Iron Metabolism

From Molecular Mechanisms to Clinical Consequences

Robert Crichton

• Publisher: Wiley
• Year: 2016
• Print ISBN: 9781118925614
• eBook ISBN: 9781118925638
• Edition: 4
• Topic: Medicine
• Subtopic: Endocrinology & Metabolism

Iron is indispensable for the growth, development and well-being of almost all living organisms. Biological systems from bacteria, fungi and plants to humans have evolved systems for the uptake, utilisation, storage and homeostasis of iron. Its importance for microbial growth makes its uptake systems  a natural target for  pathogenic microorganisms and parasites. Uniquely, humans  suffer from both iron deficiency and iron overload, while the capacity of iron to generate highly reactive free radicals, causing oxidative stress, is  associated with a wide range of human pathologies, including many neurodegenerative diseases. Whereas some essential metal ions like copper and zinc are closely linked with iron metabolism, toxic metals like aluminium and cadmium can interfere with iron metabolism. Finally, iron metabolism and homeostasis are key targets for the development of new drugs for human health.
The 4^th edition of Iron Metabolism  is written in a lively style by one of the leaders in the field, presented in colour  and covers the latest discoveries in this exciting area. It will be essential reading for researchers and students in biochemistry, molecular biology, microbiology, cell biology, nutrition and medical sciences. Other interested groups include biological inorganic chemists with an interest in iron metabolism, health professionals with an interest in diseases of iron metabolism, or of diseases in which iron uptake systems are involved (eg. microbial and fungal infections, cancer, neurodegenerative disorders), and researchers in the pharmaceutical industry interested in developing novel drugs targeting iron metabolism/homeostasis.

1
Solution Chemistry of Iron

1.1 Iron Chemistry

In the Earth’s crust, iron is the fourth most abundant element and the second most abundant metal (the most abundant is aluminium). Situated in the Periodic Table in the middle of the first transition series (characterised by having incompletely filled d orbitals), iron has access to a number of oxidation states (from –II to +VI), the principal being II (d⁶) and III (d⁵). A number of iron-dependent monooxygenases are able to generate high-valent Fe(IV) or Fe(V) reactive intermediates during their catalytic cycle. Whereas, Fe²⁺ is extremely water-soluble, Fe³⁺ is quite insoluble in water (K_sp = 10^–39 M and at pH 7.0, [Fe³⁺] = 10^–18 M) and significant concentrations of water-soluble Fe³⁺ species can be attained only by strong complex formation with appropriate ligands.

The interaction between Fe²⁺ and Fe³⁺ and ligand donor atoms will depend on the strength of the chemical bond formed between them. An idea of the strength of such bonds can be found in the concept of ‘hard’ and ‘soft’ acids and bases (HSAB) (Pearson, 1963). ‘Soft’ bases have donor atoms of high polarisability with empty, low-energy orbitals; they usually have low electronegativity and are easily oxidised. In contrast, ‘hard’ bases have donor atoms of low polarisability, and only have vacant orbitals of high energy; they have high electronegativity and are difficult to oxidise. Metal ions are ‘soft’ acids if they are of low charge density, have a large ionic radius. and have easily excited outer electrons. ‘Hard’ acid metal ions have high charge density, a small ionic radius, and no easily excited outer electrons. In general, ‘hard’ acids prefer ‘hard’ bases and ‘soft’ acids form more stable complexes with ‘soft’ bases (Pearson, 1963). Fe(III) with an ionic radius of 0.067 nm and a charge of 3⁺ is a ‘hard’ acid and will prefer ‘hard’ oxygen ligands such as phenolate and carboxylate, compared to imidazole or thiolate. In contrast, Fe(II) with an ionic radius of 0.083 nm and a charge of only 2⁺ is on the borderline between ‘hard’ and ‘soft,’ favouring nitrogen (imidazole and pyrrole) and sulphur ligands (thiolate and methionine) over oxygen ligands.

The coordination number of 6 is the most frequently found for both Fe(II) and Fe(III) giving octahedral stereochemistry, although four-coordinate (tetrahedral) and particularly five-coordinate complexes (trigonal bipyramidal or square pyrimidal) are also found. For octahedral complexes, two different spin states¹ can be observed. Strong-field ligands (e.g. Fe³⁺ OH^–), where the crystal field splitting is high and hence electrons are paired, give low-spin complexes, while weak-field ligands (e.g. CO, CN^–), where crystal field splitting is low, favour a maximum number of unpaired electrons and give high-spin complexes Changes of spin state affect the ion size of both Fe(II) and Fe(III), the high-spin ion being significantly larger than the low-spin ion. As we will see in Chapter 2, this is put to good use as a trigger for the cooperative binding of dioxygen to haemoglobin. High-spin complexes are kinetically labile, while low-spin complexes are exchange-inert. For both oxidation states only high-spin tetrahedral complexes are formed, and both oxidation states are Lewis acids, particularly the ferric state.

The unique biological role of iron comes from the extreme variability of the Fe²⁺/Fe³⁺ redox potential, which can be fine-tuned by well-chosen ligands, so that iron sites can encompass almost the entire biologically significant range of redox potentials, from about –0.5 V to about +0.6 V. However, as we will see in Chapter 13, copper allows access to an even higher range of redox potentials (0 V to +0.8 V), which turned out to be of crucial importance in the Earth’s rapidly evolving aerobic environment, following the arrival of water-splitting, oxygen-generating photosynthetic organisms.

1.2 Interactions of Iron with Dioxygen and Chemistry of Oxygen Free Radicals

Molecular oxygen was not present when life began on Earth, with its essentially reducing atmosphere, and both the natural abundance of iron and its redox properties predisposed it to play a crucial role in the first stages of life on Earth. About one billion (10⁹) years ago, photosynthetic prokaryotes (Cyanobacteria) appeared and dioxygen was evolved into the Earth’s atmosphere. It probably required 200–300 million years – a relatively short time on a geological time scale – for oxygen to attain a significant concentration in the atmosphere, since at the outset the oxygen produced by photosynthesis would have been consumed by the oxidation of ferrous ions in the oceans. Once dioxygen had become a dominant chemical en...

Table of contents

• Cover
• Title Page
• Table of Contents
• Preface
• 1 Solution Chemistry of Iron
• 2 The Essential Role of Iron in Biology
• 3 Microbial Iron Uptake
• 4 Iron Acquisition by Pathogens
• 5 Iron Uptake by Plants and Fungi
• 6 Cellular Iron Uptake and Export in Mammals
• 7 Mammalian Iron Metabolism and Dietary Iron Absorption
• 8 Intracellular Iron Utilisation
• 9 Iron Storage Proteins
• 10 Cellular and Systemic Iron Homeostasis
• 11 Iron Deficiency, Iron Overload and Therapy
• 12 Iron and Immunity
• 13 Iron and Oxidative Stress
• 14 Interactions between Iron and other Metals
• 15 Iron Homeostasis and Neurodegeneration
• Concluding Remarks
• Index
• End User License Agreement