Question

Limestone (CaCO3) can be broken down at high temperatures: When heated at 1.00x103 K for a certain amount of time, it can be completely transformed into two oxides: carbon dioxide and calcium oxide (known as “burnt lime”, figure why J). The solid is then completely dissolved in enough water to form a concentrated solution of Ca(OH)2. A 750-fold diluted solution of Ca(OH)2 required 50.0mL of a HCl 0.200M solution to completely neutralise its Ca(OH)2 content. a) How many moles of CO2 were produced in the first reaction? b) What is the Volume of CO2 that was collected (assume ideal gas behaviour at 1.00 x103 K and 1.10atm) c) What is the mass of the original limestone rock?

Answer 1

CaCO3 = CaO + CO2

CaO + H2O = Ca(OH)2

Ca(OH)2 + 2HCl = CaCl2 + H2O

mmol of acid used = MV = 0.2*50 = 10 mmol

then,

mmol of Ca(OH)2 neutralized = 1/2*10 = 5 mmol of Ca(OH)2

this is 750 fold neutralized so

total Ca(OH)2 was --> 750*5 = 3750 mmol or = 3.75 mol of Ca(OH)2

this was formed from

3.75 mol of Ca(OH)2 --> 3.75 mol of CaO

ratio is 1:1 so

if 3.75 mol of CaO were formed, then 3.75 mol of CO2 evolved

b)

find volume of CO2

PV = nRT

V = nrT/P

V = (3.75 )(0.082)(1000)/1.1

V = 279.545 L of CO2

c)

mol of a CaO = 3.75

mol of CaCO3 from CaO = 3.75 mol

mass of CaCO3 = mol*MW = 3.75 *100 = 375 g of CaCO3