Combustion of a hydrocarbon in a limited supply of air produces CO in addition to CO2 and H2O. When 1.90 g of a particular hydrocarbon was burned in air, 1.83 g CO, 2.88 g CO2, and 2.95 g H2O were produced. What is the empirical formula of the hydrocarbon? empirical formula: Calculate the mass of oxygen used in the reaction. mass: Calculate the total mass of oxygen that would have been required for complete combustion. mass:

Consider the following balanced chemical reaction: Fe2O3 (s) + 2 Al (s) → Al2O3 (s) + 2 Fe (s)Suppose 9.98 g of Fe2O3 reacts with 4.02 g of Al. Calculate the total mass of reactant that remains after the reaction goes to completion

Manganese commonly occurs in nature as a mineral. The extraction of manganese from the carbonite mineral rhodochrosite, involves a two-step process. In the first step, manganese (II) carbonate and oxygen react to form manganese (IV) oxide and carbon dioxide: 2 MNCO;(s)+O,(g)→2 MnO,(s)+2 CO,(9) In the second step, manganese (IV) oxide and aluminum react to form manganese and aluminum oxide: 3 MnO2(s)+4 Al(s)→3 Mn(s)+2 Al,O3(s) Write the net chemical equation for the production of manganese from manganese (II) carbonate, oxygen and aluminum. Be sure your equation is balanced.

Methanol has tremendous potential to be used as an alternate fuel. Carbon monoxide gas reacted with hydrogen gas can be used to produce methanol. A student reacted 32.9 g of H2(9) with excess CO(9) to produce methanol. If his percent yield was 47.7%, calculate the mass (in g) of CH3OH(1) that he collected in the experiment. Hint: Write a balanced chemical reaction equation. Do not show your work in the space provided. Record only your final answer with the correct number of significant digits and the proper units.