Question

Write net ionic equations to explain the formation of (a) a red precipitate when solutions of iron(III) nitrate and sodium hydroxide are mixed. (b) two different precipitates when magnesium sulfate and barium hydroxide are mixed.

Short answer

In conclusion, when we mix (a) iron(III) nitrate and sodium hydroxide, a red precipitate of iron(III) hydroxide (Fe(OH)3) forms. When we mix (b) magnesium sulfate and barium hydroxide, two white precipitates form: magnesium hydroxide (Mg(OH)2) and barium sulfate (BaSO4).

Step-by-step solution

(a) Balanced Molecular Equation for Iron(III) Nitrate and Sodium Hydroxide Reaction:

The first step is to write the balanced molecular equation for the reaction between iron(III) nitrate and sodium hydroxide. Iron(III) nitrate has the formula Fe(NO3)3, and sodium hydroxide has the formula NaOH. The products of the reaction are likely to be an ionic compound containing Fe(III) and OH-, in this case, iron(III) hydroxide (Fe(OH)3), and a salt containing Na+ and NO3-, which is sodium nitrate (NaNO3). The balanced molecular equation is: Fe(NO3)3 (aq) + 3 NaOH (aq) -> Fe(OH)3 (s) + 3 NaNO3 (aq)

(a) Total Ionic Equation:

Now, let's write the total ionic equation by breaking the strong electrolytes into their constituent ions: Fe^{3+} (aq) + 3 NO3^- (aq) + 3 Na^+ (aq) + 3 OH^- (aq) -> Fe(OH)3 (s) + 3 Na^+ (aq) + 3 NO3^- (aq)

(a) Net Ionic Equation:

Next, we'll write the net ionic equation by eliminating spectator ions, which are the ions that do not participate in the reaction. In this case, Na+ and NO3- are spectator ions: Fe^{3+} (aq) + 3 OH^- (aq) -> Fe(OH)3 (s)

(a) Red Precipitate Identification:

Finally, we can identify the precipitate that forms during the reaction as iron(III) hydroxide (Fe(OH)3), which appears as a red solid: Red precipitate: Fe(OH)3

(b) Balanced Molecular Equation for Magnesium Sulfate and Barium Hydroxide Reaction:

For the reaction between magnesium sulfate (MgSO4) and barium hydroxide (Ba(OH)2), the balanced molecular equation is: MgSO4 (aq) + Ba(OH)2 (aq) -> Mg(OH)2 (s) + BaSO4 (s)

(b) Total Ionic Equation:

Again, we'll write the total ionic equation by breaking the strong electrolytes into their constituent ions: Mg^{2+} (aq) + SO4^{2-} (aq) + Ba^{2+} (aq) + 2 OH^- (aq) -> Mg(OH)2 (s) + BaSO4 (s)

(b) Net Ionic Equation:

By eliminating the spectator ions, which in this case are not present, we can directly write the net ionic equation: Mg^{2+} (aq) + SO4^{2-} (aq) + Ba^{2+} (aq) + 2 OH^- (aq) -> Mg(OH)2 (s) + BaSO4 (s)

(b) Precipitates Identification:

Finally, two precipitates are formed in this reaction: magnesium hydroxide (Mg(OH)2), which is white, and barium sulfate (BaSO4), which is also white: First precipitate: Mg(OH)2 Second precipitate: BaSO4

Key concepts

Chemical Precipitation

Chemical precipitation is a process where a solid forms from a solution and falls out of the solution as a precipitate. This process often occurs in reactions involving ionic compounds when two solutions are mixed and an insoluble product, the precipitate, is formed.

For example, when solutions of iron(III) nitrate and sodium hydroxide are mixed, a red precipitate of iron(III) hydroxide forms. Precipitation reactions are crucial in various fields such as water treatment, where they are used to remove harmful ions from water, or in biomedical applications for the separation of biochemical compounds.

Balancing Chemical Equations

Balancing chemical equations is vital in accurately representing what occurs in a chemical reaction. Each atom's mass must be conserved, which means the number of atoms of each element on the reactants side must equal the number on the products side.

Our example problem shows how to balance the equation for the reaction between iron(III) nitrate and sodium hydroxide, resulting in the formation of iron(III) hydroxide and sodium nitrate. Balancing ensures that the equation adheres to the Law of Conservation of Mass. Properly balanced equations are essential not just for academic purposes but also for calculating how much of each reactant is needed in industrial processes.

Spectator Ions

Spectator ions are ions in an ionic equation that do not participate in the chemical reaction and are present in identical forms on both the reactant and product sides of the equation. These ions don't undergo any chemical changes during the reaction.

In our exercise, the Na⁺ and NO₃⁻ ions are spectator ions in the reaction of iron(III) nitrate with sodium hydroxide. They appear in the total ionic equation but are not present in the net ionic equation because they do not contribute to the formation of the precipitate. This distinction is instrumental in simplifying complex reactions to their essence and in understanding the actual chemical change occurring.

Solubility Rules

Understanding solubility rules is crucial for predicting the outcome of reactions in aqueous solutions, especially those leading to precipitation. These rules provide guidance on whether an ionic compound is soluble or insoluble in water.

For instance, compounds containing alkali metal ions and ammonium (NH₄⁺) are generally soluble, while most sulfates (SO₄²⁻) are soluble with exceptions like barium sulfate. Our exercise mentioned the mixing of magnesium sulfate and barium hydroxide, which leads to the formation of two different precipitates: magnesium hydroxide and barium sulfate. Knowledge of solubility rules allows for predictions that barium sulfate is insoluble and thus will precipitate, an important consideration in chemical analyses and industries where separation of substances is needed.