Question

Every year people die from carbon monoxide poisoning because they bring a charcoal fire into their tent, house, or enclosed area. When carbon burns, it can produce either carbon dioxide, CO2(g), or carbon monoxide gas, CO(g).

a) If you start with a system containing 1 pound of charcoal briquettes (assuming pure carbon), how many grams of oxygen is needed to turn all of the carbon into the safe carbon dioxide?

b) What volume of air is required (in L)?

c) How much heat is given off as a result of the combustion to CO2(g)

Answer 1

(a) The amount of O2 needed is 2.67 pounds.

(b) The volume required is 847,509 litres.

(c) The heat given off as a result of the combustion to CO2 is 393.5 kJ.

Step-by-step explanation:

For a complete combustion of C to C02(g)

`C+O_(2) \rightarrow  CO_(2)`

(a) The molecular mass of O2 is 32 g/mol and the molecular mass of C is 12 g/mol.

We need 1 mol O2 to burn 1 mol of C.

If we need 32 g of O2 to burn 12 g of C, to burn 1 pound of pure carbon charcoal we need (32/12)*1=2.67 pounds of O2.

(b) The density of O2, at atmospheric conditions, is 1.429 g/l. The volume of 2.67 pounds of O2 is

`V=M/\rho=(2.67  lb)*((1 l)/(1.429 g))*((453,592 g)/(1lb)) =(1,211,090)/(1.429)= 847,509 l`

(c) To calculate the heat of the reaction, we have to look up in the Table of Standard Enthalpy of Formation Values and compute the following equation

`\Delta H_(reaction)^(o)=\sum\Delta H_(f)^(o) (products)-\sum\Delta H_(f)^(o)(reactants)\\\\\Delta H_(reaction)^(o)=\Delta H_{f CO_(2)}^(o)-(\Delta H_{fO_(2)}^(o)+\Delta H_(fC)^(o))\\\\\Delta H_(reaction)^(o)=(1 mol) *(-393.5 kJ/mol)-((1mol)*(0 kJ/mol)+(1mol)*(0 kJ/mol))\\\\\Delta H_(reaction)^(o)=-393.5 kJ`

The heat given off as a result of the combustion to CO2 is 393.5 kJ.