For statements I and II, state for each if they are true or false. I) The reaction of potassium phosphate and calcium nitrate solutions goes to completion. II) When calcium phosphate is formed in solution, it precipitates as it is insoluble. If both are true, state if II is a correct explanation for I.
In statement I, we have two compounds reacting together in water, that is, in aqueous conditions. Potassium phosphate and calcium nitrate are both salts. Potassium phosphate is formed by neutralizing phosphoric acid with the appropriate base. And calcium nitrate is formed by neutralizing nitric acid with the correct base. When solutions of salts are mixed in aqueous conditions, we frequently see a double-displacement, also known as a double-decomposition, reaction taking place.
In the example shown here, A and C are cations, that is, positive ions. B and D are anions; those are negative ions. So in our reaction here, two new salts will be formed. These will be potassium nitrate and calcium phosphate. When presented as a symbol equation, we should include state symbols to indicate what state the reactants and products are present in under normal conditions. s indicates a solid, l indicates a liquid, g indicates a gas, and aq indicates aqueous, which means it’s dissolved in water.
Frequently, double-displacement reactions lead to the formation of precipitates. A precipitate is a solid product formed in a reaction. Precipitates do not dissolve well at all, or they have very limited solubility in water. We can use simple solubility rules to establish which of our salts are soluble or insoluble in this case. In general, salts of group one elements are all soluble. This also includes ammonium compounds as well. Also, salts containing nitrate ions are also soluble. So we can say that potassium phosphate, calcium nitrate, and potassium nitrate are all soluble salts with a high level of confidence.
The solubility of group two metal salts is much harder to predict. However, calcium carbonate, strontium carbonate, and barium carbonate are all insoluble. Many phosphates are also insoluble. Calcium phosphate and silver phosphate are examples. So in our equation, we have soluble potassium phosphate — a group one metal salt — reacting with soluble calcium nitrate — all nitrates are soluble. Soluble potassium nitrate is formed and insoluble calcium phosphate will appear as a precipitate. In fact, it will be a white precipitate, as most group one and group two metal salts are white in color.
We can now say with confidence that statement II is true. Calcium phosphate is formed in this reaction and it does form as a precipitate because it’s insoluble. In statement I, we see a suggestion that the reaction goes to completion. Reasons for a reaction not going to completion may be that the reaction has a high activation energy. That is, it’s hard to get started. The reaction is very slow, so, kinetically, it happens at a very slow rate. Or the reaction can be reversed and reaches equilibrium. Let us explore each option to see if it may apply to our reaction here.
The reactants here are present as hydrated positive and negative ions that are mobile within the bulk solvent water. Like-charged ions will repel each other. But it will not be long before a chance collision occurs between oppositely charged ions and electrostatic attraction takes place. Successful collisions are happening rapidly and products will, therefore, form reasonably quickly. The first two reasons listed here to not expect a complete reaction, therefore, seem unlikely. For a reaction to reach equilibrium, which is a dynamic situation, we would need the rate of formation of products to equal the rate of the reverse reaction.
Importantly, at this point, the concentrations of reactants and products would remain constant. We could express this as a Kc expression for a typical reaction involving A and B in equilibrium with C and D. It’s possible that this could happen in our reaction, if all the reactants and products were in the same phase. However, in this reaction, calcium phosphate is removed from the bulk solution rapidly, as it’s an insoluble precipitate. This would disturb the position of any possible equilibrium and will tend to drive the reaction forward according to Le Chatelier’s principle.
This is also a reason why reactions producing gases that escape from the mixture will never reach equilibrium in an open system. For this reason, the reaction that we are studying here cannot be reversed. It is very likely that the reaction will go to completion. And statement I is therefore true. As we have discussed earlier, calcium phosphate is indeed an insoluble salt. We can explore why by looking at the equation for the Gibbs free energy for this precipitation process. ΔG or the Gibbs free energy for this precipitation process must be negative if this is to be a feasible or favorable reaction.
ΔG for this precipitation process is the combination of the enthalpy change, that is, the energy change for the precipitation reaction, and the entropy change for the precipitation process, which are related to the change in disorder. This is frequently negative for our precipitation reaction, where aqueous ions are going to a solid. With calcium phosphate effectively being removed from the bulk solution rapidly, the reaction never has the chance to reach equilibrium. And it does go to completion. Statement II is part of this explanation and, therefore, is a correct explanation for statement I.