Chemical structures

Chemical structures of (a) lithium carbonate, (b) potassium bicarbonate and (c) lithium citrate
than NaCl at a basic pH. Without the addition of NH3, the solution would be acidic, as HCl is produced as a
NaCl + CO2 + NH3 + H2O → NaHCO3 + NH4Cl (2.8)
NaHCO3 is filtered, and the remaining solution of NH4Cl reacts with the CaO from step 1. The produced
CaCl2 is usually used as road salts, and NH3 is recycled back to the initial reaction of NaCl (reaction step 2).
2NH4Cl + CaO → 2NH3 + CaCl2 + H2O (2.9)
In a final step, NaHCO3 is converted to Na2CO3 (sodium carbonate) by calcination (heating to 160–230 ∘C)
with the loss of water.
2NaHCO3 → Na2CO3 + H2O + CO2 (2.10)
Again, the CO2 can be recycled and reused within the process. This means that the Solvay process consumes
only a very small amount of NH3 and the only ingredients are NaCl and CaCO3 (Figure 2.5).
In terms of pharmaceutical applications, the various salts of sodium, lithium and potassium have the widest
use and are discussed in detail in the following sections. These applications range from the use of sodium and
potassium salts in dehydration solutions in order to restore replenished mineral balances to the treatment of
BD with simple lithium salts.
Figure 2.5 The Solvay process
Alkali Metals 29
2.2 Advantages and disadvantages using lithium-based drugs
The name ‘lithium’ stems from the Greek word ‘lithos’ which means stone. Lithium salts are well known for
their use in batteries, metal alloys and glass manufacture. Nevertheless, Li also has a clinical application in the
treatment of manic depression or BD. BD affects 1–2% of the population and severely reduces the quality of
life for the patients and also increases the likelihood of patients committing suicide. Research has shown that
lithium salts are very successful in the treatment of BD, and a broad research in this area has been stimulated
[3]. The main question addressed within this research is how the lithium ion (Li+) can be modified in order to
improve its activity to patent their findings, which would result in income for the manufacturer. Unfortunately,
the simple ion is the active ingredient and this makes it difficult to patent and secure any intellectual property.
Li is a member of group 1 alkali metals and is the lightest and smallest solid element. Li has an atomic
number of 3 and contains a single valence electron, which determines its redox chemistry and reactivity.
In nature, Li is found as ores, for example, spodumene [LiAl(SiO3)2], or at low concentrations as salts, for
example, in rivers. The metal itself is soft and white in appearance; it has the lowest reactivity within the
group of alkali metals. The Li+ ion has the smallest ionic radius of all known metals.
2.2.1 Isotopes of lithium and their medicinal application
Lithium occurs as mixture of two stable, naturally occurring isotopes (see Section, namely 6Li with
an occurrence of 7.59% and the major isotope 7Li (92.41%). The nucleus of 6Li contains three protons and
three neutrons, whilst 7Li contains three protons and four neutrons.
6Li and 7Li are both NMR (nuclear magnetic resonance) active nuclei, which means their presence can be
monitored via NMR technology. Using this analytical tool, it is possible to differentiate between intra- and
extracellular Li+ concentrations and therefore, the uptake of Li+ into different cells can be monitored. The use
of 6Li results in sharper NMR spectra because of its properties, but it also has a lower intensity. 6Li salts can
also be used to monitor the distribution of lithium in the tissue. A further application of 6Li is the production
of tritium atoms (3
1H) and their use in atomic reactors. In this process, 6Li is bombarded with neutrons, which
results in the production of tritium atoms and radioactive 𝛼-particles (4
2He (2.11)
Equation 2.11 shows the activation of 6Li.
2.2.2 Historical developments in lithium-based drugs
The first medical use of Li+ was described in 1859 for the treatment of rheumatic conditions and gout. The
theory at that time was based on the ability of lithium to dissolve nitrogen-containing compounds such as uric
acid. Their build-up in the body was believed to cause many illnesses such as rheumatic conditions and gout
problems. In 1880, Li+ was first reported as being used in the treatment of BD, and in 1885 lithium carbonate
(Li2CO3) and lithium citrate [Li3C3H5O(COO)3] were included in the British Pharmacopoeia. It also became
clear that there is a direct link between NaCl intake and heart diseases as well as hypertension. Therefore,
LiCl was prescribed as replacement for NaCl in the diet of affected patients [3b].
The urea hypothesis and the connection of NaCl to heart diseases stimulated the use of lithium salts in
common food. The prime example is the soft drink 7Up©, which has been marketed in 1929 under the label
30 Essentials of Inorganic Chemistry
Bib-Label Lithiated Lemon-Lime Soda. 7Up contained lithium citrate and was also marketed as a hang-over
cure. The actual Li+ was subsequently removed in 1950 [3b].
John Cade’s experiments on guinea pigs in 1949 initiated the discovery of Li+ and its sedative and
mood-control properties. Uric acid was known to have mood-controlling properties and Cade used lithium
urate as a control solution. To his surprise, he discovered that lithium urate had tranquillising properties, and
after further experiments he concluded that this was caused by the lithium ion [4].
Nevertheless, there were drawbacks, especially when the FDA banned Li+ salts following the death of four
US patients. These patients had an average intake of 14 g of lithium chloride (LiCl) per day in order to replace
NaCl. Another stumbling stone in the way of success of Li+ was the discovery of chlorpromazine, the first
antipsychotic drug, which is still used for the treatment of BD. In the early 1970s, Li+ was re-approved by
FDA and is now used in 50% of the treatment of BD [5].
2.2.3 The biology of lithium and its medicinal application
Lithium salts are used in the prophylaxis and treatment of mania, and in the prophylaxis of BD and recurrent depression. Lithium therapy is taken orally, usually as lithium carbonate (Li2CO3) or lithium citrate
[Li3C3H5O(COO)3], with a total dose of up to 30 mmol/day. Li2CO3 is the preferred lithium salt used, as it
causes the least irritation to the stomach. The treatment has to be closely monitored, and Li+ blood concentrations are measured 12 h after administration to achieve a serum lithium concentration of 0.4–1 mmol/l. The
therapeutic index (concentration window from efficacy to toxicity) for Li+ is very narrow, and plasma Li+
concentrations above 2 mM require emergency treatment for poisoning (Figure 2.6) [6].
The administered Li+ is distributed uniformly in the body tissues and in the blood plasma, with the external
cell Li+ concentration being below 2 mM. Experiments studying the lithium distribution in rats after administration of a high dose of 6Li+ showed that there was no exceptional accumulation of 6Li+ in the brain.
6Li+ distributes fairly uniformly in the body, with bones and endocrine glands showing higher concentrations
(Figure 2.7) [7].
Li+ ions are not soluble in lipids and therefore do not cross plasma membranes. The transport into cells
occurs via exchange mechanism by lithium–sodium counter-transport, anion exchange (so-called Li+/CO3
co-transport) and other unrelated transport molecules. The specific mode of action of the simple Li+ ion is
currently unknown, but it is clear that a displacement of Mg2+ by Li+ is involved. Therefore, an alteration of
the Mg2+ balance in the blood and the urine can be observed in patients treated with Li+ [3b]. This displacement is actually not surprising because the properties of Li+ and Mg2+ are similar, which can be explained
by the concept of diagonal relationship (see Section 2.2.4).
(a) (b)
Figure 2.6 Chemical structures of (a) lithium carbonate and (b) lithium citrate
Alkali Metals 31
[LI] mmol/kg wet weight
Figure 2.7 Distribution of 6Li+ in various tissues of rats after chronic administration [7] (Reproduced with
permission from [7]. Copyright © 1999, Royal Society of Chemistry.)
2.2.4 Excursus: diagonal relationship and periodicity
Within the periodic table, the element pairs Li/Mg, Be/Al, B/Si and others form a so-called diagonal relationship to each other. With this concept, similarities in biological activity can be explained as the element
pairs have similar properties.
Diagonally adjacent elements of the second and third periods have similar properties – this concept is
called diagonal relationship. These pairs (Li/Mg, Be/Al, B/Si, etc.) have similarities in ion size, atomic
radius, reactive behaviour and other properties.
Diagonal relationship is a result of opposite effects, which crossing and descending within the periodic
table has. The size of an atom decreases within the same period (from left to right). The reason is that a
positive charge is added to the nucleus together with an extra electron orbital. Increasing the nucleus charge
means that the electron orbitals are pulled closer to the nucleus. The atomic radius increases when descending
within the same group, which is due to the fact that extra orbitals with electrons are added (Figure 2.8) [1].
Trends can be seen for the electronegativity and the ionisation energy; both increase when moving within
the same period from left to right and decrease within the same group. Within a small atom, the electrons are
located close to the nucleus and held tightly, which leads to a high ionisation energy. Therefore, the ionisation
energy decreases within the group and increases within the period (from left to right) – showing the opposite
trend compared to the atomic radius. A similar explanation can also be used for summarising the trends seen
for the electronegativity: small atoms tend to attract electrons more strongly than larger ones. This means
fluorine is the most electronegative element within the Periodic Tables of Elements (Figure 2.9).
This structure of trends is summarised under the term periodicity of the elements. This means that within
the periodic table, all elements follow the above-mentioned trends in a repetitive way. This allows predicting
trends for atomic and ionic radii, electronegativity and ionisation energy (Figure 2.10).
In relation to the diagonal relationship, it becomes clear by studying the described trends that these effects
cancel each other when descending within the group and crossing by one element within the PSE. Therefore,
32 Essentials of Inorganic Chemistry
Li Be
Na Mg
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe
LaLu Hf Ta W Re Os Ir Pt Au Hg TI Pb Bi Po At Rn
AcLr Rf Db Sg Bh Hs Mt Ds Rg Uub
B C N O F Ne
Rb Sr
Cs Ba
Fr Ra
Figure 2.8 Atomic radii
Li Be
Na Mg
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe
LaLu Hf Ta W Re Os Ir Pt Au Hg TI Pb
Bi Po At Rn
AcLr Rf Db Sg Bh Hs Mt Ds Rg Uub
B C N O F Ne
Rb Sr
Cs Ba
Fr Ra
Figure 2.9 Electronegativity
elements diagonally positioned within the periodic table have similar properties, such as similar atomic size,
electronegativity and ionisation energy (Figure 2.11).
The concept of the diagonal relationship is crucial for the biological activity of lithium drugs, which is
mainly due to the properties of the Li+ ion being similar to the Mg2+ ion. In comparison, the size of the Li+
ion is similar to that of Mg2+ and therefore they compete for the same binding sites in proteins. Nevertheless,
lithium has relatively specific effects, and so only proteins with a low affinity for Mg2+ are targeted. Li+
and Mg2+ salts have similar solubility, for example, CO3
2−, PO4
3−, F− salts have a low water solubility, and
halide and alkyl salts are soluble in organic solvents. Li+ and Mg2+ compounds are generally hydrated, for
example, LiCl⋅3H2O and MgCl2⋅6H2O. Similarities in ionic size, solubility, electronegativity and solubility
result in similar biological activity and therefore pharmaceutical application [3b].
Alkali Metals 33
Li Be
Na Mg
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe
LaLu Hf Ta W Re Os Ir Pt Au Hg TI Pb
ionisation energy
Atomic radii
Bi Po At Rn
AcLr Rf Db Sg Bh Hs Mt Ds Rg Uub
B C N O F Ne
Rb Sr
Cs Ba
Fr Ra
Figure 2.10 Periodicity showing the ‘metallic character’ trend (highlighted in grey) within the periodic table
Li Be
Na Mg
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe
LaLu Hf Ta W Re Os Ir Pt Au Hg TI Pb Bi Po At Rn
AcLr Rf Db Sg Bh Hs Mt Ds Rg Uub
B C N O F Ne
Rb Sr
Cs Ba
Fr Ra

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