gqi-00042 & gqi-00048 aula 04
DESCRIPTION
GQI-00048 Química Geral e Tecnológica & GQI-00042 Química Geral e Inorgânica Experimental. Aula 04.TRANSCRIPT
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GQI 00042 Química Geral e Inorg. Exp. III &
GQI 00048 Química Geral Tecnológica
Física e Engas. Civil, Elétrica, de Petróleo e de Recusos Hídricos e Meio Ambiente
Prof. Ednilsom Orestes 1º Semestre de 2014
17/02/2014 – 27/06/2014
Universidade Federal Fluminense
Instituto de Química de São Carlos
Departamento de Química Inorgânica
www.slideshare.net/Ednilsom AULA 04
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FORMA E ESTRUTURA MOLECULAR
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Estrutura de Lewis
Retrata apenas a conectividade (que átomos estão ligados), não a geometria (arranjo tridimensional).
Par isolado: Par de elétrons localizado sobre um átomo; não participa de ligação alguma.
Par ligante: Par de elétrons diretamente envolvidos numa ligação.
Regra do Octeto
Tendência de moléculas e íons poliatômicos em assumir estruturas onde cada átomo fica com 8(2) elétrons na
camada de valência.
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Procedimento para construção de Estruturas de Lewis
1) Calcule o número total de elétrons (pares) considerando os íons.
2) Desenhe a estrutura considerando átomos centrais (normalmente com baixa eletronegatividade; fazem mais de uma ligação; carbono sempre) e terminais (fazem uma ligação somente; hidrogênio sempre), além da simetria.
3) Forme ligações simples.
4) Distribua restante dos elétrons. Primeiro nos átomos terminais como pares isolados. Se átomos centrais tiverem menos de 8 elétrons mova pares isolados de átomos terminais para pares ligados com átomos centrais.
5) Confira se todos átomos tem 8 elétrons.
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OH2 COH2 ClO2-
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Escreva as Estruturas de Lewis para o íon cianeto, CNO-
(carbono é central).
Escreva as Estruturas de Lewis para a amônia, NH3.
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Escreva as Estruturas de Lewis para o ácido acético, CH3COOH.
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Escreva as Estruturas de Lewis para a uréia, (NH2)2CO. Escreva as Estruturas de Lewis para a hidrazina, H2NNH2. Desenhe o híbrido de ressonância para a molécula de ozônio (O3) sabendo que as duas ligações tem o mesmo comprimento.
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Valence-Shell Electron Pair Repulsion
Modelo VSEPR
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∡ 𝐻𝐶𝐻 = 109,5°
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∡ 𝐹𝑆𝐹 = 90° 𝑒 180°
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∡ 𝐶𝑙𝑃𝐶𝑙 = 90°, 120° 𝑒 180°
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Regra no. 1
Regiões com altas concentrações elétrons se
repelem e afastam-se o máximo possível para reduzir
o efeito da repulsão
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Prediga a forma da molécula de etino (acetileno), HC≡CH.
Prediga a forma das moléculas do pentafluoreto de arsênio, AsF5 (bipiramide trigonal), e do formaldeído, CH2O.
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Regra no. 2
Não há distinção entre ligações simples e múltiplas
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Prediga a forma da molécula de trifluoreto de nitrogênio, NF3.
Prediga a forma das moléculas IF5 (octaédrico) e SO2.
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Regra no. 3
Elétrons isolados também repelem elétrons ligados e
são incluídos na descrição do arranjo
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Prediga a forma da molécula de tetrafluoreto de enxofre, SF4.
Prediga a forma das moléculas I3- (linear), e XeF4.
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Regra no. 4
Repulsão entre pares isolados > pares isolado-ligado > pares
ligado
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MOLÉCULAS POLARES
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CCl3H
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Diga se as moléculas de (a) BF3 e (b) O3 são polares ou apolares.
(c) SF4, [polar] (d) SF6, [apolar] (e) PCl5 e (f) IF5.
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TEORIA DA LIGAÇÃO DE VALÊNCIA
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Teoria da Ligação de Valência
• 1ª. Teoria Mecânico-Quântica para a ligação química a ser desenvolvida.
• 2 elétrons localizados entre 2 átomos.
• Envolve somente orbitais dos átomos ligados.
• Conceitos persistem (emparelhamento de spins, ligações σ e π, hibridização).
Ex: Caso mais simples: H2.
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OU
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HIBRIDIZAÇÃO
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Diga qual é a hibridização do enxofre no tetrafluoreto de fósforo, PF5.
Resp.: sp3d
Repita para trifluoreto de cloro [sp3d] e do BrF4-.
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Diga quais são ligações 𝜎 e 𝜋 e quais os ângulos entre elas para a molécule de ácido fórmico HCOOH.
Repita para C3O2, (OCCCO). Linear, hibridização sp, C´s ligação 𝜎 e 𝜋 com C e O.
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• Híbridos = mistura de orbitais atômicos.
• Possuem forma diferente dos orbitais originais.
• No. híbridos = No. orbitais originais.
• Promovem a superposição estabilidade.
• Ligações: híbridos + híbridos; híbridos + puros; puros + puros.
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O = 1s2 2s2 2p4
↑↓ ↑↓ ↑↓ ↑ ↑
Mas,...
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TEORIA DOS ORBITAIS MOLECULARES
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Teoria do Orbital Molecular
• Elétrons deslocalizados pela molécula
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Teoria do Orbital Molecular
• Elétrons deslocalizados pela molécula Ex: Íon molecular H2
+.
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Teoria do Orbital Molecular
• Elétrons deslocalizados pela molécula Ex: Íon molecular H2
+.
𝐻 = −ℏ2
2𝑚𝑒𝛻𝑒2 −
𝑒2
4𝜋𝜖0
1
𝑟𝑎+
1
𝑟𝑏−1
𝑅
𝐻 Ψ = 𝐸Ψ
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• Interferência destrutiva. • Orbitais ocupados reduzem força de coesão entre os átomos. • Efeito desestabilizante – fora da região internuclear (região
ligante) afastando os elétrons. • 𝐸− − 𝐸𝐻1𝑠 > 𝐸+ − 𝐸𝐻1𝑠
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He 1𝑠2 He 1𝑠2
He2
E
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Moléculas diatômicas heteronucleares
• Elétrons da ligação não são verdadeiramente covalentes.
• Densidade deslocada – ligação polar.
• Átomos adquirem cargas parciais negativas, 𝛿−, e positivas, 𝛿+.
• Coeficientes da LCAO tem pesos diferentes.
Ψ = 𝑐𝐴𝜓𝐴 + 𝑐𝐵𝜓𝐵
Ex.: HF, CO
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Ordem de ligação
𝑏 =𝑛 − 𝑛∗
2
• Cada par de elétrons eleva ordem em 1 unidade.
• Quanto maior a ordem, menor comprimento da ligação.
• Quanto maior a ordem, maior a força da ligação.
Ligação Ordem Comprim. (Å) En. diss. (kJ/mol)
HH 1 0,74 432,1
NN 3 1,097 941,7
HCl 1 1,274 427,7
CH 1 1,14 435
CC 1 1,54 368
CC 2 1,34 720
CC 3 1,20 962
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