Module 5 Chemistry > Chapter 24- Transition Elements > Flashcards

Chapter 24- Transition Elements Flashcards

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1
Q

what are the d block elements

A

they are the elements located between group 2 and group 13

they have their highest energy electron in a D-subshell

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2
Q

what are 4 physical properties that they all have

A
  • high mpt and bpt
  • metallic
  • shiny
  • electrically conductive
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3
Q

what is the important to note about the way that the d block elements fill their electron shells

A

the 4s subshell fills before the 3d subshell as it ends up being at a lower energy level

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4
Q

how do the d block elements lose their electrons

A

outside inwards, the 4s electrons are lost before the 3d electrons

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5
Q

define what a transition element is

A

“Transition elements are d-block elements which form one or more stable ions which have an incompletely filled d subshell”

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6
Q

why are Zn2+ and Sc3+ not transition elements

A

zinc only forms a 2+ ion, this 2+ ion has a completely filled 3d10 subshell - this breaks both rules

Scandium only forms a 3+ ion, this has a completely empty 3d0 subshell- this breaks both rules

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7
Q

what are the exceptions on the filling rules for electron shells

A

Chromium and Copper:
Cr: 1s2, 2s2, 2p6, 3s2, 3p6, 4s1, 3d5
Cu: 1s2, 2s2, 2p6, 3s2, 3p6, 4s1, 3d10

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8
Q

why do these exceptions on electron configuration occur

A

it’s more stable to have:

  • half filled 4s and half filled 3d
  • half filled 4s and filled 3d
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9
Q

name the three key properties of transition metals

A
  • variable oxidation states
  • good catalysts
  • colourful compounds
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10
Q

what is a species with a transition element in its highest oxidation state usually

A
  • a good oxidising agent
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11
Q

what does the system ‘want’ to happen when forming an ionic compound

A
  • it’s most stable if more energy is released

- think lattice enthalpy and Born-Haber cycles

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12
Q

What are the two main processes we should consider when explaining why transition elements have variable oxidation states

A
  • lattice enthalpy/enthalpy of hydration– a greater charge will mean a greater lattice enthalpy- more energy released
  • energy required to ionise the element– a greater charge will mean more energy is required to remove/add electrons - think electron affinities and ionisation energies
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13
Q

what is the toss-up when considering variable oxidation states

A
  • forming a higher oxidation state means more energy is required to form a higher charged ion
    BUT
  • this means more energy is released on forming the compound or dissolving
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14
Q

Why does this toss-up not have as much of an effect with the transition elements leading to multiple oxidation states

A
  • The 4s and 3d orbitals have very similar energies
  • There isn’t a huge jump in the amount of energy required to remove a 3rd election
  • then any further ionisations in the d orbital don’t require much energy
  • The net effect of this is it often doesn’t matter which oxidation state occurs in a compound, the net energy release is similar
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15
Q

what is a good phrase to learn when explaining variable oxidation states

A

” transition elements have variable oxidation states because the energy difference between the 4s and 3d subshells is small, this means different numbers of electrons can be lost or gained for a similar net energy change”

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16
Q

Give 3 examples of colourful transition element compounds and why (briefly) does this occur

A
  • K2Cr2O7 = orange
  • CoCl2 = pink-purple
  • NiSO4 = green
  • the reason for these bright colours is linked to partially filled d subshells
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17
Q

define a complex ion

A

“A complex ion is a metal ion surrounded by coordinately bonded ligands”

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18
Q

Define what a ligand is

A

“A ligand is a molecule or Ion that donates a pair of electrons to a central metal ion to form a coordinate bond. They have an active lone pair in their highest energy level”

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19
Q

define coordination number

A

“The coordination number is the number of coordinate bonds with the metal ion”

NOT the number of ligands

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20
Q

what to remember when drawing complex ions

A

the bonds are coordinate, make them arrows

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21
Q

how do we work out the charge on a complex ion

A

charge on the complex ion = charge on the metal ion + (sum of) charges on ligands

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22
Q

what is the shape and angles of a complex ion with a coordination number of 6 and which ligands does this usually occur with

A
  • octahedral shape
  • all angles are 90 degrees
  • usually occurs with small ions e.g. H2O, CN-, NH3
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23
Q

what is the shape of a complex ion with a coordination number of 2 and what are the examples of when this happens

A
  • forms a linear shape
  • rare
  • occurs only with Cu+ and Ag+ because they’re small ions
  • angles are 180 degrees
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24
Q

what is the most common shape for complex ions with a coordination number of 4, why does this happen

A
  • tetrahedral
  • angles of 109.5 degrees
  • usually occurs with bigger ligands e.g. Cl-
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25
what is the less common shape for complex ions with a coordination number of 4, why does this occur
- square planar - four 90 degree angles - occurs only with Ni, Pt, Pd(II), Au(III)
26
what is a monodentate ligand
- a ligand which only forms 1 coordinate bond with the central metal ion - occasionally they have more than 1 lone pair e.g. H2O but this still only forms one bond
27
give some examples of monodentate ligands
H2O, NH3, Cl-, CN-, OH-
28
what is a bidentate ligand
- a ligand which forms 2 coordinate bonds with the central metal ion
29
give 2 examples of bidentate ligands
1,2, diaminoethane (NH2)(CH2)(CH2)(NH2) ethandioate ion COOCOO
30
define stereoisomerism
"stereoisomers are molecules with the same structural formula but different arrangements in space"
31
what are the 2 possible types of stereoisomerism in complex ions
- cis/trans isomerism | - optical isomerism
32
what is an optical isomer
"Optical isomers are non-superimposible mirror images of each other which contain no lines of symmetry"
33
when are optical isomers able to happen
they only occur where there are: - an octahedral molecule - 2 bidentate ligands and 2 monodentate ligands - 3 bidentate ligands - only occurs in the cis formation of the isomer
34
when can cis/trans isomerism occur
- only in square planar and octahedral molecules - only if there are 2 of the same type of ligand - it cannot occur in tetrahedral molecules
35
what is the cis isomer and what is the trans isomer in complex ions
- the cis isomer occurs where the 2 identical ligands are at 90 degrees to each other - the trans isomer occurs where the 2 identical ligands are at 180 degrees to each other
36
what are the key features of optical isomers
- non-superimposable - mirror images - no lines of symmetry - one isomer rotates plane polarised light one way and the other isomer rotates the light the other way
37
what is a ligand substitution reaction
- a reaction where one or more the ligands of a complex ion are substituted for ligands of a different type
38
what 3 things does Cu(II) (aq) mainly react with
- NH3, Cl-, NaOH
39
what is the overall equation for Cu(II) (aq) with an excess of NH3 and what is the colour change
with an excess of ammonia: ([Cu(H2O)6] aq 2+) + 4NH3(aq) ---> 4H2O(aq) + [Cu(NH3)4(H2O)2]2+(aq) pale blue to deep blue
40
what is the split version of the reaction of Cu(II) (aq) with NH3
1) a pale blue precipitate is formed as the ion loses its charge and is no longer soluble, here the ammonia is acting as a base and accepts the hydrogen ions from 2 H2O ligands ([Cu(H2O)6]2+aq) + 2NH3 ---> [Cu(OH)2(H2O)4](S) + 2(NH4+) 2) the Cu(OH)2 precipitate dissolves in excess ammonia to form a dark blue [Cu(OH)2(H2O)4](s) + 4NH3 ---> ([Cu(NH3)4(H2O)2]2+aq) + 2H2O + 2OH-
41
what is the ligand substitution reaction of Cu(II) (aq) with Cl- and colour change
([Cu(H2O)6]2+aq) + 4Cl-(aq) [CuCl4]2-(aq) + 6H2O(l) - pale blue to yellow - this reaction is in dynamic equilibrium and will react as such, colour is therefore often green - initially the molecule is octahedral, after it's tetrahedral
42
what must we remember when doing ligand substitution reaction with chloride ions
use conc. HCl not dilute HCl, otherwise the water content pushes equilibrium back in the other direction
43
how does Cr(III) (aq) react with NH3 overall and include colour changes
overall ([Cr(H20)6]3+aq) + 6NH3(aq) ---> ([Cr(NH3)6]3+aq) + 6(H2O)(l) -purple to purple with grey precipitate
44
what is the split reaction for Cr(III) aq with NH3
1) ammonia acts as a base in the first stage as before, taking hydrogen ions from the water molecules ([Cr(H2O)6]3+aq) + 3NH3(aq) ---> [Cr(OH)3(H2O)3](s) + 3NH4+(aq) 2) ([Cr(OH)3(H2O)3])(s) + 6NH3(aq) ---> ([Cr(NH3)6]3+aq) + 3(H2O) + 3OH-
45
how does Cobalt (II) (aq) react with Cl-
([Co(H2O)6]aq2+) + 4Cl- ([CoCl4](aq)2-) + 6H2O(l) - pink to blue - octahedral to tetrahedral - still use conc. HCl and not dilute HCl
46
how does ligand substitution occur in the blood
- blood can carry oxygen due to haemoglobin - haemoglobin contains 4 protein chains held together by weak intermolecular forces - this contains an Fe2+ ion which can bind to both O2 and CO2 - CO can also form even stronger bonds with it and this cannot be reversed
47
how does Cu2+ react with NaOH and and excess
- first forms light blue precipitate of Cu(OH)2 Cu2+(aq) + 2OH-(aq) ---> Cu(OH)2(S) - doesn't react in an excess
48
how does Fe 2+ react with NaOH and an excess
- pale green solution to green precipitate of Fe(OH)2 Fe2+(aq) + 2OH-(aq) ---> Fe(OH)2(S) - insoluble in excess but does oxidise to Fe(OH)3
49
how does Fe 3+ react with NaOH and an excess
- pale yellow solution to orange-brown precipitate of Fe(OH)3 Fe3+(aq) + 3OH-(aq) ---> Fe(OH)3(S) - insoluble in excess
50
how does Mn 2+ react with NaOH and an excess
- pale pink solution to light brown precipitate of Mn(OH)3 Mn2+(aq) + 2OH-(aq) ---> Mn(OH)2(S) - insoluble in excess
51
how does Cr 3+ react with NaOH and an excess
- Violet solution to green-grey precipitate of Cr(OH)3 Cr3+(aq) + 3OH-(aq) ---> Cr(OH)3(S) - soluble in excess NaOH to form a dark green solution Cr(OH)3(s) + 3OH- ---> [Cr(OH)6]3-(aq)
52
how does Co 2+ react with NaOH and an excess
- Pink solution forms Blue precipitate Co2+(aq) + 2OH-(aq) ---> Co(OH)2(S) - soluble in XS NaOH to form a darker blue Co(OH)2(s) + 4OH-(aq) ---> [Co(OH)6]4-(aq)
53
what are the only 2 metal hydroxide's in these examples to redissolve in an excess of NH3
Cr, and Cu | the others won't
54
what colour does the main oxidation state of Titanium have in aqueous solution
Ti3+ = Lilac
55
what colours are the main oxidation states of Chromium in solution
``` Cr2+ = Blue Cr3+ = Green Cr6+ = Orange ```
56
what colours are the main oxidation states of Vanadium in aqueous solution
``` V2+ = purple V3+ = Green V4+ = Blue V5+ = yellow ```
57
What colours are the main oxidation states of Manganese in aqueous solutions
``` Mn2+ = pink Mn7+ = purple ```
58
what colours are the main Oxidation states of Iron in aqueous solutions
``` Fe2+ = green Fe3+ = Yellow or orange-brown ```
59
what colour is the main oxidation state of cobalt in solution
Co2+ = pink
60
what colour is the main oxidation state of Nickel in aqueous solution
Ni2+ = Green
61
What colour is the main oxidation state of Copper in aqueous solution
Cu2+ = blue
62
why does zinc not tend to form coloured ionic solutions
it is not a transition metal
63
what is the redox reaction equation for the oxidation of Fe2+ to Fe3+ with manganate(VII) to (II) and what is the colour change
MnO4- + 8H+ + 5Fe2+ ---> Mn2+ + 5Fe3+ + 4H2O - the Mn is reduced from +7 to +2 - the Fe is oxidised from +2 to +3 colour change of purple to colourless
64
what is the redox reaction equation and colour change for the reduction of Fe3+ to Fe2+
2Fe3+ + 2I- ---> 2Fe2+ + I2 - Fe3+ is reduced to Fe2+ - I- is oxidised to I2 colour change of orange-brown to pale green and brown
65
what is an explanation in terms of electrode potentials for why these redox reactions of iron take place
``` redox system (Iodide) = +0.54V redox system (Iron) = +0.77V redox system (manganate) = +1.33V ``` so as the standard electrode potential becomes more positive they become better oxidising agents so can gain electrons more easily so manganate will always gain electrons forcing iron to lose electrons (oxidise) but iron will gain electrons more easily than iodine so it will reduce with iodine b
66
what is the redox equation for the reduction of Cr2O7(2-) to Cr(3+) with zinc, what is the colour change
Cr2O7(2-) + 14H(+) + 3Zn ---> 2Cr(3+) + 7H2O + 3Zn(2+) - Cr is reduced from +6 to +3 - Zn is oxidised from 0 to +2 the colour change is orange to green (think acidified potassium dichromate)
67
what happens when Cr(3+) is added with an excess of zinc, equation, redox and colour change
Zn + 2Cr(3+) ---> Zn(2+) + 2Cr(2+) - Cr is reduced from +3 to +2 - Zn is oxidised from 0 to +2 colour change is green to pale blue
68
explain in terms of standard electrode potentials why these reductions of Chromium take place
``` redox system (Zn) = -0.76V redox system (Cr3+) = -0.41V redox system (Cr2O7(2-)) = +1.33V ``` - they become more positive and thus better oxidising agents as you go down, therefore they are better at gaining electrons so Cr2O7(2-) will gain electrons when with zinc to reduce the Chromium and oxidise the zinc and Cr3+ will gain electrons when with zinc to reduce to Cr2+ and oxidise zinc
69
what is the redox reaction for the oxidation of Cr3+ to CrO42- with hydrogen peroxide
3H2O2 + 2Cr(3+) + 10OH- ---> 2CrO4(2-) + 8H2O - oxygen is reduced from 1- to 2- - Cr is oxidised from +3 to +6 this is because H2O2 is a powerful oxidising agent
70
what is the disproportionation reaction of Cu+ ions, what are the conditions and colour change
Cu2O(s) + H2SO4 ---> Cu(s) + CuSO4 (aq) + H2O - hot dilute H2SO4 - colour change of colourless to blue solution with brown solid - one Cu+ is oxidised to Cu2+ - one Cu+ is reduced to Cu
71
what is the reduction of Cu2+ ions to Cu+ ions with iodine
2Cu2+ + 4I- ---> 2CuI + I2(s) - Cu2+ is reduced to Cu+ - I- is oxidised to I2 colour change of pale blue to brown with a white precipitate
72
what to remember if there's any mention of hydrochloric acid with a hydrated copper ion
[Cu(Cl)4]2-(aq)
73
how to write salts where one of the ions is a complex ion
waters of crystallisation (usually) e.g. Cr2(SO4)3 ¬ 6H2O NOT ([Cr(H2O)6]3+)2(SO42-)3
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