Ferro

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Ferro,  26 Fe
Pure iron chips with a high purity iron cube
Ferro
Aparênciametálico brilhante com um tom acinzentado
Peso atômico padrão A r, std (Fe) 55,845 (2) [1]
Ferro na tabela periódica
Hidrogênio Hélio
Lítio Berílio Boro Carbono Azoto Oxigênio Flúor Néon
Sódio Magnésio Alumínio Silício Fósforo Enxofre Cloro Argônio
Potássio Cálcio Escândio Titânio Vanádio Cromo Manganês Ferro Cobalto Níquel Cobre Zinco Gálio Germânio Arsênico Selênio Bromo Krypton
Rubídio Estrôncio Ítrio Zircônio Nióbio Molibdênio Tecnécio Rutênio Ródio Paládio Prata Cádmio Índio Lata Antimônio Telúrio Iodo Xenon
Césio Bário Lantânio Cério Praseodímio Neodímio Promécio Samário Europium Gadolínio Térbio Disprósio Holmium Erbium Túlio Itérbio Lutécio Háfnio Tântalo Tungstênio Rênio Ósmio Iridium Platina Ouro Mercúrio (elemento) Tálio Liderar Bismuto Polônio Astatine Radon
Francium Rádio Actínio Tório Protactínio Urânio Neptúnio Plutônio Americium Curium Berquélio Californium Einsteinium Fermium Mendelévio Nobelium Lawrencium Rutherfordium Dubnium Seabórgio Bohrium Hassium Meitnerium Darmstádio Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson
-

Fe

Ru
manganêsferrocobalto
Número atômico ( Z )26
Grupogrupo 8
Períodoperíodo 4
Bloquear  bloco d
Configuração de elétron[ Ar ] 3d 6 4s 2
Elétrons por camada2, 8, 14, 2
Propriedades físicas
Fase em  STPsólido
Ponto de fusão1811  K (1538 ° C, 2800 ° F)
Ponto de ebulição3134 K (2862 ° C, 5182 ° F)
Densidade (próximo à  rt )7,874 g / cm 3
quando líquido (em  mp )6,98 g / cm 3
Calor de fusão13,81  kJ / mol
Calor da vaporização340 kJ / mol
Capacidade de calor molar25,10 J / (mol · K)
Pressão de vapor
P  (Pa) 1 10 100 1 mil 10 k 100 k
em  T  (K) 1728 1890 2091 2346 2679 3132
Propriedades atômicas
Estados de oxidação−4, −2, −1, 0, +1, [2] +2 , +3 , +4, +5, [3] +6 , +7 [4] (um  óxido anfotérico )
Eletro-negatividadeEscala de Pauling: 1,83
Energias de ionização
  • 1o: 762,5 kJ / mol
  • 2º: 1561,9 kJ / mol
  • 3o: 2957 kJ / mol
  • ( mais )
Raio atômicoempírico: 126  pm
Raio covalenteGiro baixo: 132 ± 3 pm
Giro alto: 152 ± 18 h
Raio de Van der Waals194 [1]  pm
Color lines in a spectral range
Linhas espectrais de ferro
Outras propriedades
Ocorrência naturalprimordial
Estrutura de cristalcúbica de corpo centrado (BCC)
Body-centered cubic crystal structure for iron

a = 286,65 pm
Estrutura de cristalcúbica de face centrada (FCC)
Face-centered cubic crystal structure for iron

entre 1185–1667 K; a = 364,680 pm
Velocidade do som haste fina5120 m / s (à  temperatura ambiente ) (eletrolítico)
Expansão térmica11,8 µm / (m⋅K) (a 25 ° C)
Condutividade térmica80,4 W / (m⋅K)
Resistividade elétrica96,1 nΩ⋅m (a 20 ° C)
Ponto Curie1043 K
Ordenação magnéticaferromagnético
Módulo de Young211 GPa
Módulo de cisalhamento82 GPa
Módulo de massa170 GPa
Coeficiente de Poisson0,29
Dureza de Mohs4
Dureza Vickers608 MPa
Dureza Brinell200-1180 MPa
Número CAS7439-89-6
História
Descobertaantes de 5000 aC
Símbolo"Fe": do latim ferrum
Principais isótopos de ferro
Isótopo Abundância Meia-vida ( t 1/2 ) Modo de decaimento produtos
54 Fe 5,85% estábulo
55 Fe syn 2,73 anos ε 55 Mn
56 Fe 91,75% estábulo
57 Fe 2,12% estábulo
58 Fe 0,28% estábulo
59 Fe syn 44,6 d β - 59 Co
60 Fe vestígio 2,6 × 10 6  y β - 60 Co
Category Categoria: Ferro
| referências

Ferro ( / ər n / ) é um elemento químico com símbolo Fe (a partir Latina : Ferrum ) e número atómico 26. É um metal de que pertence à primeira série de transição e o grupo 8 da tabela periódica . É, em massa, o elemento mais comum na Terra , bem na frente do oxigênio (32,1% e 30,1%, respectivamente), formando grande parte do núcleo externo e interno da Terra. É o quarto elemento mais comum na crosta terrestre .

Em seu estado metálico, o ferro é raro na crosta terrestre , limitado principalmente à deposição de meteoritos . Os minérios de ferro , por outro lado, estão entre os mais abundantes na crosta terrestre, embora a extração de metal utilizável deles requeira fornos ou fornos capazes de atingir 1.500 ° C (2.730 ° F) ou mais, cerca de 500 ° C (900 ° F) mais do que o necessário para fundir cobre . Os humanos começaram a dominar esse processo na Eurásia por volta de 2000 AC, [ não verificado no corpo ] e o uso de ferramentas e armas de ferro começou a deslocar ligas de cobre, em algumas regiões, apenas por volta de 1200 aC. Esse evento é considerado a transição da Idade do Bronze para a Idade do Ferro . No mundo moderno , as ligas de ferro, como aço , aço inoxidável , ferro fundido e aços especiais são, de longe, os metais industriais mais comuns, devido às suas propriedades mecânicas e baixo custo.

As superfícies de ferro puro lisas e imaculadas são cinza-prateadas espelhadas. No entanto, o ferro reage prontamente com o oxigênio e a água para dar óxidos de ferro hidratados de marrom a preto , comumente conhecidos como ferrugem . Ao contrário dos óxidos de alguns outros metais, que formam camadas passivantes , a ferrugem ocupa mais volume do que o metal e, portanto, descama, expondo as superfícies novas à corrosão. Embora o ferro reaja prontamente, o ferro de alta pureza, denominado ferro eletrolítico , tem melhor resistência à corrosão.

O corpo de um ser humano adulto contém cerca de 4 gramas (0,005% do peso corporal) de ferro, principalmente em hemoglobina e mioglobina . Essas duas proteínas desempenham papéis essenciais no metabolismo dos vertebrados , respectivamente no transporte de oxigênio pelo sangue e no armazenamento de oxigênio nos músculos . Para manter os níveis necessários, o metabolismo do ferro humano requer um mínimo de ferro na dieta. O ferro é também o metal no local activo de muitos importantes redox enzimas relacionadas com a respiração celular e de oxidação e redução em plantas e animais. [5]

Quimicamente, os estados de oxidação mais comuns do ferro são ferro (II) e ferro (III) . O ferro compartilha muitas propriedades de outros metais de transição , incluindo os outros elementos do grupo 8 , rutênio e ósmio . O ferro forma compostos em uma ampla gama de estados de oxidação , -2 a +7. O ferro também forma muitos compostos de coordenação ; alguns deles, como ferroceno , ferrioxalato e azul da Prússia , têm aplicações industriais, médicas ou de pesquisa substanciais.

Características

Alótropos

Volume molar vs. pressão para ferro α em temperatura ambiente

São conhecidos pelo menos quatro alótropos de ferro (diferentes arranjos de átomos no sólido), convencionalmente denotados como α, γ, δ e ε.

Diagrama de fase de baixa pressão de ferro puro

As três primeiras formas são observadas em pressões normais. À medida que o ferro fundido esfria além de seu ponto de congelamento de 1538 ° C, ele se cristaliza em seu alótropo δ, que tem uma estrutura cristalina cúbica centrada no corpo (bcc) . À medida que esfria ainda mais para 1394 ° C, ele muda para seu alótropo de ferro γ, uma estrutura cristalina cúbica de face centrada (fcc) ou austenita . A 912 ° C e abaixo, a estrutura cristalina torna-se novamente o alótropo de ferro α bcc. [6]

As propriedades físicas do ferro em pressões e temperaturas muito altas também foram estudadas extensivamente, [7] [8] por causa de sua relevância para as teorias sobre os núcleos da Terra e outros planetas. Acima de aproximadamente 10 GPa e temperaturas de algumas centenas de kelvin ou menos, o ferro α muda para outra estrutura hexagonal compactada (hcp), também conhecida como ferro ε . A fase γ de alta temperatura também se transforma em ferro ε, mas o faz com pressão mais alta.

Existem algumas evidências experimentais controversas de uma fase β estável a pressões acima de 50 GPa e temperaturas de pelo menos 1500 K. Supõe-se que ela tenha uma estrutura ortorrômbica ou dupla hcp. [9] (É confuso, o termo "β-ferro" às vezes também é usado para se referir ao α-ferro acima de seu ponto de Curie, quando ele muda de ferromagnético para paramagnético, embora sua estrutura cristalina não tenha mudado. [6] )

O núcleo interno da Terra é geralmente considerado como consistindo de uma liga de ferro- níquel com estrutura ε (ou β). [10]

Pontos de fusão e ebulição

Os pontos de fusão e ebulição do ferro, junto com sua entalpia de atomização , são mais baixos do que os dos elementos 3d anteriores de escândio a cromo , mostrando a contribuição diminuída dos elétrons 3d para ligações metálicas à medida que são atraídos cada vez mais para o inerte núcleo pelo núcleo; [11] no entanto, eles são maiores do que os valores do elemento anterior manganês porque esse elemento tem uma sub-camada 3d preenchida pela metade e, conseqüentemente, seus elétrons-d não são facilmente deslocalizados. Essa mesma tendência aparece para o rutênio, mas não para o ósmio. [12]

O ponto de fusão do ferro é experimentalmente bem definido para pressões menores que 50 GPa. Para maiores pressões, os dados publicados (a partir de 2007) ainda variam em dezenas de gigapascals e mais de mil Kelvin. [13]

Propriedades magneticas

Curvas de magnetização de 9 materiais ferromagnéticos, mostrando saturação. 1.  Folha de aço, 2.  Aço silício, 3.  Aço fundido, 4.  Aço tungstênio, 5.  Aço magnético, 6.  Ferro fundido, 7.  Níquel, 8.  Cobalto, 9.  Magnetita [14]

Abaixo de seu ponto Curie de 770 ° C, o ferro α muda de paramagnético para ferromagnético : os spins dos dois elétrons desemparelhados em cada átomo geralmente se alinham com os spins de seus vizinhos, criando um campo magnético geral . [15] Isso acontece porque os orbitais desses dois elétrons (d z 2 e d x 2 - y 2 ) não apontam para átomos vizinhos na rede e, portanto, não estão envolvidos em ligações metálicas. [6]

Na ausência de uma fonte externa de campo magnético, os átomos ficam espontaneamente particionados em domínios magnéticos , cerca de 10 micrômetros de diâmetro, [16] de modo que os átomos em cada domínio têm spins paralelos, mas alguns domínios têm outras orientações. Assim, um pedaço macroscópico de ferro terá um campo magnético total quase nulo.

A aplicação de um campo magnético externo faz com que os domínios magnetizados na mesma direção geral cresçam às custas dos domínios adjacentes que apontam em outras direções, reforçando o campo externo. Esse efeito é explorado em dispositivos que precisam canalizar campos magnéticos, como transformadores elétricos , cabeças de gravação magnética e motores elétricos . Impurezas, defeitos de rede ou limites de grãos e partículas podem "fixar" os domínios nas novas posições, de modo que o efeito persiste mesmo depois que o campo externo é removido - transformando assim o objeto de ferro em um ímã (permanente) . [15]

Comportamento semelhante é exibido por alguns compostos de ferro, como as ferritas, incluindo a magnetita mineral , uma forma cristalina da mistura de óxido de ferro (II, III) Fe
3
O
4
(embora o mecanismo de escala atômica, ferrimagnetismo , seja um pouco diferente). Pedaços de magnetita com magnetização permanente natural ( magnetitas ) forneceram as primeiras bússolas para navegação. Partículas de magnetita foram amplamente utilizadas em mídias de gravação magnética, como memórias centrais , fitas magnéticas , disquetes e discos , até serem substituídas por materiais à base de cobalto .

Isótopos

O ferro tem quatro isótopos estáveis : 54 Fe (5,845% do ferro natural), 56 Fe (91,754%), 57 Fe (2,119%) e 58 Fe (0,282%). 20-30 isótopos artificiais também foram criados. Destes isótopos estáveis, apenas 57 Fe tem spin nuclear ( -12 ). O nuclídeo 54 Fe teoricamente pode sofrer dupla captura de elétrons para 54 Cr, mas o processo nunca foi observado e apenas um limite inferior de meia-vida de 3,1 × 10 22 anos foi estabelecido. [17]

60 Fe é um radionuclídeo extinto de meia-vida longa (2,6 milhões de anos). [18] Não é encontrado na Terra, mas seu produto de decomposição final é sua neta, o nuclídeo estável 60 Ni . [17] Muitos dos trabalhos anteriores sobre a composição isotópica do ferro enfocaram a nucleossíntese de 60 Fe por meio de estudos de meteoritos e formação de minério. Na última década, os avanços na espectrometria de massa permitiram a detecção e quantificação de variações mínimas e naturais nas proporções dos isótopos estáveis de ferro. Muito deste trabalho é impulsionado peloComunidades de ciência terrestre e planetária , embora aplicações para sistemas biológicos e industriais estejam surgindo. [19]

Nas fases dos meteoritos Semarkona e Chervony Kut, uma correlação entre a concentração de 60 Ni, neta de 60 Fe, e a abundância de isótopos estáveis ​​de ferro forneceu evidências da existência de 60 Fe no momento da formação do Sistema Solar. . Possivelmente, a energia liberada pela decomposição do 60 Fe, junto com a liberada pelo 26 Al , contribuiu para a refusão e diferenciação dos asteróides após sua formação há 4,6 bilhões de anos. A abundância de 60 Ni presente em extraterrestreso material pode trazer mais informações sobre a origem e a história inicial do Sistema Solar . [20]

O isótopo de ferro mais abundante 56 Fe é de interesse particular para cientistas nucleares porque representa o ponto final mais comum da nucleossíntese . [21] Uma vez que 56 Ni (14 partículas alfa ) é facilmente produzido a partir de núcleos mais leves no processo alfa em reações nucleares em supernovas (ver processo de queima de silício ), é o ponto final das cadeias de fusão dentro de estrelas extremamente massivas , desde a adição de outra alfa partícula, resultando em 60 Zn, requer muito mais energia. Este 56O Ni, que tem meia-vida de cerca de 6 dias, é criado em quantidade nessas estrelas, mas logo decai por duas emissões de pósitrons sucessivas dentro dos produtos de decaimento da supernova na nuvem de gás remanescente da supernova , primeiro para 56 Co radioativo e depois para estável 56 Fe. Como tal, o ferro é o elemento mais abundante no núcleo dos gigantes vermelhos e é o metal mais abundante nos meteoritos de ferro e nos densos núcleos metálicos de planetas como a Terra . [22] Também é muito comum no universo, em relação a outros metais estáveis de aproximadamente o mesmo peso atômico . [22] [23]O ferro é o sexto elemento mais abundante no universo e o elemento refratário mais comum . [24]

Embora um pequeno ganho de energia adicional pudesse ser extraído sintetizando 62 Ni , que tem uma energia de ligação marginalmente maior do que 56 Fe, as condições nas estrelas são inadequadas para este processo. A produção e distribuição de elementos em supernovas na Terra favorecem enormemente o ferro em relação ao níquel e, em qualquer caso, o 56 Fe ainda tem uma massa menor por nucléon do que o 62 Ni devido à sua maior fração de prótons mais leves. [25] Conseqüentemente, os elementos mais pesados ​​que o ferro requerem uma supernova para sua formação, envolvendo a captura rápida de nêutrons iniciando núcleos de 56 Fe. [22]

No futuro distante do universo, supondo que o decaimento do próton não ocorra, a fusão a frio ocorrendo via tunelamento quântico faria com que os núcleos leves da matéria comum se fundissem em núcleos de 56 Fe. A fissão e a emissão de partículas alfa , então, causariam a decomposição de núcleos pesados ​​em ferro, convertendo todos os objetos de massa estelar em esferas frias de ferro puro. [26]

Origem e ocorrência na natureza

Cosmogênese

A abundância de ferro em planetas rochosos como a Terra se deve à sua produção abundante durante a fusão e explosão descontroladas das supernovas do tipo Ia , que espalham o ferro no espaço. [27] [28]

Ferro metálico

Uma peça polida e quimicamente gravada de um meteorito de ferro, que se acredita ser semelhante em composição ao núcleo metálico da Terra, mostrando cristais individuais da liga de ferro-níquel ( padrão Widmanstatten )

O ferro metálico ou nativo raramente é encontrado na superfície da Terra porque tende a se oxidar. No entanto, acredita-se que tanto o núcleo interno quanto o externo da Terra , responsáveis ​​por 35% da massa de toda a Terra, consistam em grande parte de uma liga de ferro, possivelmente com níquel . Acredita-se que as correntes elétricas no núcleo externo do líquido sejam a origem do campo magnético da Terra . Acredita-se que os outros planetas terrestres ( Mercúrio , Vênus e Marte ), bem como a Lua, tenham um núcleo metálico consistindo principalmente de ferro. Os asteróides tipo M também são considerados parcialmente ou em sua maioria feitos de liga de ferro metálica.

Os raros meteoritos de ferro são a principal forma de ferro metálico natural na superfície da Terra. Itens feitos de ferro meteorítico trabalhado a frio foram encontrados em vários sítios arqueológicos que datam de uma época em que a fundição de ferro ainda não havia sido desenvolvida; e os Inuit na Groenlândia usam ferro do meteorito de Cape York para ferramentas e armas de caça. [29] Cerca de 1 em 20 meteoritos consistem nos únicos minerais de ferro-níquel taenita ( 35–80 % de ferro) e kamacita (90–95% de ferro). [30]O ferro nativo também é raramente encontrado em basaltos que se formaram a partir de magmas que entraram em contato com rochas sedimentares ricas em carbono, o que reduziu a fugacidade de oxigênio o suficiente para que o ferro se cristalizasse. Isso é conhecido como ferro telúrico e é descrito em algumas localidades, como a Ilha Disko no oeste da Groenlândia, Yakutia na Rússia e Bühl na Alemanha. [31]

Minerais de manto

A ferropericlase (Mg, Fe) O , uma solução sólida de periclase (MgO) e wüstita (FeO), representa cerca de 20% do volume do manto inferior da Terra, o que a torna a segunda fase mineral mais abundante naquela região. após perovskita de silicato (Mg, Fe) SiO
3
; também é o principal hospedeiro do ferro no manto inferior. [32] Na parte inferior da zona de transição do manto, a reação γ- (Mg, Fe)
2
[SiO
4
] ↔ (Mg, Fe) [SiO
3
] + (Mg, Fe) O
transforma a γ-olivina em uma mistura de perovskita de silicato e ferropericlase e vice-versa. Na literatura, essa fase mineral do manto inferior também é freqüentemente chamada de magnesiowüstita. [33] A perovskita de silicato pode formar até 93% do manto inferior, [34] e a forma de ferro magnésio, (Mg, Fe) SiO
3
, é considerado o mineral mais abundante da Terra, representando 38% do seu volume. [35]

crosta da terrra

Caminho ocre em Roussillon .

Embora o ferro seja o elemento mais abundante na Terra, a maior parte desse ferro está concentrada nos núcleos interno e externo . [36] [37] A fração de ferro que está na crosta terrestre atinge apenas cerca de 5% da massa total da crosta e é, portanto, apenas o quarto elemento mais abundante nessa camada (depois do oxigênio , silício e alumínio ). [38]

A maior parte do ferro na crosta é combinada com vários outros elementos para formar muitos minerais de ferro . Uma classe importante são os minerais de óxido de ferro , como hematita (Fe 2 O 3 ), magnetita (Fe 3 O 4 ) e siderita (FeCO 3 ), que são os principais minérios de ferro . Muitas rochas ígneas também contêm minerais de sulfeto, pirrotita e pentlandita . [39] [40] Durante a meteorização, o ferro tende a lixiviar dos depósitos de sulfeto como o sulfato e dos depósitos de silicato como o bicarbonato. Ambos são oxidados em solução aquosa e precipitam mesmo em pH ligeiramente elevado como óxido de ferro (III) . [41]

Formação de ferro em faixas em McKinley Park, Minnesota.

Grandes depósitos de ferro são formações de ferro em faixas , um tipo de rocha que consiste em camadas finas repetidas de óxidos de ferro alternando-se com faixas de xisto pobre em ferro e sílex . As formações de ferro bandadas foram estabelecidas entre 3.700  milhões de anos atrás e 1.800  milhões de anos atrás . [42] [43]

Materiais contendo óxidos de ferro (III) finamente moídos ou óxidos hidróxidos, como ocre , têm sido usados ​​como pigmentos amarelos, vermelhos e marrons desde os tempos pré-históricos. Eles também contribuem para a cor de várias rochas e argilas , incluindo formações geológicas inteiras como Painted Hills no Oregon e Buntsandstein ("arenito colorido", Bunter Britânico ). [44] Através de Eisensandstein (um 'arenito de ferro' jurássico , por exemplo, de Donzdorf na Alemanha) [45] e Pedra de Bathno Reino Unido, os compostos de ferro são responsáveis ​​pela cor amarelada de muitos edifícios e esculturas históricas. [46] A proverbial cor vermelha da superfície de Marte é derivada de um regolito rico em óxido de ferro . [47]

Quantidades significativas de ferro ocorrem no mineral pirita de sulfeto de ferro (FeS 2 ), mas é difícil extrair ferro dele e, portanto, não é explorado. Na verdade, o ferro é tão comum que a produção geralmente se concentra apenas em minérios com grandes quantidades dele.

De acordo com o Painel de Recursos Internacional de Stocks Metal no relatório da Sociedade , o estoque mundial de ferro em uso na sociedade é 2,200 kg per capita. Os países mais desenvolvidos diferem neste aspecto dos países menos desenvolvidos (7.000–14.000 vs 2.000 kg per capita). [48]

Oceanos

A ciência oceânica demonstrou o papel do ferro nos mares ancestrais tanto na biota marinha quanto no clima. [49]

Química e compostos


Estado de oxidação
Composto representativo
-2 (d 10 ) Tetracarbonilferrato dissódico (reagente de Collman)
-1 (d 9 ) Fe
2
(CO)2−
8
0 (d 8 ) Pentacarbonil de ferro
1 (d 7 ) Dímero de dicarbonil ciclopentadienil ferro ("Fp 2 ")
2 (d 6 ) Sulfato ferroso , ferroceno
3 (d 5 ) Cloreto férrico , tetrafluoroborato de ferrocênio
4 (d 4 ) Fe (diários)
2
Cl2+
2
, Tetrafluoroborato de ferrila
5 (d 3 ) FeO3−
4
6 (d 2 ) Ferrato de potássio
7 (d 1 ) [FeO 4 ] - (isolamento de matriz, 4K)

O ferro mostra as propriedades químicas características dos metais de transição , ou seja, a capacidade de formar estados de oxidação variáveis ​​que diferem por etapas de um e uma coordenação e química organometálica muito grandes: na verdade, foi a descoberta de um composto de ferro, o ferroceno , que revolucionou o último campo na década de 1950. [50] O ferro é às vezes considerado um protótipo para todo o bloco de metais de transição, devido à sua abundância e ao imenso papel que desempenhou no progresso tecnológico da humanidade. [51] Seus 26 elétrons estão dispostos na configuração [Ar] 3d 6 4s 2, dos quais os elétrons 3d e 4s são relativamente próximos em energia e, portanto, pode perder um número variável de elétrons e não há um ponto claro onde a ionização posterior se torne inútil. [12]

O ferro forma compostos principalmente nos estados de oxidação +2 ( ferro (II) , "ferroso") e +3 ( ferro (III) , "férrico"). O ferro também ocorre em estados de oxidação mais elevados , por exemplo, o ferrato de potássio púrpura (K 2 FeO 4 ), que contém ferro em seu estado de oxidação +6. Embora o óxido de ferro (VIII) (FeO 4 ) tenha sido reivindicado, o relatório não pôde ser reproduzido e tal espécie a partir da remoção de todos os elétrons do elemento além da configuração anterior de gás inerte (pelo menos com o ferro em seu estado de oxidação +8 ) foi considerado improvável em termos computacionais. [52] No entanto, uma forma de aniônico [FeO 4 ]- com ferro em seu estado de oxidação +7, juntamente com um isômero de ferro (V) -peroxo, foi detectado por espectroscopia de infravermelho a 4 K após cocondensação de átomos de Fe ablacionados a laser com uma mistura de O 2 / Ar. [53] O ferro (IV) é um intermediário comum em muitas reações de oxidação bioquímica. [54] [55] Numerosos compostos de organo-ferro contêm estados de oxidação formais de +1, 0, -1 ou mesmo -2. Os estados de oxidação e outras propriedades de ligação são frequentemente avaliados usando a técnica de espectroscopia Mössbauer . [56] Muitos compostos de valência mistos contêm centros de ferro (II) e ferro (III), como magnetita eAzul da Prússia ( Fe
4
(Fe [CN]
6
)
3
) [55] Este último é usado como o tradicional "azul" nos projetos . [57]

O ferro é o primeiro dos metais de transição que não pode atingir seu estado de oxidação de grupo de +8, embora seus congêneres mais pesados ​​rutênio e ósmio possam, com o rutênio tendo mais dificuldade do que o ósmio. [6] O rutênio exibe uma química catiônica aquosa em seus estados de baixa oxidação semelhantes aos do ferro, mas o ósmio não, favorecendo os estados de alta oxidação nos quais forma complexos aniônicos. [6] Na segunda metade da série de transição 3d, as semelhanças verticais nos grupos competem com as semelhanças horizontais do ferro com seus vizinhos cobalto e níquel na tabela periódica, que também são ferromagnéticas à temperatura ambientee compartilham uma química semelhante. Como tal, ferro, cobalto e níquel às vezes são agrupados como a tríade de ferro . [51]

Ao contrário de muitos outros metais, o ferro não forma amálgama com o mercúrio . Como resultado, o mercúrio é comercializado em frascos padronizados de 76 libras (34 kg) feitos de ferro. [58]

O ferro é de longe o elemento mais reativo em seu grupo; é pirofórico quando finamente dividido e se dissolve facilmente em ácidos diluídos, dando Fe 2+ . No entanto, não reage com o ácido nítrico concentrado e outros ácidos oxidantes devido à formação de uma camada de óxido impermeável, que pode, no entanto, reagir com o ácido clorídrico . [6] O ferro de alta pureza, denominado ferro eletrolítico , é considerado resistente à ferrugem, devido à sua camada de óxido.

Compostos binários

Óxidos e hidróxidos

Óxido ferroso ou de ferro (II), FeO
Férrico ou óxido de ferro (III) Fe
2
O
3
Ferrosoférrico ou óxido de ferro (II, III) Fe
3
O
4

O ferro forma vários compostos de óxido e hidróxido ; os mais comuns são óxido de ferro (II, III) (Fe 3 O 4 ) e óxido de ferro (III) (Fe 2 O 3 ). O óxido de ferro (II) também existe, embora seja instável à temperatura ambiente. Apesar de seus nomes, eles são, na verdade, todos compostos não estequiométricos cujas composições podem variar. [59] Esses óxidos são os principais minérios para a produção de ferro (ver floração e alto-forno). Eles também são usados ​​na produção de ferritas , armazenamento magnético útilmídia em computadores e pigmentos. O sulfeto mais conhecido é a pirita de ferro (FeS 2 ), também conhecida como ouro de tolo devido ao seu brilho dourado. [55] Não é um composto de ferro (IV), mas na verdade é um polissulfeto de ferro (II) contendo Fe 2+ e S2−
2
íons em uma estrutura distorcida de cloreto de sódio . [59]

Diagrama de Pourbaix de ferro

Halides

Some canary-yellow powder sits, mostly in lumps, on a laboratory watch glass.
Cloreto de ferro hidratado (III) (cloreto férrico)

Os haletos ferrosos e férricos binários são bem conhecidos. Os halogenetos ferrosos surgem tipicamente do tratamento do metal de ferro com o ácido hidrohálico correspondente para dar os sais hidratados correspondentes. [55]

Fe + 2 HX → FeX 2 + H 2 (X = F, Cl, Br, I)

O ferro reage com o flúor, o cloro e o bromo para dar os haletos férricos correspondentes, sendo o cloreto férrico o mais comum. [60]

2 Fe + 3 X 2 → 2 FeX 3 (X = F, Cl, Br)

O iodeto férrico é uma exceção, sendo termodinamicamente instável devido ao poder oxidante do Fe 3+ e ao alto poder redutor do I - : [60]

2 I - + 2 Fe 3+ → I 2 + 2 Fe 2+ (E 0 = +0,23 V)

Ferric iodide, a black solid, is not stable in ordinary conditions, but can be prepared through the reaction of iron pentacarbonyl with iodine and carbon monoxide in the presence of hexane and light at the temperature of −20 °C, with oxygen and water excluded.[60] Complexes of ferric iodide with some soft bases are known to be stable compounds.[61][62]

Solution chemistry

Comparison of colors of solutions of ferrate (left) and permanganate (right)

The standard reduction potentials in acidic aqueous solution for some common iron ions are given below:[6]

Fe2+ + 2 e ⇌ Fe E0 = −0.447 V
Fe3+ + 3 e ⇌ Fe E0 = −0.037 V
FeO2−
4
+ 8 H+ + 3 e
⇌ Fe3+ + 4 H2O E0 = +2.20 V

The red-purple tetrahedral ferrate(VI) anion is such a strong oxidizing agent that it oxidizes nitrogen and ammonia at room temperature, and even water itself in acidic or neutral solutions:[60]

4 FeO2−
4
+ 10 H
2
O
→ 4 Fe3+
+ 20 OH
+ 3 O2

The Fe3+ ion has a large simple cationic chemistry, although the pale-violet hexaquo ion [Fe(H
2
O)
6
]3+
is very readily hydrolyzed when pH increases above 0 as follows:[63]

[Fe(H
2
O)
6
]3+
[Fe(H
2
O)
5
(OH)]2+ + H+
K = 10−3.05 mol dm−3
[Fe(H
2
O)
5
(OH)]2+
[Fe(H
2
O)
4
(OH)
2
]+ + H+
K = 10−3.26 mol dm−3
2[Fe(H
2
O)
6
]3+
[Fe(H
2
O)
4
(OH)]4+
2
+ 2H+ + 2H
2
O
K = 10−2.91 mol dm−3
Blue-green iron(II) sulfate heptahydrate

As pH rises above 0 the above yellow hydrolyzed species form and as it rises above 2–3, reddish-brown hydrous iron(III) oxide precipitates out of solution. Although Fe3+ has a d5 configuration, its absorption spectrum is not like that of Mn2+ with its weak, spin-forbidden d–d bands, because Fe3+ has higher positive charge and is more polarizing, lowering the energy of its ligand-to-metal charge transfer absorptions. Thus, all the above complexes are rather strongly colored, with the single exception of the hexaquo ion – and even that has a spectrum dominated by charge transfer in the near ultraviolet region.[63] On the other hand, the pale green iron(II) hexaquo ion [Fe(H
2
O)
6
]2+
does not undergo appreciable hydrolysis. Carbon dioxide is not evolved when carbonate anions are added, which instead results in white iron(II) carbonate being precipitated out. In excess carbon dioxide this forms the slightly soluble bicarbonate, which occurs commonly in groundwater, but it oxidises quickly in air to form iron(III) oxide that accounts for the brown deposits present in a sizeable number of streams.[64]

Coordination compounds

Due to its electronic structure, iron has a very large coordination and organometallic chemistry.

The two enantiomorphs of the ferrioxalate ion

Many coordination compounds of iron are known. A typical six-coordinate anion is hexachloroferrate(III), [FeCl6]3−, found in the mixed salt tetrakis(methylammonium) hexachloroferrate(III) chloride.[65][66] Complexes with multiple bidentate ligands have geometric isomers. For example, the trans-chlorohydridobis(bis-1,2-(diphenylphosphino)ethane)iron(II) complex is used as a starting material for compounds with the Fe(dppe)
2
moiety.[67][68] The ferrioxalate ion with three oxalate ligands (shown at right) displays helical chirality with its two non-superposable geometries labelled Λ (lambda) for the left-handed screw axis and Δ (delta) for the right-handed screw axis, in line with IUPAC conventions.[63] Potassium ferrioxalate is used in chemical actinometry and along with its sodium salt undergoes photoreduction applied in old-style photographic processes. The dihydrate of iron(II) oxalate has a polymeric structure with co-planar oxalate ions bridging between iron centres with the water of crystallisation located forming the caps of each octahedron, as illustrated below.[69]

Crystal structure of iron(II) oxalate dihydrate, showing iron (gray), oxygen (red), carbon (black), and hydrogen (white) atoms.
Blood-red positive thiocyanate test for iron(III)

Iron(III) complexes are quite similar to those of chromium(III) with the exception of iron(III)'s preference for O-donor instead of N-donor ligands. The latter tend to be rather more unstable than iron(II) complexes and often dissociate in water. Many Fe–O complexes show intense colors and are used as tests for phenols or enols. For example, in the ferric chloride test, used to determine the presence of phenols, iron(III) chloride reacts with a phenol to form a deep violet complex:[63]

3 ArOH + FeCl3 → Fe(OAr)3 + 3 HCl (Ar = aryl)

Among the halide and pseudohalide complexes, fluoro complexes of iron(III) are the most stable, with the colorless [FeF5(H2O)]2− being the most stable in aqueous solution. Chloro complexes are less stable and favor tetrahedral coordination as in [FeCl4]; [FeBr4] and [FeI4] are reduced easily to iron(II). Thiocyanate is a common test for the presence of iron(III) as it forms the blood-red [Fe(SCN)(H2O)5]2+. Like manganese(II), most iron(III) complexes are high-spin, the exceptions being those with ligands that are high in the spectrochemical series such as cyanide. An example of a low-spin iron(III) complex is [Fe(CN)6]3−. The cyanide ligands may easily be detached in [Fe(CN)6]3−, and hence this complex is poisonous, unlike the iron(II) complex [Fe(CN)6]4− found in Prussian blue,[63] which does not release hydrogen cyanide except when dilute acids are added.[64] Iron shows a great variety of electronic spin states, including every possible spin quantum number value for a d-block element from 0 (diamagnetic) to 52 (5 unpaired electrons). This value is always half the number of unpaired electrons. Complexes with zero to two unpaired electrons are considered low-spin and those with four or five are considered high-spin.[59]

Iron(II) complexes are less stable than iron(III) complexes but the preference for O-donor ligands is less marked, so that for example [Fe(NH
3
)
6
]2+
is known while [Fe(NH
3
)
6
]3+
is not. They have a tendency to be oxidized to iron(III) but this can be moderated by low pH and the specific ligands used.[64]

Organometallic compounds

Iron penta-
carbonyl

Organoiron chemistry is the study of organometallic compounds of iron, where carbon atoms are covalently bound to the metal atom. They are many and varied, including cyanide complexes, carbonyl complexes, sandwich and half-sandwich compounds.

Prussian blue

Prussian blue or "ferric ferrocyanide", Fe4[Fe(CN)6]3, is an old and well-known iron-cyanide complex, extensively used as pigment and in several other applications. Its formation can be used as a simple wet chemistry test to distinguish between aqueous solutions of Fe2+ and Fe3+ as they react (respectively) with potassium ferricyanide and potassium ferrocyanide to form Prussian blue.[55]

Another old example of an organoiron compound is iron pentacarbonyl, Fe(CO)5, in which a neutral iron atom is bound to the carbon atoms of five carbon monoxide molecules. The compound can be used to make carbonyl iron powder, a highly reactive form of metallic iron. Thermolysis of iron pentacarbonyl gives triiron dodecacarbonyl, Fe
3
(CO)
12
, a complex with a cluster of three iron atoms at its core. Collman's reagent, disodium tetracarbonylferrate, is a useful reagent for organic chemistry; it contains iron in the −2 oxidation state. Cyclopentadienyliron dicarbonyl dimer contains iron in the rare +1 oxidation state.[70]

Structural formula of ferrocene and a powdered sample

A landmark in this field was the discovery in 1951 of the remarkably stable sandwich compound ferrocene Fe(C
5
H
5
)
2
, by Paulson and Kealy[71] and independently by Miller and others,[72] whose surprising molecular structure was determined only a year later by Woodward and Wilkinson[73] and Fischer.[74] Ferrocene is still one of the most important tools and models in this class.[75]

Iron-centered organometallic species are used as catalysts. The Knölker complex, for example, is a transfer hydrogenation catalyst for ketones.[76]

Industrial uses

The iron compounds produced on the largest scale in industry are iron(II) sulfate (FeSO4·7H2O) and iron(III) chloride (FeCl3). The former is one of the most readily available sources of iron(II), but is less stable to aerial oxidation than Mohr's salt ((NH
4
)
2
Fe(SO
4
)
2
·6H2O
). Iron(II) compounds tend to be oxidized to iron(III) compounds in the air.[55]

Etymology

"iren," an Old English word for 'iron'

As iron has been in use for such a long time, it has many names. The source of its chemical symbol Fe is the Latin word ferrum, and its descendants are the names of the element in the Romance languages (for example, French fer, Spanish hierro, and Italian and Portuguese ferro).[77] The word ferrum itself possibly comes from the Semitic languages, via Etruscan, from a root that also gave rise to Old English bræs "brass".[78] The English word iron derives ultimately from Proto-Germanic *isarnan, which is also the source of the German name Eisen and Dutch ijzeren. It was most likely borrowed from Celtic *isarnon, which ultimately comes from Proto-Indo-European *is-(e)ro- "powerful, holy" and finally *eis "strong", referencing iron's strength as a metal.[79] Kluge relates *isarnon to Illyric and Latin ira, 'wrath').[citation needed] The Balto-Slavic names for iron (e.g. Russian железо [zhelezo], Polish żelazo, Lithuanian geležis) are the only ones to come directly from the Proto-Indo-European *ghelgh- "iron".[80] In many of these languages, the word for iron may also be used to denote other objects made of iron or steel, or figuratively because of the hardness and strength of the metal.[81] The Chinese tiě (traditional 鐵; simplified 铁) derives from Proto-Sino-Tibetan *hliek,[82] and was borrowed into Japanese as 鉄 tetsu, which also has the native reading kurogane "black metal" (similar to how iron is referenced in the English word blacksmith).[83]

History

Development of iron metallurgy

Iron is one of the elements undoubtedly known to the ancient world.[84] It has been worked, or wrought, for millennia. However, iron objects of great age are much rarer than objects made of gold or silver due to the ease with which iron corrodes.[85] The technology developed slowly, and even after the discovery of smelting it took many centuries for iron to replace bronze as the metal of choice for tools and weapons.

Meteoritic iron

Iron harpoon head from Greenland. The iron edge covers a narwhal tusk harpoon using meteorite iron from the Cape York meteorite, one of the largest iron meteorites known.

Beads made from meteoric iron in 3500 BC or earlier were found in Gerzah, Egypt by G.A. Wainwright.[86] The beads contain 7.5% nickel, which is a signature of meteoric origin since iron found in the Earth's crust generally has only minuscule nickel impurities.

Meteoric iron was highly regarded due to its origin in the heavens and was often used to forge weapons and tools.[86] For example, a dagger made of meteoric iron was found in the tomb of Tutankhamun, containing similar proportions of iron, cobalt, and nickel to a meteorite discovered in the area, deposited by an ancient meteor shower.[87][88][89] Items that were likely made of iron by Egyptians date from 3000 to 2500 BC.[85]

Meteoritic iron is comparably soft and ductile and easily cold forged but may get brittle when heated because of the nickel content.[90]

Wrought iron

A circle, with a short, simple arrow shape extending diagonally upwards and rightwards from its edge
The symbol for Mars has been used since antiquity to represent iron.
A pillar, slightly fluted, with some ornamentation at its top. It is black, slightly weathered to a dark brown near the base. It is around 7 meters (23 feet) tall. It stands upon a raised circular base of stone, and is surrounded by a short, square fence.
The iron pillar of Delhi is an example of the iron extraction and processing methodologies of early India.

The first iron production started in the Middle Bronze Age, but it took several centuries before iron displaced bronze. Samples of smelted iron from Asmar, Mesopotamia and Tall Chagar Bazaar in northern Syria were made sometime between 3000 and 2700 BC.[91] The Hittites established an empire in north-central Anatolia around 1600 BC. They appear to be the first to understand the production of iron from its ores and regard it highly in their society.[92] The Hittites began to smelt iron between 1500 and 1200 BC and the practice spread to the rest of the Near East after their empire fell in 1180 BC.[91] The subsequent period is called the Iron Age.

Artifacts of smelted iron are found in India dating from 1800 to 1200 BC,[93] and in the Levant from about 1500 BC (suggesting smelting in Anatolia or the Caucasus).[94][95] Alleged references (compare history of metallurgy in South Asia) to iron in the Indian Vedas have been used for claims of a very early usage of iron in India respectively to date the texts as such. The rigveda term ayas (metal) refers to copper, while iron which is called as śyāma ayas, literally "black copper", first is mentioned in the post-rigvedic Atharvaveda.[96]

Some archaeological evidence suggests iron was smelted in Zimbabwe and southeast Africa as early as the eighth century BC.[97] Iron working was introduced to Greece in the late 11th century BC, from which it spread quickly throughout Europe.[98]

Iron sickle from Ancient Greece.

The spread of ironworking in Central and Western Europe is associated with Celtic expansion. According to Pliny the Elder, iron use was common in the Roman era.[86] In the lands of what is now considered China, iron appears approximately 700–500 BC.[99] Iron smelting may have been introduced into China through Central Asia.[100] The earliest evidence of the use of a blast furnace in China dates to the 1st century AD,[101] and cupola furnaces were used as early as the Warring States period (403–221 BC).[102] Usage of the blast and cupola furnace remained widespread during the Song and Tang Dynasties.[103]

During the Industrial Revolution in Britain, Henry Cort began refining iron from pig iron to wrought iron (or bar iron) using innovative production systems. In 1783 he patented the puddling process for refining iron ore. It was later improved by others, including Joseph Hall.[104]

Cast iron

Cast iron was first produced in China during 5th century BC,[105] but was hardly in Europe until the medieval period.[106][107] The earliest cast iron artifacts were discovered by archaeologists in what is now modern Luhe County, Jiangsu in China. Cast iron was used in ancient China for warfare, agriculture, and architecture.[108] During the medieval period, means were found in Europe of producing wrought iron from cast iron (in this context known as pig iron) using finery forges. For all these processes, charcoal was required as fuel.[109]

Coalbrookdale by Night, 1801. Blast furnaces light the iron making town of Coalbrookdale.

Medieval blast furnaces were about 10 feet (3.0 m) tall and made of fireproof brick; forced air was usually provided by hand-operated bellows.[107] Modern blast furnaces have grown much bigger, with hearths fourteen meters in diameter that allow them to produce thousands of tons of iron each day, but essentially operate in much the same way as they did during medieval times.[109]

In 1709, Abraham Darby I established a coke-fired blast furnace to produce cast iron, replacing charcoal, although continuing to use blast furnaces. The ensuing availability of inexpensive iron was one of the factors leading to the Industrial Revolution. Toward the end of the 18th century, cast iron began to replace wrought iron for certain purposes, because it was cheaper. Carbon content in iron was not implicated as the reason for the differences in properties of wrought iron, cast iron, and steel until the 18th century.[91]

Since iron was becoming cheaper and more plentiful, it also became a major structural material following the building of the innovative first iron bridge in 1778. This bridge still stands today as a monument to the role iron played in the Industrial Revolution. Following this, iron was used in rails, boats, ships, aqueducts, and buildings, as well as in iron cylinders in steam engines.[109] Railways have been central to the formation of modernity and ideas of progress[110] and various languages (e.g. French, Spanish, Italian and German) refer to railways as iron road.

Steel

Steel (with smaller carbon content than pig iron but more than wrought iron) was first produced in antiquity by using a bloomery. Blacksmiths in Luristan in western Persia were making good steel by 1000 BC.[91] Then improved versions, Wootz steel by India and Damascus steel were developed around 300 BC and AD 500 respectively. These methods were specialized, and so steel did not become a major commodity until the 1850s.[111]

New methods of producing it by carburizing bars of iron in the cementation process were devised in the 17th century. In the Industrial Revolution, new methods of producing bar iron without charcoal were devised and these were later applied to produce steel. In the late 1850s, Henry Bessemer invented a new steelmaking process, involving blowing air through molten pig iron, to produce mild steel. This made steel much more economical, thereby leading to wrought iron no longer being produced in large quantities.[112]

Foundations of modern chemistry

In 1774, Antoine Lavoisier used the reaction of water steam with metallic iron inside an incandescent iron tube to produce hydrogen in his experiments leading to the demonstration of the conservation of mass, which was instrumental in changing chemistry from a qualitative science to a quantitative one.[113]

Symbolic role

"Gold gab ich für Eisen" – "I gave gold for iron". German-American brooch from WWI.

Iron plays a certain role in mythology and has found various usage as a metaphor and in folklore. The Greek poet Hesiod's Works and Days (lines 109–201) lists different ages of man named after metals like gold, silver, bronze and iron to account for successive ages of humanity.[114] The Iron Age was closely related with Rome, and in Ovid's Metamorphoses

The Virtues, in despair, quit the earth; and the depravity of man becomes universal and complete. Hard steel succeeded then.

— Ovid, Metamorphoses, Book I, Iron age, line 160 ff

An example of the importance of iron's symbolic role may be found in the German Campaign of 1813. Frederick William III commissioned then the first Iron Cross as military decoration. Berlin iron jewellery reached its peak production between 1813 and 1815, when the Prussian royal family urged citizens to donate gold and silver jewellery for military funding. The inscription Gold gab ich für Eisen (I gave gold for iron) was used as well in later war efforts.[115]

Production of metallic iron

Iron furnace in Columbus, Ohio, 1922

Laboratory routes

For a few limited purposes when it is needed, pure iron is produced in the laboratory in small quantities by reducing the pure oxide or hydroxide with hydrogen, or forming iron pentacarbonyl and heating it to 250 °C so that it decomposes to form pure iron powder.[41] Another method is electrolysis of ferrous chloride onto an iron cathode.[116]

Main industrial route

Iron production 2009 (million tonnes)[117]
Country Iron ore Pig iron Direct iron Steel
 China 1,114.9 549.4 573.6
 Australia 393.9 4.4 5.2
 Brazil 305.0 25.1 0.011 26.5
 Japan 66.9 87.5
 India 257.4 38.2 23.4 63.5
 Russia 92.1 43.9 4.7 60.0
 Ukraine 65.8 25.7 29.9
 South Korea 0.1 27.3 48.6
 Germany 0.4 20.1 0.38 32.7
World 1,594.9 914.0 64.5 1,232.4

Nowadays, the industrial production of iron or steel consists of two main stages. In the first stage, iron ore is reduced with coke in a blast furnace, and the molten metal is separated from gross impurities such as silicate minerals. This stage yields an alloy—pig iron—that contains relatively large amounts of carbon. In the second stage, the amount of carbon in the pig iron is lowered by oxidation to yield wrought iron, steel, or cast iron.[118] Other metals can be added at this stage to form alloy steels.

17th century Chinese illustration of workers at a blast furnace, making wrought iron from pig iron[119]
How iron was extracted in the 19th century

Blast furnace processing

The blast furnace is loaded with iron ores, usually hematite Fe
2
O
3
or magnetite Fe
3
O
4
, together with coke (coal that has been separately baked to remove volatile components). Air pre-heated to 900 °C is blown through the mixture, in sufficient amount to turn the carbon into carbon monoxide:[118]

2 C + O2 → 2 CO

This reaction raises the temperature to about 2000 °C The carbon monoxide reduces the iron ore to metallic iron[118]

Fe2O3 + 3 CO → 2 Fe + 3 CO2

Some iron in the high-temperature lower region of the furnace reacts directly with the coke:[118]

2 Fe2O3 + 3 C → 4 Fe + 3 CO2

A flux such as limestone (calcium carbonate) or dolomite (calcium-magnesium carbonate) is also added to the furnace's load. Its purpose is to remove silicaceous minerals in the ore, which would otherwise clog the furnace. The heat of the furnace decomposes the carbonates to calcium oxide, which reacts with any excess silica to form a slag composed of calcium silicate CaSiO
3
or other products. At the furnace's temperature, the metal and the slag are both molten. They collect at the bottom as two immiscible liquid layers (with the slag on top), that are then easily separated.[118] The slag can be used as a material in road construction or to improve mineral-poor soils for agriculture.[107]

This heap of iron ore pellets will be used in steel production.

Steelmaking

A pot of molten iron being used to make steel

In general, the pig iron produced by the blast furnace process contains up to 4–5% carbon, with small amounts of other impurities like sulfur, magnesium, phosphorus, and manganese. The high level of carbon makes it relatively weak and brittle. Reducing the amount of carbon to 0.002–2.1% by mass-produces steel, which may be up to 1000 times harder than pure iron. A great variety of steel articles can then be made by cold working, hot rolling, forging, machining, etc. Removing the other impurities, instead, results in cast iron, which is used to cast articles in foundries; for example stoves, pipes, radiators, lamp-posts, and rails.[118]

Steel products often undergo various heat treatments after they are forged to shape. Annealing consists of heating them to 700–800 °C for several hours and then gradual cooling. It makes the steel softer and more workable.[120]

Direct iron reduction

Owing to environmental concerns, alternative methods of processing iron have been developed. "Direct iron reduction" reduces iron ore to a ferrous lump called "sponge" iron or "direct" iron that is suitable for steelmaking.[107] Two main reactions comprise the direct reduction process:

Natural gas is partially oxidized (with heat and a catalyst):[107]

2 CH4 + O2 → 2 CO + 4 H2

Iron ore is then treated with these gases in a furnace, producing solid sponge iron:[107]

Fe2O3 + CO + 2 H2 → 2 Fe + CO2 + 2 H2O

Silica is removed by adding a limestone flux as described above.[107]

Thermite process

Ignition of a mixture of aluminium powder and iron oxide yields metallic iron via the thermite reaction:

Fe2O3 + 2 Al → 2 Fe + Al2O3

Alternatively pig iron may be made into steel (with up to about 2% carbon) or wrought iron (commercially pure iron). Various processes have been used for this, including finery forges, puddling furnaces, Bessemer converters, open hearth furnaces, basic oxygen furnaces, and electric arc furnaces. In all cases, the objective is to oxidize some or all of the carbon, together with other impurities. On the other hand, other metals may be added to make alloy steels.[109]

Applications

As structural material

Iron is the most widely used of all the metals, accounting for over 90% of worldwide metal production. Its low cost and high strength often make it the material of choice material to withstand stress or transmit forces, such as the construction of machinery and machine tools, rails, automobiles, ship hulls, concrete reinforcing bars, and the load-carrying framework of buildings. Since pure iron is quite soft, it is most commonly combined with alloying elements to make steel.[121]

Mechanical properties

Characteristic values of tensile strength (TS) and Brinell hardness (BH) of various forms of iron.[122][123]
Material TS
(MPa)
BH
(Brinell)
Iron whiskers 11000
Ausformed (hardened)
steel
2930 850–1200
Martensitic steel 2070 600
Bainitic steel 1380 400
Pearlitic steel 1200 350
Cold-worked iron 690 200
Small-grain iron 340 100
Carbon-containing iron 140 40
Pure, single-crystal iron 10 3

The mechanical properties of iron and its alloys are extremely relevant to their structural applications. Those properties can be evaluated in various ways, including the Brinell test, the Rockwell test and the Vickers hardness test.

The properties of pure iron are often used to calibrate measurements or to compare tests.[123][124] However, the mechanical properties of iron are significantly affected by the sample's purity: pure, single crystals of iron are actually softer than aluminium,[122] and the purest industrially produced iron (99.99%) has a hardness of 20–30 Brinell.[125] The pure iron (99.9 %~99.999 %), especially called electrolytic iron, is industrially produced by electrolytic refining

An increase in the carbon content will cause a significant increase in the hardness and tensile strength of iron. Maximum hardness of 65 Rc is achieved with a 0.6% carbon content, although the alloy has low tensile strength.[126] Because of the softness of iron, it is much easier to work with than its heavier congeners ruthenium and osmium.[12]

Iron-carbon phase diagram

Types of steels and alloys

α-Iron is a fairly soft metal that can dissolve only a small concentration of carbon (no more than 0.021% by mass at 910 °C).[127] Austenite (γ-iron) is similarly soft and metallic but can dissolve considerably more carbon (as much as 2.04% by mass at 1146 °C). This form of iron is used in the type of stainless steel used for making cutlery, and hospital and food-service equipment.[16]

Commercially available iron is classified based on purity and the abundance of additives. Pig iron has 3.5–4.5% carbon[128] and contains varying amounts of contaminants such as sulfur, silicon and phosphorus. Pig iron is not a saleable product, but rather an intermediate step in the production of cast iron and steel. The reduction of contaminants in pig iron that negatively affect material properties, such as sulfur and phosphorus, yields cast iron containing 2–4% carbon, 1–6% silicon, and small amounts of manganese.[118] Pig iron has a melting point in the range of 1420–1470 K, which is lower than either of its two main components, and makes it the first product to be melted when carbon and iron are heated together.[6] Its mechanical properties vary greatly and depend on the form the carbon takes in the alloy.[12]

"White" cast irons contain their carbon in the form of cementite, or iron carbide (Fe3C).[12] This hard, brittle compound dominates the mechanical properties of white cast irons, rendering them hard, but unresistant to shock. The broken surface of a white cast iron is full of fine facets of the broken iron carbide, a very pale, silvery, shiny material, hence the appellation. Cooling a mixture of iron with 0.8% carbon slowly below 723 °C to room temperature results in separate, alternating layers of cementite and α-iron, which is soft and malleable and is called pearlite for its appearance. Rapid cooling, on the other hand, does not allow time for this separation and creates hard and brittle martensite. The steel can then be tempered by reheating to a temperature in between, changing the proportions of pearlite and martensite. The end product below 0.8% carbon content is a pearlite-αFe mixture, and that above 0.8% carbon content is a pearlite-cementite mixture.[12]

In gray iron the carbon exists as separate, fine flakes of graphite, and also renders the material brittle due to the sharp edged flakes of graphite that produce stress concentration sites within the material.[129] A newer variant of gray iron, referred to as ductile iron, is specially treated with trace amounts of magnesium to alter the shape of graphite to spheroids, or nodules, reducing the stress concentrations and vastly increasing the toughness and strength of the material.[129]

Wrought iron contains less than 0.25% carbon but large amounts of slag that give it a fibrous characteristic.[128] It is a tough, malleable product, but not as fusible as pig iron. If honed to an edge, it loses it quickly. Wrought iron is characterized by the presence of fine fibers of slag entrapped within the metal. Wrought iron is more corrosion resistant than steel. It has been almost completely replaced by mild steel for traditional "wrought iron" products and blacksmithing.

Mild steel corrodes more readily than wrought iron, but is cheaper and more widely available. Carbon steel contains 2.0% carbon or less,[130] with small amounts of manganese, sulfur, phosphorus, and silicon. Alloy steels contain varying amounts of carbon as well as other metals, such as chromium, vanadium, molybdenum, nickel, tungsten, etc. Their alloy content raises their cost, and so they are usually only employed for specialist uses. One common alloy steel, though, is stainless steel. Recent developments in ferrous metallurgy have produced a growing range of microalloyed steels, also termed 'HSLA' or high-strength, low alloy steels, containing tiny additions to produce high strengths and often spectacular toughness at minimal cost.[130][131][132]

Alloys with high purity elemental makeups(as alloys of electrolytic iron ) have specifically enhanced properties such as ductility, tensile strength, toughness, fatigue strength, heat resistance, and corrosion resistance.

Apart from traditional applications, iron is also used for protection from ionizing radiation. Although it is lighter than another traditional protection material, lead, it is much stronger mechanically. The attenuation of radiation as a function of energy is shown in the graph.[133]

The main disadvantage of iron and steel is that pure iron, and most of its alloys, suffer badly from rust if not protected in some way, a cost amounting to over 1% of the world's economy.[134] Painting, galvanization, passivation, plastic coating and bluing are all used to protect iron from rust by excluding water and oxygen or by cathodic protection. The mechanism of the rusting of iron is as follows:[134]

Cathode: 3 O2 + 6 H2O + 12 e → 12 OH
Anode: 4 Fe → 4 Fe2+ + 8 e; 4 Fe2+ → 4 Fe3+ + 4 e
Overall: 4 Fe + 3 O2 + 6 H2O → 4 Fe3+ + 12 OH → 4 Fe(OH)3 or 4 FeO(OH) + 4 H2O

The electrolyte is usually iron(II) sulfate in urban areas (formed when atmospheric sulfur dioxide attacks iron), and salt particles in the atmosphere in seaside areas.[134]

Iron compounds

Although the dominant use of iron is in metallurgy, iron compounds are also pervasive in industry. Iron catalysts are traditionally used in the Haber–Bosch process for the production of ammonia and the Fischer–Tropsch process for conversion of carbon monoxide to hydrocarbons for fuels and lubricants.[135] Powdered iron in an acidic solvent was used in the Bechamp reduction the reduction of nitrobenzene to aniline.[136]

Iron(III) oxide mixed with aluminium powder can be ignited to create a thermite reaction, used in welding large iron parts (like rails) and purifying ores. Iron(III) oxide and oxyhidroxide are used as reddish and ocher pigments.

Iron(III) chloride finds use in water purification and sewage treatment, in the dyeing of cloth, as a coloring agent in paints, as an additive in animal feed, and as an etchant for copper in the manufacture of printed circuit boards.[137] It can also be dissolved in alcohol to form tincture of iron, which is used as a medicine to stop bleeding in canaries.[138]

Iron(II) sulfate is used as a precursor to other iron compounds. It is also used to reduce chromate in cement. It is used to fortify foods and treat iron deficiency anemia. Iron(III) sulfate is used in settling minute sewage particles in tank water. Iron(II) chloride is used as a reducing flocculating agent, in the formation of iron complexes and magnetic iron oxides, and as a reducing agent in organic synthesis.[137]

Biological and pathological role

Iron is required for life.[5][139][140] The iron–sulfur clusters are pervasive and include nitrogenase, the enzymes responsible for biological nitrogen fixation. Iron-containing proteins participate in transport, storage and used of oxygen.[5] Iron proteins are involved in electron transfer.[141]

Structure of Heme b; in the protein additional ligand(s) would be attached to Fe.

Examples of iron-containing proteins in higher organisms include hemoglobin, cytochrome (see high-valent iron), and catalase.[5][142] The average adult human contains about 0.005% body weight of iron, or about four grams, of which three quarters is in hemoglobin – a level that remains constant despite only about one milligram of iron being absorbed each day,[141] because the human body recycles its hemoglobin for the iron content.[143]

Microbial growth may be assisted by oxidation of iron(II) or by reduction of iron (III).[144]

Biochemistry

Iron acquisition poses a problem for aerobic organisms because ferric iron is poorly soluble near neutral pH. Thus, these organisms have developed means to absorb iron as complexes, sometimes taking up ferrous iron before oxidising it back to ferric iron.[5] In particular, bacteria have evolved very high-affinity sequestering agents called siderophores.[145][146][147]

After uptake in human cells, iron storage is precisely regulated.[5][148] A major component of this regulation is the protein transferrin, which binds iron ions absorbed from the duodenum and carries it in the blood to cells.[5][149] Transferrin contains Fe3+ in the middle of a distorted octahedron, bonded to one nitrogen, three oxygens and a chelating carbonate anion that traps the Fe3+ ion: it has such a high stability constant that it is very effective at taking up Fe3+ ions even from the most stable complexes. At the bone marrow, transferrin is reduced from Fe3+ and Fe2+ and stored as ferritin to be incorporated into hemoglobin.[141]

The most commonly known and studied bioinorganic iron compounds (biological iron molecules) are the heme proteins: examples are hemoglobin, myoglobin, and cytochrome P450.[5] These compounds participate in transporting gases, building enzymes, and transferring electrons.[141] Metalloproteins are a group of proteins with metal ion cofactors. Some examples of iron metalloproteins are ferritin and rubredoxin.[141] Many enzymes vital to life contain iron, such as catalase,[150] lipoxygenases,[151] and IRE-BP.[152]

Hemoglobin is an oxygen carrier that occurs in red blood cells and contributes their color, transporting oxygen in the arteries from the lungs to the muscles where it is transferred to myoglobin, which stores it until it is needed for the metabolic oxidation of glucose, generating energy.[5] Here the hemoglobin binds to carbon dioxide, produced when glucose is oxidized, which is transported through the veins by hemoglobin (predominantly as bicarbonate anions) back to the lungs where it is exhaled.[141] In hemoglobin, the iron is in one of four heme groups and has six possible coordination sites; four are occupied by nitrogen atoms in a porphyrin ring, the fifth by an imidazole nitrogen in a histidine residue of one of the protein chains attached to the heme group, and the sixth is reserved for the oxygen molecule it can reversibly bind to.[141] When hemoglobin is not attached to oxygen (and is then called deoxyhemoglobin), the Fe2+ ion at the center of the heme group (in the hydrophobic protein interior) is in a high-spin configuration. It is thus too large to fit inside the porphyrin ring, which bends instead into a dome with the Fe2+ ion about 55 picometers above it. In this configuration, the sixth coordination site reserved for the oxygen is blocked by another histidine residue.[141]

When deoxyhemoglobin picks up an oxygen molecule, this histidine residue moves away and returns once the oxygen is securely attached to form a hydrogen bond with it. This results in the Fe2+ ion switching to a low-spin configuration, resulting in a 20% decrease in ionic radius so that now it can fit into the porphyrin ring, which becomes planar.[141] (Additionally, this hydrogen bonding results in the tilting of the oxygen molecule, resulting in a Fe–O–O bond angle of around 120° that avoids the formation of Fe–O–Fe or Fe–O2–Fe bridges that would lead to electron transfer, the oxidation of Fe2+ to Fe3+, and the destruction of hemoglobin.) This results in a movement of all the protein chains that leads to the other subunits of hemoglobin changing shape to a form with larger oxygen affinity. Thus, when deoxyhemoglobin takes up oxygen, its affinity for more oxygen increases, and vice versa.[141] Myoglobin, on the other hand, contains only one heme group and hence this cooperative effect cannot occur. Thus, while hemoglobin is almost saturated with oxygen in the high partial pressures of oxygen found in the lungs, its affinity for oxygen is much lower than that of myoglobin, which oxygenates even at low partial pressures of oxygen found in muscle tissue.[141] As described by the Bohr effect (named after Christian Bohr, the father of Niels Bohr), the oxygen affinity of hemoglobin diminishes in the presence of carbon dioxide.[141]

A heme unit of human carboxyhemoglobin, showing the carbonyl ligand at the apical position, trans to the histidine residue[153]

Carbon monoxide and phosphorus trifluoride are poisonous to humans because they bind to hemoglobin similarly to oxygen, but with much more strength, so that oxygen can no longer be transported throughout the body. Hemoglobin bound to carbon monoxide is known as carboxyhemoglobin. This effect also plays a minor role in the toxicity of cyanide, but there the major effect is by far its interference with the proper functioning of the electron transport protein cytochrome a.[141] The cytochrome proteins also involve heme groups and are involved in the metabolic oxidation of glucose by oxygen. The sixth coordination site is then occupied by either another imidazole nitrogen or a methionine sulfur, so that these proteins are largely inert to oxygen – with the exception of cytochrome a, which bonds directly to oxygen and thus is very easily poisoned by cyanide.[141] Here, the electron transfer takes place as the iron remains in low spin but changes between the +2 and +3 oxidation states. Since the reduction potential of each step is slightly greater than the previous one, the energy is released step-by-step and can thus be stored in adenosine triphosphate. Cytochrome a is slightly distinct, as it occurs at the mitochondrial membrane, binds directly to oxygen, and transports protons as well as electrons, as follows:[141]

4 Cytc2+ + O2 + 8H+
inside
→ 4 Cytc3+ + 2 H2O + 4H+
outside

Although the heme proteins are the most important class of iron-containing proteins, the iron-sulfur proteins are also very important, being involved in electron transfer, which is possible since iron can exist stably in either the +2 or +3 oxidation states. These have one, two, four, or eight iron atoms that are each approximately tetrahedrally coordinated to four sulfur atoms; because of this tetrahedral coordination, they always have high-spin iron. The simplest of such compounds is rubredoxin, which has only one iron atom coordinated to four sulfur atoms from cysteine residues in the surrounding peptide chains. Another important class of iron-sulfur proteins is the ferredoxins, which have multiple iron atoms. Transferrin does not belong to either of these classes.[141]

The ability of sea mussels to maintain their grip on rocks in the ocean is facilitated by their use of organometallic iron-based bonds in their protein-rich cuticles. Based on synthetic replicas, the presence of iron in these structures increased elastic modulus 770 times, tensile strength 58 times, and toughness 92 times. The amount of stress required to permanently damage them increased 76 times.[154]

Nutrition

Diet

Iron is pervasive, but particularly rich sources of dietary iron include red meat, oysters, lentils, beans, poultry, fish, leaf vegetables, watercress, tofu, chickpeas, black-eyed peas, and blackstrap molasses.[5] Bread and breakfast cereals are sometimes specifically fortified with iron.[5][155]

Iron provided by dietary supplements is often found as iron(II) fumarate, although iron(II) sulfate is cheaper and is absorbed equally well.[137] Elemental iron, or reduced iron, despite being absorbed at only one-third to two-thirds the efficiency (relative to iron sulfate),[156] is often added to foods such as breakfast cereals or enriched wheat flour. Iron is most available to the body when chelated to amino acids[157] and is also available for use as a common iron supplement. Glycine, the least expensive amino acid, is most often used to produce iron glycinate supplements.[158]

Dietary recommendations

The U.S. Institute of Medicine (IOM) updated Estimated Average Requirements (EARs) and Recommended Dietary Allowances (RDAs) for iron in 2001.[5] The current EAR for iron for women ages 14–18 is 7.9 mg/day, 8.1 for ages 19–50 and 5.0 thereafter (post menopause). For men the EAR is 6.0 mg/day for ages 19 and up. The RDA is 15.0 mg/day for women ages 15–18, 18.0 for 19–50 and 8.0 thereafter. For men, 8.0 mg/day for ages 19 and up. RDAs are higher than EARs so as to identify amounts that will cover people with higher than average requirements. RDA for pregnancy is 27 mg/day and, for lactation, 9 mg/day.[5] For children ages 1–3 years 7 mg/day, 10 for ages 4–8 and 8 for ages 9–13. As for safety, the IOM also sets Tolerable upper intake levels (ULs) for vitamins and minerals when evidence is sufficient. In the case of iron the UL is set at 45 mg/day. Collectively the EARs, RDAs and ULs are referred to as Dietary Reference Intakes.[159]

The European Food Safety Authority (EFSA) refers to the collective set of information as Dietary Reference Values, with Population Reference Intake (PRI) instead of RDA, and Average Requirement instead of EAR. AI and UL defined the same as in United States. For women the PRI is 13 mg/day ages 15–17 years, 16 mg/day for women ages 18 and up who are premenopausal and 11 mg/day postmenopausal. For pregnancy and lactation, 16 mg/day. For men the PRI is 11 mg/day ages 15 and older. For children ages 1 to 14 the PRI increases from 7 to 11 mg/day. The PRIs are higher than the U.S. RDAs, with the exception of pregnancy.[160] The EFSA reviewed the same safety question did not establish a UL.[161]

Infants may require iron supplements if they are bottle-fed cow's milk.[162] Frequent blood donors are at risk of low iron levels and are often advised to supplement their iron intake.[163]

For U.S. food and dietary supplement labeling purposes the amount in a serving is expressed as a percent of Daily Value (%DV). For iron labeling purposes 100% of the Daily Value was 18 mg, and as of May 27, 2016 remained unchanged at 18 mg.[164][165] Compliance with the updated labeling regulations was required by 1 January 2020 for manufacturers with US$10 million or more in annual food sales, and by 1 January 2021 for manufacturers with lower volume food sales.[166][167] A table of the old and new adult daily values is provided at Reference Daily Intake.

Deficiency

Iron deficiency is the most common nutritional deficiency in the world.[5][168][169][170] When loss of iron is not adequately compensated by adequate dietary iron intake, a state of latent iron deficiency occurs, which over time leads to iron-deficiency anemia if left untreated, which is characterised by an insufficient number of red blood cells and an insufficient amount of hemoglobin.[171] Children, pre-menopausal women (women of child-bearing age), and people with poor diet are most susceptible to the disease. Most cases of iron-deficiency anemia are mild, but if not treated can cause problems like fast or irregular heartbeat, complications during pregnancy, and delayed growth in infants and children.[172]

Excess

Iron uptake is tightly regulated by the human body, which has no regulated physiological means of excreting iron. Only small amounts of iron are lost daily due to mucosal and skin epithelial cell sloughing, so control of iron levels is primarily accomplished by regulating uptake.[173] Regulation of iron uptake is impaired in some people as a result of a genetic defect that maps to the HLA-H gene region on chromosome 6 and leads to abnormally low levels of hepcidin, a key regulator of the entry of iron into the circulatory system in mammals.[174] In these people, excessive iron intake can result in iron overload disorders, known medically as hemochromatosis.[5] Many people have an undiagnosed genetic susceptibility to iron overload, and are not aware of a family history of the problem. For this reason, people should not take iron supplements unless they suffer from iron deficiency and have consulted a doctor. Hemochromatosis is estimated to be the cause of 0.3 to 0.8% of all metabolic diseases of Caucasians.[175]

Overdoses of ingested iron can cause excessive levels of free iron in the blood. High blood levels of free ferrous iron react with peroxides to produce highly reactive free radicals that can damage DNA, proteins, lipids, and other cellular components. Iron toxicity occurs when the cell contains free iron, which generally occurs when iron levels exceed the availability of transferrin to bind the iron. Damage to the cells of the gastrointestinal tract can also prevent them from regulating iron absorption, leading to further increases in blood levels. Iron typically damages cells in the heart, liver and elsewhere, causing adverse effects that include coma, metabolic acidosis, shock, liver failure, coagulopathy, adult respiratory distress syndrome, long-term organ damage, and even death.[176] Humans experience iron toxicity when the iron exceeds 20 milligrams for every kilogram of body mass; 60 milligrams per kilogram is considered a lethal dose.[177] Overconsumption of iron, often the result of children eating large quantities of ferrous sulfate tablets intended for adult consumption, is one of the most common toxicological causes of death in children under six.[177] The Dietary Reference Intake (DRI) sets the Tolerable Upper Intake Level (UL) for adults at 45 mg/day. For children under fourteen years old the UL is 40 mg/day.[178]

The medical management of iron toxicity is complicated, and can include use of a specific chelating agent called deferoxamine to bind and expel excess iron from the body.[176][179][180]

Cancer

The role of iron in cancer defense can be described as a "double-edged sword" because of its pervasive presence in non-pathological processes.[181] People having chemotherapy may develop iron deficiency and anemia, for which intravenous iron therapy is used to restore iron levels.[182] Iron overload, which may occur from high consumption of red meat,[5] may initiate tumor growth and increase susceptibility to cancer onset,[182] particularly for colorectal cancer.[5]

Marine systems

Iron plays an essential role in marine systems and can act as a limiting nutrient for planktonic activity.[183] Because of this, too much of a decrease in iron may lead to a decrease in growth rates in phytoplanktonic organisms such as diatoms.[184] Iron can also be oxidized by marine microbes under conditions that are high in iron and low in oxygen.[185]

Iron can enter marine systems through adjoining rivers and directly from the atmosphere. Once iron enters the ocean, it can be distributed throughout the water column through ocean mixing and through recycling on the cellular level.[186] In the arctic, sea ice plays a major role in the store and distribution of iron in the ocean, depleting oceanic iron as it freezes in the winter and releasing it back into the water when thawing occurs in the summer.[187] The iron cycle can fluctuate the forms of iron from aqueous to particle forms altering the availability of iron to primary producers.[188] Increased light and warmth increases the amount of iron that is in forms that are usable by primary producers.[189]

See also

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Bibliography

Further reading

  • H.R. Schubert, History of the British Iron and Steel Industry ... to 1775 AD (Routledge, London, 1957)
  • R.F. Tylecote, History of Metallurgy (Institute of Materials, London 1992).
  • R.F. Tylecote, "Iron in the Industrial Revolution" in J. Day and R.F. Tylecote, The Industrial Revolution in Metals (Institute of Materials 1991), 200–60.

External links