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Group 1 The Alkali metals Li, Na, K, Rb, Cs, Fr

Group 1 The Alkali metals Li, Na, K, Rb, Cs, Fr. Electronic configuration One electron in the outermost shell in an s orbital. This electron is very easily lost and they are good reducing agents. M M + + e -

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Group 1 The Alkali metals Li, Na, K, Rb, Cs, Fr

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  1. Group 1 The Alkali metalsLi, Na, K, Rb, Cs, Fr • Electronic configuration One electron in the outermost shell in an s orbital. This electron is very easily lost and they are good reducing agents. M M+ + e- Thus the chemistry is dominated by the metals in the +1 state. The M2+ ions are not known as this would involve breaking into an inner core.

  2. Properties and reactions 1 They are soft metals that can be easily cut with a pen-knife and rapidly tarnish in air. 2 Reactions with oxygen. They all burn in oxygen and the products depend upon the alkali metals. 4Li(s) + O2(g) 2Li2O(s) normal oxide 2Na(s) + O2(g) Na2O2(s) peroxide K(s) + O2(g) KO2(s) superoxide • They react rapidly with water. The vigour of the reaction increases as you go down the group. 2M(s) + 2H2O(l) 2MOH(aq) + H2(g) Note all group 1 hydroxides are very soluble in water.

  3. Properties and reactions (continued) 4 The metals and their compounds give characteristic colours to flames. When a platinum wire is dipped into a solution of these compounds, and heated in a Bunsen flame a characteristic colour is given to the flame. (This was why Bunsen developed his Bunsen burner so that he and his colleague Kirchoff could identity new elements using this technique). Alkali metal Colour of the Flame Li red Na yellow K lilac Rb red Cs blue

  4. Properties and reactions (continued) • All the carbonates are soluble in water. All the carbonates are thermally stable except for lithium carbonate which decomposes on heating. • All the hydrogencarbonates (except for LiHCO3) can be isolated as solids They all decompose on heating e.g. 2NaHCO3(s) → Na2CO3(s) + H2O(l) + CO2(g) • The nitrates all decompose on heating to form the nitrite except lithium which forms the oxide 2MNO3(s) → MNO2(s) + O2(g) LiNO3(s) → Li2O(s) + NO2(g) + ½O2(g)

  5. Properties and reactions (continued) • 8 The salts are usually white or colourless ionic solids which just dissolve in water i.e. they are not hydrolysed e.g. NaCl(s) → Na+(aq) + Cl-(aq) H2O

  6. Reactivity Thermal Solubility of with air Stability of carbonates and H2O [CO32-] OH- [NO3-] Li decomp decomp decomp to Li2O soluble Na more stable stable decomp and K and and to increasingly Rb vigorous increa- increa- MNO2 ing sing soluble Cs stablity stabilty

  7. Group 2 The alkaline EarthsBe, Mg, Ca, Sr, Ba, Ra • Chemistry (except for Be) is dominated by the tendency to lose 2 electrons and form M2+ cations. And thus compounds (except for those of Be which tend to be covalent) tend to be ionic. M → M2+ + 2e- They are harder than the elements of group 1 (two electrons involved in metallic bonding as opposed to one) and tend to be less reactive. (it takes more energy to remove 2 electrons as opposed to one electron).

  8. Reactions and properties • They burn in air and form the normal oxides. 2M(s) + O2(g) → 2MO(s) • Ca, Sr and Ba react with H2O to form M(OH)2 and H2 2M(s) + H2O(l) → 2 MO(s) Mg reacts only slowly with steam to form the oxide. Mg(s) + H2O(g) → 2MO(s) + H2(g) • The metals react with aqueous acids to liberate H2 M(s) + 2 HCl(aq) → MgCl2(aq) + H2(g) • Some of the group 2 metals give characteristic flame tests Ca brick red Sr scarlet Ba apple green

  9. Reactions and properties (continued) • The carbonates decompose with strong heating to give the oxides. MCO3(s)→ MO(s)+CO2(g) • The nitrates decompose to the oxides, NO2 and O2 2M(NO3)2→2MO(s) + 4NO2(g) + O2(g) • The hydroxides are not as thermally stable or as soluble as those of group 1 (BeOH)2 and Mg(OH 2 are not soluble in water) 2M(OH)2(s) →MO(s) + H2O(g) • Solubility of sulphates and carbonates decrease down the group • Basic strength of oxides and hydroxides increases down the group.

  10. The Beryllium Anomaly • Be has properties which make it anomalous to the other members of the groups. These include 1 The oxide is amphoteric 2 The Chloride, bromide and iodide are covalent compounds

  11. Summary of group 2 Reactivity Basic with air Solubility Solubility strength Solubility water and [SO42-] [CO32-] OH- of OH- acidsMetalBe poor amphoteric insolubleMg soluble weak insolubleCa increases insolubleSr increasingly increasinglyincreasinglyincreasinglyBa insolubleinsolublestrong soluble

  12. Group 13 The Boron Group • This group contains both non-metals and metals. Boron is the only non-metal, whilst the rest are metals. • Reactions of the elements (E) • Oxygen 4E(s) + 3O2(g) → 2E2O3(s) E = B,Al,Ga,In.Tl • Nitrogen 2E(s) + N2(g) → 2EN(s) E = B, Al • Halogen X2 2B(s) + 3X2(g.l.s) → 2BX3(g,l) 2E(s) + X2(g,l,s) → E2X6(s) E = Al,Ga,N 2Tl + X2(g,l,s) → 2TlX(s)

  13. Acid 2E(s) + 6H3O+(aq) → 2E3+(aq) + 6H2O(l) + 3 H2(g) E = Al, Ga, Tl Base 2E(s) + 6 H2O(l) + 2 OH-(aq) → 2E(OH)4-(aq) + 3 H2(g) E = Al, Ga

  14. Reactions and properties of compounds (Boron and Aluminium only) • Oxides • Boron Oxide B2O3 reacts with water to form Boric acid B(OH)3 which is a weak acid. B2O3(s) + 3H2O(l) → 2B(OH)3 • Aluminium oxide (alumina) is an amphoteric oxide and reacts with both acid and base. Al2O3(s) + 2OH-(aq) + 3 H2O(l) → 2Al(OH)4-(aq) • Al2O3(s) + 6H3O+(aq) + 3 H2O(l → 2[Al(H2O)6]3+(aq)

  15. Halides Boron The boron halides are trigonal planar molecules as predicted by VSEPR. Because of the empty p orbital on the boron atom they are all Lewis Acids. • Aluminium These are all ionic solids BUT the chloride, bromide and iodide sublime as the dimers with the formula Al2X6 (X=Cl,Br,I) These dimers are also found in non-coordinating organic solvents. The halides can acts as Lewis acids. • Cl Cl Cl Al Al Cl Cl Cl

  16. Group 14 The Carbon Group • This group has the widest range of elements within it. These range from the non-metal carbon, the metalloids, silicon and germanium, to the metals tin and lead. The electron structure of the valence electrons is ns2 np2. Since these elements are 4 away from an octet (either by losing or sharing) we would normally expect the oxidation state in these compounds to be +4 (remember oxidation state does not imply anything about a compound being ionic or not). But the +2 state is also known, this involves just losing the p electrons ns2np2→ [ns2]2+ + 2e- This lower oxidation state becomes more stable as you go down the group and is a consequence of the inert pair effect.

  17. Group 14 The Carbon Group continued • Carbon • The chemistry of carbon forms a subject in its own right. The great number of carbon compounds is the result of carbon forming strong covalent bonds with itself. This ability of an element to form bonds to itself is called is called catenation. The extensive chemistry of carbon is also due in part because carbon can form strong stable bonds with other elements particularly hydrogen. • Carbon is now known to exist in several allotropes the most well known being diamond, graphite and C60.

  18. Giant Structures • In diamond each carbon atom is covalently bonded to 4 other carbon atoms. Since covalent bonds are (usually) very difficult to break this makes diamond very hard with a high melting point. Silicon and Germanium the next two elements in the group also have the same structure as diamond, whilst at room temperature Tin and lead have metallic structures.

  19. The reactions of the elements • Hydrogen C(s) + 2H2(g) → CH4(g) and other hydrocarbons • Oxygen E(s) + O2(g) → EO2(g,s) E= C, Si, Ge Sn 2Pb(s) + O2 (g) → 2PbO(s) • Halogen E(s) + 2X2(g,l,s) → EX4(g,l,s) E=C,Si,Ge,Sn Pb(s) + X2(g,l,s) → PbX2(s) • Acid E(s) + 2H3O+(aq) ) → E2+(aq) +2H2O(l)+H2(g) E = Sn , Pb • Base E(s) + 2H2O(l) + 2OH-(aq)→ E(OH)42- + H2(g) E = Sn, Pb

  20. Properties of compounds • Oxides • Carbon • Carbon dioxide is very stable, it dissolves in water to give carbonic acid H2CO3 which is a weak acid. Carbon dioxide is one of the gases most associated with global warming. • Carbon monoxide is a highly toxic gas • Silicon • Silicon if found as its dioxide quartz. It is slightly acidic

  21. Properties of compounds oxides continued • Germanium, Tin and Lead are amphoteric e.g SnO2 (s) + 2H2SO4(aq) → Sn(SO4)2(s)+ 2H2O(l) SnO2(s) + 2H2O(l) + 2OH-(aq) → Sn(OH)62- (aq) PbO(s) + 2HCl(aq) → PbCl2(s) + H2O(l) PbO(s) + H2O + 2OH- →Pb(OH)42--

  22. Properties of Halides • Halides of carbon CX4 are stable to hydrolysis but the halides of the rest undergo hydrolysis e.g. SiCl4(l)+ 2H2O(l) → 4HCl(aq) + SnO2xH2O(s) • The lower halides become more stable as you go down the group so PbX2 is found for the lead halides.

  23. The hydrides • Element Hydrides • C Alkanes, Alkenes, Alkynes Arenes • Si Silanes SinH2n+2; n = 1- 10 • Ge Germanes GenH2n+2; n = 1- 6 • Sn Stannane SnH4 only • Pb Plumbane PbH4 only • Stability decreases as you go down the group because the bonds get weaker. Also not the lesser inability of the heavier members in the group to catenate.

  24. The greenhouse effect • (Taken from Ramsden without permission)

  25. The greenhouse effect • The temperature of the earth is a balance between the heat absorbed from the sun and the heat radiated back into space. This is in a steady state or equilibrium if everything was working correctly The radiation from the sun that reaches the earth ranging from UV 200nm to IR 3000nm , with the maximum being in the visible spectrum at 500 nm This radiation passes through the atmosphere and reaches the earth. Here is re-emitted but at a longer wavelength i.e. in the infrared part of the spectrum. The green house effect is the absorption of this radiation by gasses in the atmosphere such as CO2 and H2O. This trapped radiation cause the temperature of the atmosphere to rise.

  26. What traps where in the greenhouse effect • Taken from Ramsden without permission

  27. Greenhouse effect continued. • Water vapour does most of the tapping BUT if we look at the range of reemitted radiation we can see that there is a gap in it, in which the water vapour does not trap the radiation. This gap is partially plugged by CO2, so increasing CO2 will reduce the amount of reemitted radiated heat lost. • Note that if there was enough CO2 in the atmosphere to absorbed all the heat re-radiated from the earth increasing the CO2 concentration would have no further effect, but what would have happened is that the temperature would have reached a new steady state but at a higher temperature.

  28. The importance of the green house effect • The greenhouse effect is vital for life. If it did not occur the average temperature of the atmosphere would be about -20 ºC and not 15 ºC.

  29. History of Greenhouse effect. • The green house effect and global warming was recognised as long ago as 1890 by Arrhenius who predicted that doubling the CO2 concentration would cause global temperatures to rise by 5 C. This value is in line with some recent values put forward by the UK Met. Office.

  30. Other contributing factors • Other greenhouse gases include CH4, and Chlorofluorocarbons CFC’s These can have a particularly harmful effect if they absorb in the IR region which is not covered by either H2O or CO2 • Deforestation also cause CO2 level to rise in the atmosphere as the trees that have been removed are no longer able to trap CO2 during photosynthesis.

  31. Group 15 The pnictides N, P, As, Sb Bi • This groups contains elements that range from the non-metals nitrogen, and phosphorous, through arsenic and antimony to bismuth which is predominantly metallic in character. The non-metals N and P can form the 3- anion but only with some of the more electropositive metal e.g. Ca but is more usual for them to form, covalent bonds instead. The heavier less electronegative elements can often be found in the +3 oxidation state where the s electrons in the valence shell are not used in bonding.

  32. The elements • The structures of the elements show the transition from non-metals through metalloid to metal. Nitrogen exists as very stable diatomic N2 molecules. There are several allotropes of phosphorous, the most reactive being white phosphorous which consists of P4 molecules. Red phosphorous which used in matches is more stable. Arsenic (As) and Antimony (Sb) have sheet structures and bismuth is three dimensional metal. The changes in structure are reflected in the meting points of the elements eg. Since there are only weak van der waals forces between N2 molecules it is a gas at room tempearture.

  33. Reactions of the elements • Hydrogen N2(g) + 3H2(g) → 2NH3(g) P4(s)+6H2(g) → 4PH3(g) • Oxygen N2(g) + xO2→ NOx(g) P4(s)+ 3 or 5 O2(g) → P4O6(s) or P4O10(s) 4As(s) + 3O2(g) → As4O6(s) 4E(s) + 3O2(g) → 2E2O3 E = Sb, Bi • Water no reaction • Halogen 2E(s) + 3X2(s,l,g) → 2EX3(s,l) E = P, As, Sb, Bi 2 E(s) + 5 X2(s,l,g) → 2EX5(s,l) P, As, Sb

  34. The reactions of the elements show the non-metallic character of the earlier elements, the increasing stability of the lower oxidation states down the group and the lower reactivity of the elements down the group.

  35. Properties of compounds. • Hydrides • NH3 isa stable molecules and is manufactured in large quantities by the Haber processes, and is used as feedstock in the manufacture of many chemicals, e.g. fertilizers, explosives and plastics. It acts as a base in aqueous solution with the lone on the nitrogen “picking” up a proton from a water molecules • The other hydrides in the group are much less table. PH3 readily inflames in air if it is not absolutely pure. It is a much weaker base than NH3.

  36. Properties of compounds Continued. • Halides The most important halides in this group are those of phosphorous. They behave in a manner which is typical of many non-metals undergoing a hydrolysis reaction to form oxoacids • PCl3(l) + 3H2O(l) → H3PO3(l) + 3HCl(g) PCl5(s) + 4H2O(l) → H3PO4(l) + 5HCl(g) • Nitrides and phosphides. The Metallic nitrides and phosphides are strong bases in water reacting to give basic solution as the nitride or phosphide ion readily removes protons from water molecules e.g. Ca3N2(s) + 6H2O(l) → 3Ca(OH)2(aq) + 2NH3(aq)

  37. Properties of compounds Continued. • Oxides Oxidation State Formula Oxoacid 5 N2O5 HNO3 4 NO2 N2O4 3 N2O3 HNO2 2 NO 1 N2O The oxides of nitrogen are given the general name of NOx Several of them and the salts of their oxo acids are are important in both atmospheric and terrestrial pollution.

  38. Atmospheric Pollution • NO NO can be produced by from N2 and O2 in hot aeroplane and car engines. NO contributes to acid rain, the formation of smog and the destruction of the ozone layer. • NO2 NO2 is a choking poisonous gas it exists in an equilibrium with its dimer N2O4. Which gives rise to the characteristic brown colour of smog. N2O4 2NO2 Colourless Brown When it dissolves in water it disproportionates to give HNO3 and NO thus also adding the problem of acid rain. 3NO2(g) + H2O(l) → 2HNO3(aq) + NO(g)

  39. Oxides continued • N2O3 N2O3 dissolves to give HNO2 (nitrous acid) This acid can not be isolated but its salts the nitrites find widespread use despite they fact that they are mildly toxic and have been implicated as carcinogens in food preservation where they inhibit meat turning brown. • N2O5 is the anhydride of HNO3 and dissolves in water to give HNO3 which is a strong acid. N2O5(l) + H2O(l) → 2HNO3(aq) • P4O6 and P4O10 These are typical acidic main group oxides which dissolve in water to produce the oxo acids P4O6(s) + 6H2O(l) → 4H3PO3(aq) P4O10(s)+ 6H2O(l) → 4H3PO4(aq) These oxo acids are weak acids

  40. Pollution by the salts of the oxo acids Too much of a good thing • The salts of the oxo acids of both nitrogen and phosphorous are extensively used used in fertilizers where they are used to supply soluble nitrogen and phosphorous which are essential for plant growth, However those salts which remain used in fields can easily be dissolved, transported by water into streams rivers, and lakes. The nitrogen salts are mildly toxic but their biggest impact to pollution occurs through the proceeds of eutrophication This is when aquatic plant life becomes too well nourished, this leads to algal blooms, This increase in plant life leads to the lakes becoming depleted in oxygen as this is used up in the respiration of the plants This loss of oxygen leads to the death of fish.

  41. What are the sources of eutrophication. • 1 Natural run-off nutrients form soil caused by weathering • 2 Run-off of inorganic fertilisers (Nitrate and phosphates) • 3 Run-off Manure (containing phosphates, nitrates urea, and ammonia) • 4 discharge of phosphates used in detergents as water softeners • 5 discharges of raw or partially treated sewage containing nitrates and phosphates

  42. Group 16 The chalcogens O, S,Se, Te, Po • Introduction Most of the elements in this group are non metals. Since they are two electrons short of reaching an inert gas configuration they can do this by either gaining two electrons to form the anion with a charge of 2- or sharing two electrons in a covalent molecule. Oxygen is used widely mainly in steel production. Sulphur is also used widely in industrial processes. It is obtained either as a by-product form the extraction of metals, or by mining using the Frasch process. • Oxygen exits in two allotropic forms dioxygen O2 or ozone O3 There are several allotropes of sulphur the two main one are rhombic and monoclinic sulphur which consist of S8 molecules arranged in a ring. Several allotropes are known for Se and Te the most stable of which consist of long chains. Po has a metallic structure

  43. Compounds of group 16 elements • Hydrides. • Oxygen. There are two hydrides of oxygen. Dihyrogen oxide (water), whose properties we have extensively encountered before and hydrogen peroxide H2O2, in which there is an O-O single bond. • The other hydrides of the elements in the group are all toxic gasses with foul smells.

  44. Oxides and sulphides • The oxides and sulphides of metal and non-metals are known. In general the oxides of metals are basic, metalloids amphoteric and non-metals are acidic. • The most important compounds which we will look at are the oxides of sulphur. • SO2 is formed when Sulphur burns in air S(s) + O2(g) → SO2(g) • It dissolves in water to form sulphurous acid H2SO3 • This is a weak acid forming the hydrogensulphite ion HSO3- ions in solution. H2SO3(aq) + H2O(l) HSO3-(aq)

  45. Sulphur trioxide SO3 SO3 is formed by the oxidation of SO2 2SO2(g) + O2(g) → 2SO3(g) This reaction is slow and SO2 in the atmosphere can survive for several days. The reaction can be catalysed in heavily polluted air where tiny particles of metal, are present. Also ozone, hydrogen peroxide and ammonia can catalyse the reaction. Industrially SO3 is prepared by the contact process in which SO2 is oxidised over a V2O5 catalyst. SO3can dissolve in water to produce sulphuric acid which is a strong acid. Both SO2 and SO3 are responsible in part for acid rain.

  46. Acid rain • Rain water is naturally acidic because of carbon dioxide in the atmosphere which form H2CO3. Without this natural acidity weathering of geological features such as limestone would be immeasurably slower. This natural rainwater has a pH of about 5.6 however in rain in central Europe the pH is 4.1 This acid rain can not only cause disfigurement of building made but attacks the environment. • Part of the problem is that when acid rain containing H2SO4 reaches the ground it can make soluble some elements in the ground e.g. aluminium salts can be solublise as aluminium sulphate

  47. Acid rain • (Taken from Ramsden without permission)

  48. Element Effect solubilsed • Al Aluminium in lakes kills fish by clogging gills • poisoning of trees • Ca, Mg Deforestation by removing essential nutrients for plants

  49. Sources of SO2 • SO2 occurs naturally in the atmosphere from volcanic eruptions. , sea spray, and rotting vegetation. However the biggest source of SO2 is the combustion of fossil fuels , with power station accounting for the largest component of this.

  50. Group 17 The halogens F, Cl, Br, I, Rn • Introduction The halogens have seven electrons in their outer shell and so react by either gaining one electron to form ionic compounds in which they have a closed shell, or by sharing one electron in a covalent compound. • Properties. • They are non-metals and are found as diatomic molecules X2 • The molecules are volatile. At room temperature F2,Cl2 are gases, Br2 is a liquid, and I2 sublimes upon heating. This is because there are only weak van der waals forces between the molecules

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