**Iron-poor blood?**

Iron is an important component of red cells. Patients who have low iron will usually be anemic and have a lower than normal number of red blood cells. One way to assess serum iron concentration is with the use of Ferrozine, a complex organic molecule. Ferrozine forms a product with Fe^{3+}, producing a pink color. In order to determine factors affecting the reaction, we need to measure the equilibrium constant. If the equilibrium does not lie far in the direction of products, precautions need to be taken when using this material to measure iron in serum.

**Calculations with Equilibrium Constants**

The general value of the equilibrium constant gives us information about whether the reactants or the products are favored at equilibrium. Since the product concentrations are in the numerator of the equilibrium expression, a \begin{align*}K_{eq} > 1\end{align*}

Though it would often seem that the \begin{align*}K_{eq}\end{align*}

#### Sample Problem: Calculating an Equilibrium Constant

Equilibrium occurs when nitrogen monoxide gas reacts with oxygen gas to form nitrogen dioxide gas.

\begin{align*}2\text{NO}(g)+\text{O}_2(g) \rightleftarrows 2\text{NO}_2(g)\end{align*}

At equilibrium at 230°C, the concentrations are measured to be [NO] = 0.0542 M, [O_{2}] = 0.127 M, and [NO_{2}] = 15.5 M. Calculate the equilibrium constant at this temperature.

*Step 1: List the known values and plan the problem*.

Known

- [NO] = 0.0542 M
- [O
_{2}] = 0.127 M - [NO
_{2}] = 15.5 M

Unknown

- \begin{align*}K_{eq}\end{align*}
Keq value

The equilibrium expression is first written according to the general form in the text. The equilibrium values are substituted into the expression and the value calculated.

*Step 2: Solve*.

\begin{align*}K_{eq}=\frac{\left [\text{NO}_2 \right ]^2}{\left [\text{NO} \right ]^2 \left [\text{O}_2 \right ]}\end{align*}

Substituting in the concentrations at equilibrium:

\begin{align*}K_{eq}=\frac{\left ( 15.5\right )^2}{\left (0.0542 \right )^2 \left (0.127\right )}=6.44 \times 10^5\end{align*}

*Step 3: Think about your result*.

The equilibrium concentration of the product NO_{2} is significantly higher than the concentrations of the reactants NO and O_{2}. As a result, the \begin{align*}K_{eq}\end{align*}

The equilibrium expression only shows those substances whose concentrations are variable during the reaction. A pure solid or a pure liquid does not have a concentration that will vary during a reaction. Therefore, an equilibrium expression omits pure solids and liquids and only shows the concentrations of gases and aqueous solutions. The decomposition of mercury(II) oxide can be shown by the following equation, followed by its equilibrium expression.

\begin{align*}2\text{HgO}(s) \rightleftarrows 2\text{Hg}(l)+\text{O}_2(g) \qquad \quad K_{eq}=[\text{O}_2]\end{align*}

The stoichiometry of an equation can also be used in a calculation of an equilibrium constant. At 40°C, solid ammonium carbamate decomposes to ammonia and carbon dioxide gases.

\begin{align*}\text{NH}_4\text{CO}_2\text{NH}_2(s) \rightleftarrows 2\text{NH}_3(g)+\text{CO}_2(g)\end{align*}

At equilibrium, the [CO_{2}] is found to be 4.71 × 10^{-3} M. Can the \begin{align*}K_{eq}\end{align*}

\begin{align*}K_{eq}=[\text{NH}_3]^2 [\text{CO}_2]\end{align*}

The stoichiometry of the chemical equation indicates that as the ammonium carbamate decomposes, 2 mol of ammonia gas is produced for every 1 mol of carbon dioxide. Therefore, at equilibrium, the concentration of the ammonia will be twice the concentration of carbon dioxide. So [NH_{3}] = 2 × (4.71 × 10^{-3}) = 9.42 × 10^{-3} M. Substituting these values into the \begin{align*}K_{eq}\end{align*}

\begin{align*}K_{eq}=(9.42 \times 10^{-3})^2 (4.71 \times 10^{-3})=4.18 \times 10^{-7}\end{align*}

#### Using Equilibrium Constants

The equilibrium constants are known for a great many reactions. Hydrogen and bromine gases combine to form hydrogen bromide gas. At 730°C, the equation and \begin{align*}K_{eq}\end{align*}

\begin{align*}\text{H}_2(g)+\text{Br}_2(g) \rightleftarrows 2\text{HBr}(g) \qquad \ \ K_{eq}=2.18 \times 10^6\end{align*}

A certain reaction is begun with only HBr. When the reaction mixture reaches equilibrium at 730°C, the concentration of bromine gas is measured to be 0.00243 M. What is the concentration of the H_{2} and the HBr at equilibrium?

Since the reaction begins with only HBr and the mole ratio of H_{2} to Br_{2} is 1:1, the concentration of H_{2} at equilibrium is also 0.00243 M. The equilibrium expression can be rearranged to solve for the concentration of HBr at equilibrium.

\begin{align*}K_{eq} & = \frac{[\text{HBr}]^2}{[\text{H}_2][\text{Br}_2]}\\
\left [ \text{HBr} \right ] & = \sqrt{K_{eq}[\text{H}_2][\text{Br}_2]} \\
& = \sqrt{2.18 \times 10^6 (0.00243)(0.00243)} = 3.59 \text{ M}\end{align*}

Since the value of the equilibrium constant is very high, the concentration of HBr is much greater than that of H_{2} and Br_{2} at equilibrium.

### Summary

- Calculation of an equilibrium constant is described.

### Review

- What are the units for \begin{align*}K_{eq}\end{align*}
Keq ? - Why is the temperature specified in equilibrium problems?
- Why don’t we include solids or liquids in equilibrium calculations?