This is likely a case of size matching between the silver and iodide ions Explanation: We interrogate the equilibrium Truong-Son N. Atomic and ionic radii are listed below. The extent of ionic character should be examined, to determine the validity of using covalent radii or ionic radii. Electronegativities are: "Ag": 1.
There are few exceptions to this rule. The difference in electronegativity between silver and fluorine is much greater than that of silver and iodine. Because of this large electronegative dissimilarity, the bond is considered polar-ionic.
But what about NaI and AgF? The two cases are hard-soft and soft-hard interactions respectively. Pearson did not say anything about this in his HSAB theory. However, the theory can be extended to include these species.
If the cation is very hard, the anion has no power to influence the cation to form a covalent bond in most cases, due to its large size exception - all nitrates including silver nitrate is soluble in water! But the reason is different. However, if the electrohpile is borderline hard, or soft , while the anion is extremely hard, things can get reversed.
Consequently, AgF is soluble in water. The equation for the equilibrium between solid silver chloride, silver ion, and chloride ion is:. The solubility product is 1. AgCl will precipitate if the reaction quotient calculated from the concentrations in the mixture of AgNO 3 and NaCl is greater than K sp.
The volume doubles when we mix equal volumes of AgNO 3 and NaCl solutions, so each concentration is reduced to half its initial value. The reaction quotient, Q , is momentarily greater than K sp for AgCl, so a supersaturated solution is formed:. Since supersaturated solutions are unstable, AgCl will precipitate from the mixture until the solution returns to equilibrium, with Q equal to K sp.
Will KClO 4 precipitate when 20 mL of a 0. Remember to calculate the new concentration of each ion after mixing the solutions before plugging into the reaction quotient expression. Thus, if we know the concentration of one of the ions of a slightly soluble ionic solid and the value for the solubility product of the solid, then we can calculate the concentration that the other ion must exceed for precipitation to begin.
To simplify the calculation, we will assume that precipitation begins when the reaction quotient becomes equal to the solubility product constant. Blood will not clot if calcium ions are removed from its plasma. For this reaction Table E3 :. CaC 2 O 4 does not appear in this expression because it is a solid.
Water does not appear because it is the solvent. If a solution contains 0. Neglect any increase in volume upon adding the solid silver nitrate.
It is sometimes useful to know the concentration of an ion that remains in solution after precipitation. We can use the solubility product for this calculation too: If we know the value of K sp and the concentration of one ion in solution, we can calculate the concentration of the second ion remaining in solution. However, the concentrations are different; we are calculating concentrations after precipitation is complete, rather than at the start of precipitation. From that, we calculate the pH.
At equilibrium:. If the person doing laundry adds a base, such as the sodium silicate Na 4 SiO 4 in some detergents, to the wash water until the pH is raised to The first step in the preparation of magnesium metal is the precipitation of Mg OH 2 from sea water by the addition of Ca OH 2.
Due to their light sensitivity, mixtures of silver halides are used in fiber optics for medical lasers, in photochromic eyeglass lenses glass lenses that automatically darken when exposed to sunlight , and—before the advent of digital photography—in photographic film. When two anions form slightly soluble compounds with the same cation, or when two cations form slightly soluble compounds with the same anion, the less soluble compound usually, the compound with the smaller K sp generally precipitates first when we add a precipitating agent to a solution containing both anions or both cations.
When the K sp values of the two compounds differ by two orders of magnitude or more e. This is an example of selective precipitation , where a reagent is added to a solution of dissolved ions causing one of the ions to precipitate out before the rest. Specifically, selective precipitation is used to remove contaminants from wastewater before it is released back into natural bodies of water.
An abundance of phosphate causes excess algae to grow, which impacts the amount of oxygen available for marine life as well as making water unsuitable for human consumption. One common way to remove phosphates from water is by the addition of calcium hydroxide, known as lime, Ca OH 2. The lime is converted into calcium carbonate, a strong base, in the water.
As the water is made more basic, the calcium ions react with phosphate ions to produce hydroxylapatite, Ca 5 PO4 3 OH , which then precipitates out of the solution:.
The precipitate is then removed by filtration and the water is brought back to a neutral pH by the addition of CO 2 in a recarbonation process.
0コメント