Silver chloride solid solute

The space charge potential at the silver chloride-aqueous solution interface can be derived from thermodynamic arguments following the concept of Grimley and Mott with the assumptions that the phase boundary potential X remains constant and that there are no adsorbed surface charges. For the silver chloride-aqueous solution system the silver ion is the most mobile one in the solid and is therefore the potential-determining ion. Let and be the electrochemical potentials of the silver ion in the solution and in the crystal, respectively. Assuming ideal solutions. 

Since our calculations agree well with the experimental data, we may conclude that the assumptions made previously are reasonable and that a space charge distribution is operative in the solid side of the silver chloride-aqueous solution interface. This space charge distribution results from a negative potential Vc(0) in the solid and corresponds to a deficiency in silver ion vacancies, the maximum occurring near the equivalence point. This conclusion will be further supported in the following section on AC conductance and capacitance measurements. 

Returning now to silver chloride, let us apply these ideas to its saturated aqueous solution at 25°. From the value given in Table 42, we see that in solid AgCl the entropy per ion pair is almost exactly 1 milli-electron-volt per degree, which is equivalent to 23.0 cal/deg/mole. It makes no difference whether we express the entropies per ion pair in electron-volts per degree or in the equivalent calories per degree per mole. In the electrochemical literature the calorie per degree per mole is used and is called one entropy unit. (This is abbreviated e.u.) ... 

Equilibrium curve for silver chloride. Silver chloride (s) is in contact with Ag+ and Cl- ions in aqueous solution. The product Q of the concentration of ions [Ag+] X [Cl-] is equal to Ksp (curved line)when equilibrium exists. If 0 > K,p, AgCI(s) tends to precipitate out until equilibrium is reached. If 0 < Ksp, additional solid dissolves. 

Silver chloride is a solid that shows this effect. This solid does not dissolve readily in water. When solid silver chloride is placed in water, very little solid enters the solution and there is only a very slight increase in the conductivity of the solution. Yet there is a real and measurable increase—ions are formed. Careful measurements show that even though silver chloride is much less soluble in water than sodium chloride, it is like sodium chloride in that all the solid that does dissolve forms aqueous ions. The reaction is... 

Though both silver nitrate and sodium chloride have high solubility in water, silver chloride is very slightly soluble. What will happen if we mix a solution of silver nitrate and sodium chloride Then, we will have a solution that includes the species present in a solution of silver chloride, Ag+(aq) and Cl (ag), but now they are present at high concentration The Ag+(agJ came from reaction (8) and the Cl (aq) came from reaction (6) and their concentrations far exceed the solubility of silver chloride. The result is that solid will be formed. The formation of solid from a solution is called precipitation ... 

Some ionic compounds are soluble, others are not. Consider what happens when we pour a solution of sodium chloride (a strong electrolyte) into a solution of silver nitrate (another strong electrolyte). A solution of sodium chloride contains Na+ cations and Cl anions. Similarly, a solution of silver nitrate, AgNO, contains Ag+ cations and NO, anions. When we mix these two aqueous solutions, a white precipitate, a cloudy, finely divided solid deposit, forms immediately. Analysis shows that the precipitate is silver chloride, AgCl, an insoluble white solid. The... 

At this point, any Hg22+ has precipitated and any Ag+ present is in solution. The solution is separated from the solid, and the presence of silver ion in the solution is verified by addition of nitric acid. This acid pulls the ammonia out of the complex as NH4+, allowing white silver chloride to precipitate ... 

Show how a silver-silver chloride electrode (silver in contact with solid AgCI and a solution of Cl- ions) and a hydrogen electrode can be used to measure (a) pH (b) pOH. 

In this cell, the following independent phases must be considered platinum, silver, gaseous hydrogen, solid silver chloride electrolyte, and an aqueous solution of hydrogen chloride. In order to be able to determine the EMF of the cell, the leads must be made of the same material and thus, to simplify matters, a platinum lead must be connected to the silver electrode. It will be seen in the conclusion to this section that the electromotive force of a cell does not depend on the material from which the leads are made, so that the whole derivation could be carried out with different, e.g. copper, leads. In addition to Cl- and H30+ ions (further written as H+), the solution also contains Ag+ ions in a small concentration corresponding to a saturated solution of silver chloride in hydrochloric acid. Thus, the following scheme of the phases can be written (the parentheses enclose the species present in the given phase) ... 

Since it is useful to know what state each reagent is in, we often designate the state in the equation. The modern practice is to add to the formula the designation in parentheses (s) for solid, (l) for liquid, (g) for gas, and (aq) for aqueous solution. Thus, a reaction of silver nitrate with sodium chloride in aqueous solution, yielding solid silver chloride and aqueous sodium nitrate, may be written as... 

Metals which form sparingly soluble salts will also respond to changes in the activity of the relevant anion provided the solution is saturated with the salt, e.g. for silver in contact with a saturated solution of silver chloride and containing solid silver chloride the electrode reaction is AgCl + e = Ag + Cl, and the electrode potential is given by ... 

Energy is needed to break the ionic bonds in the solid salt and energy is liberated forming hydration complexes like VI. We also break some of the natural hydrogen bonds in the water. The overall change in enthalpy is termed the enthalpy of solution, A// olutioni. Typical values are —207 kJmol-1 for nitric acid 34 kJmol-1 for potassium nitrate and —65.5 kJmol-1 for silver chloride. 

The silver-silver chloride electrode (Ag AgCl) is easily and cheaply made. Two silver electrodes are cleaned (see Section 9.1.1 above) and immersed in aqueous KCl solution (a concentration of 0.1 mol dm is convenient). Next, a potential of about 2 V is applied across them for c. 10 min, causing a thin outer film of silver chloride to develop on the positive electrode. Solid AgCl is formed by a two-step reaction, involving first the electro-formation of silver ion ... 

Silver-Silver Chloride Electrode. This reference electrode consists of a pure silver wire in a solution of KCl saturated with solid silver chloride. The electrode reaction is... 

The effect of the cyanine dye and of gelatin on the reaction rate shows that reduction of silver ions from solution is not the rate-controlling process. These influences of adsorbed components on the reaction rate speak against the concept that solution of the silver halide is the rate controlling process. Hence, the silver catalyzed reduction of silver chloride by hydroxylamine takes place substantially at the solid silver/ silver halide interface. 

The essential difference between the hydroxylamine reaction and the hydrazine reaction appears to be that silver nuclei are formed in the solution much more readily by hydrazine than by hydroxylamine. At sufficiently low pH and in the absence of copper, hydroxylamine does not readily form nuclei in the solution, and the catalytic reduction of the silver chloride occurs essentially at a solid interface with the silver nuclei. Hydrazine, on the other hand, readily forms nuclei in the solution and an important fraction of the total reaction involves the catalytic reduction of dissolved silver chloride. This would account for the well-known photographic properties of the two agents. Hydroxylamine is a cleanworking developer which, under proper conditions, yields little fog. Hydrazine shows much less selectivity and, although it develops an image, it also yields a relatively high fog density. 

Monammino-silver Chloride, [Ag(NH3)]Cl.—Biltz and Stollen-werk5 report the existence of the monammine, and state that sesqui-annuino- and nionammino-silver chloride form mixed crystals at 30° C., whilst Bodlilnder and Fittig8 have shown that a solution of silver chloride in aqueous ammonia contains the diammine in solution even when solid sesquiammine has separated from the liquid. A certain amount of doubt, therefore, surrounds the existence of the monammine and the diammine of silver chloride, although both appear to exist in aqueous ammoniaeal solution. 

The substance so produced, even after repeated crystallisation from water, still contains sodium chloride, with which it appears to be iso-morphous. The pure chloride is best prepared by treating a solution of the crude salt with solid sodium iodide, whereby it is transformed into the iodide. The iodide is recrystallised, and is then converted into the chloride by shaking the aqueous solution with freshly precipitated silver chloride. It crystallises in transparent cubes or in small glistening needles, and loses 2J molecules of water on heating to 120° C.3... 

After addition of 10 mL of dry dichloromethane, the yellow suspension is stirred magnetically for 3h. The solution is filtered away from the silver chloride through a Schlenk frit under an argon atmosphere into another dry Schlenk tube. To the yellow solution trilluoromethane sulfonic acid (0.058 mL, 0.66 mmol) (distilled at 43 °C in vacuo with an oil pump and stored under argon) is added by means of a micropipette. The mixture is stirred for 1 h. A colorless precipitate forms. Precipitation of the product is completed by addition of 10 mL of pentane. The mixture is centrifuged and the solution decanted. The remaining solid is washed twice with 10 mL of pentane, and dried under high vacuum for 4h. Yield 620 mg (89%). 

Crystals result when a solid substance is precipitated, under appropriate conditions, from a vapour or solution. It was suspected for a long time that the geometrical forms of crystals were due to a regular arrangement of atoms or molecules. Experimental verification of this idea was provided in 1912 by von Laue, who demonstrated that crystals showed interference phenomena with x-rays and, further, that it was possible to determine the structure of the crystal so accurately that sometimes the positions of atoms in the crystal could be given in one part in 100,000. It thus became possible to find the distance between the atoms in crystals, and it is important that it is not necessary to have well-formed crystals for this purpose. Nearly all substances are crystalline, even the silver chloride flocculent precipitate which is obtained when Cl" and Ag+ ions are brought together. 

Quantitative Determination.— Dissolve. 1 gm. of pnlits.sinin cyanide in water and dilute to IOO cc. Dilute 10 ee, of this solution with 90 cc. of water, add a grannie of sodium chloride, and titrate with dcciriormal silver nitrate solid inn until a permanent, whitish turbidity appears. 

Salts therefore, are prepared (1) from solutions of acids and bases by neutralization and separation by evaporation and crystallization (2) from solutions of two salts by precipitation where the solubility of the salt formed is slight (e.g., silver nitrate solution plus sodium chloride solution yields silver chloride precipitate [almost all as sulid], and sodium nitrate present in solution as sodium cations and nitrate anions [recoverable as sodium nitrate, solid by separation of silver chlondc and subsequent evaporation of the solution]) (3) from fusion of a basic oxide (or its suitable compound—sodium carbonate above) and an acidic oxide (or its suitable compound—ammonium phosphate), since ammonium and hydroxyl are volatilized as ammonia and water. Thus, sodium ammonium hydrogen phosphate... 

In other words, a net ionic equation shows only the net chemical change that occurs (Fig. 1.7). It shows that the Ag+ ions supplied by one solution combine with the Cl- ions supplied by the other solution and that these ions precipitate as solid silver chloride, AgCl. A net ionic equation focuses our attention on the actual process taking place and emphasizes the change. 

One simple experiment is to prepare a saturated solution of silver chloride. Then add a little solid silver chloride containing radioactive 112Ag to the solution. After a while, the solution itself is found to be radioactive. That observation indicates that exchange of silver ions is still taking place between the solid and the solution even though the solution is saturated. 

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