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A2 CHEM UNIT3 > 3.4 D-block > Flashcards

3.4 D-block Flashcards

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1
Q

Define a d-block element

A
  • group of elements whose outer electrons are found in d-orbitals
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2
Q

Why do chromium and copper fill their d orbitals before s?

A
  • completely full or half full d-sub shell is more stable than a partially filled d sub level
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3
Q

Define a transition metal

A
  • d-block element that forms a stable ion with an incomplete d-sub shell
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4
Q

What is a transition metal complex?

A
  • central metal ion surrounded by coordinately bonded ligands
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5
Q

Define a ligand

A
  • small molecule with a lone pair that can bond to a transition metal ion
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6
Q

Define a complex ion

A
  • a transition metal bonded to one or more ligands
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7
Q

Define a coordinate bond

A
  • a bind where both electrons are provided for by the same atom
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8
Q

Define the coordinate number

A
  • number of coordinate bonds formed
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9
Q

Features of transition metal complexes

A
  • small with very large positive charges
  • many empty bonding orbitals
  • electron-rich molecules have lone pairs so they can from coordinate bonds with empty orbitals in transition metal ion
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10
Q

Monodentate and bidentate ligands

A

Mono: one atom that can bond to metal ion
Bi: 2 atoms bonding to metal ion (ligand able to donate 2 lone pairs)

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11
Q

[Fe(H2O)6]2+

A

Pale green

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12
Q

[Fe(H2O)6]3+

A

Yellow

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13
Q

[Cr(H2O)6]3+

A

Dark green

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14
Q

[Co(H2O)6]2+

A

Pink

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15
Q

[Cu(H2O)6]2+

A

Blue

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16
Q

[Cu(NH3)4(H2O)2]3+

A

Royal blue

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17
Q

[CoCl4]2-

A

Blue

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18
Q

[CuCl4]2-

A

Yellow green

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19
Q

What are stereoisomers

A
  • pairs of molecules with same formula but different orientation of functional groups in 3D space
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20
Q

Ligand names

A

1 - mono
2 - di
3 - tri
4 - tetra
5 - penta
6 - hexa

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21
Q

Ligand group names

A

H2O - aqua
NH3 - ammine
OH- - hydroxo
Cl- - chloro
F- - fluoro
CN- - cyano

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22
Q

Cis isomer

A

Same side

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23
Q

Trans isomer

A

Diagonal

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24
Q

Where do the colours of transitional metal complexes come from?

A
  • there is repulsion between electrons present in ligands and electrons in d orbitals
  • which destabilises the orbitals causing them to split into 3 low energy orbitals and 2 high energy
  • electrons move from low to high energy gaining energy by the absorption of light to move to higher energy level
  • colours seem are due to unabsorbed frequencies
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25
Q

Properties of d-block metals

A
  • good conductors of heat and electricity
  • high melting and boiling point
  • malleable
  • ductile
  • shiny
  • strong with high tensile and compressive strength
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26
Q

Properties of transition metals

A
  • variable oxidation states
  • have catalytic activity
  • form coloured complexes
  • other metal properties
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27
Q

What is meant by oxidation state?

A
  • hypothetical charge of an atom if all of its bonds to other atoms were fully ionic
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28
Q

Cr3+

A

Green

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29
Q

Cr2O7 2-

A

Orange

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30
Q

CrO4 2-

A

Yellow

31
Q

Co2+

A

Pink

32
Q

Fe2+

A

Pale green

33
Q

Fe3+

A

Red brown

34
Q

Cu2+

A

Pale blue

35
Q

Explain the nature of ligands

A
  • different ligands have different effects on the energies of the d-orbitals of the central ion
  • some ligands have strong electrical fields which cause a large energy gap when d-orbitals split
36
Q

Order of smallest to largest split

A

Cl-
F-
OH-
H2O
NH3
CN-

37
Q

Impact of oxidation of metal on transition metal complexes

A
  • as oxidation increases so does the amount of splitting of the d-orbital
  • changes of oxidation state change the colour of light absorbed so therefore colour of light seen
38
Q

Explain how the coordination of the ion impacts its splitting

A
  • splitting is greater in octahedral than tetrahedral
  • only normally change coordination if change ligand, which will change colour too
    —> change cannot be isolated
39
Q

Explain ligand exchange with an example

A
  • when a transition metal is exposed to a mixture of ligands they can be exchanged without impacting coordination number

[Co(H2O)6]2+ + 6NH3 —> [Co(NH3)6]2+ + 6H2O

(Similar sizes exchange)

40
Q

Explain ligands and Le Chatelier

A
  • addition of ammonia forces equilibria to the right
  • addition of water forces equilibria left
  • leads to changes in geometry
41
Q

Explain the addition of ammonia to Cu2+

A
  • excess ammonia causes ligand substitution
    —> ppt dissolve in solution and ammonia exchange with water ligands
    —> add sulphuric acid to reverse
  • small amount of ammonia to act as a base and deprotonate some water ligands forming a ppt
42
Q

Explain addition of HCl to Cu2+

A
  • ligand substitution reactiin
43
Q

Explain addition of ammonia to Cu2+

A
  • excess results in ligand substitution reaction
    —> ppt dissolves and ammonia exchange with water

Copper compound + ammonia = pale blue to royal blue

44
Q

Colour change when cobalt solution is oxidised

A
  • rapid darkening of straw coloured to deep red brown when cobalt solution oxidised
45
Q

Examples of transition metal catalysts

A
  • iron in the Haber process
  • nickel to make margarine in the dehydrogenation of vegetable oils
  • vanadium oxide in the contact process
  • manganese oxide in the catalytic decomposition of hydrogen peroxide
46
Q

Define a homogenous catalyst and how they work

A
  • same physical state as reactions
    —> use variable oxidation states to oxidise/reduce a reactant, making it more reactive
    —> the transition metal can then be converted back to its original oxidation state by reaction with another molecule
47
Q

Define a heterogenous catalyst and how they work

A
  • different physical state to reactants
    —> partially filled d-orbitals
    —> catalyst provides a solid surface on which reactants can be absorbed and brought closer together for more opportunity to react
    —> molecules with lone pairs can form coordinate binds to the metal atom because there are available empty d-orbitals to bond ti
48
Q

Reactions of Cr3+ with OH- ions

A
  • green
  • grey-green ppt when OH- added
  • ppt dissolve to deep green solution when excess OH- added
49
Q

Reaction of Fe2+ with OH- ions

A
  • pale green
  • dark green ppt when OH- added
  • no further change when excess OH- added
50
Q

Reaction of Fe3+ with OH- ions

A
  • yellow
  • red-brown ppt when OH- added
  • no further change when excess OH- added
51
Q

Reaction of Cu2+ with OH- ions

A
  • pale blue
  • pale blue ppt when OH- ions added
  • no further change when excess OH- ions added
52
Q

Ionic equation for Cr3+ with OH- ions

A

Cr3+ (aq) + 3OH- (aq) —> Cr(OH)3 (s)

Cr(OH)3 (s) + 3OH- (aq) —> [Cr(OH)6]3- (aq)

53
Q

Ionic equation for Fe2+ with OH-

A

Fe2+ (aq) + 2OH- (aq) —> Fe(OH)2 (s)

54
Q

Ionic equation for Fe3+ with OH-

A

Fe3+ (aq) + 3OH- (aq) —> Fe(OH)3 (s)

55
Q

Ionic equation for Cu2+ with OH-

A

Cu2+ (aq) + 2OH- (aq) —> Cu(OH)2 (s)

56
Q

Reaction of Mn2+ with OH-

A
  • pale pink solution to off white ppt
  • does not dissolve in excess
57
Q

Zn2+ reaction with OH-

A
  • colourless solution to white ppt
  • ppt dissolves in excess to give colourless solution
58
Q

Cr3+ with NH3

A
  • blue/purple solution to muddy green ppt
59
Q

Mn2+ with NH3

A
  • pale pink solution to off white ppt
  • does not dissolve in excess
60
Q

Fe2+ with NH3

A
  • pale green solution to dark green ppt
  • turns brown by aerial oxidation
61
Q

Fe3+ with NH3

A
  • yellow solution to reddy-brown ppt
62
Q

Ni2+ with NH3

A
  • pale green solution to pale green ppt
  • ppt dissolve in excess to give blue solution
63
Q

Cu2+ with NH3

A
  • blue solution to pale blue ppt
  • ppt dissolves in excess to give deep blue solution
64
Q

Zn2+ with NH3

A
  • colourless solution to white ppt
65
Q

State how enthalpy of formation values give an indication of stability

A
  • the more negative the enthalpy change the more stable the oxide
66
Q

Describe a test to show the presence of iron (III) ions in a solution of Fe2SO4.7H2O

A
  • add aqueous NaOH
  • give red brown ppt
67
Q

Describe how aqueous sodium hydroxide can be used to distinguish between aqueous iron (II) sulphate and iron (II) sulphate

A
  • iron (II) ions give green ppt
  • iron (III) ions give red brown ppt
68
Q

Construct the equation for the oxidation of acidified iron (II) ions by oxygen

A

4Fe2+ + O2 + 4H+ —> 4Fe3+ + 2H2O

69
Q

Explain why the H-O-H bind angle in the water ligand is 107° rather than 104.5°

A
  • water molecules have 2 lone pairs and 2 bond pairs
  • water ligands have 1 lone pair and 3 bond pairs
  • lone pairs repel more than bond
70
Q

Explain why the complex ions [Fe(H2O)6]3+ and [Fe(H2O)5(OH)]2+ are not the
same colour.

A

Different ligands cause different splitting therefore different frequencies of light are absorbed

71
Q

Suggest how you would select an appropriate wavelength to find the
concentration of [Fe(H2O)6]3+ in the equilibrium mixture.

A

Find the wavelength absorbed by [Fe(H2O)6] but not [Fe(H2O)5(OH)]

72
Q

Explain why CO is a good reducing agent

A

It has carbon in its 2+ state but its most stable state ins 4+

73
Q

Colour change when aqueous NaOH added to potassium dichromate

A

Orange to yellow

A2 CHEM UNIT3 (9 decks)

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