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Vinegar, the commercial name for acetic acid, HC2Hs02, is a monoprotic organic acid. A 5% (w/v) solution of vinegar is used to titrate a sample of an antacid that contains CaCO3 is the only basic component. If one antacid tablet contains 800 mg of CaCOs in it, then calculate how many milliliters of vinegar should be required to completely neutralize the CaCO3 present in one tablet of antacid?

Respuesta :

cristisp93

19.2 g of vinegar solution

Explanation:

Here we have the chemical reaction between acetic acid (CH₃COOH) and calcium carbonate (CaCO₃):

2 CH₃COOH +  CaCO₃ → (CH₃COO)₂Ca + CO₂ + H₂O

number of moles = mass / molecular weight

number of moles of CaCO₃ = 0.8 / 100 = 0.008 moles

Knowing the chemical reaction, we devise the following reasoning:

if       2 moles of CH₃COOH react with 1 moles of CaCO₃

then X moles of CH₃COOH react with 0.008 moles of CaCO₃

X = (2 × 0.008) / 1 = 0.016 moles of CH₃COOH

mass = number of moles × molecular weight

mass of acetic acid (CH₃COOH) = 0.016 × 60 = 0.96 g

Now to find the volume of vinegar acid (solution of acetic acid) with a concentration of 5% (weight/volume) we use the following reasoning:

if there are         5 g of acetic acid in 100 mL of vinegar solution

then there are   0.96 g of acetic acid in Y mL of vinegar solution

Y = (0.96 × 100) / 5 = 19.2 g of vinegar solution

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weight/volume concentration

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