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!
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Metals have been extracted and used for many thousands of years
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):
Great Britain Bronze Age (c. 2100-to 700 BC)
China – Xia Dynasty (c. 2100-700 BC)
(c. 1500-500 BC)
Central Europe (c. 1800-500 BC)
Map of early Iron Age Vedic India. This map shows the North-western portion of modern-day India.
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).
2CuO(s) + C(s) 2Cu(l) + CO2(g)
2Fe2O3(s) + 3C(s) 4Fe(l) + 3CO2(g)
Focus 2: Metals differ in their reactivity with other chemicals and this influences their uses
Reactions with oxygen (combustion)
All metals form oxides except Ag, Au and Pt
Metal + oxygen metal oxide
e.g. 2Mg + O2 2MgO
Tendency to form metal oxides:
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
Reactions with dilute acid
More metals react with acid than water
Metal + acid salt + hydrogen gas
Zn + 2HCl ZnCl2 + H2
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:
most reactiveleast reactive
Grp 1>Grp 2> Grp 3>some TM (Zn, Fe)>Grp 4>more TM
N.B. TM = transition metals
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 e-
Reduction 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 States (some rules)
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?
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.
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)
Many metals react with acids. Some also react with alkalis. Some of these are:
Al, Cr, Zn, Pb, Sn
Zn(s) + 2NaOH(aq) Na2ZnO2(aq) + H2(g)
2Al(s) + 2NaOH(aq) 2NaAlO2(aq) + 3H2(g)
Note: the two complex ions formed are ZnO22- (zincate) and AlO2– (aluminate). Zincate is formed by the combination of Zn2+ and 2O2- ions. Aluminate is formed by the combination of Al3+ and 2O2- ions.
M(g) + energy M+(g) + e-
1st ionisation energy is removal of 1st e-
2nd 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
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.
Choosing a metal for a specific purpose often involves the consideration of the reactivity of the metal. Below are some examples:
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
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.
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
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 aforerunner 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
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
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.
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.History of the Periodic Table
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.
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.
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
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
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)
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 elementTrends in the Periodic Table
Focus 4: For efficient resource use, industrial chemical reactions must use measured amounts of each reactant
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.
Some relative masses (atomic weights) found on the periodic table
Silver 107.9Relative Atomic Mass (atomic weight)
All atomic weights are relative to the mass of carbon -12 which is set at exactly 12.0000
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
Molecular mass or weight is the sum of the atomic weights of the atoms in a molecular formula.
The molecular weight of sucrose (table sugar) C12H22011 is calculated as:
M.W. = (12XAC) + (22XAH) + (11XAO)
M.W. = (12X12.0) + (22X1.01) + (11X16.0)
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.
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)
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.
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.
Avogadro’s number = 6.022x1023= 1mole
Molar mass = mass of 1 mole of any substance
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
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.
1 mol Fe2O3
4 mol Fe
2 mol Fe2O3
1 mol Fe
XCalculations using moles
2Fe2O3(s) + 3C(s) 4Fe(l) + 3CO2(g)
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 g of unknown
Convert to moles
12 X 1 X 4 X 55.9
159.8 X 2 X 1
8.4 g Fe
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)
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
Focus 5: The relative abundance and ease of extraction of metals influences their value and breadth of use in the community
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 are non-renewable mineral deposits that are economically feasible to extract metals from.
Typical Australian ores
Metal prices can be affected by many factors including:
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)
Metal ores are non-renewable natural resources that need to be conserved.
Less waste to dispose of in rubbish dumps.
Steps in Al recycling
Collection of used products from homes and businesses
Transport to recycling facility
Separate the Al from impurities (labels, etc.)
Re-smelt the Al into ingots for transport to product manufacturersRecycling Metals - Al
Copper ore mainly consists of the following minerals:
Chalcocite - Cu2S
Chalcopyrite - CuFeS2
Malachite - CuCO3,Cu(OH)2
Cu is usually present as 1-3%
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)Extraction of Copper from ore
“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)
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)
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)
Silica and lime are added to produce iron (II) silicate.FeO + SiO2 FeSiO3The remaining product is 50-70% copper. (Energy requirement –MEDIUM)
Compiled by: Robert Slider (2006)
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