Origin of the elements

Cosmic abundances of the elements. Even \(Z\) more abundant than odd \(Z\).

Synthesis of the elements:

  • exothermic processes in stars
    • hydrogen burning
    • helium burning
    • carbon burning
    • \(\alpha\)-process
    • e-process
  • neutron capture processes
  • other processes

\[\begin{equation} \ce{4{}^1H -> ^4He + 2e+ + 2\nu_e}\quad Q = \SI{26.72}{\mega\eV}~\mathrm{(exothermic~process)} \end{equation}\]

\(\SI{26.22}{\mega\eV}\) for radiation, or \(\SI{4.20}{\pico\joule}\) per \(\ce{He}\) atom, or \(\SI{2.5E9}{\kilo\joule\per\mole}\), i.e., \(\SI{2.5}{\tera\joule}\) per mole \(\ce{He}\).

The abundance of the elements is commonly given as a mass-fraction of the Earth’s crust. So what is the mass of the crust?

Hydrogen, deuterium, tritium

https://www.compoundchem.com/2019/01/29/iypt001-hydrogen/

Molecular hydrogen is quite unreactive at normal temperatures. It only reacts immediately with fluorine, and under illumination with chlorine, and only after ignition with oxygen. The latter is the well-known hydrogen oxidation reaction:

\[\begin{equation} \ce{H2(g) + 1/2O2(g) -> H2O(l)}; \quad\Delta H^0=\SI{-286}{\kJ} \end{equation}\]

At higher temperatures hydrogen is much more reactive, forming hydrides with a multitude of elements.

Hydrides

  • The noble gases do form hydrides.
  • The alkali and alkaline earth metal ions form salt-like hydrides with
    1. high melting points, ii) evolve hydrogen gas at the anode, and
    2. are strong reducing agents, e.g., \(\ce{NaH}\) and \(\ce{CaH2}\). These compounds have high conductivity at the melting point. They are white at high purity. \(\ce{H-}\) and its salts react immediately with substances which has or can form protons, e.g., \(\ce{NaH + H+ -> Na+ + H2(g)}\).
  • Covalent hydrides, simple such as \(\ce{BH3}\), \(\ce{CH4}\), \(\ce{NH3}\), \(\ce{H2O}\) and \(\ce{HF}\), or complex such as \(\ce{NaBH4}\), \(\ce{LiAlH4}\), \(\ce{HCo(OH)4}\), \(\ce{ReH9^{2-}}\).
  • Hydrides with metallic properties: i) lanthanum and neodymium reacts with \(\ce{H2}\) forming black graphite-like crystals, which react violently with water - similar properties as the salt-like hydrides, ii) transition-metal hydrides; \(\ce{Ti}\), \(\ce{Zr}\) \(\ce{Hf}\) absorb hydrogen and form non-stoichiometric compounds,
    1. uranium hydride, \(\ce{UH3}\), is an important starting material for the production of other uranium compounds. Palladium hydride has a remarkably high diffusion rate of hydrogen (it is not an ionic hydride but rather an alloy of palladium with metallic hydrogen).
  • Other hydrides, e.g., \(\ce{BeH2}\), \(\ce{MgH2}\), and \(\ce{CuH}\).

Bonds of hydrogen

  • Withdrawing of the electron leaving a proton, \(\ce{H+}\), which is always solvated.
  • Uptake of an electron forming hydride ion, \(\ce{H-}\).
  • Formation of a covalent bond, e.g., \(\ce{H-H}\).
  • Hydrogen bonds.

Industrial production of hydrogen

Elemental hydrogen can be produced from practically all hydrogen compounds. However, the most convenient are water together with methane, and other hydrocarbons (coal, petroleum and natural gas).

The dominant large-scale process in integrated plants is the catalytic steam hydrocarbon reforming process. Reactants are natural gas or oil-refinery feedstock. The reactants are desulfurised to protect the catalyst, then mixed with process steam and passed over a nickel-based catalyst at \(\SIrange{700}{1000}{\degreeCelsius}\).

\[\begin{equation} \ce{C3H8 + 3H2O ->[{Ni catalyst}][{$700\unicode{x2103}-1000\unicode{x2103}$}] 3CO + 7H2} \end{equation}\]

Another common process is the water gas shift reaction.

Group 1: the alkali metals (except for hydrogen)

General properties of the alkali metals

  • Has a single valence electron (\(s\)-orbital) outside the noble gas shell.
  • The first ionisation energy is low, while subsequent I.E. are high – only \(\ce{M+}\) ions are formed.
  • The \(\ce{M+}\) ion is spherical.
  • The chemistry of the alkali metals are generally ionic (electrostatic).
  • The tendency for formation of covalent interactions decreases down the group.
  • All isotopes of francium are radioactive.

The following properties decrease down the group:

  • The melting point and the heat of sublimation of the metals.
  • The lattice energies of the salts, except for those with the smallest anions.
  • The hydration energies.
  • The easiness to thermally decompose alkali carbonates and nitrates.
  • The strength of the covalent in \(\ce{M2}\) molecules.
  • The heat of formation of alkali fluorides, hydrides, oxides, and carbides.
  • The reactivity of the alkali metals versus all chemical reagents, except nitrogen.
  • The electropositivity (the opposite of electronegativity).

The reactions of alkali metals with oxygen:

  • Lithium burns in both air and oxygen, forms \(\ce{Li2O}\) and traces of \(\ce{Li2O2}\).
  • \(\ce{K}\), \(\ce{Rb}\), \(\ce{Cs}\): these metals burn in air and oxygen, forming \(\ce{M2O}\) which is oxidised further to \(\ce{MO2}\).

Other ions which have similar properties as the alkali metal ions:

  • The \(\ce{NH4^+}\) salts have properties similar to the potassium salts.
  • The chemistry of \(\ce{Tl^+}\) is in many ways similar to that of \(\ce{Rb^+}\).
  • The chemistry of the \(\ce{(\eta^5-C5H5)Co(III)}\) ion and its sandwich analogs is in many cases similar to that of \(\ce{Cs^+}\).

Dissolution of alkali metals in liquid ammonia as well as in amines gives solvated electrons.

Group 2: the alkaline earth metals

  • Their simple compounds, complexes, and organometallic compounds
  • comparisons with elements of group 1
  • how beryllium differs from the other group 2 elements

General properties of the alkaline earth metals

Beryllium has unique chemical properties with a dominating covalent chemistry, however due to its small ionic radius the aqua ion \(\ce{Be(OH2)4^{2+}}\) is formed.

All alkaline earth metals are electropositive metals, and they have much higher melting points, boiling points and densities than the alkali metals.

There are similarities between the chemistry of beryllium and aluminium, and in some aspects also with zinc.

All the isotopes of radium are radioactive.

Some systematic tendencies from Mg–Ra

  • The strength of the hydration of the alkaline earth metal decreases.
  • The solubility of the sulfates, nitrates, and halides (except fluorides) decreases.
  • The thermal stability of carbonates, nitrates, and peroxides increases.
  • The reaction rate with hydrogen increases.

Beryllium and beryllium compounds

https://www.compoundchem.com/2019/02/16/iypt-2019-elements-004-beryllium-emeralds-and-nasa-telescopes/

  • Beryllium is normally two-coordinated in linear configuration, and the maximal coordination number of beryllium is four.
  • Beryllium is a gray, light, hard and rigid metal.
  • Beryllium has low absorption of x-rays, and is therefore often used as x-ray window material.
  • \(\ce{BeO}\) and \(\ce{Be(OH)2}\) are ampholytic (meaning they will dissolve in both acid and base): \[\begin{align} \ce{BeO + 2H3O+ + H2O &-> Be(OH2)4^{2+}}\\ \ce{BeO + 2OH- + H2O &-> Be(OH)4^{2-}} \end{align}\]
  • Binary compounds: i) \(\ce{Be + O2 -> BeO}\), ii) \(\ce{3Be + N2 -> Be3N2}\),
    1. \(\ce{Be + C2H2 -> BeC2 + 2H2}\).
  • Complexes: i) the beryllate ion, \(\ce{Be(OH)4^{2-}}\), ii) \(\ce{BeF4^{2-}}\),
    1. \(\ce{BeCl4^{2-}}\), iv) the water in \(\ce{Be(OH2)4^{2+}}\) us very strongly coordinated, e.g., it is not possible to extract water from \(\ce{[Be(OH2)4]Cl2}\) with \(\ce{P_4O_{10}}\), v) ammine complex, \(\ce{Be(NH3)4^{2+}}\).
  • The hydrated beryllium ion is easily hydrolysed: \(\ce{Be(OH2)4^{2+} + H2O -> BeOH(OH2)3^+ + H3O+}\) The hydrolysed beryllium ion can have a complex configuration, depending on
    1. the nature of the anion, ii) temperature, iii) concentration and iv) pH. Possible species are \(\ce{Be2(OH)^{3+}}\), \(\ce{Be3(OH)_3^{3+}}\) (probably cyclic) and \(\ce{Be5(OH)7^{3+}}\). Note that water is also coordinated giving CN=4 around beryllium.

Magnesium

https://www.compoundchem.com/2019/03/15/iypt012-magnesium/

The following also applies to the rest of the group (calcium, strontium, barium and radium):

  • Oxides, \(\ce{MO}\), are obtained by heating the carbonates. These compounds are white with \(\ce{NaCl}\) structure.
  • Hydroxides, \(\ce{M(OH)2}\).
  • Carbides, \(\ce{MC2}\).
  • Other binary compounds; direct reaction between the metal and other elements gives binary compounds as borides, nitrides, silicides, arsenides, sulfides, etc.
  • Oxosalts
  • Complexes: only a few \(\ce{Mg}\) and \(\ce{Ca}\) complexes are known. The known ones are chelate complexes with oxygen donors (EDTA type). What about chlorophyll?
  • Grignard compounds, \(\ce{RMgX}\).

Group 14: the carbon group (Karin Larsson, 2019-10-21)

Occurrence

Carbon found in nature as diamond and graphite, both commonly mined but can also be synthesised.

Silicon is very abundant (in fact, it is the second most abundant element in Earth’s crust by mass after oxygen).

Element Mass-% in Earth’s crust/%
C 0.1800
Si 27.0000
Ge 0.0001
Sn 0.0002
Pb 0.0010
Fl 0.0000

Crystal structure

All group 14 elements have at least one solid phase with the diamond structure.

Carbon

https://www.compoundchem.com/2019/02/22/iypt-2019-elements-006-carbon-diamonds-pencils-and-life/

Common allotropes: diamond, graphite.

Graphite is more stable than diamond, but the phase-transformation does not occur to any observable rate at NTP.

Graphite can form intercalation compounds.

Other forms of carbon:

  • fullerenes, very inert (cf. the boron balls)
  • graphene
  • nanotubes
  • nanobuds
  • graphenated carbon nanotubes
  • U-carbon

Fullerenes can be reduced to form fullerene salts, \(\ce{C_{60}^{n-}}\), \(1<n<12\).

Carbon nanotubes are closely related to both fullerenes and graphene (it is essentially a stretched-out buckyball or a rolled-up sheet of graphene).

Single-wall nanotubes or multi-wall nanotubes. SWNT have a diameter close to $.

Partially crystalline carbon

There are many forms of crystallinity, e.g., carbon black, activated carbon, and carbon fibres.

Carbon black (soot) is a very finely divided form of carbon. Planar stacks (like in graphite) and multilayer balls have been proposed for its structure. Used on a huge scale as a pigment, in printer’s ink, and as a filler for rubber goods.

Activated carbon has a very large surface area (in some cases exceeding \(\SI{1000}{\square\metre\per\gram}\). It is a very efficient adsorbent for molecules, including organic pollutants in water, noxious gases in air, and impurities from reaction mixtures. Usually produced via pyrolysis.

Use of group 14 elements

Coal or coke Graphite, lubricant, pencils. Diamond in industrial cutting surfaces. Silicon as a semiconductor. Germanium is a better intrinsic semiconductor.

PbO is added to glass to raise its refractive index (to form so-called “crystal” glass).

Compounds

All form simple compounds with hydrogen, oxygen, the halogens, and nitrogen.

Only carbon and silicon can form chain-like structures.

Catenation(?)

Hydrides with carbon

The orbitals overlap and form very strong bonds.

Hydrides with silicon

Si forms compounds analogous to the alkanes, silanes.

The longest such chain contains only seven Si: \(\ce{Si_7H_{16}}\).

Si-Si double bonds are uncommon, due to larger size, less p-p overlap.

Hydrides with the rest of them

Weaker as we go down the group.

Halides with carbon

Tetrahalomethanes and partially halogenated alkanes Used because it is possible to replace the halogen with another nucleophilic element/group.

The rate of nucleophilic displacement increases greatly \(F \ll Cl < Br < I\).

Halides of Si, Ge, Sn, and Pb

Mixed valence halides are also known (for di- or tetrahalides).

Oxygen compounds

Uses of \(\ce{CO2}\).

Why is the silicon-oxygen system so rich?

Layered aluminosilicate Kaolinite (china clay). No other metals.

If we instead use magnesium in place of \(\ce{Al}\) we get talc. The repeating layers are neutral and so they readily cleave.

Prominent examples of aluminosilicate compounds are molecular sieves and zeolites.

The sodalite cage is an important building block of zeolites. Describe the structure of the sodalite cage…

Nitrogen compounds

The cyanide ion. isoelectronic with carbon monoxide.

Compounds with metals

Numerous binary compounds between carbon and metal and metalloids, classified into three groups:

  • saline carbides
  • metal carbides
  • metalloid carbides

Silicon-metal compounds (silicides) contain isolated Si, tetrahedral Si4 units, or hexagonal nets of Si.

Lead

According to the International Lead Association, around 5 million tons of lead ores are mined per year. Lead is found in various minerals within Earth’s crust, and has a wide range of industrial uses.

Lead plays an important role in crystalline silicon module manufacturing where it is used in cell interconnects.

Overview on lead from CI (use it here)

https://hackaday.com/2019/10/24/the-blessings-and-destruction-wrought-by-lead-over-millennia/

https://www.pv-magazine.com/2019/10/26/the-weekend-read-a-lead-free-future-for-solar-pv/

Excerpt from the NNNS chemistry blog:

A hexagonal planar transition metal complex is reported from the Crimmin Group (Marti Garçon et al. DOI). In textbook chemistry any metal with 6 ligands either adopts a trigonal prismatic molecular geometry or an octahedral molecular geometry geometry so this arrangement is unusual. The new compound has a central palladium atom with three hydride ligands and three magnesium ligands in a planar alternating fashion.

Group 15: the nitrogen group (pnictogens)

https://en.wikipedia.org/wiki/Pnictogen

https://cen.acs.org/food/agriculture/Periodic-Graphics-elements-fertilizers/98/i11

Nitrogen

https://www.compoundchem.com/2019/02/26/iypt-2019-elements-007-nitrogen-the-element-that-feeds-the-world/

Nitrogen, with five valence electrons, can get a filled octet in several ways.

  • Nitrogen can capture electrons to form the nitride ion, \(\ce{N^{3-}}\).
  • Formation of electron-pair bonds: the octet can be completed by formation of three single bonds (e.g., \(\ce{NH3}\), \(\ce{NF3}\)) or through formation of double or triple bonds (e.g., azo compounds \(\ce{-N=N-}\), nitrogen gas \(\ce{N#N}\)).
  • Formation of electron-pair bonds by capturing electrons to form ions: \(\ce{NH2^-}\) and \(\ce{NH^{2-}}\) (are these the only examples of such behaviour?).
  • Formation of electron-pair bonds by surrendering electrons and forming positively charged ions, e.g., ammonium \(\ce{NH4^+}\), \(\ce{NR4^+}\), and \(\ce{N2R5^+}\).

Nitrides

  • Ionic salt-like nitrides are formed by \(\ce{Mg}\) – \(\ce{Ba}\), \(\ce{Zn}\), \(\ce{Cd}\), \(\ce{Li}\) and \(\ce{Th}\).
  • Covalent nitrides such as \(\ce{BN}\), \(\ce{S4N4}\), and \(\ce{P3N5}\).
  • Interstitial nitrides are formed with the transition metals and have properties very similar to the carbides.

Nitrogen hydrides

  • Ammonia
  • Hydrazine
  • Diimine, \(\ce{HN=NH}\).
  • Hydrogen azid, \(\ce{HN3}\), and azides \(\ce{N_3^-}\).
  • Hydroxylamine

Nitrogen oxides

  • Dinitrogen oxide
  • Nitrogen monoxide
  • Dinitrogen trioxide
  • Nitrogen dioxide
  • Dinitrogen pentoxide

Nitrogen oxo acids

  • Hyponitrous acid, \(\ce{H2N2O2}\) (\(\ce{H-O-N-N-O-H}\)).
  • Nitrous acid, \(\ce{HNO2}\).
  • Nitric acid, \(\ce{HNO3}\).
  • Nitryl ion, \(\ce{NO_2^+}\).

Other nitrogen compounds

Nitrogen trichloride and nitroglycerine

Group 16: chalcogens

Oxygen

https://www.compoundchem.com/2019/03/01/iypt008-oxygen/

The electronic configuration of the free \(\ce{O}\) atom is \(\mathrm{1s^22s^22p^4}\), leading to a \(^3P_2\) ground state.

The ionic radius of \(\ce{O^{2-}}\) is assigned the standard value of \(\SI{140}{\pm}\) and all other ionic radii are derived from this. (Greenwood and Earnshaw 1997, 605)

The oxidation state of oxygen in isolable compounds can vary widely, and includes \(+\frac{1}{2}\), \(0\), \(-\frac{1}{3}\), \(-\frac{1}{2}\), \(-1\), and \(-2\) in species such as \(\ce{O2^+}\), \(\ce{O3}\), \(\ce{O3^-}\), \(\ce{O2^-}\), \(\ce{O2^{2-}}\), and \(\ce{O^{2-}}\), respectively. The coordination number of oxygen in its compounds also varies widely, and stable compounds are known for each CN from 1 to 8 (except possibly 7).

Oxygen compounds of all the elements all known except for \(\ce{He}\), \(\ce{Ne}\), \(\ce{Ar}\) and possibly \(\ce{Kr}\). Oxygen follows the octet rule which is fulfilled in one of the following ways

  • capturing two electrons to give oxide ions, \(\ce{O^{2-}}\),
  • formation of two single bonds, as in water, alcohols (\(\ce{R-O-H}\)) and ethers (\(\ce{R-O-R}\)), or of a double bond (e.g., \(\ce{O=C=O}\)),
  • capturing of an electron and formation of a single bond, e.g., \(\ce{OH-}\),
  • formation of three (or four) covalent interactions, e.g., \(\ce{H3O+}\), \(\ce{R2OH+}\).

When an element forms several oxides the oxide with the highest oxidation state is the most acidic one, e.g., \(\ce{CrO}\) is alkaline, \(\ce{Cr2O3}\) is amphoteric, and \(\ce{CrO3}\) is acidic.

The oxygen fluorides are \(\ce{OF2}\) and \(\ce{O2F2}\).

The dioxogenyl ion, \(\ce{O2^+}\) is known in the compounds \(\ce{O2PtF6}\), \(\ce{O2PF6}\), \(\ce{O2AsF6}\), and \(\ce{O2SbF6}\).

There are two oxygen hydrides: water and hydrogen peroxide, \(\ce{H2O2}\). Hydrogen peroxide is more acidic than water. The oxygen formed at oxidation reactions with hydrogen peroxide comes from the peroxide. Reactions including \(\ce{H2O2}\) are often included in free radical reactions, e.g., \(\ce{H2O^.}\) and \(\ce{OH^.}\).

Ionic peroxides.

Superoxides. Yellow-orange crystals with the formula \(\ce{MO2}\).

Ozonides. Interaction between ozone and \(\ce{KOH}\), \(\ce{RbOH}\) or \(\ce{CsOH}\) gives ozonides, \(\ce{MO3}\).

Other peroxo compounds.

Types of oxides

  • Alkaline oxides. [get back to this]
  • Acidic oxides. Covalent of non-metals are normally acidic and soluble in water, forming acidic solutions, e.g., \(\ce{N2O5(s) + 3H2O -> 2H3O+ + 2NO3^-}\). Insoluble oxides of some less electropositive (i.e., more electronegative) metals in this class are generally soluble in alkaline solution, e.g., \(\ce{Sb2O5(s) + 2OH- + 5H2O -> 2Sb(OH)6^-}\).
  • Amphoteric oxides. These oxides behave as acids in strong bases and as bases in strong acids, e.g., \(\ce{ZnO(s) + 2H3O+ -> Zn^{2+}(aq) + H2O}\) vs \(\ce{ZnO(s) + 2OH- + H2O -> 2Zn(OH)4^{2-}}\).
  • Inert oxides. A number of oxides are relatively inert and they are neither soluble in acids nor bases. Examples are \(\ce{CO}\) and \(\ce{NO}\).
  • Other oxides. Non-stoichiometric oxides or mixed-metal oxides. [get back to this]

Oxygen allotropes

Oxygen gas and liquid oxygen is produced on a vast scale for use in industry (primarily in steel production) by the fractional distillation of liquid air at temperatures near \(\SI{-183}{\degreeCelsius}\). (Greenwood and Earnshaw 1997, 603)

Ozone and oxozone

Sulfur

https://www.compoundchem.com/2019/03/29/iypt016-sulfur/

You can produce hydrogen gas from sulfuric acid using the iodine cycle (iodine acts as a catalyst):

\[\begin{align} \ce{2H2SO4 &-> 2SO2 + 2H2O + O2}~(\SI{830}{\degreeCelsius})\\ \ce{I2 + SO2 + H2O &-> 2HI + H2SO4}~(\SI{120}{\degreeCelsius})\\ \ce{2HI(g) &-> I2 + H2}~(\SI{320}{\degreeCelsius}) \end{align}\]

Group 18: the noble gases

Helium is the only substance remaining liquid down to the absolute zero temperature when cooled under its saturated vapour pressure. Thus, in order to obtain solid helium, elevated pressures have to be applied. Low Temperature Laboratory, Helsinki University of Technology

Helium diverges from the trend in the molar volume (in the solid state) by having a higher-than-normal molar volume. This is due to weak bonding, essentially only induced dipole-induced dipole.

Detailed article on mining and refining helium and whether we are running out of sources of helium.

The \(d\)-block elements

The central role of the d-block metals in the periodic table (RSC, themed collection of papers)

The copper group

Copper, silver, gold, (and roentgenium). All three occur in elemental form in nature. They are probably some of the first metals known to man.

As expected, abundance drops off down the group. Copper is found at 0.0068% by mass or 68 ppm (note that the abundance values in the table below are given in percent). Silver is found at 0.079 ppm and gold at 0.0031 ppm.

Element Mass-% in Earth’s crust/%
Cu 6.8e-03
Ag 7.9e-06
Au 3.0e-07
Rg 0.0e+00

Silver is a slightly better electrical conductor than copper, yet the latter finds more use as such due to its higher abundance and lower cost (the difference in conductivity being only about 5%).

Element Electrical conductivity/(S/m) Relative conductivity
Cu 5.9e+07 1.0000
Ag 6.2e+07 1.0508
Au 4.5e+07 0.7627
Rg NA NA

Next, a transposed table.

Table 1: Some properties of the elements copper, silver and gold. (Could really use the units here).
Property Cu Ag Au
Atomic_Number 29 47 79
Atomic_Weight 63.546 107.8682 196.96655
Boiling_Point 2835.15 2435.15 3129.15
Density 8960 10490 19300
Electrical_Conductivity 5.9e+07 6.2e+07 4.5e+07
Electron_Configuration [Ar]4s13d10 [Kr]5s14d10 [Xe]6s14f145d10
Electronegativity 1.9 1.93 2.54
Melting_Point 1357.77 1234.93 1337.33

Note the significantly higher electronegativity of gold compared to the other elements.

Oxidation states of copper, silver and gold with their most stable oxidation states circled in red.

Figure 3: Oxidation states of copper, silver and gold with their most stable oxidation states circled in red.

Copper

Copper pyrite \(\ce{CuFeS2}\) represents 50% of world deposits of copper. Other important copper minerals are copper glance \(\ce{Cu2S}\), cuprite \(\ce{Cu2O}\), and malachite \(\ce{Cu2CO3(OH)2}\).

Production of metallic copper from copper pyrite: \[\ce{CuFeS2 + SiO2 + 2.5O2 -> Cu + FeSiO3 + 2SO2}\] Copper is used as an electrical conductor (other uses?).

https://www.compoundchem.com/2019/05/21/iypt029-copper/

Silver

Important ores of silver are silver glance \(\ce{Ag2S}\) and silver chloride \(\ce{AgCl}\). Silver is commonly a byproduct of copper, zinc and lead production.

Silver finds use in photography (to a limited extent these days), silverware and jewellery, as an electrical conductor and in batteries.

https://www.compoundchem.com/2019/07/19/iypt047-silver/

Gold

Gold is found in its elemental form in nature, and in tellurides or associated with quartz and pyrite.

Commonly produced via the so-called cyanide process or via amalgamation. Mainly used for corrosion-free contacts (in electronics), the settling of international debts, and jewellery.

https://www.compoundchem.com/2019/10/17/iypt079-gold/

Dubnium

Bohrium

The \(f\)-block elements

New molecules and materials from the f-block (themed collection, RSC)

Complex and coordination chemistry (Gunnar Westin)

Coord compounds known since 150 years.

1893 A Werner published thesis on coordination chemistry. 1913 Nobel prize to Werner for his studies on the structure of complexes.

Werner’s theory: two types of valences.

Related to the development of organic chem at the time (cf. Fischer).

The main characterisation method available at the time was optical rotation spectroscopy, and of course chemical analysis (step-wise precipitation to switch or remove ligands).

\(\ce{Co3+}\) and \(\ce{Cr3+}\) are the two ions that stand out for their slow ligand exchange.

Coordination compound (or complex). Terms

  • coordination number (inner sphere ligands)
  • outer sphere complex
  • inner sphere complex
  • crystal water
  • ligand, donating atom (Lewis base)
  • central atom, acceptor (Lewis acid)

Structural characterisation methods for complex determination

Single-crystal XRD Powder XRD EXAFS WAXS (wide-angle X-ray scattering)

NMR

Spectroscopy: IR Raman UV/Vis/NIR

Gunnar is of course an authority on alkoxides. So you might want to read up on that.

Coordination number (CN) is determined by + centralatom size and charge + steric hindrance of the ligands + electronic interactions

Complexes with \(d\) and \(f\) centralatoms are the most interesting, as their ligand exchange rates can be slow enough to allow stable compounds. So we can build, construct stuff. For \(p\)-block central atoms, the ligand exchange rates are usually very fast.

CN 1 and 2

CN 3: unusual for \(d\)-block metals. Common in the \(p\)-block.

CN 4: two different coordinations possible: tetrahedral, or square planar. More details in powerpoint.

CN 5: two geometries, planpyramidal or trigonal bipyramidal. Interesting biological examples (porphyrin, heme).

CN 6: the most common CN. Three possible geometries: planar, trigonal prismatic, octahedral.

Octahedral complexes exist for all the \(d\)-elements, and a large part of the periodic table.

CN 7: quite unusual for the 3\(d\) elements, but more common for the early 4\(d\) and 5\(d\) elements.

CN 8: quite unusual. Earliest among the early \(d\)-elements, particularly 4\(d\) and 5\(d\). Cubic, or square antiprismatic, or dodecahedral.

CN 9

CN >9

Nomenclature: Flerkärniga komplex innehåller mer än en metallatom. See more in the IUPAC rules.

Isomerism Many types of isomers, e.g.,

Conformational isomers

Geometric isomers

Ionisation isomers

Hydratiseringsisomerer

Bindingsisomerer Koordinationsisomerer Ligandisomerer

Optical isomers

https://en.wikipedia.org/wiki/Metal_aquo_complex

Crystal field theory

It is a simple model, very intuitive, but limited in that it only considers ionic contributions.

Ligands considered as negative point charges. Repels the electrons in the d-orbitals, but differently for different d-orbitals causing a split in their energy levels.

Two types of \(d\)-orbitals:

  • \(e_\text{g}\): along x, y, z axes. \(d_{x^2-y^2}\), \(d_{z^2}\)
  • \(t_\text{2g}\): between x, y, z axes. \(d_{xz}\), \(d_{yz}\), \(d_{xy}\)

Högspinn: “as many spins as possible” Lågspinn: the opposite

Which of low- or high-spin that occurs is determined by the pairing repulsion vs ligand field strength.

Expand on this: Tanabe-Sugane diagrams.

Olika ligander. Different Lewis base strength (how much electrons can each donate?) Den spektrokemiska serien. Bör kunnas utantill.

\[\begin{equation} \ce{I < Br- < S2- < SCN- < Cl- < NO3- < F- < OH- < C2O42- < H2O < NCS- < C3CN < NH3 < en < biby < phen < NO2- < PPh3 < CO} \end{equation}\]

And vice versa for the central atoms:

\[\begin{equation} \ce{Mn2+ < Ni2+ < Co2+ < Fe2+ < V2+ < Fe3+ < Co3+ < Mn4+ < Mo3+ < Rh3+ < Ru3+ < Pd4+ < Ir3+ < Pt4+} \end{equation}\]

for each increase in the oxidation number you double the ligand field if you go down a group, typically you increase the ligand field by 50%

Ligandfältsstabilisering (LFSE)

\(\mathrm{LFSE} = \Delta_O\times(-0.4x + 0.6y)\)

We will come back to discuss ligand exchange rate (Jahn Teller distortion etc.)

Heterogeneous catalysis (lektion av Martin Häggblad)

A catalyst is a substance that increases the rate of a reaction but is not itself consumed. (Shriver et al. 2014, 728)

A catalyst increases the rates of processes by introducing new pathways with lower Gibbs energies of activation, \(\Delta^\ddagger G\). A catalyst does not affect the Gibbs energy of the overall reaction, \(\Delta_\text{r}G^\circ\), because \(G\) is a state function. Reactions that are thermodynamically unfavourable cannot be made favourable by a catalyst.

A substance that increases the rate of a reaction without modifying the overall standard Gibbs energy change in the reaction; the process is called catalysis. The catalyst is both a reactant and product of the reaction. […] Catalysis can be classified as homogeneous catalysis, in which only one phase is involved, and heterogeneous catalysis, in which the reaction occurs at or near an interface between phases. Catalysis brought about by one of the products of a reaction is called autocatalysis. — IUPAC (2012), p. 220

Some terminology:

  • catalytic cycle
  • catalytic efficiency and lifetime
  • selectivity
  • active site (IUPAC 2012, 29)
  • poison (IUPAC 2012, 1141). An inhibitory substance characterized by its propensity to attach very strongly, by a true chemical bond (e.g., covalent) to the surface atoms or ions constituting the catalytically active sites. Poisons act in minute quantities. A product of the catalysed reaction reaction may cause poisoining or inhibition. This phenomenon is called self-poisoning.
  • fouling agent (IUPAC 2012)
  • turnover number, defined as molecules reacting per active site in unit time. (IUPAC 2012, 1574)

An example reaction (with applicability to internal combustion engines):

\[\ce{N2(g) + O2(g) <=> 2NO(g)}\]

with equilibrium constant, \(K\)

\[K = \frac{P^2_\ce{NO}}{P_\ce{N2}P_\ce{O2}}\]

where \(K\) depends on temperature, being very small at RT and increasing with temperature. (insert table if wanted)

An important concept in catalysis: selectivity.

A given reaction might have many possible products. The definition of selectivity: the selectivity of process \(i\), \(S_i\), is the number of moles

Another concept is reactivity. That is, how good is the catalyst? A particular measure of that is the specific turnover frequency (TN). Or, how many reactions per second?

\[TN = \frac{number of mole reacted A}{number of mole reaction sites \times time}\]

The importance of the catalyst area

Basically, the larger the area the more available reaction sites (per gram of catalyst).

Many ways to get a large surface area for a given mass of catalyst:

  • porous catalyst
  • highly disperse catalyst

A classic example of a porous catalyst is Raney nickel. Used for the hydrogenation of benzene into cyclohexane, for example.

If we assume spherical particles with \(\rho=\SI{2}{\gram\per\cubic\cm}\) and diameter \(d\). Total surface area as a function of diameter: as \(d\) decreases, the surface area increases as the cube of the diameter.

\(d = \SI{1}{\m} \longrightarrow A = \SI{3E-3}{\square\metre\per\gram}\) \(d = \SI{1}{\micro\m} \longrightarrow A = \SI{3}{\square\metre\per\gram}\) \(d = \SI{100}{\angstrom} \longrightarrow A = \SI{300}{\square\metre\per\gram}\)

Common choice: a transitiion metal

Catalysis on metallic surfaces is very know.

Rule of thumb: the bond strength (M-X) will decrease with the number of \(d\)-electrons.

\[\ce{CO(g) + O2(g) <=> CO2(g)}\]

On we will form a strong bond between the metal surface and the adsorbed species, making the reaction above slow. For something like , the interaction between the adsorbate and the surface will be much weaker, again making for a poor catalyst.

Important concept to master: volcano plots.

Sabatier’s principle: an optimum bond strength exists that maximises the reaction rate.

Important industrial applications

Carbon monoxide + oxygen to carbon dioxide, and also unreacted hydrocarbons that react with oxygen to form carbon dioxide and water. Catalyst is usually Pt/Rh on some carrier.

Important consideration: a high oxygen partial pressure leads to a better oxidation reaction (and also a higher temperature in the engine), but is bad for the reaction of to and . So this requires optimisation…

Haber-Bosch synthesis of ammonia https://rootsofprogress.org/turning-air-into-bread

An exothermal reaction.

  • increase pressure leads to more ammonia
  • decrease temperature leads to more ammonia

So that sounds like easy? Just increase pressure and decrease temperature? Nope, does not work because the activation energy is too high. The common solution to increase the reaction rate, to increase temperature, won’t work at all in this case as the reaction is exothermic.

So the only way around this is to use a catalyst. In the beginning of the 1900s a concerted effort (around 20000 experiments) was made to find the best catalyst. Iron was selected.

Not all surfaces are the same

Which facet

Steps, kinks, corners

Defects and other imperfections

Questions

  • Do salt-like hydrides have high melting points in general, or is it particular to alkali and alkaline earth metal hydrides?

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Notes and references

In the creation of this text I am indebted to my teachers and colleagues at the department of chemistry at Uppsala university and the department of chemistry at SLU. In particular, I would like to thank Rickard Eriksson, on whose slides I based some of this text, and

https://www.compoundchem.com/2020/02/10/women-periodic-table/

Greenwood, N. N., and A. Earnshaw. 1997. Chemistry of the Elements. Second. Butterworth-Heinemann.
IUPAC. 2012. Compendium of Chemical Terminology: IUPAC Gold Book. IUPAC.
Shriver, Duward, Mark Weller, Tina Overton, Jonathan Rourke, and Fraser Armstrong. 2014. Inorganic Chemistry. Sixth. Oxford University Press.