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 \(\ce{He}\).

Electronic configuration

VSEPR valence shell electron pair repulsion theory

Formally applies only to molecular compounds and complex ions in the gaseous and liquid states. It applies only to main group elements. The VSEPR approach is limited to \(d^0\), \(d^{10}\), \(f^0\) and \(f^{14}\) atoms.

Hybridisations and their configuration (around central atom):

  • \(sp\) linear
  • \(sp^2\) triangular
  • \(sp^3\) tetrahedral
  • \(dsp^3\) trigonal bipyramidal
  • \(d^2sp^3\) octahedral
  • \(dsp^2\) square-planar
  • \(d^5sp^3\) tri-capped trigonal prism (CN 9)

Hund’s rule

No orbital can contain two electrons until each orbital of the same energy is singly occupied.

Shielding and penentration

Radii (atomic, ionic)

Lanthanide contraction.

Slater’s rules

Slater’s rules, for determining effective charge, \(Z_\text{eff}\).

Note: does not take into consideration if next electron enters occupied orbital (pairing, electron-pair repulsion), half-filled or filled shells, etc.

Ionisation energy, electron affinity

Electronegativity

“The power of an atom in a molecule or compound to attract an electron to itself.”

Mulliken’s definition:

\begin{equation} \chi_\text{Mulliken} = \frac{\mathrm{IE} - \mathrm{EA}}{2} \end{equation}

Pauling’s definition of electronegativity. (More complicated).

Hydrogen, deuterium, tritium

Hydrides

  • The noble gases form hydrides.
  • The alkali and alkaline earth metal ions form salt-like hydrides with i) high melting points, ii) evolve hydrogen gas at the anode, and iii) 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, iii) 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.

Industral 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

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 \(\ce{Li2O2}\).
  • Sodium burns in both air and oxygen, forms \(\ce{Na2O}\) which is further oxidised to \(\ce{Na2O2}\).
  • \(\ce{K}\), \(\ce{Rb}\), \(\ce{Cs}\): 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

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.

Beryllium and beryllium compounds

  • 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) \(\ce{3Be + N2 -> Be3N2}\), iii) \(\ce{Be + C2H2 -> BeC2 + 2H2}\).
  • Complexes: i) the beryllate ion, \(\ce{Be(OH)4^{2-}}\), ii) \(\ce{BeF4^{2-}}\), iii) \(\ce{BeCl4^{2-}}\), 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 i) 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.

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.

Magnesium, 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 13: boron, aluminium, gallium, indium, thallium, nihonium

Aluminium, gallium, indium and thallium

Oxidation numbers

  • Boron, only \(+\)III

Group 15: the nitrogen group

Nitrogen

Nitrogen 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, \(\ce{NO_2^+}\).

Group 16: chalcogens

Oxygen

The electronic configuration of the free 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 & Earnshaw, 1997, p. 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 He, Ne, Ar and possibly 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 CO and NO.
  • Other oxides. Non-stoichiometric oxides or mixed-metal oxides. [get back to this]

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 & Earnshaw, 1997, p. 603)

Sulfur

You can produce hydrogen gas 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: 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.

Heterogeneous catalysis

A catalyst is a substance that increases the rate of a reaction but is not itself consumed. (Shriver, Weller, Overton, Rourke, & Armstrong, 2014, p. 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)

Some terminology:

  • catalytic cycle
  • catalytic efficiency and lifetime
  • selectivity
  • active site (IUPAC, 2012, p. 29)
  • poison (IUPAC, 2012, p. 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, p. 1574)

Crystal field theory

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. \(d_{xz}\), \(d_{yz}\), \(d_{xy}\)

  • http://chemistry.stackexchange.com/a/41967

Notes and questions

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

References

  1. Greenwood, N. N., & Earnshaw, A. (1997). Chemistry of the Elements (second). Butterworth-Heinemann.

  2. Shriver, D., Weller, M., Overton, T., Rourke, J., & Armstrong, F. (2014). Inorganic Chemistry (sixth, p. 901). Oxford University Press.

  3. IUPAC. (2012). Compendium of Chemical Terminology – IUPAC Gold Book. IUPAC. Retrieved from http://goldbook.iupac.org/PDF/goldbook.pdf

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