AH Chemistry Unit 1: Inorganic Chemistry

Evidence for wave-particle dual nature of electrons 

The discrete lines observed in atomic spectra can be explained if electrons, like photons, also display the properties of both particles and waves.  

Standing waves 

• Electrons behave as standing (stationary) waves in an atom. 

• These are waves that vibrate in time but do not move in space.  

• There are different sizes and shapes of standing wave possible around the nucleus, known as orbitals.  

Quanta 

fixed amounts of energy, e.g. possessed by electrons   

The angular momentum quantum number l can have values: 

 from zero to n − 1

The magnetic quantum number m_l determines the orientation of the orbital can have values between: 

 − l and  + l  

Aufbau Principle

Electrons fill orbitals in order of increasing energy 

Hund’s Rule

When degenerate orbitals are available, electrons fill each singly, keeping their spins parallel before spin pairing starts

Pauli exclusion principle  

No two electrons in one atom can have the same set of four quantum numbers, therefore, no orbital can hold more than two electrons and these two electrons must have opposite spins In an isolated atom the orbitals within each subshell are degenerate.  

Dative covalent bond 

Type of covalent bond where both of the electrons in the shared pair are from the same atom, i.e. one atom “donates” the pair of electrons

Resonance structure 

Way of describing the structure of a molecule or ion when typical Lewis structures do not apply. The presence of delocalised electrons results in the structure of the molecule or ion to be halfway, or “resonate”, between that of two or more Lewis structures. For example, in benzene and ozone, the bond length is halfway between a single and a double bond.

Using VSEPR (valence shell electron pair repulsion) theory to predict the shapes of molecules and polyatomic ions 

1. taking the total number of valence (outer) electrons on the central atom and adding one for each atom attached  

2. adding an electron for every negative charge  

3. removing an electron for every positive charge  

4. dividing the total number of electrons by two to give the number of electron pairs 

 

Arrangement of electron pairs in an atom

Electron pairs are negatively charged and repel each other and are arranged to minimise repulsion and maximise separation.

Electron pair repulsions decrease in strength in the order:

non-bonding pair/non-bonding pair > non-bonding pair/bonding pair > bonding pair/bonding pair 

Name for the molecular shape with three electron pairs 

trigonal planar

Name for the molecular shape with five electron pairs 

trigonal bipyramidal

d-block transition metals  

metals with an incomplete d subshell in at least one of their ions

The filling of the d orbitals follows the aufbau principle. State the two elements which are exceptions to this rule and explain why. 

Chromium and copper atoms. These exceptions are due to the special stability associated with the d subshell being halffilled or completely filled.  

State which orbital electrons are lost from first when atoms from the first row of the transition elements form ions

4s

Compounds containing metals in high oxidation states are often  

oxidising agents

Compounds with metals in low oxidation states are often  

reducing agents

Ligands

negative ions or molecules with non-bonding pairs of electrons that they donate to the central metal atom or ion, forming dative covalent bonds

Coordination number  

The total number of bonds from the ligands to the central transition metal

Transition metal complex 

Central transition metal ion surrounded by and bonded to a number of molecules or ions  

Rules for naming transition metal ion complexes  

1. The ligands come first followed by the metal 

2. Ligands are listed in alphabetical order 

3. Negatively charged metal ions end in” –ate” ,  cuprate is used for copper and ferrate is used for iron 

4. The oxidation state of the central metal ion is stated in roman numeral 

Name the ligand with the formula NH₃

Ammine

Name the ligand with the formula OH^-

Hydroxido

Name the ligand with the formula Cl^-

Chlorido 

Name the ligand with the formula CN^-

Cyanido 

Name the ligand with the formula H₂O 

Aqua

Name the ligand with the formula C₂O₄^2- 

Oxalato

Explain why in a complex of a transition metal, the d orbitals are no longer degenerate 

• When the electrons present in the ligands approach the central metal ion, this causes the electrons in the orbitals lying along the axes to be repelled

• The d orbitals that lie along the axes now have higher potential energy. This results in the splitting of d orbitals to higher and lower energies 

 

Strong field ligands

• Ligands that cause a large difference in energy between subsets of d orbitals

• Examples, in increasing order of strength, include H₂O, NH₃ and  CN^-

Weak field ligands  

• Ligands that cause a small energy difference between subsets of d orbitals

• Examples, include the ions of group 7 elements. The strength of the ligand increases going down the group

Spectrochemical series 

The ordering of ligands based on their ability to split d orbitals. I- < Br- < Cl- < F- < H2O < NH3 < CN- 

Explain how colour arises in transition metal complexes 

• Light is absorbed when electrons in a lower energy d orbital are promoted to a d orbital of higher energy

• If light of one colour is absorbed, then the complementary colour will be observed  

• Note that no photon emission is involved for d-d transitions

If the ligands surrounding the transition metal ion are strong field ligands (those that cause the greatest splitting of the d orbitals) d-d transitions are more likely to occur in the {UV/visible} region of the electromagnetic spectrum.

UV region