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8.3 Metals

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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8.3 Metals

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  1. 8.3 Metals Focus 1: Metals have been extracted and used for many thousands of years

  2. 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)

  3. 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

  4. 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)

  5. 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.

  6. 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) Scandinavia (c. 1500-500 BC) Central Europe (c. 1800-500 BC)

  7. 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.

  8. 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 modern-day India.

  9. 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.

  10. 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)

  11. 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: 2CuO(s) + C(s) 2Cu(l) + CO2(g) Iron: 2Fe2O3(s) + 3C(s) 4Fe(l) + 3CO2(g) heat heat

  12. 8.3 Metals Focus 2: Metals differ in their reactivity with other chemicals and this influences their uses

  13. Reactions of metals 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: • 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) heat

  14. 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)

  15. 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)

  16. 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 reactiveleast reactive Grp 1>Grp 2> Grp 3>some TM (Zn, Fe)>Grp 4>more TM N.B. TM = transition metals

  17. 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 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.

  18. Oxidising agent: Accepts electrons Causes the oxidation of another substance Is always itself reduced Has its oxidation state decreased Reducing agent: Donates electrons Causes the reduction of another substance Is always itself oxidised Has its oxidation state increased Oxidation - Reduction

  19. Oxidation - Reduction Oxidation States (some rules) • 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)

  20. 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?

  21. 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.

  22. 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)

  23. 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) 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. (sodium zincate) (sodium aluminate)

  24. Ionisation Energy is the amount of energy required to remove the most loosely bound e- from an atom or ion in the gaseous state. 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

  25. 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.

  26. 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.

  27. 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

  28. 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.

  29. Antoine Lavoisier~1770-1789 (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

  30. 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 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

  31. 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

  32. William Ramsay1894 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 Moseley1914 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 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.

  33. 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

  34. 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

  35. 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 Trends in the Periodic Table

  36. 8.3 Metals Focus 4: For efficient resource use, industrial chemical reactions must use measured amounts of each reactant

  37. 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. 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 Relative Atomic Mass (atomic weight) All atomic weights are relative to the mass of carbon -12 which is set at exactly 12.0000

  38. 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

  39. 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

  40. 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

  41. 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. 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

  42. Mass of substance (g) Molar mass (MM) (g/mol) Number of moles (n) = Example: How many moles are in 25 g of CO? n CO = 25g CO/(12+16)g/mol = 0.89 moles Calculations using moles From the previous example using C-12, we now have a mathematical relationship between mass and moles, which is:

  43. 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

  44. 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.

  45. 12g Fe2O3 X 1 mol Fe2O3 159.8g Fe2O3 4 mol Fe 2 mol Fe2O3 55.9g Fe 1 mol Fe X X 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 g of unknown Convert to moles Molar ratio 12 X 1 X 4 X 55.9 159.8 X 2 X 1 8.4 g Fe = =

  46. 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) Notice that the volume ratio is equal to the coefficients in the reaction and there is no conservation of volumes. 2 volumes 1 volume 2 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

  47. 8.3 Metals Focus 5: The relative abundance and ease of extraction of metals influences their value and breadth of use in the community

  48. 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

  49. Economic considerations of mining Metal prices can be affected by many factors including: • The abundance of the metal ore (more abundant metals will generally be less expensive) • The relative cost of production including the amount of energy required (costs often passed on to the consumer) • The cost of transporting the metal or the ore from sometimes remote locations • Demand for the metal

  50. Why recycle? 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 manufacturers Recycling Metals - Al

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