Examining the relative molecular mass of washing soda powders which includes Na2CO3 by titration, use acid-base titration to examine how many waters of crystallization is found in washing soda powders.
Many molecules are hygroscopic, meaning that they are able to absorb water from their environment, including water vapor in the air. This means that hygroscopic molecules, unless completely dried and sealed in a moisture-free container, would be expected to include a certain amount of water. When performing quantitative chemistry with these compounds, the water must be included in our considerations because it can significantly affect the results; for example, by significantly increasing the molar mass of the compound.
We can't really answer how much water is normally found in washing soda powders without knowing the molecular composition of those washing powders. However, we can devise a general procedure for sodium carbonate (`Na_2 CO_3` ) and utilize titration to solve it.
When a molecule absorbs water, it is referred to as a hydrate; the waters of crystallization are the number of water molecules associated with each molecule of the primary compound. Since it is common for more than one water to be associated per primary molecule, a distinction needs to be made between the two; this comes in the form of an interpunct, which looks like a large dot. For example: `Na_2 CO_3 ` · `10 H_2 O` means that each molecule of sodium carbonate is accompanied by ten molecules of water. This is referred to as "hydrated" sodium carbonate, specifically sodium carbonate decahydrate (deca=10, hydrate=water). Decahydrate is the most common form of hydrated sodium carbonate, but the exact hydration is dependent upon the temperature and humidity of the environment, so it's difficult to say that washing powder would always contain sodium carbonate in its decahydrate form; we would need to test it in a given environment to confirm.
Since sodium carbonate acts as a base, we can utilize titration with a known concentration of acid to determine the moles of sodium carbonate present in a sample. For example, each carbonate ion would be capable of absorbing two hydrogen ions, forming carbonic acid, which would quickly decompose into one water and one carbon dioxide:
`2H^+ + CO_3 ^-2 -> H_2 CO_3 -> H_2 O + CO_2`
Adding an indicator, as is typical in titration, would allow us to determine when all of the carbonate has reacted. By determining how many moles of acid were required to do so, we can use stoichiometry to determine the moles of sodium carbonate that were present.
For example, if we used HCl, the reaction would be:
`2HCl + Na_2 CO_3 -> CO_2 + H_2 O + 2NaCl`
If we used 100mL of a 1M HCl solution to neutralize, we would have consumed 0.1 moles of HCl. Since every mole of sodium carbonate requires two moles of HCl to react, we cut the moles of HCl in half to find the moles of sodium carbonate in the sample: 0.05.
Assuming that the sodium carbonate was hydrated, we can then convert moles to mass and determine how much of the original mass of the hydrated sodium carbonate sample was composed of sodium carbonate, and assume that the remaining mass was water. This water mass is then converted to moles to determine the degree of water crystallization. In my example above, the 0.05 moles of sodium carbonate would result in a mass of about 5.3 grams. If the sodium carbonate was in decahydrate form, this means the total sample would have been about 14.3 grams.
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