Module 1-6 Flashcards
Ground state
Most stable electron configuration
Electrons in lowest energy state
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
Tranisition metal ions
Lewis acids
Cations
Have vacant d orbitals
Electron acceptors
Effective nuclear charge
Electrons in inner shell have shielding effect on electrostatic attraction between protons and electrons in outer shells
Zeff trend
Increases across the row
Decreases in order of s>p>d>f
Atomic radii
Half distance between metallic nuclei
Increases down a column
Decreases across row
Anion atomic radii
Larger than parent atom
Cation ionic radii
Smaller than parent atoms
Electronegativity
Increases across row and decreases down column
Ionisation energy
removal of electron from highest occupied orbital of neutral atom
Follows rules same as atomic radii
Chlorophyll
Green compound found in leaves and stems
Channels sunlight into chemical energy which drives biochemical reactions through the process of photosynthesis
Chlorophyll structure
Tetrapyrrolic ring
Conjugated double bond system
Central magnesium atom
Hard Lewis bases
Donor atom is small and highly electronegative
Soft Lewis base
Donor atom is larger less electronegative making its electrons more polarisable
Hard Lewis acids
Either noble gas configuration or high charge and strongly held electrons
Soft Lewis acids
Large number of d electrons and low charge
More polarisable electrons
Ligand field stabilisation energy
Pt(2) > Ni(2) > Co (2) > Cu(2) > Fe(2) > Zn (2) > Mg(2)
Lewis bases
Substance that act as electron pair donors
Lewis acids
Substances that act as electron pair acceptors
Transition metal complexes
Complex is a combination of a Lewis acids and Lewis base
Ligand forms coordinate covalent bond to central metal ion
Neutral ligands
H20
NH3
CO
Anionic ligand
OH - Cl- Br- I- CN-
Bidentate ligands
En
Polydentate
EDTA 4-
Ligands
Lewis bases
Have atleast one non bonded pair of electrons
Chelate
Claw agents
Can bind to metal in two or more places
Counter ions
Written on outside of square brackets and not bonded to metal ion
Negatively charged - right
Positively charged - left
Coordination atoms
Number of ligand donor atoms bonded to metal ion
Charge of metal complex
Sum of charge on metal ion and sum of ligand charges
Oxidation state
Charges on metal atom
Complexes containing neutral ligands the oxidation state is equal to the net charge on complex
Linear
Coordination number is 2
Tetrahedral or square planar
Coordination number is 4
Octahedral
Six coordination complexes
Factors influencing stability of metal complexes
Irving Williams series of stability
Hard soft acid base theory
Chelate effect
KF
Formation constant
Hard acids
Small size
High charge
Low electronegativity
Soft acid
Large size
Low charge
Easily polarisable
High electronegativity
Hard bases
Don’t give up electrons
High electronegativity
Small donor atoms
Soft bases
Intermediate electronegativity
Large size
K»1
More stable state is to the right
Chelate effect
Ligands that form chelate rings generally bind more strongly to a metal ion then monodentate ligands
Chelate effect kF
Larger stability so kF is bigger
Effect of chelate ring on kF
Chelate effect weakens as ring size increases
KF decreases as size increases
Octahedral complexes
5 membered ring
Lobes that point towards negative charge
Dz2 dx2-dy2
Lobes that point between point charges
Dyz
Dxz
Dxy
Crystal splitting energy
Metal ion placed in middle causes d electrons to experience repulsion and raise the energy level
Crystal filed splitting energy
Increases with increasing oxidation number
Increases down a group
Ligand effect
Halide < oxygen < nitrogen < carbon
Weak field ligands
Smaller splitting energy
High spin
Absorbs linger wavelengths
Strong field ligands
Larger splitting energy
Low spin
Short wavelengths
High spin
Max number of unpaired electrons
Low spin
Max number of parallel spins
Irving Williams series of stability
Mn 2+ < Fe2+ < Co2+ < Ni2+ < Cu2+ > Zn2+
Metal cations in water
Dissolve and are hydrated with water molecules to form squares ions
Aqua acids
Acids where acidic proton in on water molecule attached to metal cation
Acidity of aqua acids
Increases with increasing positive charge of central atom
Increases with decreasing atomic radius of central atom
Weak aqua acids
Large cations with low charge
Strong aqua acids
Small cations with high charge
Ka>1
Strong acids
Ka<1
Weak acids
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
Donor atoms - positive
Positively charged atom in side chain is unable to coordinate to metal
Reducing agent
Oxidised
Oxidising agent
Reduced
Gibbs free energy
G = delta H - temperature x delta S
Latimer diagrams
Standard reduction potentials for species
Highly oxidised form on left
Reduction potentials for species above line
E cell
Electrons x e cell + electrons 2 x e cell / electrons 1 + electrons 2
Frost diagram
Reverse of Latimers
From least oxidised to highly oxidised
Reduction potentials reversed
Lower part of frost diagram
More thermodynamically stable
Trend in reduction potential
Cu> heme> Fe-S
Acids
Generally are metals on periodic table
Chelation
Prevents iron from being captured by other insoluble compounds
Amino acids
Oh dependent
Complex can fall apart if you change pH
To find y of frost diagram
Number of electrons to get to the end x reduction potential to get to most oxidised
If reduction potential Is favouring left hand side
E cell will become less and less if it’s positive
Increasing ph
Decreases e cell value
Process is unfavourable
Acidity
Decreases as atomic radii increases
Smaller radius = stronger acid
Atoms closer to top left of periodic table
Acidity - double bonds
More double bonds attached to central atom = greater electron withdrawal
= greater acidity
Acidity - negative charges
More negative charges is stable
Less negative charges is more acidic
conjugate acid
acid formed from protonating a base
conjugate base
base formed from deprotonating an acid
weak acid strength
low positive charge
strong acid strength
high positive charge
aqua acid precipitate in water
strong acidic aqua acids form precipitates in water
more double bonds
withdraws electron density towards the central atom resulting in greater polarisation of OH group, making OH bond weaker
more negative atoms
requires greater delocalisation over the molecule, reducing stability
stability of chelate complexes
increased entropy due to increased number of molecules
