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S  68 `P   P*    6 `   R*  H  0޽h ? ̙3380___PPT10. 0X(   $+ X X 04  P    X*  X 00     Z*  X 6x `P   X*  X 6 `   Z* H X 0޽h ? ̙3380___PPT10.1 0L0 P/(    0h -AN INTRODUCTION TO TRANSITION METAL COMPLEXEST$(2 82 f0f<f + <̙? x @i 0 0 , <̙? x @e 0 0 - 0 x m OKNOCKHARDY PUBLISHING(2 fJ . C "A KTRE  / 0@"`>B 2008 SPECIFICATIONS<G$G   B  s *޽h ? ̙33y___PPT10Y+D=' = @B +  0L0   0 (  dB  <D)P0Q  0Pp@ 0 0dB @ <D)PXP  0X0@ 0 0c   0o0Pa INTRODUCTION This Powerpoint show is one of several produced to help students understand selected topics at AS and A2 level Chemistry. It is based on the requirements of the AQA and OCR specifications but is suitable for other examination boards. Individual students may use the material at home for revision purposes or it may be used for classroom teaching if an interactive white board is available. Accompanying notes on this, and the full range of AS and A2 topics, are available from the KNOCKHARDY SCIENCE WEBSITE at... www.knockhardy.org.uk/sci.htm Navigation is achieved by... either clicking on the grey arrows at the foot of each page or using the left and right arrow keys on the keyboard (2 2(2 2  75, dB   <D?     0d{0P( OKNOCKHARDY PUBLISHING(2    0tX _TRANSITION METALS2(2$f<fB  s *޽h ? ̙33y___PPT10Y+D=' = @B +C 0L0 RJ` (    0䏍S ?S CONTENTS Aqueous metal ions Acidity of hexaaqua ions - stability constants Introduction to the reactions of complexes Reactions of cobalt Reactions of copper Reactions chromium Reactions on manganese Reactions of iron(II) Reactions of iron(III) Reactions of silver and vanadium Reactions of aluminiumx 2<`2 4!P:4 dB  <D)P0Q  0Pp@ 0 0dB  @ <D)PXP   0X0@ 0 0  0<X _TRANSITION METALS2(2$f<f  0% @ 0 0  6% @ 0 0  6%H@ 0 0  6% h@ 0 0  6%@ @ 0 0  6 -@ @  0 0  6 -@8@( 0 0  6 = W @ 0 0  6 -x O @! 0 0  6 - g @ 0 0  6= B@/ 0 0  6 @r 0 0B  s *޽h ? ̙33y___PPT10Y+D=' = @B +  0L0  M( w dB  <D)P0Q  0Pp@ 0 0dB @ <D)PXP  0X0@ 0 0  0q p~ x*THE AQUEOUS CHEMISTRY OF IONS - HYDROLYSIS2+(2*ffI   0  Y when salts dissolve in water the ions are stabilised this is because water molecules are polar hydrolysis can occur and the resulting solution can become acidic the acidity of the resulting solution depends on the cation present the greater the charge density of the cation, the more acidic the solution cation charge ionic radius reaction with water / pH of chloride Na 1+ 0.095 nm Mg 2+ 0.065 nm Al 3+ 0.050 nm the greater charge density of the cation... the greater the polarising power and the more acidic the solutionZ@3J?1  t  0  /    X P   C (Aaqionsg ; = dB   <DԔ  B  s *޽h ? ̙33y___PPT10Y+D=' = @B +  0L0  L(  LdB L <D)P0Q L 0Pp@ 0 0dB L@ <D)PXP L 0X0@ 0 0 L 0Ѝq p~ x*THE AQUEOUS CHEMISTRY OF IONS - HYDROLYSIS2+(2*ff L 0Ս  $ when salts dissolve in water the ions are stabilised this is because water molecules are polar hydrolysis can occur and the resulting solution can become acidic the acidity of the resulting solution depends on the cation present the greater the charge density of the cation, the more acidic the solution cation charge ionic radius reaction with water / pH of chloride Na 1+ 0.095 nm dissolves 7 Mg 2+ 0.065 nm slight hydrolysis Al 3+ 0.050 nm vigorous hydrolysis the greater charge density of the cation... the greater the polarising power and the more acidic the solution@3Jz1  t  0  /    X P L C (Aaqionsg ; = dB  L <DԔ  B L s *޽h ? ̙33y___PPT10Y+D=' = @B +~ 0L0 0 #(  dB  <D)P0Q  0Pp@ 0 0dB @ <D)PXP  0X0@ 0 0  0=q~ kTHE AQUEOUS CHEMISTRY OF IONS2(2ff   0=F 4fTheory aqueous metal ions attract water molecules many have six water molecules surrounding them these are known as hexaaqua ions they are octahedral in shape water acts as a Lewis Base  a lone pair donor water forms a co-ordinate bond to the metal ion metal ions accept the lone pair - Lewis Acids 4@e  &v   T   C ,Ahexaaqua2C?B  s *޽h ? ̙33y___PPT10Y+D=' = @B +  0L0   @  (  dB  <D)P0Q  0Pp@ 0 0dB @ <D)PXP  0X0@ 0 0  0]=q~ kTHE AQUEOUS CHEMISTRY OF IONS2(2ff  0Po=FX  PTheory aqueous metal ions attract water molecules many have six water molecules surrounding them these are known as hexaaqua ions they are octahedral in shape water acts as a Lewis Base  a lone pair donor water forms a co-ordinate bond to the metal ion metal ions accept the lone pair - Lewis Acids Acidity as charge density increases, the cation has a greater attraction for water the attraction extends to the shared pair of electrons in water s O-H bonds the electron pair is pulled towards the O, making the bond more polar this makes the H more acidic (more d+) it can then be removed by solvent water molecules to form H3O+(aq).@+?  @v    & T  C ,Ahexaaqua2C?R   C *A2+hydrol PB  s *޽h ? ̙33y___PPT10Y+D=' = @B +y 0L0 P(  R  C *A2+hydrol3 dB  <D)P0Q  0Pp@ 0 0dB @ <D)PXP  0X0@ 0 0  0x=q~ vHYDROLYSIS - EQUATIONSD(2 f f  0= P` M2+ ions [M(H2O)6]2+(aq) + H2O(l) [M(H2O)5(OH)]+(aq) + H3O+(aq) a@ 333333 33333333 33333 3  a FF  j   < tT @  #  *B   TD3V?@@B   TD3V?@tT @  # F jB   TD3V?@@B  TD3V?@2  0<=v xpL the resulting solution will now be acidic as there are more protons in the water this reaction is known as hydrolysis - the water causes the substance to split up Stronger bases (e.g. CO32- , NH3 and OH ) can remove further protons...,H   *   FF  j  t tT @ #  *B  TDV?@@B  TDV?@tT @ # F jB  TDV?@@B  TDV?@B  s *޽h ? ̙33y___PPT10Y+D=' = @B +s 0L0 z`(  dB  <D)P0Q  0Pp@ 0 0dB @ <D)PXP  0X0@ 0 0  0=q~ vHYDROLYSIS - EQUATIONSD(2 f f  0= 1U M3+ ions [M(H2O)6]3+(aq) + H2O(l) [M(H2O)5(OH)]2+(aq) + H3O+(aq) V@  333333 33 333333 33333  V FF  j   D tT @  #  *B   TD3V?@@B   TD3V?@tT @  # F jB   TD3V?@@B  TD3V?@R  C *A3+hydrol-' FF  j  t tT @ #  *B  TDV?@@B  TDV?@tT @ # F jB  TDV?@@B  TDV?@K  0$=v xpL the resulting solution will also be acidic as there are more protons in the water this SOLUTION IS MORE ACIDIC due to the greater charge density of 3+ ions Stronger bases (e.g. CO32- , NH3 and OH ) can remove further protons..., "   *   B  s *޽h ? ̙33y___PPT10Y+D=' = @B +' 0L0 && 8:z&(  dB  <D)P0Q  0Pp@ 0 0dB @ <D)PXP  0X0@ 0 0  0=q~ iHYDROLYSIS OF HEXAAQUA IONS2(2ff   0 0 A Lewis bases can attack the co-ordinated water molecules. Theoretically, a proton can be removed from each water molecule turning the water from a neutral molecule to a negatively charged hydroxide ion. This affects the overall charge on the complex ion. [M(H2O)6]2+(aq) [M(OH)(H2O)5]+(aq) [M(OH)2(H2O)4](s) [M(OH)2(H2O)4](s) [M(OH)3(H2O)3](aq) [M(OH)4(H2O)2]2-(aq) [M(OH)4(H2O)2]2-(aq) [M(OH)5(H2O)]3-(aq) [M(OH)6]4-(aq) When sufficient protons have been removed the complex becomes neutral and precipitation of a hydroxide or carbonate occurs. e.g. M2+ ions [M(H2O)4(OH)2](s) or M(OH)2 M3+ ions [M(H2O)3(OH)3](s) or M(OH)3 @ ffffffffffffffffffffffffffffffffffffffffffffffffffff fffff     FF  j   TcmtT @  #  *B   TDV?@@B   TDV?@tT @ # F jB  TDV?@@B  TDV?@FF  j  HtT @ #  *B  TDV?@@B  TDV?@tT @ # F jB  TDV?@@B  TDV?@  0 p U tOH H+RHA  0tF U tOH H+RHAFF  j  TcmtT @ #  *B  TDV?@@B  TDV?@tT @ # F jB  TDV?@@B   TDV?@FF  j ! HtT @ "#  *B # TDV?@@B $ TDV?@tT @ %# F jB & TDV?@@B ' TDV?@ ( 0Op U tOH H+RHA ) 08X U tOH H+RHAFF  j * D cm tT @ +#  *B , TDV?@@B - TDV?@tT @ .# F jB / TDV?@@B 0 TDV?@FF  j 1 8  tT @ 2#  *B 3 TDV?@@B 4 TDV?@tT @ 5# F jB 6 TDV?@@B 7 TDV?@ 8 0``p E  tOH H+RHA 9 0e E  tOH H+RHA : VAParchmentS" ?c0B  s *޽h ? ̙33y___PPT10Y+D=' = @B +' 0L0 &&P88,z&(  ,dB , <D)P0Q , 0Pp@ 0 0dB ,@ <D)PXP , 0X0@ 0 0 , 0nq~ iHYDROLYSIS OF HEXAAQUA IONS2(2ff , 0 0 A Lewis bases can attack the co-ordinated water molecules. Theoretically, a proton can be removed from each water molecule turning the water from a neutral molecule to a negatively charged hydroxide ion. This affects the overall charge on the complex ion. [M(H2O)6]2+(aq) [M(OH)(H2O)5]+(aq) [M(OH)2(H2O)4](s) [M(OH)2(H2O)4](s) [M(OH)3(H2O)3](aq) [M(OH)4(H2O)2]2-(aq) [M(OH)4(H2O)2]2-(aq) [M(OH)5(H2O)]3-(aq) [M(OH)6]4-(aq) When sufficient protons have been removed the complex becomes neutral and precipitation of a hydroxide or carbonate occurs. e.g. M2+ ions [M(H2O)4(OH)2](s) or M(OH)2 M3+ ions [M(H2O)3(OH)3](s) or M(OH)3 @ ffffffffffffffffffffffffffffffffffffffffffffffffffff fffff     FF  j , TcmtT @  ,#  *B  , TDV?@@B  , TDV?@tT @  ,# F jB  , TDV?@@B , TDV?@FF  j , HtT @ ,#  *B , TDV?@@B , TDV?@tT @ ,# F jB , TDV?@@B , TDV?@ , 0p U tOH H+RHA , 0p U tOH H+RHAFF  j , TcmtT @ ,#  *B , TDV?@@B , TDV?@tT @ ,# F jB , TDV?@@B , TDV?@FF  j , HtT @  ,#  *B !, TDV?@@B ", TDV?@tT @ #,# F jB $, TDV?@@B %, TDV?@ &, 0dp U tOH H+RHA ', 0 U tOH H+RHAFF  j (, D cm tT @ ),#  *B *, TDV?@@B +, TDV?@tT @ ,,# F jB -, TDV?@@B ., TDV?@FF  j /, 8  tT @ 0,#  *B 1, TDV?@@B 2, TDV?@tT @ 3,# F jB 4, TDV?@@B 5, TDV?@ 6, 0(p E  tOH H+RHA 7, 0  E  tOH H+RHA 8, VAParchmentS" ?U0 B , s *޽h ? ̙33y___PPT10Y+D=' = @B +S( 0L0 b'Z'09;&(  dB  <D)P0Q  0Pp@ 0 0dB @ <D)PXP  0X0@ 0 0  0q~ iHYDROLYSIS OF HEXAAQUA IONS2(2ff[   0x 0 QLewis bases can attack the co-ordinated water molecules. Theoretically, a proton can be removed from each water molecule turning the water from a neutral molecule to a negatively charged hydroxide ion. This affects the overall charge on the complex ion. [M(H2O)6]2+(aq) [M(OH)(H2O)5]+(aq) [M(OH)2(H2O)4](s) [M(OH)2(H2O)4](s) [M(OH)3(H2O)3](aq) [M(OH)4(H2O)2]2-(aq) [M(OH)4(H2O)2]2-(aq) [M(OH)5(H2O)]3-(aq) [M(OH)6]4-(aq) In some cases, if the base is strong, further protons are removed and the precipitate dissolves as soluble anionic complexes such as [M(OH)6]3- are formed. Very weak bases H2O remove few protons Weak bases NH3, CO32- remove protons until precipitation Strong bases OH can remove all the protonszR@ ffffffffffffffffffffffffffffffffffffffffffffffffffff fffff    )X R FF  j   TcmtT @  #  *B   TDV?@@B   TDV?@tT @ # F jB  TDV?@@B  TDV?@FF  j  HtT @ #  *B  TDV?@@B  TDV?@tT @ # F jB  TDV?@@B  TDV?@  06p U tOH H+RHA  0LL U tOH H+RHAFF  j  TcmtT @ #  *B  TDV?@@B  TDV?@tT @ # F jB  TDV?@@B   TDV?@FF  j ! HtT @ "#  *B # TDV?@@B $ TDV?@tT @ %# F jB & TDV?@@B ' TDV?@ ( 0Vp U tOH H+RHA ) 0p_ U tOH H+RHAFF  j * D cm tT @ +#  *B , TDV?@@B - TDV?@tT @ .# F jB / TDV?@@B 0 TDV?@FF  j 1 8  tT @ 2#  *B 3 TDV?@@B 4 TDV?@tT @ 5# F jB 6 TDV?@@B 7 TDV?@ 8 0ip E  tOH H+RHA 9 0xr E  tOH H+RHA : 0w \ Precipitated4 HA  ; VAParchmentS" ?P0 B  s *޽h ? ̙33y___PPT10Y+D=' = @B +' 0L0 &&p894L&(  4dB 4 <D)P0Q 4 0Pp@ 0 0dB 4@ <D)PXP 4 0X0@ 0 0 4 0 q~ iHYDROLYSIS OF HEXAAQUA IONS2(2ff[ 4 0 0 QLewis bases can attack the co-ordinated water molecules. Theoretically, a proton can be removed from each water molecule turning the water from a neutral molecule to a negatively charged hydroxide ion. This affects the overall charge on the complex ion. [M(H2O)6]2+(aq) [M(OH)(H2O)5]+(aq) [M(OH)2(H2O)4](s) [M(OH)2(H2O)4](s) [M(OH)3(H2O)3](aq) [M(OH)4(H2O)2]2-(aq) [M(OH)4(H2O)2]2-(aq) [M(OH)5(H2O)]3-(aq) [M(OH)6]4-(aq) In some cases, if the base is strong, further protons are removed and the precipitate dissolves as soluble anionic complexes such as [M(OH)6]3- are formed. Very weak bases H2O remove few protons Weak bases NH3, CO32- remove protons until precipitation Strong bases OH can remove all the protonszR@ ffffffffffffffffffffffffffffffffffffffffffffffffffff fffff    )X R FF  j 4 TcmtT @  4#  *B  4 TDV?@@B  4 TDV?@tT @  4# F jB  4 TDV?@@B 4 TDV?@FF  j 4 HtT @ 4#  *B 4 TDV?@@B 4 TDV?@tT @ 4# F jB 4 TDV?@@B 4 TDV?@ 4 0dp U tOH H+RHA 4 0с U tOH H+RHAFF  j 4 TcmtT @ 4#  *B 4 TDV?@@B 4 TDV?@tT @ 4# F jB 4 TDV?@@B 4 TDV?@FF  j 4 HtT @  4#  *B !4 TDV?@@B "4 TDV?@tT @ #4# F jB $4 TDV?@@B %4 TDV?@ &4 0Dہp U tOH H+RHA '4 0 U tOH H+RHAFF  j (4 D cm tT @ )4#  *B *4 TDV?@@B +4 TDV?@tT @ ,4# F jB -4 TDV?@@B .4 TDV?@FF  j /4 8  tT @ 04#  *B 14 TDV?@@B 24 TDV?@tT @ 34# F jB 44 TDV?@@B 54 TDV?@ 64 0p E  tOH H+RHA 74 0$ E  tOH H+RHA 84 0 \ Precipitated4 HA B 4 s *޽h ? ̙33y___PPT10Y+D=' = @B +2 0L0 $11KK0(  dB  <D)P0Q  0Pp@ 0 0dB @ <D)PXP  0X0@ 0 0  0q~ iHYDROLYSIS OF HEXAAQUA IONS2(2ff  0 0 ! Lewis bases can attack the co-ordinated water molecules. Theoretically, a proton can be removed from each water molecule turning the water from a neutral molecule to a negatively charged hydroxide ion. This affects the overall charge on the complex ion. [M(H2O)6]2+(aq) [M(OH)(H2O)5]+(aq) [M(OH)2(H2O)4](s) [M(OH)2(H2O)4](s) [M(OH)3(H2O)3](aq) [M(OH)4(H2O)2]2-(aq) [M(OH)4(H2O)2]2-(aq) [M(OH)5(H2O)]3-(aq) [M(OH)6]4-(aq) AMPHOTERIC CHARACTER Metal ions of 3+ charge have a high charge density and their hydroxides can dissolve in both acid and alkali. [M(H2O)6]3+(aq) [M(OH)3(H2O)3](s) [M(OH)6]3-(aq) @ ffffffffffffffffffffffffffffffffffffffffffffffffffff fffff  _fffffffffffffffff  FF  j  TcmtT @  #  *B   TDV?@@B   TDV?@tT @  # F jB   TDV?@@B  TDV?@FF  j  HtT @ #  *B  TDV?@@B  TDV?@tT @ # F jB  TDV?@@B  TDV?@  0Vp U tOH H+RHA  0a U tOH H+RHAFF  j  TcmtT @ #  *B  TDV?@@B  TDV?@tT @ # F jB  TDV?@@B  TDV?@FF  j  HtT @  #  *B ! TDV?@@B " TDV?@tT @ ## F jB $ TDV?@@B % TDV?@ & 0gp U tOH H+RHA ' 0o U tOH H+RHAFF  j ( D cm tT @ )#  *B * TDV?@@B + TDV?@tT @ ,# F jB - TDV?@@B . TDV?@FF  j / 8  tT @ 0#  *B 1 TDV?@@B 2 TDV?@tT @ 3# F jB 4 TDV?@@B 5 TDV?@ 6 0yp E  tOH H+RHA 7 0 E  tOH H+RHA 8 0v \ Precipitated4 HA FF  j 9 ;tT @ :#  *B ; TDV?@@B < TDV?@tT @ =# F jB > TDV?@@B ? TDV?@FF  j @ 7tT @ A#  *B B TDV?@@B C TDV?@tT @ D# F jB E TDV?@@B F TDV?@ G 0} 8  QOH2HA H 0}   RH+4HA I 0p   W Insoluble2   J 0[ USoluble2 K 0,  USoluble2B  s *޽h ? ̙33y___PPT10Y+D=' = @B +% 0L0 4,T(  TdB T <D)P0Q T 0Pp@ 0 0dB T@ <D)PXP T 0X0@ 0 0 T 0q~ aSTABILITY CONSTANTS2(2ff T 0 0 ADefinition The stability constant, Kstab, of a complex ion is the equilibrium constant for the formation of the complex ion in a solvent from its constituent ions. H ! &&   B T s *޽h ? ̙33y___PPT10Y+D=' = @B +  0L0   L8 (  8dB 8 <D)P0Q 8 0Pp@ 0 0dB 8@ <D)PXP 8 0X0@ 0 0 8 0q~ aSTABILITY CONSTANTS2(2ff 8 0@ 0{ Definition The stability constant, Kstab, of a complex ion is the equilibrium constant for the formation of the complex ion in a solvent from its constituent ions. In the reaction [M(H2O)6]2+(aq) + 6X(aq) [MX6]4 (aq) + 6H2O(l) L ! ffffffffff fff@&     FF  j 8 U_tT @  8#  *B !8 TDfV?@@B "8 TDfV?@tT @ #8# F jB $8 TDfV?@@B %8 TDfV?@B 8 s *޽h ? ̙33y___PPT10Y+D=' = @B +2 0L0 A9X (  XdB X <D)P0Q X 0Pp@ 0 0dB X@ <D)PXP X 0X0@ 0 0 X 0łq~ aSTABILITY CONSTANTS2(2ff X 0͂ 0  ,Definition The stability constant, Kstab, of a complex ion is the equilibrium constant for the formation of the complex ion in a solvent from its constituent ions. In the reaction [M(H2O)6]2+(aq) + 6X(aq) [MX6]4 (aq) + 6H2O(l) the expression for the stability constant is Kstab = [ [MX64 ](aq) ] [ [M(H2O)6]2+(aq) ] [ X(aq) ]6 C@(: ! ffffffffff fff3 ffffffffffff!!&    A    /    FF  j X U_tT @  X#  *B  X TDfV?@@B  X TDfV?@tT @  X# F jB  X TDfV?@@B X TDfV?@RB X s *DoW /B X s *޽h ? ̙33y___PPT10Y+D=' = @B +0 0L0 ?7P(  PdB P <D)P0Q P 0Pp@ 0 0dB P@ <D)PXP P 0X0@ 0 0 P 0q~ aSTABILITY CONSTANTS2(2ff P 0 0  Definition The stability constant, Kstab, of a complex ion is the equilibrium constant for the formation of the complex ion in a solvent from its constituent ions. In the reaction [M(H2O)6]2+(aq) + 6X(aq) [MX6]4 (aq) + 6H2O(l) the expression for the stability constant is Kstab = [ [MX64 ](aq) ] [ [M(H2O)6]2+(aq) ] [ X(aq) ]6 The concentration of X(aq) appears to the power of 6 because there are six of the ions in the equation. Note that the water isn t included; it is in such overwhelming quantity that its concentration can be regarded as  constant .C@(% ! ffffffffff fff3 ffffffffffff!q!}&    A    /     FF  j P U_tT @  P#  *B  P TDfV?@@B  P TDfV?@tT @  P# F jB  P TDfV?@@B P TDfV?@RB P s *DoW /B P s *޽h ? ̙33y___PPT10Y+D=' = @B +% 0L0 $$+h<w$(  <dB < <D)P0Q < 0Pp@ 0 0dB <@ <D)PXP < 0X0@ 0 0 < 09q~ aSTABILITY CONSTANTS2(2ffX < 0,M 0  >Because ligand exchange involves a series of equilibria, each step in the process has a different stability constant& Kstab / dm3 mol-1 [Co(H2O)6]2+(aq) + NH3(aq) [Co(NH3)(H2O)5]2+(aq) + H2O(l) K1 = 1.02 x 10-2 [Co(NH3)(H2O)5]2+(aq) + NH3(aq) [Co(NH3)2(H2O)4]2+(aq) + H2O(l) K2 = 3.09 x 10-2 [Co(NH3)2(H2O)4]2+(aq) + NH3(aq) [Co(NH3)3(H2O)3]2+(aq) + H2O(l) K3 = 1.17 x 10-1 [Co(NH3)3(H2O)3]2+(aq) + NH3(aq) [Co(NH3)4(H2O)2]2+(aq) + H2O(l) K4 = 2.29 x 10-1 [Co(NH3)4(H2O)2]2+(aq) + NH3(aq) [Co(NH3)5(H2O)]2+(aq) + H2O(l) K5 = 8.70 x 10-1 etc < @                  !!!!!!!! !! !!!!!!!!!! !!!   ffffffffff ff ffffffffff fff   f3f3f3f3f3f3f3f3f3f3f3 f3f3 f3f3f3f3f3f3f3f3f3f3 f3f3f3   3333333333 33 33333333 333   Z     d   8 Vp `  h<Vp H NN  j < V` <tT @  <#  *B  < TDV?@@B  < TDV?@tT @  <# F jB  < TDV?@@B < TDV?@NN  j L< ^h PtT @ M<#  *B N< TD!V?@@B O< TD!V?@tT @ P<# F jB Q< TD!V?@@B R< TD!V?@NN  j S< ^Th tT @ T<#  *B U< TDV?@@B V< TDV?@tT @ W<# F jB X< TDV?@@B Y< TDV?@NN  j Z< ^h tT @ [<#  *B \< TDf3V?@@B ]< TDf3V?@tT @ ^<# F jB _< TDf3V?@@B `< TDf3V?@NN  j a< f p ` tT @ b<#  *B c< TD3V?@@B d< TD3V?@tT @ e<# F jB f< TD3V?@@B g< TD3V?@B < s *޽h ? ̙33y___PPT10Y+D=' = @B +l* 0L0 {)s),,H)(  HdB H <D)P0Q H 0Pp@ 0 0dB H@ <D)PXP H 0X0@ 0 0 H 0<ʅq~ aSTABILITY CONSTANTS2(2ff H 0 05 *Because ligand exchange involves a series of equilibria, each step in the process has a different stability constant& Kstab / dm3 mol-1 [Co(H2O)6]2+(aq) + NH3(aq) [Co(NH3)(H2O)5]2+(aq) + H2O(l) K1 = 1.02 x 10-2 [Co(NH3)(H2O)5]2+(aq) + NH3(aq) [Co(NH3)2(H2O)4]2+(aq) + H2O(l) K2 = 3.09 x 10-2 [Co(NH3)2(H2O)4]2+(aq) + NH3(aq) [Co(NH3)3(H2O)3]2+(aq) + H2O(l) K3 = 1.17 x 10-1 [Co(NH3)3(H2O)3]2+(aq) + NH3(aq) [Co(NH3)4(H2O)2]2+(aq) + H2O(l) K4 = 2.29 x 10-1 [Co(NH3)4(H2O)2]2+(aq) + NH3(aq) [Co(NH3)5(H2O)]2+(aq) + H2O(l) K5 = 8.70 x 10-1 etc The overall stability constant is simply the equilibrium constant for the total reaction. It is found by multiplying the individual stability constants... k1 x k2 x k3 x k4 ... etc Kstab or pKstab? For an easier comparison, the expression pKstab is often used& pKstab = -log10Kstab  @.                  !!!!!!!! !! !!!!!!!!!! !!!   ffffffffff ff ffffffffff fff   f3f3f3f3f3f3f3f3f3f3f3 f3f3 f3f3f3f3f3f3f3f3f3f3 f3f3f3   3333333333 33 33333333 333         0 !!!!!!!     d      .     RB H s *Do  F Vp `   H Vp H NN  j  H V` <tT @  H#  *B  H TDV?@@B  H TDV?@tT @ H# F jB H TDV?@@B H TDV?@NN  j H ^h PtT @ H#  *B H TD!V?@@B H TD!V?@tT @ H# F jB H TD!V?@@B H TD!V?@NN  j H ^Th tT @ H#  *B H TDV?@@B H TDV?@tT @ H# F jB H TDV?@@B H TDV?@NN  j H ^h tT @  H#  *B !H TDf3V?@@B "H TDf3V?@tT @ #H# F jB $H TDf3V?@@B %H TDf3V?@NN  j &H f p ` tT @ 'H#  *B (H TD3V?@@B )H TD3V?@tT @ *H# F jB +H TD3V?@@B ,H TD3V?@B H s *޽h ? ̙33y___PPT10Y+D=' = @B +v) 0L0 (}(,PD((  DdB D <D)P0Q D 0Pp@ 0 0dB D@ <D)PXP D 0X0@ 0 0 D 0rq~ aSTABILITY CONSTANTS2(2ff D 0 0 4ZBecause ligand exchange involves a series of equilibria, each step in the process has a different stability constant& Kstab / dm3 mol-1 [Co(H2O)6]2+(aq) + NH3(aq) [Co(NH3)(H2O)5]2+(aq) + H2O(l) K1 = 1.02 x 10-2 [Co(NH3)(H2O)5]2+(aq) + NH3(aq) [Co(NH3)2(H2O)4]2+(aq) + H2O(l) K2 = 3.09 x 10-2 [Co(NH3)2(H2O)4]2+(aq) + NH3(aq) [Co(NH3)3(H2O)3]2+(aq) + H2O(l) K3 = 1.17 x 10-1 [Co(NH3)3(H2O)3]2+(aq) + NH3(aq) [Co(NH3)4(H2O)2]2+(aq) + H2O(l) K4 = 2.29 x 10-1 [Co(NH3)4(H2O)2]2+(aq) + NH3(aq) [Co(NH3)5(H2O)]2+(aq) + H2O(l) K5 = 8.70 x 10-1 etc Summary " The larger the stability constant, the further the reaction lies to the right " Complex ions with large stability constants are more stable " Stability constants are often given as pKstab " Complex ions with smaller pKstab values are more stable ( @H@@____________________________________ ______________________________ ______ ______ ______________________________ ______ ______________________________ ______ ______ ____________________________________ ______ ______________________________ ______ ______ __________________________________________ ______ ______________________________ ______ ______ ____________________________________ ______ ________________________ ______ ______ _______________! #      d    !   RB D s *Do  F Vp `  -D Vp H NN  j .D V` <tT @ /D#  *B 0D TD___V?@@B 1D TD___V?@tT @ 2D# F jB 3D TD___V?@@B 4D TD___V?@NN  j 5D ^h PtT @ 6D#  *B 7D TD___V?@@B 8D TD___V?@tT @ 9D# F jB :D TD___V?@@B ;D TD___V?@NN  j D TD___V?@@B ?D TD___V?@tT @ @D# F jB AD TD___V?@@B BD TD___V?@NN  j CD ^h tT @ DD#  *B ED TD___V?@@B FD TD___V?@tT @ GD# F jB HD TD___V?@@B ID TD___V?@NN  j JD f p ` tT @ KD#  *B LD TD___V?@@B MD TD___V?@tT @ ND# F jB OD TD___V?@@B PD TD___V?@B D s *޽h ? ̙33y___PPT10Y+D=' = @B +w 0L0  ~ @  ( w dB  <D)P0Q  0Pp@ 0 0dB @ <D)PXP  0X0@ 0 0  04q~ \REACTION TYPES2(2ffv   0"` 0 The examples aim to show typical properties of transition metals and their compounds. One typical properties of transition elements is their ability to form complex ions. Complex ions consist of a central metal ion surrounded by co-ordinated ions or molecules known as ligands. This can lead to changes in ... " colour " co-ordination number " shape " stability to oxidation or reduction Reaction types ACID-BASE LIGAND SUBSTITUTION PRECIPITATION REDOX @M@A  %1fff&      6%x x5  OA-B(2     6$' x  LLS 2     6P-} b: LOX 2     605 x  [Ppt 2    6X5} X: ORED(2    6t3} : ]REDOX( 2   B  s *޽h ? ̙33y___PPT10Y+D=' = @B +  0L0 . & P  (  dB  <D)P0Q  0Pp@ 0 0dB @ <D)PXP  0X0@ 0 0  0<@q~ \REACTION TYPES2(2ff  0F 0  @vThe examples aim to show typical properties of transition metals and their compounds. LOOK FOR... substitution reactions of complex ions variation in oxidation state of transition metals the effect of ligands on co-ordination number and shape increased acidity of M3+ over M2+ due to the increased charge density differences in reactivity of M3+ and M2+ ions with OH and NH3 the reason why M3+ ions don t form carbonates amphoteric character in some metal hydroxides (Al3+ and Cr3+) the effect a ligand has on the stability of a particular oxidation stateLZ@ H@J@Zf'f39f%ffffff5ffffffJZ     E  6 B  s *޽h ? ̙33y___PPT10Y+D=' = @B +  0L0   p h (  hdB h <D)P0Q h 0Pp@ 0 0dB h@ <D)PXP h 0X0@ 0 0 h 0 lq~ eREACTIONS OF COBALT(II)2(2ff  h 0r"`0 l" aqueous solutions contain the pink, octahedral hexaaquacobalt(II) ion " hexaaqua ions can also be present in solid samples of the hydrated salts " as a 2+ ion, the solutions are weakly acidic but protons can be removed by bases... OH [Co(H2O)6]2+(aq) + 2OH(aq)   > [Co(OH)2(H2O)4](s) + 2H2O(l) pink, octahedral blue / pink ppt. soluble in XS NaOH ALL hexaaqua ions precipitate a hydroxide with OH(aq). Some re-dissolve in excess NaOHI}O}!#  fffffffffffffffffAd 2           #       h 6ܜ"` OA-B(2  XB h 0Do8 I 8 T h C ,A2++hydrol B h s *޽h ? ̙33y___PPT10Y+D=' = @B +s  0L0  z   (  dB  <D)P0Q  0Pp@ 0 0dB @ <D)PXP  0X0@ 0 0  0q~ eREACTIONS OF COBALT(II)2(2ffl   0D"`0  @NH3 [Co(H2O)6]2+(aq) + 2NH3(aq)   > [Co(OH)2(H2O)4](s) + 2NH4+(aq) ALL hexaaqua ions precipitate a hydroxide with NH3 (aq). It removes protons ZKN}ffffffffffffffffff 2 @P  (      6χ"`&0 OA-B(2  L  C $A 2+nh38XB  0DjJ! B  s *޽h ? ̙33y___PPT10Y+D=' = @B + 0L0    P (  dB  <D)P0Q  0Pp@ 0 0dB @ <D)PXP  0X0@ 0 0  0Pq~ eREACTIONS OF COBALT(II)2(2ff  00]  NH3 [Co(H2O)6]2+(aq) + 2NH3(aq)   > [Co(OH)2(H2O)4](s) + 2NH4+(aq) ALL hexaaqua ions precipitate a hydroxide with NH3 (aq). It removes protons Some hydroxides redissolve in excess NH3(aq) as ammonia substitutes as a ligand. [Co(OH)2(H2O)4](s) + 6NH3(aq)   > [Co(NH3)6]2+(aq) + 4H2O(l) + 2OH(aq) rKNZ}Uffffffffffffffffff 2 ) *ffffffffffffffffff ftP  (  /   0  W   6%&0 OA-B(2     6) 0  LLS 2  L   C $A 2+nh38XB   0DjJ! B  s *޽h ? ̙33y___PPT10Y+D=' = @B +9 0L0 H@ ( w dB  <D)P0Q  0Pp@ 0 0dB @ <D)PXP  0X0@ 0 0  0q~ eREACTIONS OF COBALT(II)2(2ff  0(>0q l RNH3 [Co(H2O)6]2+(aq) + 2NH3(aq)   > [Co(OH)2(H2O)4](s) + 2NH4+(aq) ALL hexaaqua ions precipitate a hydroxide with NH3 (aq). It removes protons Some hydroxides redissolve in excess NH3(aq) as ammonia substitutes as a ligand. [Co(OH)2(H2O)4](s) + 6NH3(aq)   > [Co(NH3)6]2+(aq) + 4H2O(l) + 2OH(aq) but ... ammonia ligands make the Co(II) state unstable. Air oxidises Co(II) to Co(III) [Co(NH3)6]2+(aq)   > [Co(NH3)6]3+(aq) + e yellow / brown octahedral red / brown octahedralKNZ}SYFED}ffffffffffffffffff 2 ) *ffffffffffffffffff fYffffffffffff f 9P  (  /   0  g             6&0 OA-B(2     6 0  LLS 2     6d00% LOX 2  L   C $A 2+nh38XB   0DjJ! B  s *޽h ? ̙33y___PPT10Y+D=' = @B +  0L0    : (  dB  <D)P0Q  0Pp@ 0 0dB @ <D)PXP  0X0@ 0 0  0q~ eREACTIONS OF COBALT(II)2(2ff   0"`0N pCO32- [Co(H2O)6]2+(aq) + CO32-(aq)   > CoCO3(s) + 6H2O(l) mauve ppt. Hexaaqua ions of metals with charge 2+ precipitate a carbonate but heaxaaqua ions with a 3+ charge don t. :I&}l}ffffff fffff ffff kZi    <        6 [Ppt 2  ^B  6DjJHHB  s *޽h ? ̙33y___PPT10Y+D=' = @B + 0L0  $(  dB  <D)P0Q  0Pp@ 0 0dB @ <D)PXP  0X0@ 0 0  0Ìq~ eREACTIONS OF COBALT(II)2(2ffb  0\0   ~CO32- [Co(H2O)6]2+(aq) + CO32-(aq)   > CoCO3(s) + 6H2O(l) mauve ppt. Hexaaqua ions of metals with charge 2+ precipitate a carbonate but heaxaaqua ions with a 3+ charge don t. Cl [Co(H2O)6]2+(aq) + 4Cl(aq)   > [CoCl4]2-(aq) + 6H2O(l) pink, octahedral blue, tetrahedral " Cl ligands are larger than H2O " Cl ligands are negatively charged - H2O ligands are neutral " the complex is more stable if tetrahedral - less repulsion between ligands " adding excess water reverses the reactionPI&}ld}ffffff fffff ffff kffffff fffff fff'   - i    <                V  /   6Ȍ LLS 2  P   C (A cltetra L   C $A 6aqua *   6 [Ppt 2  ^B   6DjJHHB  s *޽h ? ̙33y___PPT10Y+D=' = @B +  0L0    : (  dB  <D)P0Q  0Pp@ 0 0dB @ <D)PXP  0X0@ 0 0  0q~ eREACTIONS OF COPPER(II)2(2ff   0h"`0 ntAqueous solutions of copper(II) contain the blue, octahedral hexaaquacopper(II) ion Most substitution reactions are similar to cobalt(II). OH [Cu(H2O)6]2+(aq) + 2OH(aq)   > [Cu(OH)2(H2O)4](s) + 2H2O(l) blue, octahedral pale blue ppt. insoluble in XS NaOH H8}+,#< fffffff ffffffffffffff_     1     %      6B"`0 OA-B(2  ^B  6DjJB  s *޽h ? ̙33y___PPT10Y+D=' = @B + 0L0  +(   dB   <D)P0Q   0Pp@ 0 0dB  @ <D)PXP   0X0@ 0 0   0Iq~ eREACTIONS OF COPPER(II)2(2ff   0]0S Aqueous solutions of copper(II) contain the blue, octahedral hexaaquacopper(II) ion Most substitution reactions are similar to cobalt(II). OH [Cu(H2O)6]2+(aq) + 2OH(aq)   > [Cu(OH)2(H2O)4](s) + 2H2O(l) blue, octahedral pale blue ppt. insoluble in XS NaOH NH3 [Cu(H2O)6]2+(aq) + 2NH3(aq)   > [Cu(OH)2(H2O)4](s) + 2NH4+(aq) blue ppt. soluble in excess NH3 then [Cu(OH)2(H2O)4](s) + 4NH3(aq)   > [Cu(NH3)4(H2O)2]2+(aq) + 2H2O(l) + 2OH(aq) royal blue NOTE THE FORMULA 8}y,}},#< fffffff ffffffffffffff_fffffff fffffffffff ffff   fffffffffff fffffffffffffffff     1     %  c      6c0 OA-B(2     6c0X OA-B(2     6+ 0  LLS 2  ^B   6DjJ^B   6DjJP P B   s *޽h ? ̙33y___PPT10Y+D=' = @B +  0L0  ( w dB  <D)P0Q  0Pp@ 0 0dB @ <D)PXP  0X0@ 0 0  0Pq~ eREACTIONS OF COPPER(II)2(2ff  0ڍ0 ZAqueous solutions of copper(II) contain the blue, octahedral hexaaquacopper(II) ion Most substitution reactions are similar to cobalt(II). OH [Cu(H2O)6]2+(aq) + 2OH(aq)   > [Cu(OH)2(H2O)4](s) + 2H2O(l) blue, octahedral pale blue ppt. insoluble in XS NaOH NH3 [Cu(H2O)6]2+(aq) + 2NH3(aq)   > [Cu(OH)2(H2O)4](s) + 2NH4+(aq) blue ppt. soluble in excess NH3 then [Cu(OH)2(H2O)4](s) + 4NH3(aq)   > [Cu(NH3)4(H2O)2]2+(aq) + 2H2O(l) + 2OH(aq) royal blue NOTE THE FORMULA CO32- [Cu(H2O)6]2+(aq) + CO32-(aq)   > CuCO3(s) + 6H2O(l) blue ppt.8}y,}}O$},#< fffffff ffffffffffffff_fffffff fffffffffff ffff   fffffffffff ffffffffffffffffffffffff fffffff ffff      1     %  c  -     6>0 OA-B(2     6Dc0X OA-B(2     6H+ 0  LLS 2     6@0 [Ppt 2  ^B   6DjJ^B   6DjJP P B  s *޽h ? ̙33y___PPT10Y+D=' = @B + 0L0     (  dB  <D)P0Q  0Pp@ 0 0dB @ <D)PXP  0X0@ 0 0  0Pq~ eREACTIONS OF COPPER(II)2(2ffv   0TF  Cl [Cu(H2O)6]2+(aq) + 4Cl(aq)   > [CuCl4]2-(aq) + 6H2O(l) yellow, tetrahedral " Cl ligands are larger than H2O and are charged " the complex is more stable if the shape changes to tetrahedral " adding excess water reverses the reaction I 2Cu2+(aq) + 4I(aq)   > 2CuI(s) + I2(aq) off - white ppt. " a redox reaction " used in the volumetric analysis of copper using sodium thiosulphaten}w@q*}@Hfffffff fffffff ffff  fff ffff fffQ q         E       6% LLS 2     6̍xm  OREDOX 2   ^B   6DjJPPB  s *޽h ? ̙33y___PPT10Y+D=' = @B +  0L0 $   (  dB  <D)P0Q  0Pp@ 0 0dB @ <D)PXP  0X0@ 0 0  0Hq~ dREACTIONS OF COPPER(I)2(2ff   0 "`0v  "The aqueous chemistry of copper(I) is unstable compared to copper(0) and copper (II). Cu+(aq) + e   > Cu(s) E = + 0.52 V Cu2+(aq) + e   > Cu+(aq) E = + 0.15 V subtracting 2Cu+(aq)   > Cu(s) + Cu2+(aq) E = + 0.37 V This is an example of DISPROPORTIONATION where one species is simultaneously oxidised and reduced to more stable forms. This explains why the aqueous chemistry of copper(I) is very limited. Copper(I) can be stabilised by formation of complexes.=}K8Vffffffffffffff ff fff$^7ZY  V    L B  s *޽h ? ̙33y___PPT10Y+D=' = @B + 0L0 ' d(  ddB d <D)P0Q d 0Pp@ 0 0dB d@ <D)PXP d 0X0@ 0 0 d 0ǐq~ hREACTIONS OF CHROMIUM(III)2(2ff  d 0"`0_ b Chromium(III) ions are typical of M3+ ions in this block. Aqueous solutions contain the violet, octahedral hexaaquachromium(III) ion. OH [Cr(H2O)6]3+(aq) + 3OH(aq)   > [Cr(OH)3(H2O)3](s) + 3H2O(l) violet, octahedral green ppt. soluble in XS NaOH As with all hydroxides the precipitate reacts with acid [Cr(OH)3(H2O)3](s) + 3H+(aq)   > [Cr(H2O)6]3+(aq) being a 3+ hydroxide it is AMPHOTERIC as it dissolves in excess alkali [Cr(OH)3(H2O)3](s) + 3OH(aq)   > [Cr(OH)6]3-(aq) + 3H2O(l) green, octahedral:L}P}AL}N6}# 3(  ffffff f ffffffffff ffff?Bffffffff fffffffffff  $ffffffff fffffff ffffh  _      S   d 6ِ"`F OA-B(2    d 6"` F  OA-B(2   d 6"`F OA-B(2  T d C ,A3++hydrol= B d s *޽h ? ̙33y___PPT10Y+D=' = @B +.  0L0 = 5   (  dB  <D)P0Q  0Pp@ 0 0dB @ <D)PXP  0X0@ 0 0  02q~ hREACTIONS OF CHROMIUM(III)2(2ffr   0?"`0   CO32- 2 [Cr(H2O)6]3+(aq) + 3CO32-(aq)   > 2[Cr(OH)3(H2O)3](s) + 3H2O(l) + 3CO2(g) The carbonate is not precipitated but the hydroxide is. high charge density of M3+ makes the solution too acidic to form the carbonate CARBON DIOXIDE IS EVOLVED. Z@2ffffffffffffffffffffffffffQ Q    6]"` OA-B(2  ^B  6DjJX B  s *޽h ? ̙33y___PPT10Y+D=' = @B + 0L0  (  dB  <D)P0Q  0Pp@ 0 0dB @ <D)PXP  0X0@ 0 0  0Lcq~ hREACTIONS OF CHROMIUM(III)2(2ff  0u0  | (CO32- 2 [Cr(H2O)6]3+(aq) + 3CO32-(aq)   > 2[Cr(OH)3(H2O)3](s) + 3H2O(l) + 3CO2(g) The carbonate is not precipitated but the hydroxide is. high charge density of M3+ makes the solution too acidic to form the carbonate CARBON DIOXIDE IS EVOLVED. NH3 [Cr(H2O)6]3+(aq) + 3NH3(aq)   > [Cr(OH)3(H2O)3](s) + 3NH4+(aq) green ppt. soluble in XS NH3 With EXCESS AMMONIA, the precipitate redissolves [Cr(OH)3(H2O)3](s) + 6NH3(aq)   > [Cr(NH3)6]3+(aq) + 3H2O(l) + 3OH(aq)Z@2S;}ffffffffffffffffffffffffffQ Qfffffff fffffffffffffff7  5ffffffffffffffffffffffff@u  <   V   6Dw   LLS 2     6hU J  OA-B(2     6XÑ OA-B(2  ^B   6DjJX B  s *޽h ? ̙33y___PPT10Y+D=' = @B +  0L0   (  dB  <D)P0Q  0Pp@ 0 0dB @ <D)PXP  0X0@ 0 0  0,ېq~ hREACTIONS OF CHROMIUM(III)2(2ff   0Б"`0D   Oxidation In the presence of alkali, Cr(III) is unstable and can be oxidised to Cr(VI) 2Cr3+(aq) + 3H2O2(l) + 10OH(aq)   > 2CrO42-(aq) + 8H2O(l) green yellow Acidification of the yellow chromate will produce the orange dichromate(VI) ion Reduction Chromium(III) can be reduced to the less stable chromium(II) by zinc in acid 2 [Cr(H2O)6]3+(aq) + Zn(s)   > 2 [Cr(H2O)6]2+(aq) + Zn2+(aq) green blue@~}3} Lff fffffff fff _ Lffffff"fffffffff3(  %        $   1  c    6"`^ LOX 2     6"` ^  ORED(2  ^B   6DjJppB  s *޽h ? ̙33y___PPT10Y+D=' = @B + 0L0 (  dB  <D)P0Q  0Pp@ 0 0dB @ <D)PXP  0X0@ 0 0  0q~ gREACTIONS OF CHROMIUM(VI)2(2ff   0'0  ;gOccurrence dichromate (VI) Cr2O72- orange chromate (VI) CrO42- yellow Interconversion dichromate is stable in acid solution chromate is stable in alkaline solution in alkali Cr2O72-(aq) + 2OH(aq) 2CrO42-(aq) + H2O(l) in acid 2CrO42-(aq) + 2H+(aq) Cr2O72-(aq) + H2O(l) U\ ffffffffff P ffffff ffffffffffffff fffffffff fffff&U   F ?   c 8N    fB   6DfolB   <BDfo{>T  # L)?fB  6DfolB  <BDfo{F ?   / 8N   fB  6DfolB  <BDfo{>T  # L)?fB  6DfolB  <BDfo{B  s *޽h ? ̙33y___PPT10Y+D=' = @B +N 0L0 ]U  (  dB  <D)P0Q  0Pp@ 0 0dB @ <D)PXP  0X0@ 0 0  # lN̙ ?A >  HCONTENTS$ (2  <̙? x @ 0 0  0lq~ q#OXIDATION REACTIONS OF CHROMIUM(VI)2$(2#ff   0o"`0  Being in the highest oxidation state (+6), chromium(VI) will be an oxidising agent. In acidic solution, dichromate is widely used in both organic (oxidation of alcohols) and inorganic chemistry. It can also be used as a volumetric reagent but with special indicators as its colour change (orange to green) makes the end point hard to observe. Cr2O72-(aq) + 14H+(aq) + 6e   > 2Cr3+(aq) + 7H2O(l) [ E = +1.33 V ] orange green " Its E value is lower than that of Cl2 (1.36V) so can be used in the presence of Cl ions " MnO4 (E = 1.52V) oxidises chloride in HCl so must be acidified with sulphuric acid " chromium(VI) can be reduced back to chromium(III) using zinc in acid solution6GUffffff fffffffffff'( : t>  2  -       B  s *޽h ? ̙33y___PPT10Y+D=' = @B +X  0L0 g _ 0 l (  ldB l <D)P0Q l 0Pp@ 0 0dB l@ <D)PXP l 0X0@ 0 0 l 0Dq~ iREACTIONS OF MANGANESE(VII)2(2ff  l 00"`"  ^" in its highest oxidation state therefore Mn(VII) will be an oxidising agent " occurs in the purple, tetraoxomanganate(VII) (permanganate) ion (MnO4) " acts as an oxidising agent in acidic or alkaline solution acidic MnO4(aq) + 8H+(aq) + 5e   > Mn2+(aq) + 4H2O(l) E = + 1.52 V N.B. Acidify with dilute H2SO4 NOT dilute HCl alkaline MnO4(aq) + 2H2O(l) + 3e   > MnO2(s) + 4OH(aq) E = + 0.59 V >O @ffffffffff ffff  fffffffffff fffZ,  6    f B l s *޽h ? ̙33y___PPT10Y+D=' = @B + 0L0 @ %( w dB  <D)P0Q  0Pp@ 0 0dB @ <D)PXP  0X0@ 0 0  0`Ŕq~ o!VOLUMETRIC USE OF MANGANATE(VII)2"(2!ff:   05 >Potassium manganate(VII) in acidic (H2SO4) solution is extremely useful for carrying out redox volumetric analysis. MnO4(aq) + 8H+(aq) + 5e   > Mn2+(aq) + 4H2O(l) E = + 1.52 V It must be acidified with dilute sulphuric acid as MnO4 is powerful enough to oxidise the chloride ions in hydrochloric acid. It is used to estimate iron(II), hydrogen peroxide, ethanedioic (oxalic) acid and ethanedioate (oxalate) ions. The last two titrations are carried out above 60C due to the slow rate of reaction. No indicator is required; the end point being the first sign of a permanent pale pink colour. Iron(II) MnO4(aq) + 8H+(aq) + 5Fe2+(aq)   > Mn2+(aq) + 5Fe3+(aq) + 4H2O(l) this means that moles of Fe2+ = 5 moles of MnO4 1}%  0ffffffffff ffffO Effffffffff fff fffffff  #  t    B         x RB   s *Dop  p RB   s *Dop p B  s *޽h ? ̙33y___PPT10Y+D=' = @B +q 0L0 x@ (  dB  <D)P0Q  0Pp@ 0 0dB @ <D)PXP  0X0@ 0 0  0q~ cREACTIONS OF IRON(II)2(2ff"   0x"`0 :When iron reacts with acids it gives rise to iron(II) (ferrous) salts. Aqueous solutions of such salts contain the pale green, octahedral hexaaquairon(II) ion OH [Fe(H2O)6]2+(aq) + 2OH(aq)   > [Fe(OH)2(H2O)4](s) + 2H2O(l) pale green dirty green ppt. it only re-dissolves in very conc. OH but... it slowly turns a rusty brown colour due to oxidation by air to iron(III) increasing the pH renders iron(II) unstable. Fe(OH)2(s) + OH(aq)   > Fe(OH)3(s) + e dirty green rusty brown NH3 Iron(II) hydroxide precipitated, insoluble in excess ammonia CO32- Off-white coloured iron(II) carbonate, FeCO3, precipitatedHY}P#}#}fffffff ffffffffff ffff#ffff f#>+ -  W  u  s        !    6M F  OA-B(2    6PRUFJ OA-B(2    6UF  LOX 2  ^B  6DjJ p8   6Z F [Ppt 2  ^B  6DjJZ p8Z B  s *޽h ? ̙33y___PPT10Y+D=' = @B +J  0L0 Y Q P  (  dB  <D)P0Q  0Pp@ 0 0dB @ <D)PXP  0X0@ 0 0  # l\d̙ ?A >  HCONTENTS$ (2  <̙? x @ 0 0  04iq~ cREACTIONS OF IRON(II)2(2ffp   0o0 LVolumetric Iron(II) can be analysed by titration with potassium manganate(VII) in acidic (H2SO4) solution. No indicator is required. MnO4(aq) + 8H+(aq) + 5Fe2+(aq)   > Mn2+(aq) + 5Fe3+(aq) + 4H2O(l) this means that moles of Fe2+ = 5 moles of MnO4 1 z  R  'ffffffffff ffffffffff*   &@    RB   s *Do RB   s *DoB  s *޽h ? ̙33y___PPT10Y+D=' = @B +K  0L0 Z R ` t(  tdB t <D)P0Q t 0Pp@ 0 0dB t@ <D)PXP t 0X0@ 0 0 t 0q~ eREACTIONS OF IRON(III)2(2ff  t 0"`0a $Aqueous solutions contain the yellow-green, octahedral hexaaquairon(III) ion OH [Fe(H2O)6]3+(aq) + 3OH(aq)   > [Fe(OH)3(H2O)3](s) + 3H2O(l) yellow rusty-brown ppt. insoluble in XS .-}s. fffffff ffffffffff ffff-@7  k     t 6좕"`~  OA-B(2  ^B t 6DjJp8B t s *޽h ? ̙33y___PPT10Y+D=' = @B +D 0L0 SK (  dB  <D)P0Q  0Pp@ 0 0dB @ <D)PXP  0X0@ 0 0  0@ɕq~ eREACTIONS OF IRON(III)2(2ffv  0 ڕ0   Aqueous solutions contain the yellow-green, octahedral hexaaquairon(III) ion OH [Fe(H2O)6]3+(aq) + 3OH(aq)   > [Fe(OH)3(H2O)3](s) + 3H2O(l) yellow rusty-brown ppt. insoluble in XS CO32- 2[Fe(H2O)6]3+(aq) + 3CO32-(aq)   > 2[Fe(OH)3(H2O)3](s) + 3H2O(l) + 3CO2(g) rusty-brown ppt. The carbonate is not precipitated but the hydroxide is; the high charge density of M3+ makes the solution too acidic to form a carbonate CARBON DIOXIDE EVOLVED. -}|s. fffffff ffffffffff ffff-fffffffffffffffffffffffffffX NZ7  k       6~  OA-B(2     6[~ P OA-B(2  ^B   6DjJp8^B   6DjJ p8 B  s *޽h ? ̙33y___PPT10Y+D=' = @B + 0L0  ^( w dB  <D)P0Q  0Pp@ 0 0dB @ <D)PXP  0X0@ 0 0  04ݕq~ eREACTIONS OF IRON(III)2(2ff  0+0  h Aqueous solutions contain the yellow-green, octahedral hexaaquairon(III) ion OH [Fe(H2O)6]3+(aq) + 3OH(aq)   > [Fe(OH)3(H2O)3](s) + 3H2O(l) yellow rusty-brown ppt. insoluble in XS CO32- 2[Fe(H2O)6]3+(aq) + 3CO32-(aq)   > 2[Fe(OH)3(H2O)3](s) + 3H2O(l) + 3CO2(g) rusty-brown ppt. The carbonate is not precipitated but the hydroxide is; the high charge density of M3+ makes the solution too acidic to form a carbonate CARBON DIOXIDE EVOLVED. NH3 [Fe(H2O)6]3+(aq) + 3NH3(aq)   > [Fe(OH)3(H2O)3](s) + 3NH4+(aq) rusty-brown ppt. insoluble in XS -}|s/i. fffffff ffffffffff ffff-fffffffffffffffffffffffffffX Nfffffff fffffffffff fffff !t7  k         6/~  OA-B(2     6@[~ P OA-B(2     6|N ~ C  OA-B(2  ^B   6DjJp8^B   6DjJ p8 ^B  6DjJ p8 B  s *޽h ? ̙33y___PPT10Y+D=' = @B +c 0L0 rj ( w  dB   <D)P0Q   0Pp@ 0 0dB  @ <D)PXP   0X0@ 0 0   0䅖q~ eREACTIONS OF IRON(III)2(2ff   00 X\Aqueous solutions contain the yellow-green, octahedral hexaaquairon(III) ion OH [Fe(H2O)6]3+(aq) + 3OH(aq)   > [Fe(OH)3(H2O)3](s) + 3H2O(l) yellow rusty-brown ppt. insoluble in XS CO32- 2[Fe(H2O)6]3+(aq) + 3CO32-(aq)   > 2[Fe(OH)3(H2O)3](s) + 3H2O(l) + 3CO2(g) rusty-brown ppt. The carbonate is not precipitated but the hydroxide is; the high charge density of M3+ makes the solution too acidic to form a carbonate CARBON DIOXIDE EVOLVED. NH3 [Fe(H2O)6]3+(aq) + 3NH3(aq)   > [Fe(OH)3(H2O)3](s) + 3NH4+(aq) rusty-brown ppt. insoluble in XS SCN [Fe(H2O)6]3+(aq) + SCN(aq)   > [Fe(SCN)(H2O)5]2+(aq) + H2O(l) blood-red colour Very sensitive; BLOOD RED COLOUR confirms Fe(III). No reaction with Fe(II) -}|s/iwN. fffffff ffffffffff ffff-fffffffffffffffffffffffffffX Nfffffff fffffffffff fffff !fffffff fffffffffffff 7  k      1  }        6D~  OA-B(2     6[~ P OA-B(2     6p ~ e LLS 2     6 N ~ C  OA-B(2  ^B   6DjJp8^B   6DjJ p8 ^B   6DjJ p8 B   s *޽h ? ̙33y___PPT10Y+D=' = @B + 0L0 p p~(  pdB p <D)P0Q p 0Pp@ 0 0dB p@ <D)PXP p 0X0@ 0 0 p 0Tq~ dREACTIONS OF SILVER(I)2(2ffR  p 0"`0  j " aqueous solutions contains the colourless, linear, diammine silver(I) ion " formed when silver halides dissolve in ammonia eg AgCl(s) + 2NH3(aq)   > [Ag(NH3)2]+(aq) + Cl(aq) [Ag(SO3)2]3- Formed when silver salts are dissolved in sodium thiosulphate "hypo" solution. Important in photographic fixing. Any silver bromide not exposed to light is dissolved away leaving the black image of silver as the negative. AgBr + 2S2O32-   > [Ag(S2O3)2]3- + Br [Ag(CN)2] Formed when silver salts are dissolved in sodium or potassium cyanide the solution used for silver electroplating [Ag(NH3)2]+ Used in Tollen s reagent (SILVER MIRROR TEST) Tollen s reagent is used to differentiate between aldehydes and ketones. Aldehydes produce a silver mirror on the inside of the test tube Formed when silver halides dissolve in ammonia - TEST FOR HALIDESR7} #x'&f:ff ffffffffff ff fffffffffffff fM,fGfOAD;    ?    6  D       %  *          B p s *޽h ? ̙33y___PPT10Y+D=' = @B +p  0L0  w  x (  xdB x <D)P0Q x 0Pp@ 0 0dB x@ <D)PXP x 0X0@ 0 0 x 0Deq~ cREACTIONS OF VANADIUM2(2ff  x 0l"`06  |Reduction using zinc in acidic solution shows the various oxidation states of vanadium. Vanadium(V) VO2+(aq) + 2H+(aq) + e   > VO2+(aq) + H2O(l) yellow blue Vanadium(IV) VO2+(aq) + 2H+(aq) + e   > V3+(aq) + H2O(l) blue blue/green Vanadium(III) V3+(aq) + e   > V2+(aq) blue/green lavender%}U3}64}X fffffffff% fffffffff3fffff4  B x s *޽h ? ̙33y___PPT10Y+D=' = @B +< 0L0 KC | (  |dB | <D)P0Q | 0Pp@ 0 0dB |@ <D)PXP | 0X0@ 0 0 | 0ԕq~ !OXIDATION & REDUCTION - A SUMMARYD"(2f f  | 0L"`F  * Oxidation " complex transition metal ions are stable in acid solution " complex ions tend to be less stable in alkaline solution " in alkaline conditions they form neutral hydroxides and/or anionic complexes " it is easier to remove electrons from neutral or negatively charged species " alkaline conditions are usually required e.g. Fe(OH)2(s) + OH(aq)   > Fe(OH)3(s) + e Co(OH)2(s) + OH(aq)   > Co(OH)3(s) + e 2Cr3+(aq) + 3H2O2(l) + 10OH(aq)   > 2CrO42-(aq) + 8H2O(l) " Solutions of cobalt(II) can be oxidised by air under ammoniacal conditions [Co(NH3)6]2+(aq)   > [Co(NH3)6]3+(aq) + e v Ds@}} C fff ffffffff fffffffffffff ff ffff fffffJffffffff fffffff ftf  3  x      @ B | s *޽h ? ̙33y___PPT10Y+D=' = @B + 0L0  $(  $dB $ <D)P0Q $ 0Pp@ 0 0dB $@ <D)PXP $ 0X0@ 0 0 $ 0q~ !OXIDATION & REDUCTION - A SUMMARYD"(2f f $ 0H F t 4Oxidation " complex transition metal ions are stable in acid solution " complex ions tend to be less stable in alkaline solution " in alkaline conditions they form neutral hydroxides and/or anionic complexes " it is easier to remove electrons from neutral or negatively charged species " alkaline conditions are usually required e.g. Fe(OH)2(s) + OH(aq)   > Fe(OH)3(s) + e Co(OH)2(s) + OH(aq)   > Co(OH)3(s) + e 2Cr3+(aq) + 3H2O2(l) + 10OH(aq)   > 2CrO42-(aq) + 8H2O(l) " Solutions of cobalt(II) can be oxidised by air under ammoniacal conditions [Co(NH3)6]2+(aq)   > [Co(NH3)6]3+(aq) + e Reduction " Zinc metal is used to reduce transition metal ions to lower oxidation states " It acts in acid solution as follows... Zn(s)   > Zn2+(aq) + 2e e.g. it reduces iron(III) to iron(II) vanadium(V) to vanadium (IV) to vanadium(III)@ Ds@}}se C fff ffffffff fffffffffffff ff ffff fffffJffffffff fffffff f zff fff fef  3  x        7       $ ^B $ 6Do p8 B $ s *޽h ? ̙33y___PPT10Y+D=' = @B +2 0L0 A9(  dB  <D)P0Q  0Pp@ 0 0dB @ <D)PXP  0X0@ 0 0  0|Fq~ dREACTIONS OF ALUMINIUM2(2ff   0_"` z" aluminium is not a transition metal as it doesn t make use of d orbitals " BUT, due to a high charge density, aluminium ions behave as typical M3+ ions " aqueous solutions contain the colourless, octahedral hexaaquaaluminium(III) ion OH [Al(H2O)6]3+(aq) + 3OH(aq)   > [Al(OH)3(H2O3](s) + 3H2O(l) colourless, octahedral white ppt. soluble in XS NaOH As with all hydroxides the precipitate reacts with acid [Al(OH)3(H2O)3](s) + 3H+ (aq)   > [Al(H2O)6]3+(aq) being a 3+ hydroxide it is AMPHOTERIC and dissolves in excess alkali [Al(OH)3(H2O)3](s) + 3OH(aq)   > [Al(OH)6]3-(aq) + 3H2O(l) colourless, octahedral CO32- 2 [Al(H2O)6]3+(aq) + 3CO32-(aq)   > 2[Al(OH)3(H2O)3](s) + 3H2O(l) + 3CO2(g) As with 3+ ions, the carbonate is not precipitated but the hydroxide is. NH3 [Al(H2O)6]3+(aq) + 3NH3(aq)   > [Al(OH)3(H2O)3](s) + 3NH4+(aq) white ppt. insoluble in XS NH3iTu}BJ}O&}`LZ} '- ffffff&fffffffff:9fffffffffffffffJffffff$fffffff&ffffff fffffffff ff fffMffffff ffffffff fffff&  C    ~    \        6te"`8 OA-B(2     68"`8 OA-B(2     6Ě"`8  OA-B(2     6Dǚ"` 8  OA-B(2    6\˚"`^8S OA-B(2  ^B  6DjJ p8 ^B  6DjJ p8 B  s *޽h ? ̙33y___PPT10Y+D=' = @B + 0L0 % @(   X @ @ <̙? x @D 0 0 @ 0 ښx m i/ 2009 JONATHAN HOPTON & KNOCKHARDY PUBLISHING0(20 f @ 0KfM  ATHE END(2f @ 0X -AN INTRODUCTION TO TRANSITION METAL COMPLEXEST$(2 82 f0f<fJ  @ C "A KTRE H @ 0޽h ? ̙33y___PPT10Y+D=' = @B +r@^0%Y+ l@\6  "4:1r<w}C`lU Y,?L`rk~HOW`%И8RckwUJXvѓ_$p19p>n`gYj|w!sOh+'0T hp   No Slide TitleHOPTONJONATHAN HOPTON995Microsoft PowerPoint@0$@@hNx@ 8GSg  )'    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