2.7 The periodic table groups 2 and 7

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1 2.7 The periodic table groups 2 and 7 Students will be assessed on their ability to: 1 Properties down group 2 a. explain the trend in the first ionization energy down group 2 b. recall the reaction of the elements in group 2 with oxygen, chlorine and water c. recall the reactions of the oxides of group 2 elements with water and dilute acid, and their hydroxides with dilute acid d. recall the trends in solubility of the hydroxides and sulfates of group 2 elements e. recall the trends in thermal stability of the nitrates and the carbonates of the elements in groups 1 and 2 and explain these in terms of size and charge of the cations involved f. recall the characteristic flame colours formed by group 1 and 2 compounds and explain their origin in terms of electron transitions g. describe and carry out the following: i. experiments to study the thermal decomposition of group 1 and 2 nitrates and carbonates ii. flame tests on compounds of group 1 and 2 iii. simple acid-base titrations using a range of indicators, acids and alkalis, to calculate solution concentrations in g dm -3 and mol dm -3, eg measuring the residual alkali present after skinning fruit with potassium hydroxide h. demonstrate an understanding of how to minimise the sources of measurement uncertainty in volumetric analysis and estimate the overall uncertainty in the calculated result.

2 a. explain the trend in the first ionization energy down group 2

3 Explanation of this trend Definition: The first ionisation energy is the enthalpy change when one mole of gaseous atoms loses one mole of electrons to form one mole of gaseous mono-positive ions: Going down Group 2: Nuclear charge increases. The radius of the atom increases, so the distance between the nucleus and the outer electron increases. There are more filled energy levels between the nucleus and the outer electron these shield the outer electron from the attraction of the nucleus. The first factor is not as important as the other two, therefore the force of attraction between the nucleus and outer electron is reduced, and less energy is needed to remove the outer electron.

4 With regard to successive IEs, the big jump here is between IE(2) and IE(3). Why? IEs (kjmol-1) With regard to successive IEs, the big jump here is between IE(2) and IE(3). Be 1st nd rd Why? Mg Ca Sr Ba

5 Reaction of Groups 2 #Except Be, whose oxide is amphoteric *The energy required to overcome IEs is recovered in the EAs and the lattice enthalpies. The compounds formed are almost completely ionic, except with Be. The ions formed correspond to the +2 oxidation state, simply because of energetics more stable compounds are formed with a larger net energy release*. With O2 Group 2 (except Ba) react to form the expected simple, basic oxide.# The colour of the flame on combustion is dealt with later. 2M (s) + O2 (g) 2MO (s) White solid b. recall the reaction of the elements in group 2 with oxygen, chlorine and water

6 Ba 2Ba (s) + O2 (g) solid 2BaO (s) White Ba (s) + O2 (g) BaO2 (s) White solid Both the simple oxide and the peroxide are formed.

7 REACTIONS OF THE GROUP 2 ELEMENTS WITH WATER Beryllium has no reaction with water or steam even at red heat. Magnesium burns in steam to produce magnesium oxide and hydrogen. Very clean magnesium has a very slight reaction with cold water. The reaction soon stops because the magnesium hydroxide formed is almost insoluble in water and forms a barrier on the magnesium preventing further reaction. Note: As a general rule, if a metal reacts with cold water, you get the metal hydroxide. If it reacts with steam, the metal oxide is formed. This is because the metal hydroxides thermally decompose (split up on heating) to give the oxide and water.

8 Calcium, strontium and barium These all react with cold water with increasing vigour to give the metal hydroxide and hydrogen. Strontium and barium have reactivities similar to lithium in Group 1 of the Periodic Table. The equation for the reactions of any of these metals would be: Summary of the trend in reactivity The reactions become easier as the energy needed to form positive ions falls. This is mainly due to a decrease in ionisation energy as you go down the Group. This leads to lower activation energies, and therefore faster reactions.

http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/021_sodiumchlor.

9 Reactions with Cl2 All react on heating to form predominantly ionic chlorides (except Be): Ca(s) + Cl2(g) CaCl2(s) Na(s) + Cl2(g) NaCl(s) MgCl2 shows some covalent character, due to the high charge density of the Mg2+ ion. BeCl2 is covalent, and forms a solid polymer via dative covalent bonds to give the Be an octet of electrons:

10 Solubility of Group 2 Sulphates and Hydroxides Sulphates (and carbonates) become less soluble as you go down the Group; hydroxides become more soluble. Be Mg Ca Sr Ba Sulphates More soluble Hydroxides Easy to remember, as you know that BaSO4 is a white ppt. used to identify sulphates, NB1. The ph of magnesium hydroxide NB2. Limewater d. recall the trends in solubility of the hydroxides and sulfates of group 2 elements

11 Thermal stability of the s-block carbonates and nitrates You will need to know: - the trends - the reactions - the explanation In all cases, for Gp1 or Gp2 carbonates and nitrates, the thermal stability increases down the group as the ionic radius of the cation increases, and so its polarising power decreases. For the same reason, the thermal stability decreases from Gp1 to Gp2 across a period. The carbonates of Li, and the Gp2 metals decompose according to the general equation: MCO3 MO (s) + CO2 (g) You try the Lithium eqn. The other Gp1 metal carbonates (and barium carbonate) do not decompose at bunsen temperatures. e. recall the trends in thermal stability of the nitrates and the carbonates of the elements in groups 1 and 2 and explain these in terms of size and charge of the cations involved

12 The same pattern is observed for the s-block nitrates. The nitrates of Li, and the Gp2 metals decompose according to the general equation: 2M(NO3)2 2MO (s) + 4NO2 (g) + O2(g) Brown gas is evolved Eg Again, you try the Lithium eqn. The other Gp1 metal nitrates decompose to form the nitrites: Eg 2MNO3(s) 2MNO2(s) + O2(g) The nitrates are white solids, the nitrites yellow.

13 Explanation of the thermal stability trends One view is that thermal stability increases as polarising power of the cation decreases. Basically, large polarisable anions (CO3 2-, NO3 - ) will be more stable with nonpolarising cations, but small cations with a large polarising power (high charge density) will pull the oxygen in the anion towards them, thus favouring the decomposition of the anion to form the oxide.

14 The flame test The distinctive colours that appear when we heat metals or their compounds can be used to identify them. You need to learn them metal barium calcium strontium copper potassium sodium lithium flame colour light (apple) green brick red crimson red blue/green lilac bright orange carmine red f. recall the characteristic flame colours formed by group 1 and 2 compounds and explain their origin in terms of electron transitions

15 The energy from the bunsen flame causes electrons to jump up to higher energy levels. When they return to their original ground state, they release a certain characteristic amount of energy.

http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/039_flametestsmet.mov Observation can be qualitative or via a spectrometer.

17 Observation can be qualitative or via a spectrometer. The latter method gives a series of lines of frequencies that correspond to the differences in the energy levels for a particular atom (ion). See p180. These are called line spectra. Sodium and sodium compounds emit almost monochromatic light used for distinguishing optical isomers. Other uses: Analysing body fluids, astronomy. Expt: Flame tests P Questions

19 g. describe and carry out the following: i. experiments to study the thermal decomposition of group 1 and 2 nitrates and carbonates ii. flame tests on compounds of group 1 and 2 iii. simple acid-base titrations using a range of indicators, acids and alkalis, to calculate solution concentrations in g dm -3 and mol dm -3, eg measuring the residual alkali present after skinning fruit with potassium hydroxide h. demonstrate an understanding of how to minimise the sources of measurement uncertainty in volumetric analysis and estimate the overall uncertainty in the calculated result.

Atomic and ionic radii The only intermolecular forces are weak dispersion forces.

20 Group 7: The Halogens Physical properties Not surprisingly, with the addition of an extra shell, the atomic radii increase down the group. The ionic radii are much larger than the atomic radii due to the increased repulsion forces of the extra electron in the now complete valence shell. Atomic and ionic radii The only intermolecular forces are weak dispersion forces. These increase as the halogen molecules increase in size and the greater number of electrons make the momentary dipoles more significant. What do the halogens look like? What do they look like dissolved in water and organic solvents?

21 Inorganic chemistry of group 7 (limited to chlorine, bromine and iodine) a. recall the characteristic physical properties of the elements limited to the appearance of solutions of the elements in water and hydrocarbon solvents b. describe and carry out the following chemical reactions of halogens: i. oxidation reactions with metal and non-metallic elements and ions such as iron(ii) and iron(iii) ions in solution ii. disproportionation reactions with cold and hot alkali, eg hot potassium hydroxide with iodine to produce potassium iodate(v) c. carry out an iodine/thiosulfate titration, including calculation of the results and evaluation of the procedures involved, eg determination of the purity of potassium iodate(v) by liberation of iodine and titration with standard sodium thiosulfate solution d. describe and carry out the following reactions: i. potassium halides with concentrated sulfuric acid, halogens and silver nitrate solution ii. silver halides with sunlight and their solubilities in aqueous ammonia solution iii. hydrogen halides with ammonia and with water (to produce acids) e. make predictions about fluorine and astatine and their compounds based on the trends in the physical and chemical properties of halogens.

22 Ionisation Energy Look at your table of the halogens. Sketch a graph of the first ionisation energies. Explain the pattern. On your graph, predict where the graphs for IE(2) and IE(3) would be. On your graph, predict where the graph for Group 0 would be.

23 Tests for Cl 2, Br 2 and I 2 *used to extract bromine from sea water. Cl2 tests make use of the oxidising power of chlorine Bleaches damp litmus paper (damp blue litmus paper momentarily turns red before white). Oxidises Br- and I- to Br2* and I2 respectively. Write the observations and the colour changes for the latter of these tests. More distinctive colours are seen if an organic solvent is added.

Br2 Bleaches damp litmus paper, but slowly.

solvents (in the absence of O2) Extra: Amylose in starch is responsible for the formation of a deep blue

Iodine - KI Reagent: Iodine is not very soluble in water, therefore the iodine reagent is made by dissolving

24 Br2 Bleaches damp litmus paper, but slowly. Oxidises I - (aq) I2 No change with litmus paper Blue/black colour with starch solution Purple in organic solvents (in the absence of O2) Extra: Amylose in starch is responsible for the formation of a deep blue color in the presence of iodine. The iodine molecule slips inside of the amylose coil. Iodine - KI Reagent: Iodine is not very soluble in water, therefore the iodine reagent is made by dissolving iodine in water in the presence of potassium iodide. This makes a linear triiodide ion complex which is soluble. The triiodide ion ion slips into the coil of the starch causing an intense blue-black colour.

H-X The hydrogen halides all are c-less gases.

25 H-X The hydrogen halides all are c-less gases. give misty fumes in moist air (why?) are very water soluble (large enthalpy of hydration of the ions compensates for the energy required to break the H-X bond). form strong acids with water*, stronger as the H-X bond becomes weaker (Hydroiodic acid, however, is prone to oxidation). HX(g) + H2O(l) H3O+(aq) + X-(aq) *except H-F

H-F again the exception Forms a weak acid with water due to the greater bond enthalpy HX(g) + H2O(l) H3O + (aq) + X - (aq) and due to the hydrogen bonds formed between the water molecules and the H-F

26 H-F again the exception Forms a weak acid with water due to the greater bond enthalpy HX(g) + H2O(l) H3O + (aq) + X - (aq) and due to the hydrogen bonds formed between the water molecules and the H-F molecules, and between dissociated fluoride ions and H-F molecules: H-F(aq) + F - (aq) [F H-F] - (aq) The production of H3O + (aq) is inhibited, whichever way you look at it. l

Testing for aqueous halide ions STAGE 1 The solution is acidified by adding dilute nitric acid (nitric acid reacts with, and removes, other ions such as carbonate ions that might also give a

27 Testing for aqueous halide ions STAGE 1 The solution is acidified by adding dilute nitric acid (nitric acid reacts with, and removes, other ions such as carbonate ions that might also give a confusing precipitate with silver nitrate). 2. Silver nitrate solution is then added to give: ion present F - Cl - Br - I - observation no precipitate white precipitate very pale cream precipitate very pale yellow precipitate

28 Ionic Equations: STAGE 2: Confirmation Add aqueous ammonia solution: original precipitate AgCl AgBr AgI observation precipitate dissolves to give a colourless solution precipitate is almost unchanged using dilute ammonia solution, but dissolves in concentrated ammonia solution (or in XS dilute NH3(aq)) to give a colourless solution precipitate is insoluble in ammonia solution of any concentration

29 Explanation The silver halides are only very slightly soluble (see p69). In aqueous solution, rather than the above equations showing precipitation, the following equilibrium is set up: Ag + (aq) + Cl - (aq) AgCl(s) The addition of NH 3 (aq) results in the following equilibrium: By Le Chatelier, if we remove Ag + ions from solution, the top equilibrium will shift to the left and the ppt. Redissolves. This happens readily with chloride ions, less with bromide ions and not at all with iodide ions. Reactions of the silver halides with sunlight...investigate!!

30 Halide salts with concentrated sulphuric acid An alternative test for halide ions is adding concentrated sulphuric acid to the solid salt. You see: ion present F - Cl - Br - observation steamy acidic fumes (of HF) steamy acidic fumes (of HCl) steamy acidic fumes (of HBr) contaminated with brown bromine vapour I - Some steamy fumes (of HI), but lots of purple iodine vapour (plus various red colours in the tube) The concentrated sulphuric acid can act both as an acid and as an oxidising agent, depending on how easy it is to oxidise the halide ion. With all of the salts, a proton is donated to the halide ion to form the hydrogen halide gas: NaCl(s) + H2SO4(l) NaHSO4(s) + HCl(g)

31 With fluorides and chlorides the reaction stops here, but bromides and iodides are better reducing agents and redox takes place: With Br- Overall ionic eqn:

32 With I - The sulphuric acid is reduced in 3 different ways. The iodide ions are powerful enough reducing agents to reduce it: first to sulphur dioxide (sulphur oxidation state = +4) then to sulphur itself (oxidation state = 0) and all the way to hydrogen sulphide (sulphur oxidation state = -2) Bad eggs aroma Try the other reduction 1/2 eqns and try to deduce the overall ionic equations for the 3 reactions. The brown colour is due to the formation of the tri-iodide ion: I - + I2 I3 -

33 Reactions of the Hydrogen Halides: with Ammonia: with Water

34 disproportionation reactions with cold and hot alkali, eg hot potassium hydroxide with iodine to produce potassium iodate(v) Redox and Group VII: Positive oxidation states of the halides Although the syllabus only mentions chlorate(i) and chlorate(v), it is possible that a question could refer to the equivalent ions of bromine and iodine. Give the oxidation states of chlorine in its common forms and in the above 2 ions: Everyday bleach sodium chlorate(i) is prepared by reacting chlorine water with sodium hydroxide solution: 2NaOH(aq) + Cl2(aq) NaOCl(aq) + NaCl(aq) + H2O(l) Write the ionic eqn and indicate the oxidation states of the chlorine. When the same species in a chemical reaction is both oxidised and reduced, we say that disproportionation has taken place.

35 With chlorine and hot concentrated NaOH solution: 3Cl2(aq) + 6NaOH(aq) 5NaCl(aq) + NaClO3(aq) + 3H2O(l) Ionic eqn? Oxidation states? Chlorate(I) ions on heating in solution will disproportionate: 3 OCl - (aq) ClO 3- (aq) + 2Cl (aq) Full eqn? Show the oxidation states. Full eqn? Show the oxidation states. Chlorine with water ( chlorine water ): Cl 2 (g) + H 2 O HCl(aq) + HClO(aq) Write the ionic equation and consider the effects of adding acid or alkali to the mixture.

36 Redox and Group VII: Iodine with Sodium Thiosulphate c. carry out an iodine/thiosulfate titration, including calculation of the results and evaluation of the procedures involved, eg determination of the purity of potassium iodate(v) by liberation of iodine and titration with standard sodium thiosulfate solution

Fluorine Gas. Chlorine Gas. Bromine Liquid. Iodine Solid

Fluorine Gas. Chlorine Gas. Bromine Liquid. Iodine Solid Halogens Fluorine (F 2 ): very pale yellow gas. It is highly reactive Chlorine : ( ) greenish, reactive gas, poisonous in high concentrations Bromine ( ) : red liquid, that gives off dense brown/orange