# Question #b8a9f

Apr 24, 2016

${\text{C"_ ((s)) + "O"_ (2(g)) -> "CO}}_{2 \left(g\right)}$

#### Explanation:

Carbon dioxide, ${\text{CO}}_{2}$, has two constituent elements

• carbon, $\text{C}$
• oxygen, $\text{O}$

Your job here will be write a balanced chemical equation that will describe the synthesis reaction that leads to the formation of carbon dioxide.

Now, strictly speaking, the standard state is defined exclusively in terms of pressure. More specifically, the standard state of solids, liquids, or gases makes use of a standard pressure of ${10}^{5} \text{Pa}$.

However, temperature is often included when referring to gases. More often than not, you'll see the standard state for a gas being defined as a working pressure of ${10}^{5} \text{Pa}$ and a temperature of ${25}^{\circ} \text{C}$.

In your case, carbon is a solid in its standard state, i.e. at a pressure of ${10}^{5} \text{Pa}$.

On the other hand, oxygen is a gas in its standard state. Moreover, it's crucial to remember that oxygen is one of the seven diatomic elements.

This means that you represent oxygen in its standard state as a diatomic molecule, ${\text{O}}_{2}$, not as an element, $\text{O}$.

Therefore, the equation that describes the synthesis of carbon dioxide from its constituent elements in their standard state will look like this

$\textcolor{g r e e n}{| \overline{\underline{\textcolor{w h i t e}{\frac{a}{a}} \textcolor{b l a c k}{{\text{C"_ ((s)) + "O"_ (2(g)) -> "CO}}_{2 \left(g\right)}} \textcolor{w h i t e}{\frac{a}{a}} |}}}$

As a side note, it's worth noting that the enthalpy change that accompanies this reaction is actually carbon dioxide's standard enthalpy of formation, $\Delta {H}_{f}^{\circ}$.

A substance's standard enthalpy of formation tells you the enthalpy change that occurs when one mole of that substance is formed from its constituent elements in their standard state.