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Why does dimethyl dichlorosilane undergo hydrolysis to form siloxane polymers, but its carbon analog does not?
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+ Biochemistry
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Marlene Dixon
Why does dimethyl dichlorosilane undergo hydrolysis to form siloxane polymers, but its carbon analog does not?
Actually a $\ce{Si-Cl}$ bond is much stronger than a $\ce{C-Cl}$ bond. It takes about 90 kcal/mole to break a $\ce{Si-Cl}$ bond, but only around 81 kcal/mole to break a $\ce{C-Cl}$ bond. The strengths of the corresponding $\ce{Si-O}$ (110 kcal/mole) and $\ce{C-O}$ (85 kcal/mole) bonds also need to be considered.
In the silicon reaction we need 90 kcal/mole to break the $\ce{Si-Cl}$ bond, but we get 110 kcal/mole back when we form the $\ce{Si-O}$; the reaction is exothermic by roughly 20 kcal/mole - a large amount. In the analogous carbon case, we need 81 kcal/mole to break a $\ce{C-Cl}$ bond, and we only get 85 kcal/mole back when we form the $\ce{C-O}$ bond; the reaction is exothermic by only 4 kcal/mole - not much. There is quite a difference in the overall driving force for the two reactions with the silicon case being energetically much more favorable than the carbon case!
Actually a $\ce{Si-Cl}$ bond is much stronger than a $\ce{C-Cl}$ bond. It takes about 90 kcal/mole to break a $\ce{Si-Cl}$ bond, but only around 81 kcal/mole to break a $\ce{C-Cl}$ bond. The strengths of the corresponding $\ce{Si-O}$ (110 kcal/mole) and $\ce{C-O}$ (85 kcal/mole) bonds also need to be considered.
In the silicon reaction we need 90 kcal/mole to break the $\ce{Si-Cl}$ bond, but we get 110 kcal/mole back when we form the $\ce{Si-O}$; the reaction is exothermic by roughly 20 kcal/mole - a large amount. In the analogous carbon case, we need 81 kcal/mole to break a $\ce{C-Cl}$ bond, and we only get 85 kcal/mole back when we form the $\ce{C-O}$ bond; the reaction is exothermic by only 4 kcal/mole - not much. There is quite a difference in the overall driving force for the two reactions with the silicon case being energetically much more favorable than the carbon case!
Your guess is correct. The geminal diol is unstable because carbon atom can't hold more than one $\ce{-OH}$ group. Where as Si atom can hold three $\ce{-OH}$ groups. It is this property of silicon that makes the formation of organosilicon polymers possible.
Your guess is correct. The geminal diol is unstable because carbon atom can't hold more than one $\ce{-OH}$ group. Where as Si atom can hold three $\ce{-OH}$ groups. It is this property of silicon that makes the formation of organosilicon polymers possible.
Its not that you couldnt make polyacetal. In fact Dupont will sell it to you (under the name of Delrin) and its a favourite with machinists because it holds shape extremely well.More
Actually a $\ce{Si-Cl}$ bond is much stronger than a $\ce{C-Cl}$ bond. It takes about 90 kcal/mole to break a $\ce{Si-Cl}$ bond, but only around 81 kcal/mole to break a $\ce{C-Cl}$ bond. The strengths of the corresponding $\ce{Si-O}$ (110 kcal/mole) and $\ce{C-O}$ (85 kcal/mole) bonds also need to be considered.
In the silicon reaction we need 90 kcal/mole to break the $\ce{Si-Cl}$ bond, but we get 110 kcal/mole back when we form the $\ce{Si-O}$; the reaction is exothermic by roughly 20 kcal/mole - a large amount. In the analogous carbon case, we need 81 kcal/mole to break a $\ce{C-Cl}$ bond, and we only get 85 kcal/mole back when we form the $\ce{C-O}$ bond; the reaction is exothermic by only 4 kcal/mole - not much. There is quite a difference in the overall driving force for the two reactions with the silicon case being energetically much more favorable than the carbon case!
Actually a $\ce{Si-Cl}$ bond is much stronger than a $\ce{C-Cl}$ bond. It takes about 90 kcal/mole to break a $\ce{Si-Cl}$ bond, but only around 81 kcal/mole to break a $\ce{C-Cl}$ bond. The strengths of the corresponding $\ce{Si-O}$ (110 kcal/mole) and $\ce{C-O}$ (85 kcal/mole) bonds also need to be considered.
In the silicon reaction we need 90 kcal/mole to break the $\ce{Si-Cl}$ bond, but we get 110 kcal/mole back when we form the $\ce{Si-O}$; the reaction is exothermic by roughly 20 kcal/mole - a large amount. In the analogous carbon case, we need 81 kcal/mole to break a $\ce{C-Cl}$ bond, and we only get 85 kcal/mole back when we form the $\ce{C-O}$ bond; the reaction is exothermic by only 4 kcal/mole - not much. There is quite a difference in the overall driving force for the two reactions with the silicon case being energetically much more favorable than the carbon case!
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Your guess is correct. The geminal diol is unstable because carbon atom can't hold more than one $\ce{-OH}$ group. Where as Si atom can hold three $\ce{-OH}$ groups. It is this property of silicon that makes the formation of organosilicon polymers possible.
Your guess is correct. The geminal diol is unstable because carbon atom can't hold more than one $\ce{-OH}$ group. Where as Si atom can hold three $\ce{-OH}$ groups. It is this property of silicon that makes the formation of organosilicon polymers possible.
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