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Influence of pH on precipitate crystallite sizes?
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Kevin McClear
Influence of pH on precipitate crystallite sizes?
At low pHs there is little "free" $\ce{CO3^{2-}}$ to precipitate the $\ce{Mg^{2+}}$ and $\ce{Ca^{2+}}$ cations. Most of the carbonate species are dissolved $\ce{CO2}$, $\ce{H2CO3}$ and $\ce{HCO3-}$. Thus the precipitate forms slowly and you get relatively large crystals.
At high pHs there is a lot "free" $\ce{CO3^{2-}}$ to precipitate the $\ce{Mg^{2+}}$ and $\ce{Ca^{2+}}$ cations. Thus the precipitate forms relatively rapidly and you get relatively small crystals.
I have no idea what literature would lead you to believe otherwise.
At low pHs there is little "free" $\ce{CO3^{2-}}$ to precipitate the $\ce{Mg^{2+}}$ and $\ce{Ca^{2+}}$ cations. Most of the carbonate species are dissolved $\ce{CO2}$, $\ce{H2CO3}$ and $\ce{HCO3-}$. Thus the precipitate forms slowly and you get relatively large crystals.
At high pHs there is a lot "free" $\ce{CO3^{2-}}$ to precipitate the $\ce{Mg^{2+}}$ and $\ce{Ca^{2+}}$ cations. Thus the precipitate forms relatively rapidly and you get relatively small crystals.
I have no idea what literature would lead you to believe otherwise.
Aside of influence on the crystallization rate, you have also consider the direct chemical influence.
The sodium carbonate hydrolyzes in not enough alkaline solutions.
$$\ce{Na2CO3 + H2O <=> NaHCO3 + NaOH}$$
Precipitation of carbonate is interfered by solubility of Bicarbonates.
$$\ce{CaCO3 v + H2O + HCO3- <=> Ca(HCO3)2 + OH-} $$
Addition of the hydroxide affects both processes. It shifts the equilibrium towards precipitation and also speeds up the precipitation.
The same for magnesium.
For magnesium, there is possibility of a side reaction of precipitation the insoluble hydroxide. But lower carbonate solubility will probably overrule this.
Aside of influence on the crystallization rate, you have also consider the direct chemical influence.
The sodium carbonate hydrolyzes in not enough alkaline solutions.$$\ce{Na2CO3 + H2O <=> NaHCO3 + NaOH}$$
Precipitation of carbonate is interfered by solubility of Bicarbonates.$$\ce{CaCO3 v + H2O + HCO3- <=> Ca(HCO3)2 + OH-} $$
Addition of the hydroxide affects both processes. It shifts the equilibrium towards precipitation and also speeds up the precipitation.
The same for magnesium.
For magnesium, there is possibility of a side reaction of precipitation the insoluble hydroxide. But lower carbonate solubility will probably overrule this.
At low pHs there is little "free" $\ce{CO3^{2-}}$ to precipitate the $\ce{Mg^{2+}}$ and $\ce{Ca^{2+}}$ cations. Most of the carbonate species are dissolved $\ce{CO2}$, $\ce{H2CO3}$ and $\ce{HCO3-}$. Thus the precipitate forms slowly and you get relatively large crystals.
At high pHs there is a lot "free" $\ce{CO3^{2-}}$ to precipitate the $\ce{Mg^{2+}}$ and $\ce{Ca^{2+}}$ cations. Thus the precipitate forms relatively rapidly and you get relatively small crystals.
I have no idea what literature would lead you to believe otherwise.
At low pHs there is little "free" $\ce{CO3^{2-}}$ to precipitate the $\ce{Mg^{2+}}$ and $\ce{Ca^{2+}}$ cations. Most of the carbonate species are dissolved $\ce{CO2}$, $\ce{H2CO3}$ and $\ce{HCO3-}$. Thus the precipitate forms slowly and you get relatively large crystals.
At high pHs there is a lot "free" $\ce{CO3^{2-}}$ to precipitate the $\ce{Mg^{2+}}$ and $\ce{Ca^{2+}}$ cations. Thus the precipitate forms relatively rapidly and you get relatively small crystals.
I have no idea what literature would lead you to believe otherwise.
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Aside of influence on the crystallization rate, you have also consider the direct chemical influence.
The sodium carbonate hydrolyzes in not enough alkaline solutions. $$\ce{Na2CO3 + H2O <=> NaHCO3 + NaOH}$$
Precipitation of carbonate is interfered by solubility of Bicarbonates. $$\ce{CaCO3 v + H2O + HCO3- <=> Ca(HCO3)2 + OH-} $$
Addition of the hydroxide affects both processes. It shifts the equilibrium towards precipitation and also speeds up the precipitation.
The same for magnesium.
For magnesium, there is possibility of a side reaction of precipitation the insoluble hydroxide. But lower carbonate solubility will probably overrule this.
$$\ce{MgCl2 + 2 NaOH -> Mg(OH)2 v + 2 NaCl}$$
Aside of influence on the crystallization rate, you have also consider the direct chemical influence.
The sodium carbonate hydrolyzes in not enough alkaline solutions.$$\ce{Na2CO3 + H2O <=> NaHCO3 + NaOH}$$
Precipitation of carbonate is interfered by solubility of Bicarbonates.$$\ce{CaCO3 v + H2O + HCO3- <=> Ca(HCO3)2 + OH-} $$
Addition of the hydroxide affects both processes. It shifts the equilibrium towards precipitation and also speeds up the precipitation.
The same for magnesium.
For magnesium, there is possibility of a side reaction of precipitation the insoluble hydroxide. But lower carbonate solubility will probably overrule this.
$$\ce{MgCl2 + 2 NaOH -> Mg(OH)2 v + 2 NaCl}$$
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