Module 1-6 Flashcards Preview

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Flashcards in Module 1-6 Deck (88):
1

Ground state

Most stable electron configuration
Electrons in lowest energy state

2

Hunds rule

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

3

Tranisition metal ions

Lewis acids
Cations
Have vacant d orbitals
Electron acceptors

4

Effective nuclear charge

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

5

Zeff trend

Increases across the row
Decreases in order of s>p>d>f

6

Atomic radii

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

7

Anion atomic radii

Larger than parent atom

8

Cation ionic radii

Smaller than parent atoms

9

Electronegativity

Increases across row and decreases down column

10

Ionisation energy

removal of electron from highest occupied orbital of neutral atom

Follows rules same as atomic radii

11

Chlorophyll

Green compound found in leaves and stems
Channels sunlight into chemical energy which drives biochemical reactions through the process of photosynthesis

12

Chlorophyll structure

Tetrapyrrolic ring
Conjugated double bond system
Central magnesium atom

13

Hard Lewis bases

Donor atom is small and highly electronegative

14

Soft Lewis base

Donor atom is larger less electronegative making its electrons more polarisable

15

Hard Lewis acids

Either noble gas configuration or high charge and strongly held electrons

16

Soft Lewis acids

Large number of d electrons and low charge
More polarisable electrons

17

Ligand field stabilisation energy

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

18

Lewis bases

Substance that act as electron pair donors

19

Lewis acids

Substances that act as electron pair acceptors

20

Transition metal complexes

Complex is a combination of a Lewis acids and Lewis base
Ligand forms coordinate covalent bond to central metal ion

21

Neutral ligands

H20
NH3
CO

22

Anionic ligand

OH -
Cl-
Br-
I-
CN-

23

Bidentate ligands

En

24

Polydentate

EDTA 4-

25

Ligands

Lewis bases
Have atleast one non bonded pair of electrons

26

Chelate

Claw agents
Can bind to metal in two or more places

27

Counter ions

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

28

Coordination atoms

Number of ligand donor atoms bonded to metal ion

29

Charge of metal complex

Sum of charge on metal ion and sum of ligand charges

30

Oxidation state

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

31

Linear

Coordination number is 2

32

Tetrahedral or square planar

Coordination number is 4

33

Octahedral

Six coordination complexes

34

Factors influencing stability of metal complexes

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

35

KF

Formation constant

36

Hard acids

Small size
High charge
Low electronegativity

37

Soft acid

Large size
Low charge
Easily polarisable
High electronegativity

38

Hard bases

Don’t give up electrons
High electronegativity
Small donor atoms

39

Soft bases

Intermediate electronegativity
Large size

40

K>>1

More stable state is to the right

41

Chelate effect

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

42

Chelate effect kF

Larger stability so kF is bigger

43

Effect of chelate ring on kF

Chelate effect weakens as ring size increases
KF decreases as size increases

44

Octahedral complexes

5 membered ring

45

Lobes that point towards negative charge

Dz2 dx2-dy2

46

Lobes that point between point charges

Dyz
Dxz
Dxy

47

Crystal splitting energy

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

48

Crystal filed splitting energy

Increases with increasing oxidation number
Increases down a group

49

Ligand effect

Halide < oxygen < nitrogen < carbon

50

Weak field ligands

Smaller splitting energy
High spin
Absorbs linger wavelengths

51

Strong field ligands

Larger splitting energy
Low spin
Short wavelengths

52

High spin

Max number of unpaired electrons

53

Low spin

Max number of parallel spins

54

Irving Williams series of stability

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

55

Metal cations in water

Dissolve and are hydrated with water molecules to form squares ions

56

Aqua acids

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

57

Acidity of aqua acids

Increases with increasing positive charge of central atom
Increases with decreasing atomic radius of central atom

58

Weak aqua acids

Large cations with low charge

59

Strong aqua acids

Small cations with high charge

60

Ka>1

Strong acids

61

Ka<1

Weak acids

62

Donor atoms - neutral and negative

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

63

Donor atoms - positive

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

64

Reducing agent

Oxidised

65

Oxidising agent

Reduced

66

Gibbs free energy

G = delta H - temperature x delta S

67

Latimer diagrams

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

68

E cell

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

69

Frost diagram

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

70

Lower part of frost diagram

More thermodynamically stable

71

Trend in reduction potential

Cu> heme> Fe-S

72

Acids

Generally are metals on periodic table

73

Chelation

Prevents iron from being captured by other insoluble compounds

74

Amino acids

Oh dependent
Complex can fall apart if you change pH

75

To find y of frost diagram

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

76

If reduction potential Is favouring left hand side

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

77

Increasing ph

Decreases e cell value
Process is unfavourable

78

Acidity

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

79

Acidity - double bonds

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

80

Acidity - negative charges

More negative charges is stable
Less negative charges is more acidic

81

conjugate acid

acid formed from protonating a base

82

conjugate base

base formed from deprotonating an acid

83

weak acid strength

low positive charge

84

strong acid strength

high positive charge

85

aqua acid precipitate in water

strong acidic aqua acids form precipitates in water

86

more double bonds

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

87

more negative atoms

requires greater delocalisation over the molecule, reducing stability

88

stability of chelate complexes

increased entropy due to increased number of molecules