Transcript 8.3 Metals - slider-chemistry-11

8.3 Metals
Focus 1:
Metals have been extracted and used
for many thousands of years
Metals – historical uses of copper
• Bronze (alloy of tin(10%) and copper)
used to make axe heads, other
weapons and tools. (Early Bronze age
c. 3500 BC). The very first alloy!
• Brass (alloy of copper(65%) and zinc)
used for decorative items in Roman
homes (c. 100 BC)
• Copper wires used in the transmission
of electricity (1878)
• Solar cells, using copper wiring,
generate electricity from the sun
(1954)
Metals – historical uses of iron
• Tips of spears, daggers and ornaments,
were being fashioned from iron recovered
from meteorites. (c. 4000 BC)
• The Hittites from Turkey learned to smelt
iron to make weapons that were superior to
Bronze. (c. 1500 BC)
• Steel (iron with carbon added) first made in
India c. 350 BC.
• Shipbuilding with steel began in the early
1900's, then skyscrapers started going up
around 1910.
• The first stainless steel was melted on the
13th August 1913. It contained 0.24%
carbon and 12.8% chromium. This
replaced silver coated cutlery
The Bronze Age
The Bronze Age in the Middle East (known
at this time as the Near East) is divided
into three main periods (the dates are very
approximate):
• EBA - Early Bronze Age (c.3500-2000 BC)
• MBA - Middle Bronze Age (c.2000-1600 BC)
• LBA - Late Bronze Age (c.1600-1200 BC)
The Bronze Age in the Middle East
• EBA - The mountains of
Anatolia (modern-day
Turkey) were rich in
copper and tin.
Metallurgy began here.
• EBA - Copper was also
mined throughout
Mesopotamia, from
Egypt to the Persian
Gulf.
• MBA – Increased trade
of metals among states.
• LBA – The Hittites in
Anatolia were
dominating battles by
using iron weapons.
The Bronze Age in the rest of the world
Great Britain Bronze Age
(c. 2100-to 700 BC)
China – Xia Dynasty
(c. 2100-700 BC)
Central Europe
(c. 1800-500 BC)
Scandinavia
(c. 1500-500 BC)
The Iron Age in the Middle East
• Ancient Egyptians
used iron from
4000 BC
• Widespread use
was not seen until
@1300 BC when
the Hittites used it
for weapons.
• Around 1000 BC,
iron took over
bronze in the
Middle East as the
dominant metal.
Iron Age in India
• Iron implements
found as far back
as 1800 BC in the
province of Uttar
Pradesh
• Carbon steel was
being produced
@300 BC and was
exported
throughout Asia
and Europe
Map of early Iron Age Vedic India. This map
shows the North-western portion of modernday India.
The Iron Age in Europe
• Iron working was
introduced to Europe
around 1000 BC,
probably from Asia
Minor (or Anatolia).
• In Eastern Europe, the
Iron Age begins
around 900 BC.
• The British Isles Iron
Age lasted from 500
BC to 500 AD.
Metals in the modern era
• Building materials – nails, beams,
windows, electrical wiring (Cu, Fe, Al)
• Transportation – cars, trains, planes
(Al, Ti, Cr and alloys of Fe)
• Replacement parts for the human body
- (Ti, alloys of Co, stainless steel)
• Electronics - (Sn/Pb alloy-solder, Cu,
Fe, Ni, Cr)
• Jewellery - (Au, Ag, Cu, Pd, Pt)
• Money - (Ni/Cu alloy-silver coins;
Cu/Al/Ni alloy- gold coins)
Energy is required to extract metals
In order to extract metals from their ores,
energy is required to break the existing
bonds in the minerals. In ancient times,
ores were heated with carbon (charcoal).
Copper:
heat
2CuO(s) + C(s)  2Cu(l) + CO2(g)
Iron:
heat
2Fe2O3(s) + 3C(s)  4Fe(l) + 3CO2(g)
8.3 Metals
Focus 2: Metals differ in their
reactivity with other chemicals and
this influences their uses
Reactions of metals
Reactions with oxygen (combustion)
All metals form oxides except Ag, Au and Pt
heat
Metal + oxygen  metal oxide
e.g. 2Mg + O2  2MgO
Tendency to form metal oxides:
• Li, Na, K, Ca, Ba (react at room temp)
• Mg, Al, Fe, Zn (react slowly at room temp, vigorously when heated)
• Sn, Pb, Cu (react slowly and only when heated)
Reactions of metals
Reactions with water
Reactive metals react with water or steam
Metal + water  metal hydroxide + hydrogen gas
e.g. Na + 2H2O  2NaOH + H2
Metal + steam  metal oxide + hydrogen gas
e.g. Zn + H2O  ZnO + H2
Relative reactivity:
• Li, Na, K, Ca, Ba (react with water at room temp)
• Mg, Al, Zn, Fe (react with steam at high temp)
• Sn, Pb, Cu, Ag, Au, Pt (do not react)
Reactions of metals
Reactions with dilute acid
More metals react with acid than water
Metal + acid  salt + hydrogen gas
Zn + 2HCl  ZnCl2 + H2
Relative reactivity:
• Li, Na, K, Ca, Mg, Al, Zn, Fe, Co, Ni (react readily)
• Sn, Pb (slow to react without heat)
• Cu, Ag, Hg, Pt, Au (do not react)
Reactions of metals
Based on the ease of reactions with oxygen, water
and acids, metals can be organised in order of
reactivity, known as an activity series.
Activity series for metals:
K>Na>Ba>Ca>Mg>Al>Zn>Fe>Sn>Pb>Cu>Ag>Hg>Pt>Au
most reactive
least reactive
Grp 1>Grp 2> Grp 3>some TM (Zn, Fe)>Grp 4>more TM
N.B. TM = transition metals
Oxidation - Reduction
The reactions of metals with oxygen, water and acids
involve the metals losing electrons to form +ve metal
ions.
When an atom loses one or more electrons, it is oxidised.
If an atom gains electrons, it is reduced. Therefore:
Oxidation is loss of eReduction is gain of e-
In any equation, there is no overall loss or gain of e-.
Therefore, oxidation and reduction occur
simultaneously and are known as redox reactions.
Oxidation - Reduction
Oxidising agent:
Reducing agent:
• Accepts electrons
• Causes the oxidation
of another substance
• Is always itself
reduced
• Has its oxidation state
decreased
• Donates electrons
• Causes the reduction
of another substance
• Is always itself
oxidised
• Has its oxidation state
increased
Oxidation - Reduction
Oxidation States (some rules)
1.
2.
3.
4.
5.
6.
The oxidation state of a free element (i.e. not part of a compound)
is zero (e.g. Zn(s), O2(g))
The oxidation state of an element in an ionic compound is equal
the electrical charge on its ion. (e.g. Na+ = +1)
Oxidation states of elements in covalent compounds are
calculated as if they were ionic. The most electronegative atom
(closest to F in the periodic table) is assumed to gain electrons.
(e.g. NH3; N = -3, H = +1)
The sum of the oxidation states of all the elements in a compound
is zero.
The oxidation state of oxygen in a compound is normally -2,
except for peroxides, when it is -1. (e.g. H2O2; H = 1, O = -1)
The oxidation state of hydrogen in a compound is normally +1,
except for metal hydrides, when it is -1. (e.g. NaH; Na = 1, H = -1)
Oxidation - Reduction
A metal reacting with acid is an example of a redox reaction. Consider
the following rxn:
Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g)
This reaction can be written as an ionic equation:
Zn(s) + 2H+(aq) + 2Cl-(aq)  Zn2+(aq) + 2Cl-(aq) + H2(g)
Note the two chloride ions that appear on both sides of the equation. These are
known as spectator ions. These can be removed to give us a net ionic
equation:
Zn(s) + 2H+(aq)  Zn2+(aq) + H2(g)
Which of these species has been oxidised? Which has been reduced?
Oxidation - Reduction
Zn(s) + 2H+(aq)  Zn2+(aq) + H2(g)
This net ionic equation can be written as two half reactions:
Oxidation: zinc dissolves and loses electrons
Zn(s)  Zn2+(aq) + 2e- (loss of e-)
Reduction: hydrogen ions gain electrons to form H gas
2H+(aq) + 2e-  H2(g) (gain of e-)
Note that combining these two half reactions results in a
balance of electrons. Try this process using sulfuric acid.
Oxidation - Reduction
Consider the two half reactions for the reaction of
aluminium and a dilute acid:
Al(s)  Al3+(aq) + 3e- (oxidation)
2H+ (aq) + 2e-  H2 (g) (reduction)
Notice that adding these two half reactions results in an
imbalance in the number of electrons. In this case, we
must multiply the first by 2 and the second by 3 to get:
2Al(s) + 6H+(aq)  2Al3+(aq) + 3H2 (g)
Some metals can react with acids
and alkalis
Many metals react with acids. Some also react
with alkalis. Some of these are:
Al, Cr, Zn, Pb, Sn
Some examples:
Zn(s) + 2NaOH(aq)  Na2ZnO2(aq) + H2(g)
(sodium zincate)
2Al(s) + 2NaOH(aq)  2NaAlO2(aq) + 3H2(g)
(sodium aluminate)
Note: the two complex ions formed are ZnO22- (zincate) and AlO2–
(aluminate). Zincate is formed by the combination of Zn2+ and 2O2ions. Aluminate is formed by the combination of Al3+ and 2O2- ions.
Ionisation Energy
Ionisation Energy is the amount of energy required to remove the most loosely bound efrom an atom or ion in the gaseous state.
M(g) + energy  M+(g) + e•
•
1st ionisation energy is removal of 1st e2nd ionisation energy is removal of 2nd e-
Highly reactive metals have low ionisation energies.
Less reactive metals have high ionisation energies.
Periodic table trends:
In general, the closer and more tightly bound an electron is to the nucleus, the higher the ionisation
energy.
Moving left to right  ionisation energy increases
Moving top to bottom  ionisation energy decreases
Question: Why are these trends observed?
Ionisation Energy - trends
Left to right  increase:
This trend is due to the number of
protons increasing, which
leads to a stronger force of
attraction action on the
electrons.
Top to bottom  decrease:
This trend is due to the increase
in the number of electrons in
lower shells shielding the force
of attraction between the
nucleus and the valence
electrons.
Uses of metals based on reactivity
Choosing a metal for a specific purpose often involves the consideration of the
reactivity of the metal. Below are some examples:
• Coatings – Zn is used to make galvanised iron. Iron is
dipped in molten Zn forming a protective layer and
serving as a sacrificial anode (i.e. corrodes first). This
can be a cheaper option to Al building materials (e.g.
roof gutters).
• Jewellery – Au, Ag and Pt are the least reactive of all
metals and therefore retain their lustre.
• Plumbing – Cu is often used in water pipes as it resists
corrosion better than the cheaper alternative, Fe.
• Electrical contacts – Cu (which eventually corrodes) or
more expensive Au contacts.
8.3 Metals
Focus 3: As metals and other elements were
discovered, scientists recognised that
patterns in their physical and chemical
properties could be used to organise the
elements into a Periodic Table
History of the Periodic Table
Aristotle~330 BC
Four element theory: earth, air,
fire & water.
Aristotle classified the elements
on whether they were hot or
cold and whether they were
wet or dry.
• Fire and earth were dry.
• Air and water were wet.
• Fire and air were hot.
• Earth and water were cold.
History of the Periodic Table
Antoine Lavoisier~17701789 (Father of modern
Chemistry)
• Wrote the first extensive
list of elements
containing 33 elements.
• Distinguished between
metals and non-metals.
• Some of Lavoisier's
elements were later
shown to be compounds
and mixtures.
Jöns Jakob Berzelius 1828
• Developed a table of
atomic weights.
• Introduced letters to
symbolize elements.
History of the Periodic Table
Johann Döbereiner -1829
Developed 'triads', groups of 3
elements with similar
properties.
• Lithium, sodium & potassium
formed a triad.
• Calcium, strontium & barium
formed a triad.
• Chlorine, bromine & iodine
formed a triad.
• Sulfur, selenium & tellurium
formed a triad.
Döbereiner was a forerunner to
the notion of groups.
John Newlands -1864
The known elements (>60) were
arranged in order of atomic
weights and he observed
similarities between the first
and ninth elements, the
second and tenth elements
etc.
He proposed the 'Law of Octaves‘
which identified many
similarities amongst the
elements, but also required
similarities where none
existed.
He did not leave spaces for
elements as yet undiscovered.
Forerunner to the notion of
periods.
History of the Periodic Table
Lothar Meyer -1869
Compiled a Periodic Table of 56
elements based on the
periodicity of properties such
as molar volume when
arranged in order of atomic
weight.
He produced graphs to show
the changes in physical
properties as a function of
atomic weights.
Meyer & Mendeleev produced
their Periodic Tables
simultaneously. Mendeleev
was given more credit as he
was able to make accurate
predictions about
undiscovered elements.
Dmitri Mendeleev -1869
Produced a table based on atomic weights
but arranged 'periodically' with
elements with similar properties under
each other.
Gaps were left for elements that were
unknown at that time and their
properties predicted (the elements
were gallium, scandium and
germanium).
The order of elements was re-arranged if
their properties dictated it, eg, tellurium
is heavier than iodine but comes before
it in the Periodic Table.
Mendeleev's Periodic Table was important
because it enabled the properties of
elements to be predicted by means of
the 'periodic law': properties of the
elements vary periodically with their
atomic weights.
History of the Periodic Table
William Ramsay 1894
Discovered the Noble Gases.
In 1894Ramsay removed oxygen,
nitrogen, water and carbon dioxide
from a sample of air and was left with
a gas 19 times heavier than hydrogen,
very unreactive and with an unknown
emission spectrum. He called this gas
Argon.
In 1895 he discovered helium as a decay
product of uranium and matched it to
the emission spectrum of an unknown
element in the sun that was
discovered in 1868. (helios is the
Greek for Sun).
He went on to discover neon, krypton and
xenon, and realised these represented
a new group in the Periodic Table.
Ramsay was awarded a Nobel Prize in
1904.
Henry Moseley 1914
Determined the atomic number of each of the
elements.
He modified the 'Periodic Law' to read that
the properties of the elements vary
periodically with their atomic numbers.
Moseley's modified Periodic Law puts the
elements tellerium and iodine in the right
order, as it does for argon and potassium,
cobalt and nickel.
Glenn Seaborg1940
Synthesised transuranic elements (the
elements after uranium in the periodic
table)
In 1940 uranium was bombarded with neutrons
in a cyclotron to produced neptuniun
(Z=93). Plutonium (Z=94) was produced
from uranium and deuterium. These new
elements were part of a new block of the
Periodic table called Actinides. Seaborg
was awarded a Nobel Prize in 1951.
Trends in the Periodic Table
Electrical Conductivity
Moving left to right across a
period, electrical and
thermal conductivities
tend to decrease as
metallic character
decreases.
Moving down groups,
metallic character tends
to increase as metallic
character increases.
Ionisation Energy
As stated in section
8.2.2, ionisation
energy increases
across a period.
Moving down a group,
ionisation energy
decreases.
Trends in the Periodic Table
Atomic radius
The trend moving left to right in a
period is a decrease in atomic
radius.
This decrease is due to the
additional positive charge
pulling the outermost e-, which
are in the same energy level.
The trend down a group is an
increase due to additional
energy levels.
Melting point/Boiling point
The melting and boiling points
peak in group IV.
Elements in group IV tend to form
strongly bonded covalent
network solids, which have
high mp/bp.
Noble gases have almost no
tendency to form bonds.
Groups I and II undergo metallic
bonding and therefore have
moderate mp/bp.
Group VII have weak
intermolecular forces and low
mp/bp.
Trends in the Periodic Table
Combining power (valency)
This generally refers to the
number of available
bonding sites on an
element.
The trend is a peak in group
IV.
Examples: group I (LiCl);
group II (BeCl2); group IV
(CCl4); group VII (Cl2O)
Electronegativity
This refers to the relative
power of an element to
attract e- to itself or a
drive towards a stable
octet.
The trend is an increase
from left to right and a
decrease from top to
bottom.
F is the most
electronegative element
8.3 Metals
Focus 4: For efficient resource
use, industrial chemical
reactions must use measured
amounts of each reactant
Relative Atomic Mass (atomic weight)
Atoms of particular elements have
a specific mass. Most of this
mass is associated with the
mass of the nucleus.
Since the mass of an atom is a
very small number, it is very
difficult to measure individual
masses.
For this reason, Chemists
determined the relative mass
of atoms. For example, a silver
atom has four times the mass
of a carbon atom. Since they
are relative, they have no
units.
All atomic weights are relative to
the mass of carbon -12 which
is set at exactly 12.0000
Some relative masses (atomic
weights) found on the periodic
table
Carbon 12.0
Aluminium 27.0
Chlorine 35.5
Gold 197.0
Lead 207.2
Silver 107.9
Isotopes
You may notice that many elements do not have
atomic weights that are whole numbers. This is
because most elements have more than one
isotope (different numbers of neutrons) and the
relative atomic weights are weighted averages
of these isotopes.
For example, naturally occurring chlorine is 75% Cl-35 and
25% Cl-37. Therefore:
Avg mass Cl = (75X35) + (25X37) = 35.5
100
Molecular Mass
Molecular mass or weight is the sum of the
atomic weights of the atoms in a molecular
formula.
Example:
The molecular weight of sucrose (table sugar) C12H22011 is
calculated as:
M.W. = (12XAC) + (22XAH) + (11XAO)
M.W. = (12X12.0) + (22X1.01) + (11X16.0)
= 342.2
Formula Mass
Formula mass or weight is the sum of the atomic weights of
the atoms in a compound that has no discreet molecules
(e.g. ionic compounds). These describe the ratios of the
atoms present, but are calculated the same way as
molecular weights.
Example:
The formula weight of calcium phosphate Ca3(PO4)2 is
calculated as:
F.W. = (3XACa) + (2XAP) + (8XAO)
F.W. = (3X40.1) + (2X31.0) + (8X16.0)
= 310.3
The Mole /Avogadro's number
It was eventually determined that a single
C-12 atom has a mass of 1.99X10-22 g.
Therefore, 12g of C-12 contains
6.022x1023 atoms. This is known as
Avogadro’s Number (NA) and is
equivalent to 1 mole of carbon atoms.
In fact, a mole of anything contains
6.022X1023 units.
Avogadro’s number =
6.022x1023
= 1mole
Notice that the mass of 1 mole of C-12 is
the same value as the relative atomic
mass for C-12.
In the same way, 1 mole of any element or
compound is equivalent to its atomic,
molecular or formula weight.
We can now define the relative atomic
masses from the periodic table as
molar masses with the units g/mol.
Substance
Molar mass
(g/mol)
carbon
12
chlorine
oxygen
35.5
16
water
18
Molar mass = mass of 1 mole of
any substance
Calculations using moles
From the previous example using C-12, we now
have a mathematical relationship between mass
and moles, which is:
Mass of substance (g)
Number of moles (n) =
Molar mass (MM) (g/mol)
Example: How many moles are in 25 g of CO?
n CO = 25g CO/(12+16)g/mol
= 0.89 moles
Calculations using moles
We can also convert between moles and the
number of atoms or molecules using Avogadro’s
number.
Number of atoms/molecules = moles (n) x NA
Example: How many atoms are there in a copper pipe that
weighs 2.56g?
n = 2.56/63.6 = 0.0403 moles
number of Cu atoms = 0.0403 X 6.022X1023
= 2.43X1022
Calculations using moles
We can now use the mole concept to determine
how much product to expect in a chemical
reaction. Take the following example:
2Fe2O3(s) + 3C(s)  4Fe(l) + 3CO2(g)
The coefficients in front of each species provide us with useful ratios that
we can use to calculate expected masses of products in a chemical
reaction. We previously said that these were ratios based on the
numbers of atoms. However, with Avogadro’s number, we can now say
that these are molar ratios.
We say, 2 moles of iron (III) oxide react with 3 moles of carbon to produce
4 moles of iron and 3 moles of carbon dioxide gas.
Calculations using moles
2Fe2O3(s) + 3C(s)  4Fe(l) + 3CO2(g)
Example:
How many grams of iron will we expect if we react 12g of
iron (III) oxide as in the above reaction, assuming neither
reactant is in excess?
Convert to moles
12g Fe2O3 X
=
Molar ratio
1 mol Fe2O3
X
159.8g Fe2O3
12 X 1 X 4 X 55.9
159.8 X 2 X 1
Convert to g of
unknown
=
4 mol Fe
2 mol Fe2O3
8.4 g Fe
X
55.9g Fe
1 mol Fe
Avogadro and Gay-Lussac
Gay-Lussac’s law states that gases at equal temperatures and
pressures react in whole number ratios to one another. For example:
2H2(g) + O2(g)  2H2O(g)
2 volumes
1 volume
2 volumes
Notice that the volume ratio is equal to the coefficients in the reaction
and there is no conservation of volumes.
Avogadro’s Law states that equal volumes of gases at the same
temperature and pressure contain the same numbers of molecules.
Now we have a convenient way of determining gas volumes in a
reaction since we can replace gas volumes for moles in a reaction.
e.g. 2 mol of hydrogen gas + 1 mol of oxygen gas  2 mol water gas
8.3 Metals
Focus 5: The relative abundance and
ease of extraction of metals
influences their value and breadth
of use in the community
Minerals and Ores
Minerals
Minerals are naturally occurring compounds found in the
Earth. Most metals (except Au and Ag) are found as
minerals. The most common minerals that contain
metals in Australia are oxides and sulfides.
Ores
Ores are non-renewable mineral deposits that are
economically feasible to extract metals from.
Typical Australian ores
Sphalerite
Chalcopyrite
Haematite
Bauxite
ZnS
CuFe2S2
Fe2O3
Al2O3
Economic considerations of mining
Metal prices can be affected by
many factors including:
Metal
Price
(per tonne Au$)
aluminium
2060
• The abundance of the metal
ore (more abundant metals will
generally be less expensive)
copper
2400
gold
19000000
• The relative cost of production
including the amount of energy
required (costs often passed
on to the consumer)
lead
740
tin
8800
Metal
Energy (kJ/kg)
aluminium
200
copper
70
mild steel
40
• The cost of transporting the
metal or the ore from
sometimes remote locations
• Demand for the metal
Recycling Metals - Al
Why recycle?
1. Less energy is required
to recycle a metal than
is required to extract it
from its ore. (Al
recycling requires
7kJ/kg rather than 200
kJ/kg to extract the ore)
2. Metal ores are nonrenewable natural
resources that need to
be conserved.
3. Less waste to dispose
of in rubbish dumps.
Steps in Al recycling
1. Collection of used
products from homes
and businesses
2. Transport to recycling
facility
3. Separate the Al from
impurities (labels, etc.)
4. Re-smelt the Al into
ingots for transport to
product manufacturers
Extraction of Copper from ore
Copper ore
Copper ore mainly consists of the following
minerals:
•
Chalcocite - Cu2S
•
Chalcopyrite - CuFeS2
•
Malachite - CuCO3,Cu(OH)2
Cu is usually present as 1-3%
Electrolysis
“Blister copper” is made into anodes (+) in an
electrolytic cell (Cu and other metals
oxidised). The cathode (-) is pure Cu.
Solution is CuSO4 and H2SO4 (to inhibit
oxidation of H2O to O2). Cu2+ is reduced to
Cu at the cathode and less reactive metals
fall to the bottom of the cell . (Energy
requirement –HIGH approx. 200kW/tonne)
Froth flotation (concentration of Cu)
Copper ore is mixed with water and a frothing
agent which adheres to the copper allowing
it to float to the surface. The froth containing
approx.15-25% copper is skimmed off the
surface. (Energy requirement – LOW)
Reduction of Cu2S to Cu
Copper sulfide is reduced to copper metal to
produce 98% “blister copper” as there are
SO2 blisters on the copper. The overal
reaction is Cu2S + O2  2Cu + SO2 (Energy
requirement –LOW)
Roasting
Copper ore is heated in a furnace with oxygen to
remove sulphur compounds as SO2 gas.
2CuFeS2 + 3O2  2CuS + 2FeO + 2SO2
(Energy requirement –MEDIUM as the
reaction is highly exothermic)
Iron Removal
Silica and lime are added to produce iron (II)
silicate. FeO + SiO2  FeSiO3 The remaining
product is 50-70% copper. (Energy
requirement –MEDIUM)
Metals
Compiled by: Robert Slider (2006)
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