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The infobox gives the solubility data for various hydrates as
Decahydrate: 7 g/100 mL (0 °C) 16.4 g/100 mL (15 °C) 34.07 g/100 mL (27.8 °C) Heptahydrate: 48.69 g/100 mL (34.8 °C) Monohydrate: 50.31 g/100 mL (29.9 °C) 48.1 g/100 mL (41.9 °C) 45.62 g/100 mL (60 °C) 43.6 g/100 mL (100 °C)[1]
The labels seem to say that the values are grams of various hydrates that will dissolve into 100 mL of water. However, comparing with the solubility chart in this Solvay US Inc document (after converting from mass percentage to g/dL) and another couple of sources, it seems that all the values are grams of anhydrousNa 2CO 3. What those labels actually say is which hydrate would crystallize at those temperatures. Moreover, the point "50.31 (29.9 °C)" seems to have the correct amount (the maximum solubility) but this page gives the temperature as 35.37 °C, quoting a 1907 paper.[2] Also, the values for temperatures above that maximum seem to be slightly lower than what can be read from the Solvay chart. Here is the table as I read from that chart:
However, I read out these values visually from Solvay's chart and did the conversions myself; and I cannot now check the most authoritative sources. Therefore, I will only fix the labels and remove the bad point for now, and wait in case someone has more reliable data.--Jorge Stolfi (talk) 21:35, 17 June 2018 (UTC)[reply]
The paragraph ends weirdly: "For the complete sequence of tests used for qualitative cation analysis." It seems like they were going to mention or link to something. -- PaulxSA (talk) 12:49, 13 July 2018 (UTC)[reply]
The following was removed today. Reinsert it if you think I erred. Just seems pretty niche.--Smokefoot (talk) 21:32, 23 December 2018 (UTC)
Sodium carbonate, in a solution with common salt, may be used for cleaning silver. In a nonreactive container (glass, plastic, or ceramic), aluminium foil and the silver object are immersed in the hot salt solution. The elevated pH dissolves the aluminium oxide layer on the foil and enables an electrolytic cell to be established. Hydrogen ions produced by this reaction reduce the sulfide ions on the silver restoring silver metal. The sulfide can be released as small amounts of hydrogen sulfide. Rinsing and gently polishing the silver restores a highly polished condition.[1][reply]
In taxidermy, sodium carbonate added to boiling water will remove flesh from the bones of animal carcasses for trophy mounting or educational display.
In chemistry, it is often used as an electrolyte. Electrolytes are usually salt-based, and sodium carbonate acts as a very good conductor in the process of electrolysis. In addition, unlike chloride ions, which form chlorine gas, carbonate ions are not corrosive to the anodes. It is also used as a primary standard for acid-base titrations because it is solid and air-stable, making it easy to weigh accurately.
Soda ash is used as a water softener in laundering: it competes with the magnesium and calcium ions in hard water and prevents them from bonding with the detergent being used, but does not prevent scaling.[1] Sodium carbonate can be used to remove grease, oil, and wine stains.
Soda lakes are deadlands.
Feyalite sands is common orange color sands available by density separation occurring in ferrosol etc.
then just add quartz sands powder.
You can also make Feyalite from ferrosol and quartz.
of course chemical Quartz powder and FeSiO3 is the best Wikistallion (talk) 06:38, 23 July 2019 (UTC)[reply]
I am baffled as to how to understand the section "Inexpensive, weak base" because it begins "Sodium carbonate is also used as a relatively strong base in various fields." Can anyone help to explain how to reconcile the seeming contradiction of weak and strong? Czrisher (talk) 13:19, 16 August 2019 (UTC)[reply]
Hi @Czrisher: In chemistry, an officially "weak" or "strong" base or acid is an indication of how much the substance dissociates into ions in solution. For example, sodium hydroxide (NaOH), also known as lye, is a "strong" base because, when dissolved into water, the NaOH completely splits into Na+ and OH- ions that float around separately in the water. There is no actual NaOH as a whole anywhere the solution, just the ions the NaOH has split into. On the other hand, carbonic acid (H2CO3) is a "weak" acid because, when dissolved in water, only a fraction of the carbonic acid splits into H+ and CO32- ions, while the rest stays as undivided H2CO3 that floats around in the water. Because sodium carbonate is the result of these two chemicals reacting with each other, and because sodium hydroxide is "strong" and carbonic acid is "weak", sodium carbonate behaves like a "weak" base in solution, but in comparison to other "weak" bases, the ratio of split ions to whole sodium carbonate is much greater which increases its ability to carry out reactions, but it's not a "strong" base because there is still sodium carbonate floating around in the solution. You might say it's a "relatively strong" base, but not an actual "strong" base. I know that might be a little confusing, so if you have any further questions, feel free to ask and I'll try to help. You may also find this article useful. Matt18224 (talk) 16:34, 18 August 2019 (UTC)[reply]
While researching washing soda, in particular its cleaning mechanism, I only found saponification on this Wikipedia page, with other sources explaining the effect by washing soda binding with mineral cations (Mg or Ca) and thus improving other detergents' effect.
Citation is needed for the saponification claim. BratwurstBaron (talk) 20:35, 14 January 2023 (UTC)[reply]
In the washing soda section there is this sentence, "It is one of the few metal carbonates that is soluble in water." Isn't this something that that is true of sodium carbonate in general, and not specific to washing soda? Can someone confirm? Ike9898 (talk) 20:20, 24 January 2025 (UTC)[reply]
level-4 vital article is rated B-class on Wikipedia's content assessment scale.
