What Is Really in the Baking Soda Jar?
A jar labelled 'baking soda' has stood open on a damp laboratory shelf for months, and the teacher suspects that part of it has changed to washing soda. A student weighs exactly 5.00 g of the dry powder into a pre-weighed crucible, heats it strongly, cools it in a desiccator and reweighs: the residue is 3.88 g. The teacher asks for the heating to be repeated until two successive weighings agree, but the student, in a hurry, stops as soon as the powder stops crackling and records a residue of 4.15 g. One classmate argues that the mass lost is only moisture absorbed from the air; another says it must be carbon dioxide from the baking soda. The teacher points out that the powder was kept in a stoppered bottle during weighing, that the crucible had been heated and cooled before use, and that the desiccator was freshly charged. The student must decide what the powder contains, how reliable each result is, and how to prove what actually escaped from the crucible.
1. From the 3.88 g residue, calculate the percentage of sodium hydrogencarbonate and of sodium carbonate in the sample, stating your assumptions.
Read the worked answer
On strong heating only the hydrogencarbonate decomposes:
2NaHCO3 -> Na2CO3 + H2O + CO2
So 2 mol (168 g) of NaHCO3 lose 18 + 44 = 62 g and leave 106 g of Na2CO3. Any Na2CO3 already present is unchanged and simply stays in the crucible.
Mass lost = 5.00 - 3.88 = 1.12 g. This loss comes only from NaHCO3:
mass of NaHCO3 = 1.12 x 168/62 = 3.03 g, i.e. 3.03/5.00 x 100 = 60.7%.
Original Na2CO3 = 5.00 - 3.03 = 1.97 g, i.e. 39.3%.
Check: Na2CO3 formed from the hydrogencarbonate = 3.03 x 106/168 = 1.91 g; total residue = 1.97 + 1.91 = 3.88 g, which matches.
Assumptions: the sample was dry and contained nothing else volatile; no NaHCO3 was lost by spitting; decomposition went to completion; the residue is pure Na2CO3. The 60:40 split is consistent with a jar left open, since NaHCO3 slowly converts to Na2CO3 with uptake of atmospheric moisture and CO2 loss.
2. Repeat the calculation using the hurried 4.15 g residue. In which direction is the error, and what exactly should the student have done?
Read the worked answer
Mass lost = 5.00 - 4.15 = 0.85 g. Then mass of NaHCO3 = 0.85 x 168/62 = 2.30 g, i.e. 46.1%, against the correct 60.7%. The result is an underestimate by about 15 percentage points.
The reason: decomposition was incomplete. Some NaHCO3 remained undecomposed in the residue and was counted as if it were Na2CO3, so the measured loss was only part of the water and carbon dioxide that should have escaped. The residue is therefore too heavy and the calculated hydrogencarbonate content too low.
The fix is heating to constant mass: heat, cool in the desiccator, weigh, and repeat until two successive weighings differ by no more than about 0.01 g. The crucible should also be heated gently at first to avoid spitting, and covered with a lid or watch glass during the early stage so no solid is lost. 'Stops crackling' is not an endpoint; it only means the visible water has gone.
3. How would you prove experimentally that both water vapour and carbon dioxide are evolved? Address the objection that the loss is merely moisture from the air.
Read the worked answer
Carbon dioxide: arrange the mouth of the heated test tube so the gas bubbles through freshly prepared limewater. Limewater turns milky because Ca(OH)2 + CO2 -> CaCO3 + H2O, and the milkiness clears again if excess CO2 is passed, because CaCO3 + H2O + CO2 -> Ca(HCO3)2. This is specific for CO2 among common gases from such a salt.
Water vapour: hold a cool, dry watch glass or dry test tube in the stream of gas; colourless droplets condense on it. A more definite test is to pass the vapour over anhydrous copper sulphate, which turns from white to blue, or to use cobalt chloride paper, which turns from blue to pink. Both tests need the apparatus to be dry beforehand.
The 'moisture from the air' objection fails on the teacher's evidence: the powder was stoppered during weighing, the crucible was pre-dried, and the residue was cooled in a freshly charged desiccator, so surface water had little chance to enter. More decisively, the loss is 22% of the sample mass, far more than a dry crystalline powder could carry, and an inert substance such as sodium chloride would not lose mass at all on strong heating.