Solubility Calcium carbonate

1 solubility

1.1 varying co2 pressure
1.2 varying ph, temperature , salinity: caco3 scaling in swimming pools
1.3 solubility in strong or weak acid solution

solubility
with varying co2 pressure

travertine calcium carbonate deposits hot spring

calcium carbonate poorly soluble in pure water (47 mg/l @ normal atmospheric co2 partial pressure shown below).

the equilibrium of solution given equation (with dissolved calcium carbonate on right):

where solubility product [ca][co3] given anywhere ksp = 3.7×10 ksp = 8.7×10 @ 25 °c, depending upon data source. equation means product of molar concentration of calcium ions (moles of dissolved ca per liter of solution) molar concentration of dissolved co3 cannot exceed value of ksp. seemingly simple solubility equation, however, must taken along more complicated equilibrium of carbon dioxide water (see carbonic acid). of co3 combines h in solution according to:

hco3 known bicarbonate ion. calcium bicarbonate many times more soluble in water calcium carbonate—indeed exists in solution.

some of hco3 combines h in solution according to:

some of h2co3 breaks water , dissolved carbon dioxide according to:

and dissolved carbon dioxide in equilibrium atmospheric carbon dioxide according to:

for ambient air,




p


co

2









{\displaystyle p_{\ce {co2}}}

around 3.5×10 atmospheres (or equivalently 35 pa). last equation above fixes concentration of dissolved co2 function of




p


co

2









{\displaystyle p_{\ce {co2}}}

, independent of concentration of dissolved caco3. @ atmospheric partial pressure of co2, dissolved co2 concentration 1.2×10 moles/liter. equation before fixes concentration of h2co3 function of [co2]. [co2]=1.2×10, results in [h2co3]=2.0×10 moles per liter. when [h2co3] known, remaining 3 equations with

(which true aqueous solutions), , fact solution must electrically neutral,

2[ca] + [h] = [hco3] + 2[co3] + [oh]

make possible solve simultaneously remaining 5 unknown concentrations (note above form of neutrality equation valid if calcium carbonate has been put in contact pure water or neutral ph solution; in case initial water solvent ph not neutral, equation modified).

the table on right shows result [ca] , [h] (in form of ph) function of ambient partial pressure of co2 (ksp = 4.47×10 has been taken calculation).

at atmospheric levels of ambient co2 table indicates solution alkaline maximum caco3 solubility of 47 mg/l.
as ambient co2 partial pressure reduced below atmospheric levels, solution becomes more , more alkaline. @ extremely low




p


co

2









{\displaystyle p_{\ce {co2}}}

, dissolved co2, bicarbonate ion, , carbonate ion largely evaporate solution, leaving highly alkaline solution of calcium hydroxide, more soluble caco3. note




p


co

2







=

10

−
12



a
t
m



{\displaystyle p_{\ce {co2}}=10^{-12}\mathrm {atm} }

, [ca][oh] product still below solubility product of ca(oh)2 (8×10). still lower co2 pressure, ca(oh)2 precipitation occur before caco3 precipitation.
as ambient co2 partial pressure increases levels above atmospheric, ph drops, , of carbonate ion converted bicarbonate ion, results in higher solubility of ca.

the effect of latter evident in day-to-day life of people have hard water. water in aquifers underground can exposed levels of co2 higher atmospheric. such water percolates through calcium carbonate rock, caco3 dissolves according second trend. when same water emerges tap, in time comes equilibrium co2 levels in air outgassing excess co2. calcium carbonate becomes less soluble result , excess precipitates lime scale. same process responsible formation of stalactites , stalagmites in limestone caves.

two hydrated phases of calcium carbonate, monohydrocalcite, caco3·h2o , ikaite, caco3·6h2o, may precipitate water @ ambient conditions , persist metastable phases.

with varying ph, temperature , salinity: caco3 scaling in swimming pools

in contrast open equilibrium scenario above, many swimming pools managed addition of sodium bicarbonate (nahco3) 2 mm buffer, control of ph through use of hcl, nahso4, na2co3, naoh or chlorine formulations acidic or basic. in situation, dissolved inorganic carbon (dic) far equilibrium atmospheric co2. progress towards equilibrium through outgassing of co2 slowed (i) slow reaction h2co3 ⇌ co2(aq) + h2o; (ii) limited aeration in deep water column , (iii) periodic replenishment of bicarbonate maintain buffer capacity (often estimated through measurement of ‘total alkalinity’).

in situation, dissociation constants faster reactions h2co3 ⇌ h + hco3 ⇌ 2 h + co3 allow prediction of concentrations of each dic species in solution, added concentration of hco3 (which comprises more 90% of total dic ph 7 ph 8 @ 25 ˚c in fresh water. addition of hco3 increase co3 concentration @ ph. rearranging equations given above, can see [ca] = ksp / [co3], , [co3] = ka2 × [hco3] / [h]. therefore, when hco3 concentration known, maximum concentration of ca ions before scaling through caco3 precipitation can predicted formula:

camax = (ksp / ka2) × ([h] / [hco3])

the solubility product caco3 (ksp) , dissociation constants dic species (including ka2) substantially affected temperature , salinity, overall effect camax increases fresh salt water, , decreases rising temperature, ph, or added bicarbonate level, illustrated in accompanying graphs.

the trends illustrative pool management, whether scaling occurs depends on other factors including interactions mg, b(oh)4 , other ions in pool, supersaturation effects. scaling commonly observed in electrolytic chlorine generators, there high ph near cathode surface , scale deposition further increases temperature. 1 reason pool operators prefer borate on bicarbonate primary ph buffer, , avoid use of pool chemicals containing calcium.

solubility in strong or weak acid solution

solutions of strong (hcl), moderately strong (sulfamic) or weak (acetic, citric, sorbic, lactic, phosphoric) acids commercially available. commonly used descaling agents remove limescale deposits. maximum amount of caco3 can dissolved 1 liter of acid solution can calculated using above equilibrium equations.

in case of strong monoacid decreasing acid concentration [a] = [a], obtain (with caco3 molar mass = 100 g):

where initial state acid solution no ca (not taking account possible co2 dissolution) , final state solution saturated ca. strong acid concentrations, species have negligible concentration in final state respect ca , neutrality equation reduces approximately 2[ca] = [a] yielding




[


c
a


2
+


]
≃



[


a


−


]

2





{\displaystyle \scriptstyle [\mathrm {ca} ^{2+}]\simeq {\frac {[\mathrm {a} ^{-}]}{2}}}

. when concentration decreases, [hco3] becomes non-negligible preceding expression no longer valid. vanishing acid concentrations, 1 can recover final ph , solubility of caco3 in pure water.

in case of weak monoacid (here take acetic acid pka = 4.76) decreasing total acid concentration [a] = [a]+[ah], obtain:

for same total acid concentration, initial ph of weak acid less acid 1 of strong acid; however, maximum amount of caco3 can dissolved approximately same. because in final state, ph larger pka, weak acid dissociated, yielding in end many h ions strong acid dissolve calcium carbonate.

the calculation in case of phosphoric acid (which used domestic applications) more complicated since concentrations of 4 dissociation states corresponding acid must calculated [hco3], [co3], [ca], [h] , [oh]. system may reduced seventh degree equation [h] numerical solution of gives

where [a] = [h3po4] + [h2po4] + [hpo4] + [po4] total acid concentration. phosphoric acid more efficient monoacid since @ final neutral ph, second dissociated state concentration [hpo4] not negligible (see phosphoric acid).