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Inorganic chemistry > Module 1-6 > Flashcards

Module 1-6 Flashcards

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1
Q

Ground state

A

Most stable electron configuration

Electrons in lowest energy state

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2
Q

Hunds rule

A

For degenerate orbitals, orbitals of same energy level), the lowest energy is obtained when the number of electrons with same spin is maximised

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3
Q

Tranisition metal ions

A

Lewis acids
Cations
Have vacant d orbitals
Electron acceptors

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4
Q

Effective nuclear charge

A

Electrons in inner shell have shielding effect on electrostatic attraction between protons and electrons in outer shells

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5
Q

Zeff trend

A

Increases across the row

Decreases in order of s>p>d>f

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6
Q

Atomic radii

A

Half distance between metallic nuclei
Increases down a column
Decreases across row

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7
Q

Anion atomic radii

A

Larger than parent atom

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8
Q

Cation ionic radii

A

Smaller than parent atoms

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9
Q

Electronegativity

A

Increases across row and decreases down column

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10
Q

Ionisation energy

A

removal of electron from highest occupied orbital of neutral atom

Follows rules same as atomic radii

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11
Q

Chlorophyll

A

Green compound found in leaves and stems

Channels sunlight into chemical energy which drives biochemical reactions through the process of photosynthesis

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12
Q

Chlorophyll structure

A

Tetrapyrrolic ring
Conjugated double bond system
Central magnesium atom

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13
Q

Hard Lewis bases

A

Donor atom is small and highly electronegative

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14
Q

Soft Lewis base

A

Donor atom is larger less electronegative making its electrons more polarisable

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15
Q

Hard Lewis acids

A

Either noble gas configuration or high charge and strongly held electrons

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16
Q

Soft Lewis acids

A

Large number of d electrons and low charge

More polarisable electrons

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17
Q

Ligand field stabilisation energy

A

Pt(2) > Ni(2) > Co (2) > Cu(2) > Fe(2) > Zn (2) > Mg(2)

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18
Q

Lewis bases

A

Substance that act as electron pair donors

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19
Q

Lewis acids

A

Substances that act as electron pair acceptors

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20
Q

Transition metal complexes

A

Complex is a combination of a Lewis acids and Lewis base

Ligand forms coordinate covalent bond to central metal ion

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21
Q

Neutral ligands

A

H20
NH3
CO

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22
Q

Anionic ligand

A 
OH -
Cl-
Br- 
I- 
CN-
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23
Q

Bidentate ligands

A

En

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24
Q

Polydentate

A

EDTA 4-

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25
Q

Ligands

A

Lewis bases

Have atleast one non bonded pair of electrons

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26
Q

Chelate

A

Claw agents

Can bind to metal in two or more places

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27
Q

Counter ions

A

Written on outside of square brackets and not bonded to metal ion
Negatively charged - right
Positively charged - left

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28
Q

Coordination atoms

A

Number of ligand donor atoms bonded to metal ion

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29
Q

Charge of metal complex

A

Sum of charge on metal ion and sum of ligand charges

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30
Q

Oxidation state

A

Charges on metal atom

Complexes containing neutral ligands the oxidation state is equal to the net charge on complex

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31
Q

Linear

A

Coordination number is 2

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32
Q

Tetrahedral or square planar

A

Coordination number is 4

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33
Q

Octahedral

A

Six coordination complexes

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34
Q

Factors influencing stability of metal complexes

A

Irving Williams series of stability
Hard soft acid base theory
Chelate effect

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35
Q

KF

A

Formation constant

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36
Q

Hard acids

A

Small size
High charge
Low electronegativity

37
Q

Soft acid

A

Large size
Low charge
Easily polarisable
High electronegativity

38
Q

Hard bases

A

Don’t give up electrons
High electronegativity
Small donor atoms

39
Q

Soft bases

A

Intermediate electronegativity

Large size

40
Q

K»1

A

More stable state is to the right

41
Q

Chelate effect

A

Ligands that form chelate rings generally bind more strongly to a metal ion then monodentate ligands

42
Q

Chelate effect kF

A

Larger stability so kF is bigger

43
Q

Effect of chelate ring on kF

A

Chelate effect weakens as ring size increases

KF decreases as size increases

44
Q

Octahedral complexes

A

5 membered ring

45
Q

Lobes that point towards negative charge

A

Dz2 dx2-dy2

46
Q

Lobes that point between point charges

A

Dyz
Dxz
Dxy

47
Q

Crystal splitting energy

A

Metal ion placed in middle causes d electrons to experience repulsion and raise the energy level

48
Q

Crystal filed splitting energy

A

Increases with increasing oxidation number

Increases down a group

49
Q

Ligand effect

A

Halide < oxygen < nitrogen < carbon

50
Q

Weak field ligands

A

Smaller splitting energy
High spin
Absorbs linger wavelengths

51
Q

Strong field ligands

A

Larger splitting energy
Low spin
Short wavelengths

52
Q

High spin

A

Max number of unpaired electrons

53
Q

Low spin

A

Max number of parallel spins

54
Q

Irving Williams series of stability

A

Mn 2+ < Fe2+ < Co2+ < Ni2+ < Cu2+ > Zn2+

55
Q

Metal cations in water

A

Dissolve and are hydrated with water molecules to form squares ions

56
Q

Aqua acids

A

Acids where acidic proton in on water molecule attached to metal cation

57
Q

Acidity of aqua acids

A

Increases with increasing positive charge of central atom

Increases with decreasing atomic radius of central atom

58
Q

Weak aqua acids

A

Large cations with low charge

59
Q

Strong aqua acids

A

Small cations with high charge

60
Q

Ka>1

A

Strong acids

61
Q

Ka<1

A

Weak acids

62
Q

Donor atoms - neutral and negative

A

Neutral and negatively charged atoms with lone electrons in side chain at pH 7 can be donor atoms and coordinate to metal cations

63
Q

Donor atoms - positive

A

Positively charged atom in side chain is unable to coordinate to metal

64
Q

Reducing agent

A

Oxidised

65
Q

Oxidising agent

A

Reduced

66
Q

Gibbs free energy

A

G = delta H - temperature x delta S

67
Q

Latimer diagrams

A

Standard reduction potentials for species
Highly oxidised form on left
Reduction potentials for species above line

68
Q

E cell

A

Electrons x e cell + electrons 2 x e cell / electrons 1 + electrons 2

69
Q

Frost diagram

A

Reverse of Latimers
From least oxidised to highly oxidised
Reduction potentials reversed

70
Q

Lower part of frost diagram

A

More thermodynamically stable

71
Q

Trend in reduction potential

A

Cu> heme> Fe-S

72
Q

Acids

A

Generally are metals on periodic table

73
Q

Chelation

A

Prevents iron from being captured by other insoluble compounds

74
Q

Amino acids

A

Oh dependent

Complex can fall apart if you change pH

75
Q

To find y of frost diagram

A

Number of electrons to get to the end x reduction potential to get to most oxidised

76
Q

If reduction potential Is favouring left hand side

A

E cell will become less and less if it’s positive

77
Q

Increasing ph

A

Decreases e cell value

Process is unfavourable

78
Q

Acidity

A

Decreases as atomic radii increases
Smaller radius = stronger acid
Atoms closer to top left of periodic table

79
Q

Acidity - double bonds

A

More double bonds attached to central atom = greater electron withdrawal
= greater acidity

80
Q

Acidity - negative charges

A

More negative charges is stable

Less negative charges is more acidic

81
Q

conjugate acid

A

acid formed from protonating a base

82
Q

conjugate base

A

base formed from deprotonating an acid

83
Q

weak acid strength

A

low positive charge

84
Q

strong acid strength

A

high positive charge

85
Q

aqua acid precipitate in water

A

strong acidic aqua acids form precipitates in water

86
Q

more double bonds

A

withdraws electron density towards the central atom resulting in greater polarisation of OH group, making OH bond weaker

87
Q

more negative atoms

A

requires greater delocalisation over the molecule, reducing stability

88
Q

stability of chelate complexes

A

increased entropy due to increased number of molecules

Inorganic chemistry (2 decks)

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