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Inorganic chemistry > Module 1-6 > Flashcards

Module 1-6 Flashcards

(88 cards)

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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
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

Inorganic chemistry (2 decks)

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