank spaces.
Student Name...........................................
“Minerals”
are
naturally
occurring
a).............................. which are mixed together in
rocks.
An “ore” is a b)............................. from which it
is c)................................ worthwhile to extract a
desired d)..............................
After mining, the ore is crushed, then
concentrated
by
“n)...........................
.......................................”. This process uses a
froth
of
bubbles
to
separate
the
o).............................. density copper compounds
from the worthless rock which is mainly
p)................................. minerals.
Whether it is worthwhile (or not) to mine an ore
depends on the balance between the
e)........................................ and the f).....................
................................... of mining, transporting
and g)........................... the metal.
The
“smelting”
process
involves
q)...................................... reactions. For a sulfide
ore, it reacts with r)....................... to form
s)...................... metal and t).............................
gas.
h)........................... analysis of an ore deposit is
vital to predict the i)..................................... from
the ore, to determine if it is worth mining.
The final step is to u)........................... the copper
by a process of v)............................................
There are many good reasons to w)......................
metals, especially x)............................ which
requires large amounts of y).................................
energy to extract from its ore. Producing the
electricity required is often done by burning
z).......................... fuels such as aa)......................
This contributes to the “ab).................................
Effect”, responsible for global climate changes.
Recycling aluminium requires only a fraction of
this energy.
Ores are j)........................................... resources
because
once
used,
they
cannot
k)........................................... due to the immense
time it takes for l).......................................
processes to form them.
Copper ores contain compounds such as
m)........................... and ......................................
Worksheet 14 Metals from Ores
Test Questions
Student Name...........................................
3. (8 marks)
a) Give the name and formula for a compound
commonly found in copper ores.
Multiple Choice
1.
The “smelting” of a metal ore always involves:
A. separating the metal mineral from the rock.
B. decomposing a compound of the metal.
C. purifying the extracted metal by electrolysis.
D. all of the above.
b) Name, and briefly describe the process by
which a copper ore is concentrated and
separated from the surrounding “rock”.
Longer Response Questions
Mark values shown are suggestions only, and are to
give you an idea of how detailed an answer is
appropriate. Answer on reverse if insufficient space.
2. (5 marks)
a) Differentiate between a “mineral” and an
“ore”.
c) Write a chemical equation to describe the
reaction which occurs in the smelting of the ore.
(Involving the compound you named in part (a))
b) Outline the role of Chemical Science in
assessing the economic feasibility of mining a
mineral resource.
d) Name the process by which the smelted
copper is purified, and relate the need for
purification to a common use of the metal.
c) Briefly discuss the sustainability of using the
Earth’s mineral resources, and outline a
strategy for conservation.
Preliminary Chemistry Topic 2 “Metals”
Copyright © 2005-2
2009 keep it simple science
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