calcium reactor??

derek_s · Mar 18, 2009

The problem with my seachem is that I alwars ge the same results. 1100. NA the other problem is the reference sample is...you guessed it, 1100. Boo that.

rostato · Mar 18, 2009

I use salifret and seachem. The differ about 200 PPM and the seachem is always lower... I use the two and meet them both halfway...

derek_s · Mar 18, 2009

The seachem def is not the easiest to use. THe Elos is not for the weak, but it is easier to perform. You basically do 2 titration tests and subtract the results, then multiply them by a number.

Is salifert the same method>?

glxtrix · Mar 18, 2009

Derek_S;311209 wrote: You basically do 2 titration tests and subtract the results, then multiply them by a number.

:confused2:

glxtrix · Mar 18, 2009

I've been a salifert buyer since day 1. Always does me well....so I think, lol.

derek_s · Mar 18, 2009

Yeah its wierd. Count the drops, subtract..yada

rostato · Mar 18, 2009

no salifret is just adding some drops and powder, and then adding a solution intill it turns a color. Then what ever is left in your vial is your number of MG. Really simple test kit. The seachem is a PITA but I still use it.

glxtrix · Mar 18, 2009

Yeah all Salifert tests are pretty easy. I heard a rumor though that they may be going out of business....this that true?

derek_s · Mar 18, 2009

I have no clue, but I though that rumor had been around for years. I bet Tim would know.

glxtrix · Mar 18, 2009

I hope not, I love their test kits! lol.

dannybradley · Mar 18, 2009

Derek_S;311227 wrote: Yeah its wierd. Count the drops, subtract..yada

I would say the Seachem method is more accurate than that. Because you're dealing with a known volume of titrant, not counting drops that may vary in size.

The lower Seachem level may be due to you forming smaller drops with Elos than what is required by either rushing them out of the bottle, or not holding it vertically.

Although Seachem does have drop counting, they don't affect the outcome like the titrant does. Step 2 is precipitation, where a slightly smaller amount won't affect it as severely. Step 4 is just color change indicator dye. It would change color regardless of the volume of dye added.

sailfish · Mar 18, 2009

The problem with sailfert was that they had a bad batch of test and lots of people contacted them I got not response.

There are store owners here that contacted them and got no response.

I guess I have a thing for taking responsibility and believe that if you screw up one should admit it not run from it.

I have started to believe that all mag tests suck although I have not found Elos to be wrong yet.

I just figure if I keep it around 1350 I should be ok not to low and not to high even if the test is off by 100.

Joe

glxtrix · Mar 18, 2009

That is a nice safe number Joe. I typically keep mine at 1400-ish.

grouper therapy · Mar 18, 2009

I think whichever test kit you use ( as long as there reasonably accurate) keep using that particular brand and set your dosing, reactors etc. to that brand and your ahead of the game . Stability

tdwyatt · Mar 18, 2009

I have the Pinpoint Ca++ probe, it was an expensive mistake. They are great if you are going to calibrate it, then test a bunch of different samples within a day or so, but after that, it really needs recalibration to give useful results (like if you have several systems to test for Ca++). They will not be accurate for continuous monitoring like we use pH meters.

Many issues with buffer content, as it is really a mix of both bicarbonate (~89%) and carbonate(~9%) plus other buffers, and is a reflection of pH, pCO2, Total CO2 (ZCO2), and total alkalinity ( AT ), so much so that given the value of any three factors, we can determine the others. The problem then becomes "how do we measure these variables accurately so a controller could use their input?"

The Ca++ probe/meter is around $250 USD, can you imagine how much the alk meter would have to be? Makes a Salifert kit look like a bargain.

derek_s · Mar 18, 2009

David, nowhere on here did I comment on using the CA reactor to raise your alk/CA or balancing. You false accuser....

grouper therapy · Mar 18, 2009

see ya in about 30

dannybradley · Mar 18, 2009

tdwyatt;312316 wrote:
Many issues with buffer content, as it is really a mix of both bicarbonate (~89%) and carbonate(~9%) plus other buffers, and is a reflection of pH, pCO2, Total CO2 (ZCO2), and total alkalinity ( AT ), so much so that given the value of any three factors, we can determine the others.

I considered this, that even though the pH is based on what comprises total alkalinity, wouldn't it take the same amount of a given acid to achieve the desired pH from another pH regardless of what alkalinity is comprised of?

tdwyatt · Mar 18, 2009

DannyBradley;312384 wrote: I considered this, that even though the pH is based on what comprises total alkalinity, wouldn't it take the same amount of a given acid to achieve the desired pH from another pH regardless of what alkalinity is comprised of?

pH is not based on alkalinity, although it has an effect on the pH of a given water column: pH is driven by the amount of CO2 dissolved in the water (basically the partial pressure of CO2 (pCO2)). The issue is that the different salts have different buffer capacities, and will reach a different final pH value based on its reaction with carbonic acid in marine environments. Buffers are just like acids, they have dirrerent strengths based on composition, ionization, equilibria, pKa, and concentration at equilibrium.

Much of what follows is based on the discussion by James Millero of pH in seawater and the measurement of such by indirect references to known values and dissociation constants of known buffers and acids, see Chemical Oceanography, 2nd edition, pp272-277*

When a CO2 atmosphere is in contact with seawater, the following eqilibria are established:

(1) CO2(g) ↔ CO2(aq)

(2) CO2(aq) + H2O ↔ H+ + HCO3-

(3) HCO3- ↔ H+ + CO3--

(4) Ca++ + CO3-- ↔ CaCO3(s)

The proportions of products and to what extent each equilibrium proceeds to products or reactants is directly controlled by the kinetics of these reactions. Reaction 2 is first order with respect to CO2 with a first order rate constant (k) where k[size]1[/size] = 0.03 s, or a half-timie t1/2 = 1n 2/k1 = 23 s.The reaction of OH- + CO2 ↔ HCO3- is second order with resepect to [CO2] and [ OH-] :

(5) -d[CO2]/dt = k2[CO2][OH-]

where k2 = 850 M-1 s-1. This process is particuarly important at high valuese of pH. The dehydration of hydrocarbonate ( H2CO3↔ CO2 + H2O ) is first order with respect to [ H2CO3 ] where the rate constant is k-1= 20 s-1 and t1/2 = 0.03 s.The values for a combinatioin of the forward and backwards reations (6) can then be used for a determination of the equilibrium ratio, K (7):

(6) CO2 + H2O <--k1/k2---> H2CO3

(7) K = k1 / k-1 = 0.03 / 20 = 1/670

raj wrote: So Tom, what does that mean to us???

This would indicate that at equilibrium, based on the rates of reaction, that the concentration of CO2 ( [CO2] ) is about 670 times as high as the concentration of H2CO3, or that for the dissolved CO2, the majority that is not converted to carbonic acid remains as molecules of CO2 in the water column. Our rate constants indicate that not only is the forward reaction occuring slowly, but that due to this, equilibrium will 9be reached fairly slowly, as the reverse reaction rate is rapid. In many living organisams, the hydration of carbonic acid is rapid due to the presence of carbonic anhydrase (how RBC's move CO2), so corals, etc can then move dissolved CO2 as bicarbonate and convert it back easiliy thrugh the use of carbonic anhydrase. Having bicarbonate as the primary source of inorganic carbon available to corals reduces the work they need to do to acquire and assimilate inorganic carbon.

Rather than get into the deep math of buffer intensity, I'll refer folks to http://aslo.org/lo/toc/vol_20/issue_2/0222.pdf">http://aslo.org/lo/toc/vol_20/issue_2/0222.pdf</a> where ther3e is a good bit of relivant info on the topic.

Although buffer does have an effect on system pH, especially at the extremese of the pH range of sewater, it is often a rate issue associated with why we may see swings in the closed systems we keep, not to mentionu issues with degassing and accumulation of CO2 in home aquaria (and the indoor environmental rise of [CO2]. It is the [CO2} (the cncentration of dissolkved carbon dioxide), that more directly controls our pH, and this is going to swing within a daily7 range of 7.8 to 8.4 based on how long th lights are on, how much photosynthesis is going on, the population density o8f photoautotro8phs, and the length of the photoperiod. I am excluding issuese with poorly adjusted Ca++ reactors or controllers/probese that have not beem calibrated or cleaned or are too old to give accurate results.

HTH

* I will delve into this if there is interest:

Carbonic acid is diprotic, that is it has two hydrogens (protons) which dissociate from the parent molecule, and thus there are two rates of ionization and two dissociation constants:

H2CO3 ⇌ HCO3− + H+
Ka1 = 2.5×10−4; pKa1 = 3.60 at 25 °C, for -log (2.5×10−4) = 3.60. ****

HCO3− ⇌ CO32− + H+
Ka2 = 5.61×10−11; pKa2 = 10.25 at 25 °C.

Although we would be correct in quoting the first dissociation constant, care must be taken when quoting and using the first dissociation constant to describe what occurs as carbonic acid ionizes its first proton (not photon ~ :D ). The value given above is correct for the H2CO3 molecule, and shows that it is actually a stronger acid than many organic acids (examples: formic or acetic acid), as one might expect due to the electronegative pull of the oxygen moiety. This is not a faCTOR IN SEAWATER, however, as carbonic acid only exists in equilibrium with carbon dioxide at SST&P (see equation with constants above), and the concentration of H2CO3 there is much lower than the CO2 concentration (remember, ratio of 1:670). This renders the effect of the ionization of this acid to ineffectual levels in seawater at SST, reducing the measured acidity in seawater due to carbonic acid. The equation may be rewritten as follows based on its diprotic nature (both protons dissociating and combining the constants to form one value Ka):

CO2 + H2O ⇌ HCO3− + H+
Ka = 4.30×10−7; pKa = 6.36.

This figure is quoted as the dissociation constant of carbonic acid, although this is a little ambiguous: it might better be referred to as the acidity constant of carbon dioxide, and can be used in calculating the pH of CO2 solutions of seawater. See: [IMG]http://ajpcell.physiology.org/cgi/content/abstract/260/5/C1113">http://ajpcell.physiology.org/cgi/content/abstract/260/5/C1113</a> and
[IMG]http://www.springerlink.com/content/m7354331t2885092/">http://www.springerlink.com/content/m7354331t2885092/</a>

To determine pH and composition of a carbonic acid based on concentratioin in seawater at a given temperature, the composition of a pure carbonic acid solution (or of a pure CO2 solution) is completely determined by the partial pressure, (pCO2) of carbon dioxide above the solution. To calculate this composition, account must be taken of the above equilibria between the three different carbonate forms (H2CO3, HCO3− and CO32− in formulae 2, 3, and 4) as well as of the hydration equilibrium between dissolved CO2 and H2CO3 with constant (see above) and of the following equilibrium between the dissolved CO2 and the gaseous CO2 above the solution ( see Henry’s constant, ( [IMG]http://en.wikipedia.org/wiki/Partial_pressure">http://en.wikipedia.org/wiki/Partial_pressure</a> ):

(1) CO2(gas) ↔ CO2(dissolved) with where kH=29.76 atm/(mol/L) at 25°C

The corresponding equilibrium equations together with the relation and the neutrality condition result in six equations for the six unknowns [CO2], [H2CO3], [H+], [OH−], [HCO3−] and [CO32−], showing that the composition of the solution is fully determined by pCO2 . The equation obtained for [H+] is a cubic whose numerical solution yields the following values for the pH and the different species concentrations:

pCO2** . . pH . . . [CO2]***. . . [H2CO3]***. . . [HCO3−]*** . . . [CO32−]
10−8. . . . 7.00 . 3.36 × 10-10 5.71 × 10−13 . 1.42 × 10−9. . 7.90 × 10−13
10−6. . . . 6.81 . 3.36 × 10−8. 5.71 × 10−11 . 9.16 × 10−8. . 3.30 × 10−11
10−4. . . . 5.92 . 3.36 × 10−6. 5.71 × 10−9 . . 1.19 × 10−6. . 5.57 × 10−11
3.5×10−4 5.65 . 1.18 × 10−5. 2.00 × 10−8 . . 2.23 × 10−6. . 5.60× 10−11
10−3. . . . 5.42 . 3.36 × 10−5. 5.71 × 10−8 . . 3.78 × 10−6. . 5.61 × 10−11
10−2. . . . 4.92 . 3.36 × 10−4. 5.71 × 10−7 . . 1.19 × 10−5. . 5.61 × 10−11
10−1. . . . 4.42 . 3.36 × 10−3. 5.71 × 10−6 . . 3.78 × 10−5. . 5.61 × 10−11
1 . . . . . . 3.92 . 3.36 × 10−2. 5.71 × 10−5 . . 1.20 × 10−4. . 5.61 × 10−11
2.5. . . . . 3.72 . 8.40 × 10−2. 1.43 × 10−4 . . 1.89 × 10−4. . 5.61 × 10−11
10 . . . . . 3.42 . . . 0.336. . . . 5.71 × 10−4 . . 3.78 × 10−4. . 5.61 × 10−11

[I]info courtesy "Tracer studies with radioactive oxygen-15. Exchange between carbon dioxide and water," Michael J. Welch, Judith F. Lifton, Jane A. Seck. [B]The Journal of Physical Chemistry[/B], 1969 73 (10), 3351-3356. See [IMG]http://pubs.acs.org/doi/abs/10.1021/j100844a033">http://pubs.acs.org/doi/abs/10.1021/j100844a033</a> 

Caveats:<ul>
<li>We see that in the total range of pressure, the pH is always largely lower than pKa2 so that the CO32− concentration is always negligible with respect to HCO3− concentration. In fact CO32− play no quantitive role in the present calculation (see remark below).</li>
<li>For vanishing pCO2 (at 1x10-8, or 0.00000001), the pH is close to the one of pure water (pH = 7) and the dissolved carbon is essentially in the HCO3− form.</li>
<li>For normal atmospherics conditions ( pCO2 = 3.5 x 10-4 atm), we get a slightly acid solution (pH = 5.7) and the dissolved carbon is now essentially in the CO2 form (1/670). From this pressure on, [OH−] (the concentration of hydroxyl ion) becomes also negligible so that the ionized part of the solution is now an equimolar mixture of H+ and HCO3−.</li>
<li>For a CO2 pressure typical of the one in soda drinks bottles ( pCO2 =~ 2.5 atm), we get a relatively acidic water column in the absence of buffer (as in fresh water, pH = 3.7) with a high concentration of dissolved CO2. These features contribute to the sour and sparkling taste of sparkling carbonated beverages.</li>
<li>Between 2.5 and 10 atm, the pH crosses the pKa1 value (3.60) giving a dominant H2CO3 concentration (with respect to HCO3−) at high pressures.

Remark: As noted above, [CO32−] may be neglected for this specific problem, resulting in the following very precise analytical expression for [H+]:

[IMG]http://www.atlantareefclub.org/forums/attachment.php?attachmentid=15832&stc=1&d=1237660341" alt="" /></li>
</ul>

**values are in terms of ATM
***values are in terms of mols/L
**** see: [IMG]http://actachemscand.dk/pdf/acta_vol_01_p0204-0209.pdf">http://actachemscand.dk/pdf/acta_vol_01_p0204-0209.pdf</a>

This would be easier to post if the vBull for ARC had some documentation... ...anyone have a primer?
<fieldset class="gc-fieldset">
<legend> Attached files </legend> [IMG]http://atlantareefclub.org/boards/data/uploads/attachments/312465=15832-831ffd4f77d667186b1d24cf7131235e.png>

312465=15832-831ffd4f77d667186b1d24cf7131235e.png

class="gc-images" title="831ffd4f77d667186b1d24cf7131235e.png[/IMG] style="max-width:300px" /></a> </fieldset>

calcium reactor??

derek_s

Active Member

rostato

Active Member

derek_s

Active Member

glxtrix

Active Member

glxtrix

Active Member

derek_s

Active Member

rostato

Active Member

glxtrix

Active Member

derek_s

Active Member

glxtrix

Active Member

dannybradley

Active Member

sailfish

Member

glxtrix

Active Member

grouper therapy

Active Member

tdwyatt

Member

derek_s

Active Member

grouper therapy

Active Member

dannybradley

Active Member

tdwyatt

Member

Similar threads