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CHEM2007 - General Inorganic Chemistry > Chemistry of the p-Block > Flashcards

Chemistry of the p-Block Flashcards

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1
Q

What is the definition of ionisation energy?

A

The energy required to remove completely an electron from the gaseous atom or molecule in its ‘ground state’.

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2
Q

What is the equation for the first ionisation energy?

A

M(g) → M+(g) + e-

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3
Q

What is the equation for the third ionisation energy?

A

M2+(g) → M3+(g) + e-

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4
Q

What is a feature of the values for ionisation energies?

A

They are always positive as ionisation requires an input of energy.

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5
Q

What are the trends in ionisation energy in a group?

A

Ionisation energy decreases down a group. This is because the electron is further away from the nucleus.

However, there are deviations in period 4 and other lower periods.

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6
Q

Why is there a deviation in the ionisation energy trend in period 4?

A

Periods 3 and 4 and periods 5 and 6 have very similar ionisation energies due to the d and f blocks increasing the nuclear charge but not providing much shielding.

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7
Q

What are the trends in ionisation energy across a period?

A

Generally, ionisation energy increases from left to right. This is because the atoms have a greater nuclear charge and so attract the electron more tightly.

However, there are deviations from the trend in group 13 and group 16.

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8
Q

Why is there a drop in ionisation energy in group 13 and group 16?

A

In group 13, the electron is being removed from a p orbital rather than an s orbital. p orbitals are higher energy and so are further away from the nucleus.

In group 16, the electron is being removed from a p orbital with two electrons in it. This creates extra Coulombic repulsion.

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9
Q

What is the definition of electron affinity?

A

The energy released when a gaseous atom, molecule or ion in its ‘ground state’ gains an electron.

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10
Q

What is a generic equation for the first electron affinity?

A

X(g) + e- → X-(g)

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11
Q

What is a feature of the values for electron affinity?

A

The electron affinity is positive as this is a favourable process.

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12
Q

What is the general trend for the electron affinity across a period?

A

It generally increases. However there are many deviations in the trend.

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13
Q

How do the electron affinities compare in group 1 and group 13?

A

Group 1 have larger electron affinities compared with group 13. This is because group 1 atoms will get a full s orbital which is more stable.

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14
Q

How do the electron affinities compare in group 1 to group 2?

A

Group 1 electron affinities are positive whereas group 2 electron affinities are negative. This is because in group 2 the electron is being added to a higher energy p orbital.

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15
Q

What are the electron affinities like in group 15?

A

They are anomalously low. This is because there are already in an np3 sub-shell. The extra electron added will be paired which creates Coulombic repulsion

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16
Q

How do the electron affinities compare in periods 3 and 4 and periods 5 and 6?

A

They have very similar values due to the presence of a the d and f block. There is a large increase in nuclear charge but a small increase in shielding.

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17
Q

How do the electron affinities compare between the 2nd and 3rd period?

A

The 2nd period has lower electron affinities than the 3rd period. This is because 2nd period elements are very small and so have a high charge to radius ratio. This increases Coulombic and inter-electron repulsions.

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18
Q

What is the definition of electronegativity?

A

The ability of an atom to attract electron density towards itself in a molecule.

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19
Q

How is Pauling’s electronegativity calculated?

A

It is calculated by taking the change in bond energies between the measured bond energy and the expected bond energy and applying this into an equation. If the atoms had identical electronegativities, the difference would be zero.

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20
Q

How can electronegativity vary for the same element?

A

It depends on what is bonded to the element, the oxidation state and the hybridisation.

sp > sp2 > sp3 due to the greater s character and the electrons being held more tightly.

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21
Q

What are the features of a van Arkel Ketelaar triangle?

A

The y-axis is the ionicity parameter and is the difference in electronegativity between the two atoms.

The x-axis is the covalency parameter which is the average electronegativity of the two atoms.

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22
Q

What is the definition of atomic radius?

A

The distance from the centre of the nucleus to the outermost electron. However, the outermost electron does not have a well defined position.

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23
Q

What is the covalent atomic radius defined as?

A

Half the length of the symmetrical homonuclear bond.

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24
Q

What is the metallic atomic radius defined as?

A

The equivalent distance between ions in a metal lattice.

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25
Q

What is effective nuclear charge?

A

The outermost electrons feel a nuclear charge which is less than the actual nuclear charge because of shielding effects from other electrons.

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26
Q

What are the trends in effective nuclear charge?

A

Across a period, effective nuclear charge increases substantially.

Down a group, effective nuclear charge increases (to a smaller extent).

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27
Q

What is the most important factor for ionisation energy?

A

The distance between the nucleus and the electron.

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28
Q

Why can Slater’s rules not explain the reduction in ionisation energy down a group?

A

Slater’s rules do not take into account the distance from the nucleus or penetration.

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29
Q

What happens to atomic radius across a period?

A

Radii decrease from left to right.

This is because increasing nuclear charge means electrons are more tightly held which means the orbital size and energy decrease. This causes the radius to contract.

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30
Q

What happens to atomic radius down a group?

A

Populating orbitals of the next principal quantum shell so further away from the nucleus.

Slater’s rules do not explain this.

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31
Q

How do ionic radii compare to atomic radii?

A

They follow similar trends but anions are larger due to greater inter-electron repulsion from more electrons.

Cations are smaller due to a greater Zeff.

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32
Q

What happens to radii when the oxidation state increases?

A

Radii will decrease due to a higher Zeff.

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33
Q

How does the radii change depending on coordination number?

A

More ligands around a metal centres means more electron density at the metal and an increased apparent size.

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34
Q

What happens to HOMO energies going down a group?

A

HOMO energies increase (less stable) and the s-p separation decreases.

This is because the electrons are in higher principal quantum shells so the electrons are further away from the nucleus.

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35
Q

Why are s electron energies lower than p electron energies?

A

S electrons are closer to the nucleus and more stabilised by the positive attractive force.

P electrons are less penetrating and so less stabilised by positive attractive force.

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36
Q

How does s and p orbital energy change across a period?

A

s and p orbital energies decrease across a period. This is more pronounced for s orbitals.

This is because Zeff is increased and s orbitals have a better penetration compared to p orbitals.

This also explains why the π and σ levels swap around from N2 to O2.

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37
Q

What are promotion energies?

A

The energy required to promote an electron from the ground state to the excited state.

This could be from an s orbital to a p orbital.

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38
Q

How do promotion energies change down a period?

A

Promotion energies generally increase down a group.

An exception is Ga and Ge.

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39
Q

How do forming oxides and fluorides affect the types of molecules formed?

A

Due to compounds containing fluorine having a higher covalency parameter, they are more likely to be discrete covalent molecules than in a covalent network structure.

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40
Q

What is the inert pair effect?

A

The tendency of the electrons in the outermost atomic s orbital to remain unionised or unshared in compounds of the group 13-16 elements.

It is also the observation that as a group is descended, the n-2 oxidation state becomes more favoured.

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41
Q

What are the causes of the inert pair effect?

A

Strength of covalent bonds decreases down a group due to poor orbital overlap. This means that bond enthalpy does not offset the hybridisation energy cost.

There are also relativistic effects which stabilise the 6s orbital for Tl and Pb.

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42
Q

What are relativistic effects?

A

The heavier the element, the faster the outer electrons move.

The theory of special relativity sates that objects moving near the speed of light gain mass.

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43
Q

How do relativistic effects cause stabilisation?

A

As the mass increases, the orbital contracts.

This is called direct relativistic orbital contraction.

This causes the s electrons of heavy elements to be more stable and less likely to be ionised or involved in bonding.

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44
Q

How are the different orbitals stabilised by relativistic effects?

A

S orbitals are stabilised the most as they have the best penetration.

P orbitals contact but not as much due to poorer penetration.

d and f orbitals experience indirect relativistic orbital expansion (destabilised). This is due to poor penetration and the contraction of s and p leaves them more shielded and less affected by Zeff.

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45
Q

What are trends in bond strength?

A

Generally, bonds become weaker and longer down a group. This is because there is a larger valence orbital volume but only two electrons per orbital so they are more diffuse. There is also less overlap and so sharing electron density is less effective.

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46
Q

Why do first row elements commonly form multiple bonds?

A

The 2pπ orbitals have a good overlap.

In second and later rows, they have poor overlap and so single bonds are favoured.

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47
Q

What are the features of pπ-dπ bonding?

A

A filled p orbital donates electrons to a vacant d orbital.

This happens in the P=O bond.

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48
Q

What is the valency of group 13 elements?

A

They only have three valence electrons so have a valence of 3. 3 bond pairs gives only 6 valence electrons so triel compounds are electron deficient (Lewis acids).

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49
Q

What is an exception in homonuclear bond strengths?

A

2nd period (from nitrogen to fluorine) are lower than expected due to greater electrostatic repulsions for smaller atoms which more charge-dense orbitals.

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50
Q

What are the trends for homonuclear bond energies going down a group?

A

Generally, bonds get weaker down the group due to larger valence orbitals and less effective overlap or sharing of electron density.

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51
Q

Why do group 1 and group 2 metals have low homonuclear bond energies?

A

The elements are more electropositive and so have adjacent partial positive charges. This causes electrostatic repulsion which destabilises the bonding interaction.

Hydrogen is an exception as it has a higher electronegativity along with small orbitals, good overlap, effective sharing of electron density. There are no electrostatic repulsions as there are only 2 electrons present. This is a very strong bond.

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52
Q

How do homonuclear bond energies change across a period?

A

Going from left to right (group 1 to 14), atoms get smaller and more electronegative. This leads to more effective orbital overlap and sharing of electron density causing stronger bonds.

From groups 15 to 17, lone pairs are present so the increased electrostatic repulsions weaken the bonds.

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53
Q

Why are B-F bonds very strong?

A

They have small orbitals with effective overlap and sharing of electron density. It is a polar covalent bond so there is additional electrostatic attraction. Pi donation from lone pairs on fluorine to empty p orbital on B is possible.

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54
Q

Why are Al-F bonds weaker than B-F bonds?

A

There are larger valence orbitals. This trend continues down the group.

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55
Q

What are the differences in bond strength between B-F and C-F?

A

C-F bonds are weaker. This is due to a lower difference in electronegativity causing a reduced electrostatic attraction. There is no pi-donation as all carbon orbitals are occupied.

56
Q

How do Si-F bonds compare in strength to C-F bonds?

A

Si-F bonds are stronger than C-F bonds. This is due to a higher difference in electronegativity causing an increased electrostatic attraction. Si also has low energy d-orbitals for pπ-dπ bonding.

57
Q

What are the differences in bond strength between N-F and P-F?

A

N-F bonds are weaker than P-F bonds. This is because there is a lower difference in electronegativity and therefore decreased electrostatic attraction. pπ-dπ bonding is also possible for P.

58
Q

How does the S-F bond compare in strength to the P-F bond?

A

The S-F bond is weaker than the P-F bond. This is because there is a lower difference in electronegativity which has a decreased electrostatic attraction. S and F have lone pairs that are small and the electrostatic repulsion is greater than that for P.

59
Q

How does the bond strength for the hydrides differ to the fluorides?

A

As there are no lone pairs on H, electronegativity difference plays a greater role. The strongest bonds are on the right.

60
Q

What are general trends of the triels?

A

Their chemistry is dominated by electron deficiency as they have fewer valence electrons than the number of valence orbitals. Their compounds are Lewis acids - they can accept an electron pair from Lewis bases.

The +1 oxidation state becomes more stable as the group is descended due to the inert pair effect.

61
Q

What are the features of boron halides?

A

BX3 exists for all halides.

All fluorides are trivalent and can participate in π bonding.

It is a planar molecule and donation of electrons from a full p orbital on F into the vacant p orbital on B.

Lewis acid strengths increase down the group: BI3 > BBr3 > BCl3…

62
Q

What are the features of non-fluoride halides for other triels?

A

They dimerise. This is due to poor orbital overlap between the halogen and the triel due to large diffuse orbitals. Dimerisation is preferred to π-bonding.

63
Q

What is an example of an exchange reaction of halides in Triels?

A

BX3 + ROH → B(OR)3

However, BF3 resists due to strong B-F bonds.

64
Q

What are the features of boron hydrides (boranes)?

A

B cannot form π bonds with H as there are no lone pairs so BH3 forms an electron deficient dimer - B2H6

65
Q

What is the bonding like in B2H6?

A

The B-H bonds are 2-centre, 2-electron bonds.

The B-H-B bonds are 3-centre, 2-electron bonds.

66
Q

Why do Triel hydrides become less stable as the group is descended?

A

There are larger valence orbitals which causes less effective sharing of electron density between nuclei and therefore weaker bonds.

67
Q

What terms are used when there are an equal number of vertices to BH units, one fewer BH units to vertices and two fewer?

A

Equal - Closo

One fewer - Nido

Two fewer - Arachno

68
Q

How can Wade’s rules be used for carboranes?

A

The BH unit can be used the same as a C-H+ unit as they are isolobal.

69
Q

Which is the most thermodynamically stable isomer for a carboranes?

A

The one where the carbons are placed apart.

70
Q

What are group 14 elements called?

A

The tetrels.

71
Q

Why does the stability of tetrel hydrides decrease down the group?

A

The tetrel-H bond decreases in strength due to larger more diffuse orbitals.

72
Q

Why is CX4 hydrolytically stable but SiX4 isn’t?

A

The bonds are more polar and the LUMOs are low-lying. This causes silicon to be more electrophilic and so is susceptible to nucleophilic attack.

73
Q

Why can the tetrels Ge, Sn and Pb form MX2 structures?

A

Due to the inert pair effect.

74
Q

How do tetrel halides react with alcohols?

A

The alcohol attacks as a nucleophile and displaces the halogen.

75
Q

What are silicones?

A

Siloxane polymers.

There are two R groups attached to the Si with an Si - O repeating chain.

76
Q

What is a two-step process that can form silicones?

A

Hydrolysis of silylhalides:
Me3SiCl + H2O → Me3SiOH + HCl.

Condensation reactions:
2 Me3SiOH – -H2O→ Me3Si-O-SiMe3.

77
Q

What are the three types of monomer units needed to form polymeric siloxanes?

A

A terminal group. This has a single OH connected to the silicon.

A chain forming group. This has two OH connected to the silicon.

Cross linking group. This has three OH connected to the silicon.

78
Q

How stable are silicones?

A

They are very stable due to strong Si-C and Si-O-Si bonds.

However, they will react with fluorinating agents as the Si-F bond is incredibly strong.

79
Q

Why is the Si-O bond so strong?

A

Si-O can undergo pπ-dπ bonding as the filled 2p orbital on oxygen can donate into the empty 3d of Si.

80
Q

Why is tetramethylsilane used as an NMR reference rather than SiCl4?

A

TMS is more stable and less reactive than SiCl4. This is due to steric shielding, less polar bonds and the fact that Cl- is a better leaving group than Me-.

81
Q

Why can’t silicon form a double bond with oxygen?

A

The orbital overlap in the p orbitals is poor.

82
Q

How can silicon double and triple bonds be made?

A

Using a bulky, sterically hindering group stops bonding through σ bonds and forces it to use pi bonds.

83
Q

What are Zintl ions?

A

An ion containing a group 1 or group 2 metal, along with post-transition metals (group 12) or the metalloids from groups 13-16.

84
Q

What are group 15 elements called?

A

The Pnictogens.

85
Q

What are some features of nitrogen?

A

Dinitrogen is one of the strongest covalent bonds and is very short.

There is little catenation as the N-N single bond is weak.

The N3- ions does exist and is formed when Li metal reacts with N2:
6 Li + N2 → 2Li3N.

The lattice energy of Li3N is high as Li+ is very small.

86
Q

What is catenation?

A

Linking of atoms of the same (or very similar element) together in a chain or ring.

87
Q

What are general features of Phosphorus?

A

There are a number of allotropes due to catenation.

Diphosphorus does exist but only under extreme conditions.

88
Q

What are general features of As, Sb and Bi and a trend in the Pnictogens?

A

As, Sb and Bi form more layered structures.

Elements become more metallic as the group is descended.

89
Q

What are the features of white phosphorus?

A

White phosphorus (P4 is pyrophoric which means it is very reactive towards oxygen and ignites in contact with air).

Reacts to form P4O10. This contains lots of P-O and P=O bonds. P-O and P=O bonds are strong due to pπ-dπ bonding of O into the d-orbital of P.

90
Q

What phosphorus oxides can form?

A

Phosphorus (V) - P4O10

Phosphorus (III) - P4O6

91
Q

What nitrogen oxides can form?

A

There are 7 molecular oxides.

This includes ·NO, NO2-, N2O4, ·NO2, NO2+.

92
Q

How can ·NO2 dimerise?

A

Dimerisation can occur to form N2O4.

However, the N-N bond is very weak.

93
Q

Why are there no negative oxidation states for N-oxides?

A

Oxygen is more electronegative so present as O2-.

94
Q

How are phosphates used in biology?

A

They are used in the structure of ATP which provides energy to the body.

It is also essential for plant growth.

95
Q

What is the general formula for pnictogen hydrides?

A

EH3.

96
Q

How does the E-H bond angle change for pnictogens down the group?

A

The bond angle decreases.

97
Q

What are the features of ammonia?

A

It has a relatively high boiling point due high polarity and H bonding.

It can participate in acid-base chemistry.

Liquid ammonia is a good solvent.

98
Q

What are the features of PH3?

A

PH3 is not basic.

It is highly flammable and reacts to form oxide and water.
4PH3 + 8O2 → 2P2O5 + 6H2O.

99
Q

How can phosphines act as ligands?

A

They are good ligands as π-bonding involves overlap between a σ* orbital on P and a filled metal d-orbital.

The Tolman cone angle changes based on the R group connected to the phosphorus.

As the size of the substituents increases, the angle also increases.

100
Q

What pnictogen halides are formed?

A

All pnictogens can form MX3 for all halides.

However, for MX5 not all halides are formed. As the halide gets larger, it cannot form bonds with the larger pnictogens and iodine cannot form these.

This is because the bonding would involve large valence orbitals on both elements which would form very weak bonds.

101
Q

What are the features of NF3?

A

It is a gas that has a low boiling point due to no H bonding.

It can be prepared by: 4NH3 + 3F2 → NF3 + 3NH4F.

The F-N-F bond angle is less than the H-N-H bond angle in ammonia. This is because F pulls electron density away from N which causes less electrostatic repulsion between bond pairs.

102
Q

What are the features of PF3?

A

A gas at room temperature that is pyramidal.

103
Q

What are the features of PF5?

A

A trigonal bipyramidal shape with axial bond lengths slightly longer than equatorial bond lengths.

104
Q

What are the features of PCl5?

A

In the gas phase, trigonal bipyramidal.

In the solid state it consists of [PCl4]+ cations and [PCl6]- anions.

105
Q

What are the features of PBr5?

A

Crystallises as [PBr4]+[SbF6]-.

106
Q

What are the features of AsF2 and SbF2?

A

They are strong F- acceptors.

SbF5 + 2HF → [H2F]+[SbF6]-.

This is a super acid as it results in a great increase in the acidity of HF.

107
Q

What is the general formula of an oxyacid?

A

OmE(OH)n.

108
Q

What are Pauling’s rules for oxyacids?

A

The strength of the acid is independent of the number of OH groups, n. This is becuase the OH groups do not allow delocalisation of the charge on the anion.

The acid strength increases by 105 for each E=O moiety.

109
Q

What the pKa of an oxyacid?

A

Around 8-5m.

110
Q

What are the general principles of oxyacids of phosphorus?

A

All contain 4 coordinate P; at least one P=O.

At least one P-OH (ionisable).

May contain P-H (non-ionisable).

Can polymerise into chains or rings via P-O-P or P-P.

111
Q

What happens to the pKa of Phosphorus oxyacids as it becomes ionised?

A

The pKa increases as the charge is less well stablised.

112
Q

What is the suffix for an oxyacid formed from an EO2 and an EO3?

A

EO2 = -ous.

EO3 = -ic.

113
Q

What are the features of group 17 halogen oxyacids?

A

Group 17 elements have a wide range of common oxidation states.

They are named based on the relative number of oxygen atoms.

114
Q

How are group 17 halogen oxyacids named?

A

HOX - hypohalous (e.g. hypochlorous).

HOXO - halous (e.g. chlorites).

HOXO2 - halic (e.g. chlorates).

HOXO3 - perhalic (e.g. perchlorates).

115
Q

What are the features of HOXO2 and HOXO3 oxyacids?

A

They form XO3- and XO4- and are very strong oxidising agents.

116
Q

What are group 16 elements called?

A

The Chalcogens.

117
Q

How is ozone formed in a laboratory?

A

At high electrical potential at a surface in an ozonizer and in low concentrations under UV irradiation.

O2 + O2* → O3 + O.

O2 + O + M → O3 + M*

118
Q

What is the structure of ozone?

A

It is bent with resonance between a central positive oxygen and a negative outer oxygen.

119
Q

How can chalcogens react?

A

They can dissolve in oxidising agents.

They can be attacked by halogens.

120
Q

What are the features of sulfur?

A

There are lots of allotropes because of S-S catenation.

Catenation is favoured as the S-S-S bond distance and angles can vary greatly and cycles are thermodynamically favoured.

121
Q

What are the features of chalcogen hydrides?

A

Bond angles change from 104.5 degrees (sp3 hybrid) to 90 degrees (p orbitals) after H2O.

EH2 is prepared by chalcogenide and acid.

Hydrides other than H2O are highly toxic and very smelly.

122
Q

Why is SF6 so inert compared to SF4 and SeF6?

A

Although they have similar bond strengths due to similar differences in electronegativity, SF6 is coordinatively saturated and sterically hindered. As Se is larger, it is more susceptible to attack.

Therefore the reasons are more due to kinetic factors.

123
Q

What are the main features of the halogens?

A

They are diatomic which gives rise to high volatility for the halogens.

Chains are unlikely; catenation occurs to an extent for I but not through covalent interactions.

An element’s behaviour is dominated by high electronegativity F > Cl > Br > I.

124
Q

What are the chemical properties of the halogens?

A

They are very reactive and will directly combine with most elements.

Reactions with fluorine are very exothermic and violent due to a low F-F bond energy (due to short length and electrostatic repulsions) and high E-F bond energy (large ionic contribution).

Both ionic and covalent compounds are formed.

125
Q

How does hydrogen react with halogens?

A

All react to form hydrogen halides.

The reaction becomes less vigorous as the halogen changes.

126
Q

What are the trends in reactivity and acid strength for hydrogen halides?

A

Reactivity: HF > HCl > HBr > HI.

Acid strength: HF < HCl < HBr < HI.

Both of these trends are related to bond strength.

127
Q

For group 18 elements to form, what must be present?

A

Very electronegative substituents such as F, O or Cl to stablise the compounds of the group 18 elements.

128
Q

Which group 18 element has the most extensive chemistry?

A

Xenon as it has oxidation states ranging from +2 to +8.

129
Q

What are the features of xenon halides?

A

They are all colourless, volatile solids.

They are very powerful oxidising agents and fluorinating agents.

Fluorinating ability: XeF2 < XeF4 < XeF6.

130
Q

What are the structures of XeF2, XeF4 and XeF6?

A

XeF2: Linear

XeF4: Square planar

XeF6 Monocapped octahedron with lone pair projecting from one face. The lone pair distorts the octahedron by pushing back three of the F- ligands. The structure is fluxional and lone pair shifts from face to face and the overall structure is an average of all eight structures.

131
Q

What are the solution state structure and solid state structure for XeF6?

A

Solution state structure - 129Xe NMR spectrum shows a single 25 line resonance arising from coupling of the Xe nucleus to 24 equivalent F atoms.

Solid state structure - best represented as an [XeF5]+ cation and a F- anions. One of the F’s in each molecule is partially dissociated and bridges between the [XeF5]+ cations to form a tetramer.

132
Q

What are the features of other group 18 elements?

A

They need highly electronegative elements to stabilise oxidation state.

He - few compounds known and generally formed under extreme conditions.

Ne - hexagonal array of Ne atoms in a MOF structure.

Ar - HArF is metastable.

Kr - Much more limited chemistry than Xe, but KrF2 is known.

Rn - Very low abundance and radioactive, all isotopes have short half-lives. RnF2 is known but very unstable.

133
Q

What are the features of silsesquioxanes (a type of siloxane)?

A

They form many ring, cluster and polymeric structures determined kinetically.

134
Q

What phosphorus sulfides can form?

A

P4S10 which is isostructural with P4O10. The bonds are longer because of the larger size of sulfur and smaller electronegativity difference between component atoms.

P4S3 can from which is similar to white phosphorus.

135
Q

What is the structure of germylenes and how can it be formed?

A

A Ge=Ge double bond.

It can be formed by having an electropositive substituent that favours the double bond character. This is caused by the size and electronegativity of E which affects bond polarity and relative energies of singlet vs triplet states and HOMO and LUMO.

136
Q

What are the features of borazine?

A

Boron-nitrogen analogue of benzene. Nitrogen is a lot more electronegative and smaller than boron. So, more localised charges than benzene in structures that are in resonance with the neutral representation.

CHEM2007 - General Inorganic Chemistry (6 decks)

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