A 5.60g sample of an unknown compound containing only C, H, and O combusts in an oxygen-rich environment... (2 parts to this question)

Question

A 5.60g sample of an unknown compound containing only C, H, and O combusts in an oxygen-rich environment. When the products have cooled to 20.0 degrees Celsius at 1 bar, there are 7.05 L of CO2 and 5.22 mL of H2O. The density of water at 20.0 degrees Celsius is 0.998 g/mL.

What is the empirical formula of the unknown compound?

If the molar mass is 116.2 g/mol, what is the molecular formula of the compound?

What is the

Answer 1

1) Find moles or grams of CO₂ produced (IGL) and the moles or grams of H₂O produced (g using density)

2) Since all the moles/grams of C and H come from the two products, you can use the formulas to get moles and grams for C and H.

3) You need both grams and moles of C and H because

a) 5.6 g of sample - g C - g H = g O

b) In order to find the empirical formula: Find moles C:moles H: moles O and divide through by the smallest number of moles. If the mole ratios are integral, you have your subscripts. If they are obvious decimal fractions, multiply through until you have whole numbers (1:3:2.5 --> 2:6:5)

4) The subscript multiplier is obtained using M_MF / M_EF = n (It is how many times the empirical formula goes into the molecular formula - If you have kept a lot of sig figs throughout, this should be very close to an integer) If the empirical formula was C_xH_yO_z , then the MF is C_nxH_nyO_nz

Good Luck!