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# Calculating Heat of Reaction from Heat of Formation

## An application of Hess's Law

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Practice Calculating Heat of Reaction from Heat of Formation

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Hot Enough to Melt Steel

### Hot Enough to Melt Steel

Credit: Robert Lopez, CK-12
Source: https://www.dropbox.com/s/1v622e3diamagtn/fire_wood.jpg

Burning wood is a useful way to generate heat. In the image above, the fire is safely kept within its container, because the flame is not hot enough to melt the metal. However, under some circumstances, we might want to generate molten (liquid) metals. One surprisingly simple way to create this much heat is by igniting a substance known as thermite, which is a powdered mixture of aluminum and iron (III) oxide (Fe2O3, the primary component of rust).

#### Amazing But True!

• One of the products of the reaction is neutral iron metal. Because so much heat is released so quickly, the iron is heated above its melting point of 1,538°C.
• Part of what makes this reaction so violent is that the materials are in a powdered form. Powdered materials have a very high surface area, and this solid-state reaction can only occur when the surfaces of the two reactants are in contact with one another. As a result, heat is generated much faster than it can be dissipated to the surroundings, creating very high temperatures.
• Combustion reactions like the burning of a log require oxygen as one of the reactants, so they can be stopped by smothering the resulting fire, thus preventing any more oxygen from reaching the fuel. The thermite reaction does not make use of molecular oxygen, so it cannot be stopped in this way.

#### Explore More

With the links below, learn more about the reactivity of different metals and the amount of heat absorbed or released during a given chemical reaction. Then answer the following questions.

1. The balanced equation for the thermite reaction (after the molten iron is allowed to cool and solidify) is shown below:2 Al(s)+Fe2O3(s)Al2O3(s)+2 Fe(s)\begin{align*}2 \ Al(s)+Fe_2 O_3 (s) \rightarrow Al_2 O_3 (s)+2 \ Fe(s)\end{align*}Using the standard enthalpy of formation values from the link above, calculate the amount of heat released when one mole of Fe2O3 is reacted with excess aluminum.
2. What if copper were used instead of aluminum? Would as much energy be produced? The balanced reaction in this case would be the following:
3. 3 Cu(s)+Fe2O3(s)3 CuO(s)+2 Fe(s)\begin{align*}3 \ Cu(s)+Fe_2 O_3 (s) \rightarrow 3 \ CuO(s)+2 \ Fe(s)\end{align*}The two reactions above are examples of single-replacement reactions. Another common type of single-replacement reaction involves a neutral metal reacting with an aqueous metal salt. Suppose the solid iron (III) oxide were replaced with aqueous iron (III) chloride. Use the activity series of metals to determine whether either the following two reactions would proceed spontaneously. Then relate this result to the answers for the first two questions.
4. Al(s)+FeCl3(aq)3 Cu(s)+2 FeCl3(aq)AlCl3(aq)+Fe(s)3 CuCl2(aq)+2 Fe(s)\begin{align*}Al(s)+FeCl_3 (aq) & \rightarrow AlCl_3 (aq)+Fe(s) \\ 3 \ Cu(s)+2 \ FeCl_3 (aq) & \rightarrow 3 \ CuCl_2 (aq)+2 \ Fe(s)\end{align*}Could you create molten iron by wrapping a rusty nail with aluminum foil and heating it with a lit match? Why or why not?

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