Creation and Measurement of a Saturated Solution --saturated solution of Aluminum Nitrate in Water Experiment, and Measurements: --pure, solid Al(NO3)3 is slowly added to exactly 100.0 ml of pure water until no more dissolves at exactly 25 °C. The measured density of the final solution is 1.37 g/ml, and the measured total mass of the final solution is 286.4 g. Definitions to Review: solvent, solution, solute, mass, volume, density, concentration, molarity

In Exercise 12.49, what is the allowed concentration of AgCl in g per 100 g of water? The Safe Drinking Water Act of 1974 established the maximum permitted concentration of silver ion at 0.05 ppm. What is the concentration of Ag+ in parts per million in a saturated solution of AgCl? (NOTE: 1 ppm = 1 mgofsolute/L of solution.)

If 55 mg of lead(II) sulfate is placed in 250 mL of pure water, does all of it dissolve? If not, how much dissolves?

A solution contains 0.00740 M calcium ion. A concentrated sodium fluoride solution is added dropwise to precipitate calcium fluoride (assume no volume change). a At what concentration of F does precipitate start to form? b When [F] = 9.5 104 M, what is the calcium-ion concentration? What percentage of the calcium ion has precipitated?

You make a saturated solution of NaCl at 25 C. No solid is present in the beaker holding the solution. What can be done to increase the amount of dissolved NaCl in this solution? (See Figure 13.11.) (a) Add more solid NaCl. (b) Raise the temperature of the solution. (c) Raise the temperature of the solution, and add some NaCl. (d) Lower the temperature of the solution, and add some NaCl. Figure 13.11 The temperature dependence of the solubility of some ionic compounds in wafer. The solubility of most ionic compounds increases with increasing temperature This is illustrated using NH4CI (ports b and c).

Which of the following pictures best represents an unsaturated solution of sodium chloride, NaCl? In these pictures, the dark gray spheres represent Na+ ions and the green spheres represent chloride ions. For clarity, water molecules are not shown.

A saturated solution of a slightly soluble electrolyte in contact with some of the solid electrolyte is said to be a system in equilibrium. Explain. Why is such a system called a heterogeneous equilibrium?

The weak electrolyte NH3(g) does not obey Henrys law. Why? O2(g) obeys Henrys law in water but not in blood (an aqueous solution). Why?

. Why does the amount of excess solid solute present in a solution not affect the amount of solute that ultimately dissolves in a given amount of solvent?

Given the measured final mass, and final density of the observed solution, what must have been the final volume of the solution?   Steps and explanation please.